Studies in Surface Science and Catalysis 31 PREPARATION OF CATALYSTS IV Scientific Bases for the Preparation of Heterogeneous Catalysts
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Studies in Surface Science and Catalysis Advisory Editors: B. Delman and J.T. Yates Vol. 31
PREPARATION OF CATALYSTSIV Scientific Bases for the Preparation of Heterogeneous Catalysts Proceedings of the Fourth International Symposium, Louvain-Ia-Neuve, September 1-4, 1986
Editors B. Delmon and P. Grange Catalyse et Chimie des Materiaux Divises, Groupe de Physico-Chimie Minerale et de Catalyse, Universite Catholique de Louvain, Louvain-Ia-Neuve, Belgium
P.A.Jacobs Centrum voor Oppervlaktescheikunde en Colloidale Scheikunde, Katholieke Universiteit Leuven, Heverlee, Belgium
and G. Poncelet Catalyse et Chimie des Materiaux Divises, Groupe de Physico-Chimie Minerale et de Catalyse, Universite Catholique de Louvain, Louvain-Ia-Neuve, Belgium
ELSEVIER Amsterdam - Oxford - New York - Tokyo
1987
ELSEVIER SCIENCE PUBLISHERS B.V. Sara Burgerhartstraat 25 P.O. Box 211, 1000 AE Amsterdam, The Netherlands Distributors for the United States and Canada:
ELSEVIER SCIENCE PUBLISHING COMPANY INC. 52, Vanderbilt Avenue NewYork,NY 10017,U.S.A.
Library of Congress Cataloging-in-Publication Data
Preparation of catalysts IV. (Studies in surface science and catalysis ; 31) "Proceedings of the IVth International Symposium on the Scientific Bases for the Preparation of Heterogeneous Catalysts"--Foreword. Includes bibliographies and index. 1. Catalysts--Congresses. 2. Heterogeneous catalysis --Congresses. I. Delman, Bernard. II. International Symposium on the Scientific Bases for the Preparation of Haerogeneous Catalysts (4th : 1986 : Louvain-la-Neuve, Belgium) Ill. Title: Preparation of catalysts 4. IV. Title: Preparation of catalysts four. V. Series. QD505.P68 1987 660.2'995 87-6868 ISBN 0-444-42796-1 (U.S.)
ISBN 0-444-42796-1 (Vol. 31) ISBN0-444-41801-6 (Series)
© Elsevier Science Publishers B.V., 1987 All rights reserved. No part of this publication may be reproduced, stored in a retrieval system or transmitted in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without the prior written permission of the publisher, Elsevier Science Publishers B.V.I Science & Technology Division, P.O. Box 330, 1000 AH Amsterdam, The Netherlands. Special regulations for readers in the USA - This publication has been registered with the Copyright Clearance Center Inc. (Ccq, Salem, Massachusetts. Information can be obtained from the CCC about conditions under which photocopies of parts of this publication may be made in the USA. All other copyright questions, including photocopying outside of the USA, should be referred to the copyright owner, Elsevier Science Publishers B.V., unless otherwise specified. Printed in The Netherlands
v CONTENTS Studies in Surface Science and Catalysis (other volumes in the series) Organizing Committee ~~~~
Acknowledgements Financial Support
XI XIII XV XVII XVIII
Preparation of metal distributions within catalyst supports M.S. Heise and J.A. Schwarz A study of some parameters in catalyst preparation and their influence on catalyst performance P.T. Cardew, R.J. Davey, P. Elliott, A.W. Nienow and J.P. Winterbottom
15
Impregnation of controlled-porosity silicas with hexachloroplatinic acid: parameters affecting Pt dispersion and location and support modification M.A. Martin Luengo, P.A. Sermon and K.S.W. Sing
29
Preparation and properties of the catalysts by a chemical mixing procedure F. Mizukami, S. Niwa, M. Toba, T. Tsuchiya, K. Shimizu, S. Imai and J. Imamura
45
Electrochemically controlled deposition-precipitation. A new method for the production of supported catalysts P.C.M. van Stiphout, H. Donker, C.R. Bayense and J.W. Geus
55
Preparation of highly dispersed, carbon supported, platinum catalysts D. Richard and P. Gallezot
71
Application of precursors of catalytically active materials on preshaped supports by impregnation with solutions of badly crystallizing compounds G.R. Meima, B.G. Dekker, A.J. van Dillen, J.W. Geus, J.E. Bongaarts. F.R. van Buren. K. Delcour and J.M. Wigman
83
Preparation and characterization of sol-gel based catalysts for the selective catalytic reduction of NO with NH~ H. Barten. F. Janssen, F. V.D.Kerkhof, R. Lefer1nk, E.T.C. Vogt, A.J. van Dillen and J.W. Geus
103
High selectivity of CVD Sn02/Si02 catalyst for oxidative dehydrogenation of ethyl benzene T. Hattori, S. Itoh, T. Tagawa and Y. Murakami
113
The preparation and characterization of vanadia supported rhodium catalysts Y.-J. Lin. R.J. Fenoglio. D.E. Resasco and G.L. Haller
125
Preparation of highly selective and abrasion-resistant thick shell catalysts for heterogeneously catalyzed exothermic oxidation reactions D. Arntz and G. Prescher
137
VI
The preparation and characterization of iron and vanadium oxide monolayer catalysts on Ti02 and Zr02 J.G. van Ommen, H. Bosch, P.J. Gellings and J.R.H. Ross
151
Interlayer accessibility in layered double-metal hydroxides F.A.P. Cavalcanti, A. Schutz and P. Biloen
165
Molybdenum catalyzed oxygen-transfer reactions. Heterogenization of homogeneous catalysts by using new dihydroxyboryl-substituted resins E. Tempesti, L. Giuffre, C. Mazzocchia, F. Di Renzo and P. Gronchi
175
Preparation of VO(HP04).0.5H20 and its transformation to the maleic anhydride catalyst lVO)2P207 J.W. Johnson, D.C. Johnston and A.J. Jacobson
181
Effect of ultrasonic treatment on the physico-chemical properties of Cr-Mo-O catalysts for methanol oxidation T.S. Popov, D.G. Klissurski, K.I. Ivanov and J. Pesheva
191
Modification of the textural and acid properties of A1P04 with sulfate anions J.M. Campelo, A. Garcia, D. Luna and J.M. Marinas
199
Processes of the formation of the active structures of the V-Mo-O catalysts for selective benzene oxidation M. Najbar, A. Bielanski, J. Camra, E. Bielanska, W.Wal, J. Chrzaszcz and W. Ormaniec
217
Preparation chemistry of V-Ti-O mixed oxides. Comparison of coprecipitation grafting and impregnation methods F. Cavani, G. Centi, F. Parrinello and F. Trifiro
227
Cation effects in the preparation of and catalysis by heteropoly oxometallates J.B. Moffat
241
Preparation of heteropolyvanadophosphate catalysts supported by silica and an active carbon fibre felt and their catalytic properties Y. Kera, Y. Ishihama, T. Kawashima, T. Kamada, T. Inoue and Y. Matsukaze
259
Catalyst preparation via hydrous metal oxide ion-exchangers H.P. Stephens and R.G. Dosch
271
Photo-assisted deposition of noble metals: investigation of a new route for metallic and bimetallic catalyst preparation J.-M. Herrmann, J. Disdier, P. Pichat and C. Leclercq
285
The preparation and the characterization of some ternary titanium oxide photocatalysts R.I. Bickley, T. Gonzalez-Carreno and L. Palmisano
297
Development of methods for regulating the charged surface groups of y-A1203 in aqueous solutions. L. Vordonis, A. Akratopulu, P.G. Koutsoukos and A. Lycourghiotis
309
VII
Understanding the morphological transformations that occur in the preparation of alumina supports W.C. Conner, E.L. Weist and L.A. Pedersen
323
Microporous amorphous alumina of a zeolitic type for catalytic reactions with methanol G. Tournier, M. Lacroix-Repellin. G.M. Pajonk and S.J. Teichner
333
The relevance of kneading and extrusion parameters in the manufacture of active porous aluminas from pseudoboehmites A. Danner and K.K. Unger
343
Amorphous alloys as catalysts or catalyst precursors M. Shibata and T. Masumoto
353
Nickel catalysts derived from eutectic and pro-eutectic nickel al uminum a11 oy C.S. Brooks
375
Amorphous metal alloys as precursors in catalyst preparation ammonia synthesis catalysts from amorphous Ni-Zr systems E. Armbruster. A. Baiker, H.J. Guentherodt, R. Schloegl and B. Walz
389
The production of iron- and nickel-copper alloy catalysts by means of catalytic decomposition of the gaseous metal carbonyls A.F.H. Wielers. C.M.A.M. Mesters. G.W. Koebrugge, C.J.G. van der Grift and J.W. Geus
401
Reactivity and structure of metal catalyst particles C. Lee, S. Gao and L.n. Schmidt
421
Influence of lanthanum oxide on the surface structure and CO hydrogenation activity of supported cobalt catalysts J.S. Ledford, M. Houalla, L. Petrakis and D.M. Hercules
433
Effects of the technique for the preparation of supported cobalt catalysts on selectivity in the Fischer-Tropsch synthesis F. Liu
443
Selective Fischer-Tropsch catalysts containing iron and lanthanide oxides B.G. Baker and N.J. Clark
455
Titania supported iron-ruthenium catalysts for Fischer-Tropsch synthesis Lin Liwu, F.J. Berry. Du Hongzhang. Liang Dongbai. Tang Renyuan, Wang Chengyu and Zhang Su
467
Preparation of carbon-supported K-Fe-Mn and Fe-Mn catalysts using carbonyl clusters J. Venter, M. Kaminsky, G.L. Geoffroy and M.A. Vannice
479
Catalytic activity of carbon supported catalysts for CO-hydrogenation and their preparation by oxidative decomposition of Fe(CO)5 U. Peters, H. Greb, R. Jockers and J. Klein
493
On the mechanism of formation of colloidal monodisperse metal boride particles from reversed micelles composed of CTAB I-hexanol - water I. Ravet, J. B.Nagy and E.G. Derouane
505
VIII
Preparation of nickel catalyst from nickel containing chrysotile Y. Ono, N. Kikuchi and H. Watanabe
519
Parameters influencing the preparation and characterization of sodium on zeolite catalysts L.R.M. Martens, W.J.M. Vermeiren, P.J. Grobet and P.A. Jacobs
531
Preparation of bimodal alumina and other refractory inorganic oxides-suitable supports for hydrotreating catalysts K. Onuma
543
Hydrotreating NiMo/sepiolite catalysts: influence of catalyst preparation on activity for HDS, hydrogenation and chain isomerization reactions F.V. Melo, E. Sanz, A. Corma and A. Mifsud
557
New uranium-based hydrotreatment catalysts G. Agostini, M.J. Ledoux, L. Hilaire and G. Maire
569
On the addition of various metals as inorganic salts or organometallic complexes to a MoS2-yA1203 hydroprocessing catalyst preparation, characterization and hydrogenation activity A. Wambeke, H. Toulhoat, J.P. Boutrois, J. Grimblot and J.P. Bonnelle
581
Control of concentration profiles by rational preparation of pelleted hydrodesulfurization catalysts J.L.G. Fierro, P. Grange and B. Delmon
591
Influence of the activation procedure on the nature and concentration of the active phase in HDS catalysts R. Prada Silvy, J.L.G. Fierro, P. Grange and B. Delmon
605
Palladium catalysts for selective gas-phase hydrogenation of phenol to cyclohexanone J.R. Gonzalez-Velasco, J.I. Gutierrez-Ortiz, M.A. GutierrezOrtiz, M.A. Martin. S. Mendioroz. J.A. Pajares and M.A. Folgado
619
Aluminum-oxide-pillared montmorillonite: effect of hydrothermal treatment of pillaring solution on the product structure J.P. Sterte and J.-E. Otterstedt
631
Synthesis and properties of cross-linked hydroxy-titanium bentonite Sun Guida. Van Fushan. Zhu Huihua and Liu Zhonghui
649
Effect of hydroxy-aluminum polymeric cations on acidity of crosslinked hydroxy-aluminum smectities Sun Guida. Van Fusan. Sun Dehai and Liu Zhonghui
659
Controlled preparation with three different supported bimetallic J.B. Michel and J.T.
669
of monodisperse bimetallic Pd-Au colloids microstructures and their use in preparing catalysts Schwartz
Controlled surface reactions for the preparation of different types of alumina supported Sn-Pt catalysts E. Kern-Talas, M. HegedUs, S. G6bolos, P. Szedlacsek and J. Margitfalvi
689
Preparation and properties of anchored Pt-Mo/Si0 2 bimetallic catalyst Yang Yashu, Guo Xiexian. Li Huimin. Deng Maicun and Lin Zhiyin
701
Controlled preparation of bimetallic hydrogenation catalysts O.A. Ferretti. L.C. Bettega de Pauli, J.P. Candy, G. Mabillon and J.P. Bournonville
713
Characterization and catalytic properties of cobalt FischerTropsch catalysts prepared by chemical reduction and used in a liquid phase C. Bechadergue-Labiche. S. Maille. P. Canesson. M. Blanchard and D. Vanhove
725
Comparison of the quantitative studies by STEM of hydrated hydroxycarbonates and related mixed oxides catalysts for CO hydrogenation to alcohols R. Szymanski. Ch. Travers. P. Chaumette._ Ph. Courtyand D. Durand
739
Preparation of multicomponent catalysts for the hydrogenation of carbon monoxide via hydrotalcite-like precursors S. Gusi. F. Pizzoli:lt. Trifiro. A. Vaccari and G. Del Piero
753
Preparation and characterization of copper/zinc oxide/alumina catalysts for methanol synthesis E.B.M. Doesburg. R.H. HBppener. B. de Koning. Xu Xiaoding and J.J.F. Scholten
767
The activity of coprecipitated Cu. Zn catalysts for methanol synthesis B.S. Rasmussen. P.E. H~jlund Nielsen. J. Villadsen and J.B. Hansen
785
The preparation and characterization of sequentially precipitated and coprecipitated nickel-alumina catalysts and a comparison of their properties H.G.J. Lansink Rotgerink. J.G. van Ommen and J.R.H. Ross
795
Preparation and characterization of thick layers of semiconductors oxides for gas chemisorption and detection C. Lucat. F. Menil. M. Destriau. J. Salardenne and J. Portier
809
Standardization of catalyst test methods by the Committee on Reference Catalyst of the Catalysis Society of Japan T. Hattori. H. Matsumoto and Y. Murakami
815
Europt-l : the first platinum on silica reference catalyst G.C. Bond and P.B. Wells
827
List of Participants Author Index
841 865
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XI
STUDIES IN SURFACE SCIENCE AND CATALYSIS Advisory Editors: B. Delmon, Universite Catholique de Louvain, Louvain-Ia-Neuve, Belgium J.T. Yates, University of Pittsburgh, Pittsburgh, PA, U.S.A.
Volume 1 Preparation of Catalysts I. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the First International Symposium, Brussels, October 14-17,1975 edited by B. Delmon, P.A. Jacobs and G. Poncelet Volume 2 The Control of the Reactivity of Solids. A Critical Survey of the Factors that Influence the Reactivity of Solids, with Special Emphasis on the Control of the Chemical Processes in Relation to Practical Applications by V.V. Boldyrev, M. Bulens and B. Delmon Volume 3 Preparation of Catalysts II. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the Second International Symposium, Louvain-Ia-Neuve, September 4-7, 1978 edited by B. Delmon, P. Grange, P. Jacobs and G. Poncelet Volume 4 Growth and Properties of Metal Clusters. Applications to Catalysis and the Photographic Process. Proceedings of the 32nd International Meeting of the Societe de Chimie Physique, Villeurbanne, September 24-28, 1979 edited by J. Bourdon Volume 5 Catalysis by Zeolites. Proceedings of an International Symposium, Ecully (Lyon), September 9-11, 1980 edited by B. Imelik, C. Naccache, Y. Ben Taarit, J.C. Vedrine, G. Coudurier and H. Praliaud Volume 6 Catalyst Deactivation. Proceedings of an International Symposium, Antwerp, October 13-15, 1980 edited by B. Delmon and G.F. Froment Volume 7 New Horizons in Catalysis. Proceedings of the 7th International Congress on Catalysis, Tokyo, June 30-July 4, 1980. Parts A and B edited by T. Seiyama and K. Tanabe Volume 8 Catalysis by Supported Complexes by Yu.1. Yermakov, B.N. Kuznetsov and V.A. Zakharov Volume 9 Physics of Solid Surfaces. Proceedings of a Symposium, Bechyi\e, September 29-0ctober 3, 1980 edited by M. Laznicka Volume 10 Adsorption at the Gas-Solid and Liquid-Solid Interface. Proceedings of an International Symposium, Aix-en-Provence, September 21-23, 1981 edited by J. Rouquerol and K.S. W. Sing Volume 11 Metal-Support and Metal-Additive Effects in Catalysis. Proceedings of an International Symposium, Ecully (Lyon), September 14-16,1982 edited by B. Imelik, C. Naccache, G. Coudurier, H. Praliaud, P. Meriaudeau, P. Gallezot, G.A. Martin and J.C. Vedrine Volume 12 Metal Microstructures in Zeolites. Preparation - Properties - Applications. Proceedings of a Workshop, Bremen, September 22-24, 1982 edited by P.A. Jacobs, N.I. Jaeger, P. Jiru and G. Schulz-Ekloff Volume 13 Adsorption on Metal Surfaces. An Integrated Approach edited by J. Benard Volume 14 Vibrations at Surfaces. Proceedings of the Third International Conference, Asilomar, CA, September 1-4, 1982 edited by C.R. Brundleand H. Morawitz
XII Volume 15 Heterogeneous Catalytic Reactions Involving Molecular Oxygen by G.I. Golodets Volume 16 Preparation of Catalysts III. Scientific Basesfor the Preparation of Heterogeneous Catalysts. Proceedings of the Third International Symposium, Louvain-Ia-Neuve, September 6-9, 1982 edited by G. Poncelet, P. Grange and P.A. Jacobs Volume 17 Spillover of Adsorbed Species. Proceedings of an International Symposium, LyonVilleurbanne, September 12-16, 1983 edited by G.M. Pajonk, S.J. Teichner and J.E. Germain Volume 18 Structure and Reactivity of Modified Zeolites. Proceedings of an Intenational Conference. Prague,July 9-13. 1984 edited by P.A. Jacobs, N.1. Jaeger. P. Jiru, V.B. Kazansky and G. Schulz-Ekloff Volume 19 Catalysis on the Energy Scene. Proceedings of the 9th Canadian Svmposium on Catalysis, Quebec. P.Q., September 30-0ctober 3. 1984 edited by S. Kaliaguine and A. Mahay Volume 20 Catalysis by Acids and Bases. Proceedings of an International Symposium, Villeurbanne (Lyon). September 25-27. 1984 edited by B. Imelik, C. Naccache, G. Coudurier, Y. Ben Taarit and J.C. Vedrine Volume 21 Adsorption and Catalysis on Oxide Surfaces. Proceedings of a Symposium, Uxbridge, June 28-29. 1984 edited by M. Che and G.C. Bond Volume 22 Unsteady Processes in Catalytic Reactors by Yu.Sh. Matros Volume 23 Physics of Solid Surfaces 1984 edited by J. Koukal Volume 24 Zeolites: Synthesis. Structure, Technology and Application. Proceedings of an International Symposium, Portorof-Portorose, September 3-8, 1984 edited by B. Driaj, S. Hol:evar and S. Pejovnik Volume 25 Catalytic Polymerization of Olefins. Proceedings of the International Symposium on Future Aspects of Olefin Polymerization. Tokyo, July 4-6, 1985 edited by T. Keii and K. Soga Volume 26 Vibrations at Surfaces 1985. Proceedings of the Fourth International Conference. Bowness-on-Windermere. September 15-19.1985 edited by D.A. King. N.V. Richardson and S. Holloway Volume 27 Catalytic Hydrogenation edited by L. Cerveny Volume 28 New Developments in Zeolite Science and Technology. Proceedings of the 7th International Zeolite Conference. Tokyo, August 17-22, 1986 edited by Y. Murakami, A. Iijima and J.W. Ward Volume 29 Metal Clusters in Catalysis . edited by B.C. Gates. L. Guczi and H. KnOzinger Volume 30 Catalysis and Automotive Pollution Control. Proceedings of the First International Symposium, Brussels, September 8-11, 1986 edited by A. Crucq and A. Frennet Volume 31 Preparation of Catalysts IV. Scientific Basesfor the Preparation of Heterogeneous Catalysts. Proceedings of the Fourth International Symposium, Louvain-Ia-Neuve. September 1-4. 1986 edited by B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet Volume 32 Thin Metal Films and Gas Chemisorption edited by P. Wissmann Volume 33 Synthesis of High-silica Aluminosilicate Zeolites by P.A. Jacobs and J.A. Martens
XIII
ORGANIZING COMMITTEE President
Prof. B. DELMON, Universite Catholique de Louvain
Executive Chairmen
Dr. P. GRANGE, Universite Catholique de Louvain Dr. P.A. JACOBS, Katholieke Universiteit Leuven Dr. G. PONCELET, Universite Catholique de Louvain
Scientific Committee
Dr. U. BLINDHEIM, Senter for Industriforskning, Norway Dr. G. DE CLIPPELEIR, Labofina S.A., Belgium Prof. B. DELMON, U.C.L., Belgium Prof. E. DEROUANE, Facultes Universitaires de Namur, Belgium Dr. J. DETHY, Catalysts and Chemicals Europe, Belgium Dr. P. ENGELHARD, Total - Compagnie Fran~aise de Raffinage, France Dr. V. FRANZEN, Lonza, Switzerland Prof. J. GARCIA DE LA BANDA, Instituto de Catalisis, Madrid, Spain Dr. P. GRANGE, U.C.L. Belgium Dr. W. HOLDERICH, B.A.S.F., W. Germany Dr. P. JACOBS, K.U.L., Belgium Dr. K. KOCHLOEFL, SUd-Chemie, W. Germany Dr. C. KOMODROMOS, British Gas, England Dr. A. LECLOUX, Solvay &Cie, Belgium Dr. J. MAGNUSSON, Katalistiks, The netherlands Dr. C. MASQUELIER, U.C.B., Belgium Prof. J. B.NAGY, Facultes Universitaires de Namur, Belgium Dr. S. NOTARI, Assoreni, Italy Dr. G. PONCELET, U.C.L., Belgium
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xv FOREWORD
We are pleased to present here the Proceedings of the IVth International Symposium on the "Scientific Bases for the Preparation of Heterogeneous Catalysts". The series started rather modestly in 1975, but since then it has attracted increasing scientific attention. The call for papers for this fourth symposium resulted in the submission of no less than 210 extended abstracts, an appreciable proportion of which stemmed from industrial research laboratories. This large number posed a great problem for the Scientific Committee, made up of 19 scientists with international reputations in the field and representing industry as well as academia. It seemed a hopeless task to make a selection from this wealth of proposals. It also became clear that quality could not be the sole yardstick in judging the suitability of the papers, for reasons we hope will become clear in what follows. A much appreciated feature of this series of symposia has been that the Scientific Committee really is responsible for the scientific content of the Proceedings; during a one-day meeting of the Committee, some 60 abstracts (the maximum possible number) were selected, unfortunately a mere 29% of the contributions submitted. Thanks to this selection, most of the accepted abstracts fell within the scope of the symposium as initially defined. The few exceptions were justified by their originality or outstanding quality. It had been stressed when announcing the scope that new results should have a clear bearing on catalysts either already in industrial use or under development, that contributions dealing with new preparation methods would be particularly welcome, and that new routes of catalyst manufacture should be explored. It had also been hoped that new ways of manufacturing catalysts could be presented and discussed. Not all of these expectations were fulfilled, but relevance to problems of practical applications and originality were certainly the hallmarks of the meeting. The Scientific Committee and the local Committee attemped to organise, as they had done for the previous symposia, self-contained sessions on the basis of the selected papers. Accordingly, sessions devoted to catalysts for selective oxidation, hydrodesulphurisation, Fischer-Tropsch catalysis, methanol synthesis, and metal supported catalysts, including new trends in catalyst preparation, were arranged. The programme comprised two parallel sessions, the major concern being the organisation of a well-balanced programme around the chosen topics. In the light of this, the many authors whose abstracts were not selected should not infer from their refusal that the quality of their work was being questioned.
XVI
Against.this background, 4 experts were invited to deliver plenary lectures or extended communications, and plenty of time was left for discussions. The conference chairmen and editors considered, therefore, that these discussions were worth being included in these proceedings. Approximately 350 participants attended the IVth Symposium. Over 40% were affiliated with industry, confirming again the need for this kind of event, as felt by all scientists working in the field. This will be a stimulus to the local organisers of future conferences when they start with the preparations for the Vth symposium. When initiating this series of symposia, the organisers hoped to begin bridging the gap between those scientists who are directly involved in the day-to-day manufacture or uses of catalysts and those whose contribution to this area is more on the theoretical side. Modest but substantial progress has been made in this direction during the course of these symposia. We shall strive to speed up this progress, with the help of the next scientific committee and the members of past ones, who have already helped us so much.
B. DELMON P. GRANGE P.A. JACOBS G. PONCELET
XVII
ACKNOWLEDGEMENTS The Organizing Committee thanks Mgr Ed. Massaux, Rector of the Universite Catholique de Louvain, who, for the third time, made it possible for this symposium to be held in Louvain-la-Neuve, and gave access to the various facilities provided by the University. We also thank the University Authorities and, in particular, Professor E. Buyse, for the welcome address to the participants. Our deep gratitude goes to the members of the Scientific Committee who accomplished very efficiently their difficult task of selecting the communications, and accepted to act as session chairmen during the symposium. The minisymposium on Catalyst Normalization was handled by Professor E. Derouane. We are most indebted to him and also to Professor G.C. Bond, Professor T. Hattori and Professor J.W.E. Geus for leading brilliantly this session. The organizers are very pleased to acknowledge the contributions of Professor J.A. Schwarz, Professor J.B. Moffat and Dr. Shibata who delivered most stimulating plenary lectures. The Organizing Committee also acknowledges the authors of the 200 papers submitted, both those who presented a paper and those whose paper could not be retained, mostly due to time limitation. Our congratulations for the perfect achievement and our gratitude are addressed, as usual, to the hostesses of the REUL (Relations Exterieures de l'Universite Catholique de Louvain), headed by Mrs F. Bex, and to Mr J. Therer (Service du Logernent), for their enthusiasm and efficiency. It is a special pleasu~e to thank them for this symposium after similar thanks for the preceding ones. The Organizing·Committee wishes to acknowledge all the persons from the "Groupe de Physico-chimie Minerale et de Catalyse", and from the "Laboratorium voor Oppervlakte Chemie", K.U. Leuven, who worked for the success of the symposium, in particular, A. Arteaga, D. Balloy, P. Berteau, E. Churin, Chr. Dhayer, T. Machej, M.A. Martin-Luengo, R. Prada Silvy, P. Ruiz, M. Ruwet, L. wang, B. Yasse, B. Zhou; A.Boden, J. Geerts, L. Jacobs, S. Pelgrims, J. Perez, P.Tastenhove, M. Tielen. Finally, special thanks are due to the team of secretaries, F. Somers, J. Liagre and especially P. Theys. who. right from the beginning up to the end. took care of the most thankless part of the organization of the symposium.
XVIII
FINANCIAL SUPPORT The organizers gratefully acknowledge the "Fonds National de la Recherche Scientifique" and the "Minist~re de 1'Education Nationale et de la Culture Fran9aise" for having faithfully offered their financial guarantee for the organization of the IVth Symposium. The following companies and catalyst manufacturers have largely contributed to the success of this symposium by providing financial support. Most of them had already supported one or several of the previous symposia. The organizers are much indebted to them for their generosity. AKZO Chemie, Ketjen Catalysts (The Netherlands) Catalysts and Chemicals Europe (Belgium) Condea Chemie (W. Germany) Degussa Hanau (W. Germany) Dow Chemicals (Nederland) B.V. (The Netherlands) Haldor Tops¢e (Denmark) Harshaw Chemie B.V. (The Netherlands) Imperial Chemical Industries (United Kingdom) Labofina S.A. (Belgium) Lonza A.G. (Switzerland) Metallurgie Hoboken-Overpelt (Belgium) Nederlandse Unilever Bedrijven B.V. (The Netherlands) Norton Chemical Process Products (Europe) Ltd. (United Kingdom) Shell International Petroleum Company Ltd. (United Kingdom) Societe Nationale Elf Aquitaine (France) Solvay et Cie (Belgium) SUd-Chemie A.G. (W. Germany) The P.Q. Corporation (U.S.A.) The organizers also thank Harshaw, and in particular Mr. Brull, Director of Marketing and Sales, for generously offering the conference folders.
B, Delmon, P, Grange, P,A, Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
PREPARATION OF METAL DISTRIBUTIONS WITHIN CATALYST SUPPORTS M.S. Heisel and J.A. Schwarz 2 lCurrent Address W.R. Grace and Co. 2Department of Chemical Engineering and Materials Science, Syracuse University, Syracuse, NY 13244 SUMMARY The objective of catalyst design is to obtain the optimum metal profile for a particular reaction system. This is accomplished by the addition of certain ingredients to the impregnating solution, which are selected largely on the basis of empirical evidence. This paper proposes a classification scheme for these ingredients founded on three predominant interfacial effects. The first class of ingredients affects the electrostatics of the solution near the surface of the support. Ingredients in the second class alter the pH of the solution and consequently the potential of the support surface. The third class includes ingredients that adsorb onto the support and compete with the catalytic metal for adsorption sites. The conceptual classification scheme developed allows one to predict adsorption profiles for both uniform and nonuniform metal distributions. The concepts are completely general and thus provide a theoretical as well as a practical basis for the preparation of catalysts. INTRODUCTION Studies of the relationship between catalytic performance and metal profiles have been primarily directed towards the oxidation of carbon monoxide and hydrocarbons in automobile exhaust. Improvements in the activity and poison resistance of oxidation catalysts have been achieved by modifying the depth of the metal impregnation in the catalyst support (1-4). One way to alter the distribution of metal in the support is to add ingredients to the impregnating solution. Maatman (5) showed that the impregnation profile of hexachloroplatinic acid on alumina could be changed from an eggshell profile to a uniform profile by adding HC1, HN03, or various inorganic nitrates to the impregnating solution. Similarly, Benesi, Curtis, and Studer (6) demonstrated that the adsorption profiles of metal cations could be altered by changing the pH of solution. The first comprehensive study on the effects of adding various chemical ingredients to the impregnating solution was performed by Shyr and Ernst (7). They obtained an eggshell profile for the adsorption of hexachloroplatinic acid on gamma-alumina in the absense of other ingredients. The individual addition of fourteen salts and acids produced nine distinct adsorption profiles. This illustrates the diversity which can be achieved through the
2
addition of ingredients to the impregnating solution, and reflects the need for an understanding of the physical and chemical phenomena involved in the impregnation process. EXPERIMENTAL This section describes two experimental techniques that were employed to determine the metal profile and the effects of added ingredients on this profile. These procedures will allow for a qualitative validation of the classification scheme developed in the folloWing section. Impregnation Profile Experiments The impregnation procedure shown in Figure 1 involves contacting a smooth end of a dried gamma-alumina pellet with a solution of hexach1oroplatinic acid and the ingredient under study. The solution is drawn up by capillarity and the ingredients are adsorbed onto the outer channels or pores of the pellet.
r
IMPREGNATION
~,~ SOLUTION
CALCINATIO'I
4i1O't
J
I'tCIllCJW'HIC NEGATIII£S
1
MICROQENS/lllMETER TRACINGS
.-
Fig. 1. Impregnation technique including steps (a) impregnation of alumina pellets with solution containin9 the active ingredient,(b) calcination of the platinum/alumina catalyst, (c) photography of the calcined sample, and (d) transmission results from the microdensitomer tracing. Solutions of hexach1orop1atinic acid were prepared by dissolving H2PtC16'(6H20), obtained from Eng1ehard Industries in deionized and distilled water. Experiments were performed in the dark, since significant decomposition of chloroplatinate was observed in room light in less than an hour. The support used was cylindrical gamma-alumina extrudate characterized as follows: 0.3 cm, diameter; 190 m2/g, surface area; 0.68 cm 3/g, pore volume; 2.25x10-6 cm, pore radius.
3
The adsorbed platinate colors the white support yellow, which turns dark gray upon calcination. Calcination is carried out for four hours at 450 degrees C, with a low initial heating rate to minimize metal displacement in this step. The intensity of the color of the support is directly related to the amount of platinum adsorbed. Photographic negatives of the calcined pellets were then taken, which show adsorbed platinum as white. A scanning microdensitometer was then used to analyze the negatives. The extent of transmission was plotted as a function of axial distance from the dipped end of the pellet. The percent transmission is directly related to the platinum concentration and the area under the curve is proportional to the total amount of platinum adsorbed. In all experiments, 20.0 ml of solution was contacted with the alumina pellets for four hours. Additional ingredients were checked to make sure that they did not color the support upon calcination. Ingredients which did color the support (i.e. large organic acids) were reduced in flowing hydrogen for four hours. The photography was done with a platinum standard and a control (an unimpregnated pellet). This ensured that the experimentally measured profiles would be unaffected by minor differences in the transmission results of the photographic negatives. For each experiment, four pellets were scanned and the tracings were averaged to obtain a composite profile. Adsorption Experiments Experiments were run at the pore-filling time of twenty minutes with no prior heat treatment. One hundred mg of dry alumina pellets was added to various impregnating solutions while the platinum concentration in solution was monitored with a Beckman DB-GT spectrophotometer. Ultraviolet scans were run at 262 nm for the platinum (IV) complex (8). The amount adsorbed was determined by the difference between initial and final concentrations. CLASSIFICATION SCHEME It is proposed that ingredients added to the impregnating solution can be classified according to their effect on three interfacial phenomena. The first class of ingredients consists of simple inorganic electrolytes such as NaN03' NaCl, and CaC12' which affect the electrostatics at the solution-surface interface. The second class of ingredients includes simple inorganic acids and bases such as HC1, HN03' and NaOH, which affect the pH of the system. These compounds alter the chemistry of the surface by changing the surface potential. The interfacial effects associated with the first two classes are not the result of specific adsorption. Instead, the affinity of the metal ion for the
4
surface is altered by changing the number of available surface sites in the case of class two and by changing the accessibility of the metal to those sites in the case of class one. The third type of ingredient is one that can compete with the metal ion for possible adsorption sites. Although many compounds will adsorb onto the surface, the strongest and most effective are those that contain hydroxyl, carboxyl, and phosphoryl groups. If this type of ingredient is added to the impregnating solution. it will affect the metal adsorption in a chromatographic manner. These ingredients can also introduce significant pH and electrostatic effects into the system. Class 1 Ingredients Simple inorganic salts such as NaND3. NaCl. and CaC12 do not adsorb strongly enough on alumina to compete with the platinum ion for adsorption sites (9). It is therefore apparent that the cations and anions of these salts have a higher affinity for the aqueous phase. They modify the adsorption of platinum by altering the charge distribution near the surface of the support. Adsorption experiments were run at an initial platinum concentration of 5.4xlO-4M. With no Class 1 additions, approximately 0.55 wt% platinum was deposited on the alumina. Four different Class 1 ingredients were added to the impregnating solution: sodium chloride, sodium nitrate, calcium chloride, and calcium nitrate. Figure 2 shows that the addition of NaCl and NaN03. univalent Class 1 ingredients. produces similar effects on platinum adsorption. Since at any given concentration these 1:1 electrolytes introduce the same amount of electrostatics into the solution, they should produce the same effect on the amount of platinum adsorbed. The effects of the 2:1 electrolytes, CaC12 and Ca(ND3)2, are shown in Figure 3. As expected. these ingredients show similar results. For a given concentration, the 2:1 electrolytes show much less platinum adsorption than the 1:1 electrolytes. The 2:1 electrolytes do not show any significant platinum adsorption up to 1.OxlO-1M. This discrepancy can be accounted for by Poisson-Boltzmann theory. The extension of the electric field of the surface into the bulk solution is determined by the ionic strength of the solution, varying inversely with its square root. Because the ionic strength of a 2:1 electrolyte is three times that of a 1:1 electrolyte for a given concentration. a 2:1 electrolyte is a more effective site blocking agent. Figure 4 shows the dependence of the amount of platinum adsorbed on the ionic strength of solution for Class 1 ingredients.
5
30 o NoCI I::. NoN~
o
o
I::. 0 Q>
Fig. 2. Effect of univalent Class 1 ingredients on the amount of platinum adsorption on gamma-alumina pellets. Initial hexachloroplatinic acid concentration of 5.46xlO- 4 molar.
30 • cocI 2 4CO(N~12
o o
I
2
4
6
8
10
ELECTROLYTE CONCENTRATlON,co, .IO'(MOLARI
Fig. 3. Effect of divalent Class 1 ingredients on the amount of platinum adsorption on gamma-alumina pellets.
6
o'--_.......
_ - - l . . _ - - - L_ _.L.-.......Q)~
08
0.7
0.6
0.5
0.4
0.3
PLATINUM ION ACTIVITY COEFFIClENT,op
Fig. 4. Relation of the mean activity coefficient of the platinum ion in solution and the amount of platinum adsorbed on gamma-alumina pellets. Initial hexachloroplatinic acid concentration of 5.46xlO-4 molar. The experimental profiles for Class 1 ingredients are shown in Figures Sa-Sc. The initial concentration of platinum was fixed at 2.SxlO- 3M. Figure Sa shows the platinum standard, no Class 1 ingredient added. As NaN03 is added to the solution, the amount of platinum adsorbed decreases and the profiles became uniform. The electrolytes effectively decrease the number of active sites on the alumina surface by electrically screening them from the bulk solution. This effort occurs uniformly down the length of the support pore, causing the platinum coating to be thinner and thus extend deeper into the pore.
7
1.0
1.0 (0)
(bl
~0.8
0.8
0:
UJ
§0.6
0.6
..J
oct
~ 04
04
~
u
oct
~ 0.2 0.00
0.2
01
0.0
0.4
0
0./
0.2
PORE AXIAL LENGTH
1.0 (e
~
oct
I
0.8
~
u 0.6 ..J
~ 2 0.4 I-
u
oct
a:
"- 0.2
0.2
0.3
0.4
AXIAL LENGTH
Fig. 5. Experimental platinum adsorption profiles produced by impregnating gamma-alumina pellets with an initial hexachloroplatinic acid concentration of 2.5xlO-3 molar and various initial concentrations of sodium nitrate: (a) C02 = 0.0 - platinum - platinum standard, (b) C02 = 2.5xlO-3 molar, and C02 = 2.5xlO-2 molar. Class 2 Ingredients Work done by Maatman (5) showed that uniform profiles could be obtained for platinum deposition from hexachloroplatinic acid on an alumina support by adding simple inorganic acids such as HCl and HN03' It was assumed that the anions competed with the platinum ion for adsorption sites. However, no significant chloride or nitrate ion adsorption can be measured on alumina. Since these anions are not binding to the surface, the pH and electrostatics in the system are affecting the platinum distribution.
8
Figure 6 shows the amount of platinum adsorption as a function of the pH of the solution. The amount of platinum adsorbed drops off rapidly with the addition of base to the system. The zero point of charge occurs at a pH of 8.2 (10). No adsorption occurs past this point. Although the alumina surface is positively charged below a pH of about eight, at a pH of three or less, the dissolution of alumnina is significant (11). High concentrations of acids decrease the amount of platinum adsorbed by decreasing the number of active sites on the alumina surface. As expected, the addition of HCl and HN03 show the same results since they both introduce the same pH and electrostatic effects to the solution.
E0 ~
~40 III
0 NoOH 0 HCl
Q.l
g
/:1 HN~
.5- 0 CJ:) 3 z
0
i=
ft20 0
~
0 <1 ~
:::>
10
z
i=
<1 ...J
Q.
0 0
4
6
8
10
12
pH
Fig. 6. Effect of acidic and basic Class 2 ingredients on the amount of platinum adsorbed on gamma-alumina pellets. Initial hexachloroplatinic acid concentration of 5.46xlO-4 molar.
9
Figures 7a-7d shows the effects on the platinum profile caused by the addition of base to the solution. Up to the neutralization point, a hydroxide ion concentration of 5xlO- 3M, the profiles resemble that of the platinum standard. The height and the length of penetration decrease slightly with increasing base additions. This corresponds to a gradual decrease in the area under the profile, which represents a decrease in the amount of platinum adsorbed. At the acid-base neutralization point, Figure 7c, the shape of the adsorption profile changes rapidly. The shape changes from an eggshell to a linearly decreasing profile. The length of penetration and the amount adsorbed have decreased. Well past the neutralization point, at a concentration of 7.0xlO- 3M, the height of the platinum profile has decreased by a factor of four. A very small amount of platinum has been deposited at the pore mouth. At a concentration of added hydroxide of 1.OxlO- 2M, no platinum is adsorbed.
10
1,0 (bl
(a)
wO,8
08
~ a:
w 6°6 u
0,6
.
...J
0,4
~0,4
>= ~
If 0,2
0,2
0.° 0
0,/
0,0
02
°
01
PORE AXIAL LENGTH
I.
Idl
(c)
0,8 0,6
;J. z
0,4 -
Q04 tu
.
e: 0.2
0,2
0,0
°
00 0,1 0.4 PORE AXIAL LENGTH
0,3
°
0,2
03
0,4
Fig. 7. Experimental platinum adsorption profiles produced by impregnating gamma-alumina pellets with an initial hexachloroplatinic acid concentration of 2.5xlO- 3 molar and various initial concentrations of sodium hydroxide: (a) C02 = 0.0 - platinum standard, (b) C02 = 4.5xlO- 3 molar, (c) C02 = 5.0xlO- 3 molar, and (d) C02 = 7.0xlO- 3 molar.
10
1.0
(0)
1.0
(bl
10 -
'"cr:
08
0.8
~
06
0.6
0.4
0.4
0.2
02
0
CJ
(e)
..J
z
0
;:: CJ
cr: "-
0.0
0 0.1
0.2 0.3 0.4 0.5
0.0 0
or
02
0.3
0.4 0.5
PORE AXIAL I£NGTH
Fig. 8. Experimental platinum adsorption profiles produced by impregnating gamma-alumina pellets with an initial hexachloroplatinic acid concentration of 2.5xlO-3 molar and various initial concentrations of nitric acid: (a) C02 0.0 - platinum standard, (b) C02 = 5.0xlO-3 molar, and (c) C02 = 1.OxlO- 2 molar. Figures 8a-8c show how the addition of high concentrations of nitric acid affects platinum adsorption profiles. The metal profiles are uniform with a length of penetration that increases with increasing acid concentrations. The addition of an acidic Class 2 ingredient is shown to be an efficient method of producing uniform profiles with a high degree of penetration into the pore. Class 3 Ingredients In the previous sections, it was shown that the electrostatic (Class 1) and pH (Class 2) effects could each be reduced to a single parameter that was independent of the added ingredients. Class 1 and Class 2 additions control the amount of platinum adsorption and the depth that the platinum penetrates into the support pore. These ingredients can be used to obtain eggshell, uniform, or linearly decreasing platinum profiles. However, depending on reaction conditions, a core or other types of nonuniform metal profiles may be desired. These profiles are obtained by the addition of Class 3 ingredients. which complete with the platinum ion for adsorption sites on the alumina surface.
11
Class 3 ingredients that have a higher affinity for the surface than the active ingredient will adsorb to the surface at the front of the pore. As the solution flows down the pore, the concentration ratio of the active ingredient to the added ingredient becomes large, and the active ingredient will adsorb. In this manner the amount of Class 3 ingredient added and its relative affinity for the surface controls the distribution of the active ingredient. Whether the Class 3 ingredients are added as acids, bases, or salts will affect the resulting adsorption profile by introducing significant Class 1 or Class 2 effects into the system. Figures 9a-9c show the effect of adding a competing phosphate ion to the impregnating solution. The addition of a small amount of phosphoric acid, Figure 9b, yields a platinum profile that is linearly decreasing. The amount of platinum adsorbed at the pore mouth is suppressed because the phosphate ion has taken up adsorption sites. Farther down the pellet more and more platinum is adsorbed. Here, the amount of phosphate in solution has decreased enough to allow the platinum to compete effectively for adsorption sites.
10
10
10 (0
w
08
'"
(c)
(b)
I
0.8
II:
W
0.6
~ 06 u
..J
~
04
0.4
>=
u
II:
"-
0.2
0.2 00
02 00
0
0
01
PORE AXIAL LENGTH
Fig. 9. Experimental platinum adsorption profiles produced by impregnating gamma-alumina pellets with an initial hexachloroplatinic acid concentration of 2.5xlO-3 molar and various initial concentrations of phosphoric acid: (a) C02 = 0.0 - platinum standard, (b) C02 = 3.7xlO-3 molar, (c) C02 = 7.4xlO-3 molar.
12
At a much higher phosphate concentration, Figure 9c, the length of the platinum profile is about twice as great as that of the platinum standard, Figure 9a; the platinum distribution shows a long uniform profile with a sharp peak at the end. These experimental profiles illustrate the wide variety of profiles that can be produced by a Class 3 ingredient. Changing the ratio of the amount of Class 3 ingredient to active ingredient can control the amount of adsorption. the depth of penetration. and the overall shape of the profile. CONCLUSIONS The classification scheme proposed in this study divides the ingredients in an impregnating solution according to their effects on the impregnation process. By relating the changes in metal distributions to measurable solution parameters, adsorption-transport modeling can be developed. In conclusion, the proposal developed here provides a methodology for producing a wide variety of metal-supported catalysts. The ultimate goal is to relate impregnation modeling to reaction kinetics. This would then allow for the prediction of catalytic behavior based on the composition of the impregnating solution. ACKNOWLEDGEMENT One of us (J.A.S.) acknowledges the support of the Division of Chemical Sciences, U.S. Department of Energy, Basic Energy Science under Contract DE-AC02-84ER13158 during the preparation of this manuscript. REFERENCES 1 E.R. Becker and J. Wei, J. Catal. 46, 365 (1977). 2 E.R. Becker and J. Wei, J. Catal. 46, 372 (1977). 3 D.P. McArthur, Advan. Chern. Ser. 143, 85 (1975). 4 J.C. Summers and L.L. Hegedus, J. Catal. 51, 185 (1978). 5 R.W. Maatman, Ind. Eng. Chern. 51(8), 913 (1959). 6 H.A. Benesi, R.M. Curtis and H.P. Studer, J. Catal. 10, 328 (1968). 7 Y. Shyr and W.R. Ernst, J. Catal. 63,425 (1980). 8 F.R. Harty, The Chemistry of Platinum and Palladium, Wiley-Interscience, New York, 1973. 9 H.L. Bohn, B.L. McNeal and G.A. O'Connor, Soil Chemistry, Wiley-Interscience, New York, 1979. 10 J.A. Schwarz, C.T. Driscoll and A.K. Bhanot, J. Colloid Inter. Sci., Vol. 97, No.1, Jan. 1984. 11 H.M. May, P.A. Helmke and M.L. Jackson, Geochim. Cosmochim Acta 43, 861 (1979).
13
DISCUSSION B. DELMON : This refers to your class I ingredients: Na, Ca are often unwanted. They permanently modify the support. Volatile ions are preferred NHt, substituted ammonium, N0 3, organic acid anions. Van der Waals adsorption on the support modifies its apparent charge. The effect of those ions corresponds to what you describe correctly for class I ingredients, but with additional effects. One has sort of second order effects. One might also say that those ions are somehow situated between class I and class III. I suppose you observed that influence of atomic weight (Rb, Cs, Sr, instead of Na, Cal or molecular weight. Could you comment on that remark? May I add that I like very much your classification. J.A. SCHWARZ: Your question is indeed relevant and the points you raise demonstrate the limitations of any universal classification scheme for predicting the distribution of catalyst precursors within porous supports. In addition to your comments regarding "size" effects, there are other effects which I did not have time to comment upon which further demonstrate the subtle overlapping of the induced effects of added ingredients. For example, NaBr would normally be considered a class I ingredient. However, bromide ions exhange readily with chloride in the PtC1 62 anion, thus giving rise to speciation effects; each of the chloro-bromo complexes absorb at different rates. If the catalysis community is willing to recognize that there will always be exceptions to the rule, I do think that the proposed classification scheme does provide a useful guideline for manipulating the position of catalytic metal precursor within the support structure. G.M. PAJONK : As, in general, catalysts are rarely monomodal in their pore radii distribution, do you think that this distribution will exert an influence beside your three classes of added ingredients? J.A. SCHWARZ: We have confined our studies to a single catalyst support which has a specific pore size distribution. This has eliminated any confusion that might arise when trying to compare impregnation profiles obtained from different supports under similar experimental impregnation conditions. The answer to your question is that catalysts with different pore size distributions will likely lead to qualitatively different profiles. The single pore model assumes the support is comprised of pores with a constant radius, R. The convective transport is modeled by the methods described by Washburn (Phys. Rev. 17, 273 (1921)) which assumes that the flow of the liquid column passes througn an infinite succession of steady states. The expression for the radial average velocity of the liquid front is V p
= ~ (6~)~
t-~
where the pressure drop is a constant and equals 2A/R; A is the surface tension of the solution; ~ is the solution viscosity; and R is the pore radius. Thus the fluid velocity depends upons R~ and of course will change the qualitative behavior of the metal distributions. I might mention that we have carried out a similar study of the impregnation of Pt(NH4)4(OH)2 onto a high surface area silica and have demonstrated that the effect of a cnange in average pore size did not influence the predicted metal distribution based on the proposed classification scheme.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
15
© 1987 ElsevierScience Publishers B.V., Amsterdam - Printed inThe Netherlands
A STUDY OF SOME PARAMETERS IN CATALYST PREPARATION AND THEIR INFLUENCE ON CATALYST PERFORMANCE l• 2 A.W. NIENOW . 2 AND J.M. WINTERBOTTOM 3 P.T. CARDEW 1 • R.J. DAVEy P. ELLIOTT. 1I CI• Process Technology Group. New Science Group, PO Box 11. The Heath, Runcorn. Cheshire. WA7 4QE. U.K. ZICI pIc, Petrochemicals and Plastics Division. PO Box 8, The Heath, Runcorn, Cheshire. WAl 4QD. U.K. 3Department of Chemical Engineering, The University of Birmingham, PO Box 363, Edgbaston, Birmingham. B15 2TT, U.K. ABSTRACT The results are reported of an investigation into the influence of preparative conditions on catalyst performance, using silica-supported nickel as a model system. Nickel lII) dimethylglyoximate was precipitated in the presence of silica in a stirred mixing vessel; reduction with hydrogen then yielded metallic nickel supported on the silica. Metal surface area and specific activities of the catalysts were measured. Variation of the precipitation conditions enabled the relationship between preparative parameters and catalytic performance to be examined. 2 -1 Very high metal surface areas (ZOO-230cm g Ni) were obtained by this preparative route, but these were found to be almost independent of preparative conditions. The reduction of Nickel (II) dimethylglyoximate is accompanied by sublimation; this is believed to be responsible for both the high surface areas and the lack of correlation with the preparative conditions. INTRODUCTION Industrial catalysts are manufactured traditionally by established routes based upon empirical knowledge rather than upon a scientific understanding of how the preparation conditions affect catalyst performance.
Much is known about
the behaviour of catalysts once prepared. but little has been published which attempts a systematic analysis of catalyst preparation, or an identification of how preparative conditions might influence catalyst performance.
This paper
reports the results of a study of the way in which preparative conditions may affect the physical properties and catalytic performance of a model catalyst. The catalyst studied was metallic nickel on a silica support, prepared by a recently-developed route (ref.l) in which Nickel (II) dimethylglyoximate is precipitated in the presence of silica.
Following suitable filtration, washing
and drying. the organic complex is reduced in hydrogen to metallic nickel supported on the silica.
This route gave highly dispersed nickel catalysts with metal surface areas as high as 250 m2 (gNi)-l; atomically dispersed nickel would have a surface area of 670 m2(gNirl(awrox). Of importance is the manner in which the components are brought together and a range of parameters was varied identifying those
16 of greatest importance in determining the properties & quality of the resulting catalyst. Measurements of catalytic activity, metal surface area and crystallite size, and metal-support interaction were used to characterise the catalysts thus prepared. Reported are the methods used in the preparation and characterisation of the catalysts, together with the results of efforts to correlate changes in preparative conditions with variations in catalyst properties and performance. Preliminary findings in this field have previously been reported (refs.2,3,4), while a more detailed study is currently in preparation (ref.S). CATALYST PREPARATION Supported nickel catalysts were prepared according to the following scheme: 1. Precipitation of Nickel (II) dimethylglyoximate in the presence of suspended silica. Z. Filtration, washing and drying of the precipitate and silica mixture. 3. Reduction of the nickel complex to metallic nickel by heating in hydrogen. Nickel was precipitated from an aqueous solution of Nickel (II) nitrate by the addition of an ethanolic solution of dimethylglyoxime (C4C8NZOZ) according to the following equation: Ni Z+ (aq) + ZHZDMG (EtOH) ~ Ni (HDMG)Z (s) + ZH+ (aq) dimethylglyoxime
Nickel (II)
in ethanol
dimethylglyoximate
The solubility product of Ni(HDMG)Z in water is reported to be 4.4 x 10- 18 (ref.6), so the precipitation is essentially complete under normal conditions. Indeed, dimethylglyoxime has long been the standard gravimetric reagent for the determination of nickel (ref.7).
However, the liberation of ~
ions
during precipitation necessitates the use of a pH buffer, since solubility increases significantly with decreasing pH. Precipitations were performed in a stirred, baffled mixing vessel of 9dm 3 capacity, in which the dimethylglyoxime solution was added, by means of a metering pump, to the Ni Z+ solution containing the silica in suspension. The rate of addition of dimethylglyoxime was such that precipitation was complete after 20 minutes. Agitation was provided by a 7cm, 4-blade, 45 0 impeller, pumping downards. The following parameters were varied systematically in order to investigate the influence of each in turn on catalyst performance: 1. Ni Z+ concentration, varied from 5 x 10- 5 to 5 x 10- 3 mol dm-3; Z. pH, varied at unit pH intervals from 3.0 to 10.0 by means of buffers; 3. Temperature, varied in 100e steps from ooe to 400C; 4. Impeller speed, varied in 5 steps from 60 rpm to 900 rpm. After precipiation was complete,themixtureofNilHDMG)Z precipitate and
17
suspended silica was collected by filtration, washed with water, and dried. Two drying temepratures were used, 200 e and 80°C.
The effect of washing
additionally with ethanol prior to drying was also investigated. In all cases, Davison 95Z grade silica was used. range was
The measured particle size
(by optical microscopy) and the nominal surface area of this grade is 300-350m Z/g. The ratio of nickel:silica was maintained such that 5-Z30~
complete reduction of the Ni(HDMG)Z would yield a catalyst loading of 5% w/w on the silica support.
The HZDMG concentration in ethanolwasnormallyO.lmoldm 3•
For all precipitations, an induction time was measured, this being the delay between the beginning of HZDMG addition and the first appearance of the Ni(HDMG)Z precipitate, immediately recognisable from its distinctive pink coloration. Once dry, the catalyst precursor was heated at ZOoC/min to 500 0C in a stream of hydrogen at 40 cc/min to reduce the Ni(HDMG)Z complex to metallic nickel crystallites on the silica surface.
In order to isolate the effect of changes
only at the precipitation stage, the reduction step was generally identical for all catalysts.
However, asa result of developments in the work the effect of
reducing in hydrogen diluted to only 5% v/v by an inert gas (nitrogen), and of calcination prior to reduction, was additionally investigated in some instances. CATALYST PRECURSOR CHARACTERISATION When Ni(HDMG)Z is precipitated from aqueous solution, it forms well-documented, elongated needles approximately
10~
x
O.I~
in size (refs.8,9) (fig.l).
The
effect of some precipitation parameters on the size of these crystals has previously been studied by electron microscopy (refs.IO,II).
Since the
Ni(HDMG)Z crystal size could determine the metal crystallite size (and hence the metal surface area) of the final catalyst, it is desirable to know how the Ni(HDMG)Z
crystal size varies with precipitation conditions.
Beforereducing each precipitated catalyst precursor to nickel metal, the precipitates were themselves examined for any influence of all the parameters listed under Catalyst Preparation.
Ni(HDMG)Z crystal sizes were measured by
optical microscopy, scanning electron microscopy (SEM), and by means of a Coulter Counter.
Each technique of experimental characterisation gave similar
results relating precipitation conditions to crystal size and morphology. A complex set of interactions between all the variables were found. Therefore, it is essential to quote results for the variation of only one parameter at a time. One example for varations of Ni Z+ concentration with all other parameters held constant is given in fig.Z. The effect of Ni Z+ concentration was in accordance with expectations.nam~ly crystal size increases with decreasing Ni Z+ concentration, and hence with lower
18
Fig. 1 Electron micrograph of needle-like Ni(HDMG)Z crystals. (Scale; Smm = 1~; pH = 7; 250 rpm)
..c
"2'20 .!
~OIl 16
........
~12
!i! o %:
8
:i
g .. 4 x 5.10-2 5J.10·3 5.10"'4 5.10-5 Initial concentration of Ni2'mol dm-3)
Fig. 2 Variation of mean Ni(HDMG)2 crystal length with initial Ni Z+ concentration (Z80 rpm; temp = ZOoC, pH = 7).
19
levels of supersaturation (fig.2).
This is a familiar effect in precipitation
processes. Fig. 3 shows how mean Ni(HDMG)Z needle length varied with pH.
The pH-
dependence arises from the increasing solubility of Ni(HDMG)Z in water in acidic conditions.
Indeed, below pH 3 this nickel complex is appreciably water-soluble
and is not precipitated quantitatively.
At low pH values, therefore, the lower
supersaturation results in larger Ni(HDMG)Z crystals.
e
:s....
.:,.2.... • ~14
c:
o
III
>~10
..!:"
'" 58 =. %6 c: .,.
•
•
•
..
::1:
4 3
4
5
6
7
8
9
10
pH during preparation
Fig. 3 Variation of mean Ni(HDMG)2 crystal length with pH during precipitation. (Z80 rpm; temp ZooC; Ni Z+ concentration = 5 x 10- 3 M) The effect of temperature on crystal size was far less pronounced: the difference in mean cystal size between samples precipitated at ZOOC and SOOC was barely measurable. Crystals precipitated at extremes of pH and temperature were examined by X-ray diffraction (XRD) to detect any difference in the phase of the precipitated materials.
In fact, all samples gave identical XRD profiles.
Optical microscopy before and after drying the precipitate provided information on the degree of physical association between the Ni(HDMG)Z and the silica which had been in the tank during precipitation.
Inspection of the Ni(HDMG)2 - silica
mixture before drying revealed that there was no apparent association whatsoever between the two components.
This is consistent with the observation that the
induction times for precipitations with and without silica present were identical, showing that the silica plays no part in the precipitation and fails to act as a nucleation site for the precipitation of this nickel complex.
After drying,
however, a significant proportion of the Ni(HDMG)Z was seen to be physically
20
associated with the silica particles.
No difference in association was observed
between samples dried at ZOOC and those dried at SOOC, nor between those washed with ethanol and those washed only with water.
Very rapid drying, however
(e.g. by heating a thin film on a glass slide) precluded any accompanying association process.
The drying step therefore seems to be more important
than the precipitation process in determining the association between the catalyst precursor and its silica support. CATALYST CHARACTERISATION - PHYSICAL PROPERTIES Reduced catalysts were characterised according to metal surface area, metalsupport interaction, metal crystallite size and metal loading. Surface areas were measured by hydrogen chemisorption, using apparatus and methods described by Benesi et al. (ref.13) and Falconer and Schwarz (ref.14), except that hydrogen uptake during cooling was measured, rather than hydrogen desorption during heating.
About 100mg of each catalyst was weighed accurately
into a U-shaped tube, which was then placed in a programmable oven and purged with hydrogen. With a HZ flow rate over the sample of 40 cc/min, the oven was heated to 500 oC, reducing the nickel complex to metallic nickel. The gas flow was then changed to 5% HZ in Argon and the system allowed to equilibrate. A thermal conductivity detector (TCD) measured the composition of the gas at the exit from the tube, which was then cooled rapidly.
Hydrogen chemisorption onto
the metal surface depleted the HZ content of the exist gas passing through the TeD.
Integration of the TCD signal output determined the total HZ uptake by
the catalyst, from which the metal surface area was calculated using the BET equation, assuming monolayer coverage.
Calibration was achieved by injecting
pulses of the accurately known volume of fl2into the gas flow upstream of the TeD. Specific metal surface areas can be calculated if the exact loading of the catalyst is known.
Loadings were measured by dissolving away the Ni(HDMG)Z
from an accurately weighed sample of catalyst precursor using nitric acid, and measuring the Ni Z+ concentration of the resulting solution by atomic absorption spectroscopy.
Owing to the quantitative precipitation of Ni(HDMG)Z, measured
loadings were generally in good agreement with theoretical values. Metal-support interactions were investigated by temperature-programmed reduction (TPR) (refs.15,16). measurements.
The apparatus was the same as for HZ chemisorption
50mg of catalyst precursor was weighed into the U-tube, through
which 5% HZ in Argon was passed at 30 cc/min.
As the temperature was raised
at ZOoC/min, the progressive reduction of the catalyst removed HZ from the gas flow.
This change in composition was measured as before by the TCD.
Recording
the TeD output signal as a chart record resulted in a profile exhibiting two peaks, corresponding to a two-step reduction mechanism.
The temperature at
which these peaks occur reflects the ease of reducibility of the catalyst; any
21
shift in peak temperatures relative to unsupported Ni(HDMG)2 suggestive of metal-support interaction, rendering the reduction of the catalyst precursor more, or less, facile than for the unsupported material. Nickel crystallite sizes in the reduced catalysts were measured by transmission electron micrscopy (TEM), using magnification of 150 000 times. CATALYST CHARACTERISATION - CATALYTIC ACTIVITY Specific activities of catalysts were measured using a standard toluene hydrogenation reaction in a tubular furnace reactor operated in differential mode.
100mg of catalyst precursor was weighed accurately into a 1cm diameter
pyrex tube, forming a bed 4-5mm deep between two glass wool plugs. was flushed with HZ and heated to 5000C to reduce the catalyst.
The tube
The reactor
was then allowed to cool before switching the gas flow to 5% HZ in NZ previously saturated with toluene vapour.
The gas leaving the reactor was sampled and
analysed by gas chromatography. The fraction of toluene hydrogenated to methylcyclohexane (MCH) is depended on furnace temperature, and the relationship between catalyst temperature and % conversion (expressed as a reaction rate) gives a measure of the activation energy of the catalysed reaction, and hence of the catalyst activity. Conversions were maintained below 5% in order to employ the relatively simple kinetic analysis of the differential mode reactor. RESULTS Despite the known wide variations in Ni(HDMG)Z crystal size distribution (CSD) under different precipitation conditions, no significant differences were found between the properties and characteristics of the catalysts after reduction in HZ'
Metal surface areas were found to be ZOO-ZZO HZ (gNi)-l, but these were
similar regardless of precipitation conditions (table 1). Table 1 - HZ Chemisorption and TRP results for catalysts prepared at different impeller speeds Impeller Speed rpm
60 120 250 500 900 Unsupported Ni(HDMG)2
Measured
loading %w/w Ni
Lower TPR Peak Temp °C
Upper TPR Area Under Peak Temp TPR Profile Arbitrary °c Units
Metal Surface Area by H2 Chemisorption (g-1 Ni)
4.92 4.94 4.84 4.96 4.81
334 335 333 334 337
451 460 457 465 460
3.18 3.05 3.12 3.24 3.17
215 214 207 214 224
N/A
300
350
2.15
16.2
Temperature programmed reduction revealed a shift in the position of the two reduction peaks for supported relative to unsupported samples (fig.4), but the TPR profiles were similar for all catalysts tested (table 1). Fig. 5 shows nickel crystallites on the silica support, imaged by rEM, with
22
Ill-I o ~
." s-,
:t: 0'
c
. .s "iii d
u
In)
40
40Q 350 Temperature • C
' 300
Fig. 4 Temperature programmed reduction profiles for Ni(HDMG)Z. (a) with no silica present. (b) with added silica to give a loading of 5% w/w nickel on silica catalyst.
Fig. 5 TEM photograph of a reduced catalyst. Nickel crystallites are clearly seen against the silica support. This sample was reduced firstly in 5%' HZ diluted with NZ before further reduction in pure HZ'
23
sizes ranging from lOA and 1501.
Again, no difference in size distributions
was apparent between the differently prepared samples. Likewise, activity data from the catalytic hydrogenation experiments showed no clear correlation with preparation conditions, and no pattern of activation energies emerged from the samples examined. Despite this lack of correlation, two potentially valuable relationships did emerge from this study.
Firstly, catalysts reduced in pure HZ exhibited
consistently lower surface areas than catalysts previously heated in 5% HZ in NZ' Prolonged calcination at 400°C in NZ prior to HZ-reduction, however, left the surface area unchanged. Ni(HDMG)Z
Secondly, the metal surface areas and TPR profile of
physically mixed with silica were the same as for samples where the
silica had been present during precipitation.
Hence no advantage is gained by
adding the silica at the precipitation stage.
Of particular significance,
mixtures of Ni(HDMG)2 and silica exhibited TPR profiles identical to those of supported Ni(HDMG)2, and quite different from the characteristic profile of unsupported Ni)HDMG)Z (fig.4), which might have been expected. DISCUSSION The lack of correlation between precipitation conditions and catalyst properties can be explained by the observation that, whereas the reduction of Ni(HDMG)Z in hydrogen commences at 3000C, Ni(HDMG)2 sublimes above 230 0C. Therefore, any variations in Ni(HDMG)2 CSD arising from different precipitation conditions will be eradicated during the reduction stage by the vapour-phase redistribution of the Ni(HDMG)2 at temperature just below the reduction temperature.
This also explains why physical mixtures of Ni(HDMG)2 and silica
produce similar TPR profiles, surface areas and catalytic activities to samples in which the silica was present during precipitation. The sublimation of Ni(HDMG)2 was examined by hot-stage microscopy: when rough mixtures of Ni(HDMG)2 + silica were heated on the microscope stage, the Ni(HDMG)2 was seen clearly to sublime above 230 0C, and the silica particles become well coated with Ni(HDMG)2 at 275°C.
The effect was observed again when
portions of this mixture, heated separately to 175, 2Z5, 250 and 275OC, were examined by SEM.
The progressive sublimation of the unassociated Ni(HDMG)2
crystals onto the surface of the silica support was clearly visible on comparing these samples. l~is
also seems to explain why physical mixtures of dry Ni(HDMG)2 crystals
+ silica powder exhibit TPR profiles identical to those of supported samples. of Ni(HDMG)2 precipitated in the presence of silica, and quite unlike the characteristic profile of unsupported Ni(HDMG)2: the Ni(HDMG)Z sublimes onto the silica surface and thereby becomes supported before reduction occurs. The measurement of lower surface areas for catalyst reduced in pure hydrogen,
24 rather than in 5% HZ in NZ' seems to be due to the sintering of small nickel crystallites when heated to 500 0C in pure hydrogen.
This is supported by TEM
results: the sample in fig. 5 was "pre-reduced" in 5% HZ -NZ before switching briefly to pure HZ' (below 100
R)
In samples reduced only in pure HZ, the small crystallites
are no longer seen by TEM, presumably having sintered into the
larger aggregates which are then clearly visible (fig.6).
The sintering and
mobility of Ni crystallites in HZ has been reported (refs.18,19,ZO), but it is not known whether in this case the process is one of genuine nickel crystallite migration, or a redistribution of the material during reduction by sublimation owing to the evolution of local heat, due to the exothermic reaction, while reducing in pure hydrogen.
Fig. 6 TEM photograh of a catalyst reduced in pure HZ at 500 oC. Crystallites of less than 100 A diameter are entirely absent, apparently due to sintering. CONCLUSIONS Supported nickel catalysts are difficult to prepare in a form which is both highly dispersed and highly reduced, on account of (a) sintering and (b) metalsupport interaction.
Nickel dimethylglyoximate as a catalyst precursor has been
reported to yield catalyst with a high degree of both reduction and dispersion (ref.Zl).
It has been found, however, that during reduction to metallic nickel
a vapour-phase redistribution of the catalyst precursor takes place.
Whereas
this precludes the modification of the catalyst by varying the precipitation
25
conditions, it also appears that this sublimation of the catalyst precursor prior to reduction to supported nickel is a vital factor in achieving the desired degree of dispersion in the final catalyst. Further work in this area might investigate the performance of catalysts prepared from different oxime complexes which might be more volatile than Ni(HDMG)2, or else examine alternative reduction conditions which specifically suppress, or enhance, the sublimation of the nickel complex.
A possible method
will be via sodium tetrahydroborate reduction in the liquid phase or via hydrazine vapour at temperatures less than 2000C i.e. below the sublimation temperature. REFERENCES 1 A.I. Thompson and J.M. Winterbottom, unpublished work. 2 A.J.S. Anderson and I.K. Minto, 3rd year research project, University of Birmingham. 3 G. Cambanis and M. Polack, 3rd year research project, University of Birmingham, 1982. 4 A.J.S. Anderson, G.C. Cambanis, P. Elliott, I.K. Minto, A.W. Nienow, M. Polack, A.I. Thompson and J.M. Winterbottom, Actas Simp. Iberoam. Catal., 9th, 1984, 2, 1609-10. 5 P. Eliott, Ph.D. Thesis, University of Birmingham (in preparation). 6 H. Christopherson and E.B. Sandell, Anal. Chim. Acta 10 (1954), 1-9. 7 A.F. Vogel, A Textbook of Quantiative Inorganic Analysis, 3rd ed. (1964),125. 8 K. Takiyoma and L. Gordon, Talanta 10 (1963), 1165-1167. 9 J.L. Jones and L.C. Hawick, Talanta 11 (1964), 757-60. 10 R.B. Fischer and S.H. Simonsen, Anal. Chem. 20 (1948), 1107-1109. 11 M. Ishibashi, E. Suito , K. Takayima and E. Sekido, Bull. Inst. Chem. Res. Kyoto Univ. 31 (1953), 365-367. 12 P.P. von Weirmarn, Chem. Rev. 2 (1926), 217-242. 13 H.A. Benesi, L.T. Atkins and R.B. Mosely, J.Catal. 23 (1971), 211-213. 14 J.L. Falconer and J.A. Schwarz, Catal. Rev. - Sci. Eng. 25(2), (1983), 141-227. 15 J.W. Jenkins, B.D. McNicol and S.D. Robertson, Chemtech., May 1977, 316-320. 16 S.D. Robertson, B.D. McNicol, J.H. de Baas and S.C. Kloet, J.Catal. 37 (1975), 424-431. 17 R.B. Anderson, "Kinetics of Catalytic Reactions" in B.R. Anderson and P.T. Davison, Experimental Methods in Catalytic Research (1976), 1-43. 18 t. Nakayama, M. Arai and Y. Nishiyama, J.Catal. 79 (1983), 497-500. 19 C.H. Bartholomew and W.L. Sorensen, J.Catal. 81 (1983), 131-141. 20 T. Nakayama, M. Arai and Y. Nishiyama, J.Catal. 87 (1984), 108-115. 21. E. Zogli and J.L. Falconer, Applied Catalysis 4 (1982), 135-143.
26
DISCUSSION L. GUCZI : Have you tried to calcine your catalyst precursor prior to hydrogenation? It may affect the sublimation of Ni(HDMG)2 simply accelerating the decomposition into Ni oxide. P. ELLIOTT: Brief calcination of samples at 300°C in nitrogen reduced the surface area measured by hydrogen chemisorption by 40%. The TPR profile exhibited two peaks in the same area ratio as uncalcined samples, but the total area under the TPR profile was also reduced by 40%. Calcination at 400°C gave only one TPR peak, and a surface area 30% lower than uncalcined samples. Prolonged calcination at these temperatures failed to produce a catalyst superior in surface area to material that has not been calcined. V. PERRI CHON : The 20°C min- l heating rate used for the reduction of the sample seems to me rather high. Don't you think that water vapour produced during the reduction can favour the sintering of the nickel particles? Did you try a lower heating rate to limit this phenomenon? P. ELLIOTT: The exact nature of the reaction between hydrogen and Ni(HDMG)2 is not yet known; the presence of water as a reaction product, although likely, has not been confirmed. Heating rates from SOC/min to 40°C/min were tried, and the heating rate was found not to have an effect upon catalyst surface area within this range. Very low heating rates would be expected to promote extensive sublimation of NlTHDMG)2 prior to reduction and, in a gas flow, result in removal of this volatile material altogether from the system. G.M. PAJONK : The fact that your Ni dimethyl glyoximate sublimates when on your silica support reminds me of course of the CVD (chemical vapor deposition) method. Do you expect to observe identical results if your silica was directly treated by your Ni l I complex in the vapour phase? P. ELLIOTT: Ni(HOMG)2-silica systems mixed mechanically by hand showed TPR profiles and metal surface areas (after reduction) which were identical to samples in which the nickel had been precipitated in the presence of the silica. Since we were concerned with the influence of mixing and precipitation conditions upon catalyst performance, no attempt was made to sublime bulk Ni(HDMG)2 onto the silica support in the conventional CVD manner. Nevertheless, this remains an option for future investigation, provided that dispersion and loading can be adequately controlled. J. B.NAGY : Did you use surfactant molecules to stabilize the Ni(II) ions or Ni particles? How the presence of surfactants influenced the size of the particles? Were the particles monodisperse in size? P. ELLIOTT: The presence of a surfactant at the precipitation stage resulted in Ni(HDMG)2 crystals 3-4 times smaller than in the absence of a surfactant. However, differences in Ni(HDMG)2 CSD at this stage are eradicated by the sublimation step which precedes reduction. The Ni(HDMG)2 size distribution was monodisperse with a clearly defined mean size. The nickel crystallites following reduction covered a broad size range, and the exact nature of the size distribution has not yet been elucidated. P. LANCASTER: How do nickel crystal sizes vary with time (activity variation with time), after reduction? Do different preparation nethods Wiich give specific crystal sizes show similar catalytic activity after reduction and usage for a short period of time?
27
P. ELLIOTT: The variation of nickel crystallite size or catalytic activity with time was not specifically investigated. Data reported above were measured immediately after catalyst preparation. However, when metal surface areas, determined by hydrogen chemisorption, were measured repeatedly on the same samples, the values initially decayed by about 10% before reaching stable values after five or six determinations. This behaviour was independent of preparation conditions.
J. KIWI: You buffer your solutions for Ni precipitation on 5i0 2 using dimethylglyoxime? Did you assess the competition of ions of the buffer (-ions) with the precursor of the catalyst at different pH ? If so, did you use straight acid and base to adjust the pH values during Ni-precipitation and compare the final catalyst produced? P. ELLIOTT: The use of cation buffer solutions is precluded by adverse chemical reactions; for example, phosphate-containing buffers precipitate nickel phosphate, whereas some organic molecules are said to complex with H2DMG. Beyond this, specific ionic contributions and interactions were not investigated, although the characteristics of catalysts prepared from unbuffered systems (adding base to maintain neutral pH) did not differ from those prepared in buffered systems.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
29
IMPREGNATION OF CONTROLLED-POROSITY SILICAS WITH HEXACHLOROPLATINIC ACID: PARAMETERS AFFECTING Pt DISPERSION AND LOCATION AND SUPPORT MODIFICATION M.A. MARTIN LUENGO, P.A. SERMON* and K.S.W. SING Department of Chemistry, Brunel University, Uxbridge, Middlesex UBB 3PH, UK
SUMMARY The surface area rather than the average pore size distribution of silica supports appears to define the size and dispersion of Pt produced by impregnation with hexachloroplatinic acid solutions. Partial pore blocking as a result of impregnation is to some extent reversed by calcination and reduction. The silica is not significantly depolymerised-dissolved despite the low pH of the impregnating solution (i .e. pH,0.5). However, the temperature of the impregnating solution does appear to be important in defining the location of the supported Pt particles; this useful preparative parameter has not been considered previously. It may also affect the chemical state and reducibility of the silica-supported Pt species. INTRODUCTI ON Porous silica is often used as a catalyst support with the active phase being introduced by adsorption or impregnation from aqueous solution. It generally has a wide range of pore radii; thus mesoporous Davison 70 has been used with particular frequency to support metals (ref. 1) and oxides (ref. 2) but its distribution of pore diameters lies within the range 5 nm to 15 ~m (although at least 50% of pores detected by mercury porosimetry have diameters of 7.5-11.5 nm (see Fig. la)). The average pore diameter a por e is reported (ref. 3) to affect the growth and ultimate size of supported metal particles. Certainly one assumes that the average metal particle size apart (i.e. 2rpar t ) cannot exceed the average pore size (see Fig. lb), unless the pore structure is ruptured; when the two approach one another there may be some pore blocking. Consider a Pt particle blocking a pore in 5i0 2 (i.e. r par t = r por e) after reduction of a precursor salt. Raising the temperature by ~T causes strain on the support where a is the linear coefficient of (apt - asilica)·~T expansion. The resulting stress intensity is Esilica·(apt - asilica)~~T, where Esilica is the Young's modulus of the support (in this case 0.7 kg.cm ). Then the maximum stress which the silica support can withstand before rupture is its tensile strength Fsilica (700 kg.cm- 2). Therefore rupture of the local support pore structure would occur at Fsilica/Esilica.(aPt - asilica) or 111 K. In other words when the Pt/Si0 2 was heated a little above its temperature of preparation or the temperature when r par t = r pore then the local support pore
30
Fig. 1 Pore size distribution in Davison 70 silica defined by dpore (2rpore) (la) and the relationship between the average size of supported metal particles dDart (=2rpart) measured as a function of concentration (lb) and temperature (lC).
% pore volume (V) 0.3
a 100 f-
dV d(d
<E-<--_.
80 f-
_
0.2
-
0.1
;
60 -
I
I I
40 '-
•
••
I
•
I:~
20
......
I
-~
II
. I
loo active phase
8 _
o
I 1 I I
o 4 _
b
cPl c9 cP:O
%
o
I
op
~
01-...;:..----71..--"'------.......I ---1 1000 I-
o
I
r600 r-
800 T(K)
400 I- 0 2
10 I 0
,
0
Fo~ ~f
Q)
C
o 0 0
0 0 0
0
tl' II
I
10
20
dpore (-2r pore) (nm) (a) dpart (=2rpart) (nm) (b.c)
100
pore
)
3 -1 -1 cm.g .nm
31
structure will be liable to fracture and subsequent further Pt particle aggregation and growth. However, below this critical temperature it is assumed that r por e for the support pores defines the maximum value of r part for the supported metal. Fig. 1c shows that r part exceeds r pore at higher temperatures (but not at higher metal loadings (see Fig. 1b)) for Pt upon Davison 70 meso-porous silica. However, because of the wide range of pore sizes present in this and similar silicas it is not easy to discern the role of the support pore size on the size of the particles of the supported metal or conversely the effect of the impregnation-reduction-calcination on the silica support porosity. Therefore the present work was undertaken with two silicas of well controlled porosity within very narrow size ranges and into which Pt could be introduced, to determine the extent of this mutual relationship between the support pore size and the size (and location) of the supported metal crystallites. This work is now reported. EXPERIMENTAL Materials Two silica gels (Shell International Chemical Co Ltd) were selected because of their high purity, homogeneity of pore size and surface area. These consisted of spheres (2.5 mm diamter: S-980A-2.5/3 denoted A in Table 1; 1.5 mm diameter: S-980A-1.5 denoted B in Table 1) which were particularly easy to use for the preparation of supported metal catalysts. The physical characteristics of these two silicas are given in Table 1. TABLE 1 Properties of silica supports SN
Sample
A Shell B Shell C Shell D Shell
* t
2 (m 2g-1) 980-2.5 980-1.5 980-2.5 t 980-1.5 t
211 56 215
pores BET % (nm) of 6-10 C (20-30) 120 243 176
83 (21) 81 (25)
in range r por e 10-20 (30-60) 6 (70) 7 (65)
total pore volwn e 1) 6-10 10-20 (cm '9(20-30) (30-60) Hg N2 vpore*
0.77 (0.20 0.70 (0.23)
0.06 (0.67) 0.06 (0.59)
0.93 0.95 0.87 0.91
0.95 0.94
3 -1 pore volume (cm.g ) in range (nm) of r por e' samples C and D were produced from silicas A and B respectively by treatment with HCl solution (at the same pH as those of hexach1orop1atinic acid used in catalyst preparation)at 373 K for Sh.
32
Hexachloroplatinic acid H2PtC1 6.xH20 (Johnson Matthey; >99.9% purity) was used for the preparation of all catalysts. Catalyst Preparation All silica-supported platinum catalysts were prepared using the incipient wetness technique in the following manner. An aqueous solution of hexachloroplatinic acid was used which was of suitable strength and volume to just fill the support pore volume and producE 3% Pt on drying. Supports were impregnated under isothermal conditions at one of two temperatures (273 Kand 373 K). Samples containing the supported Pt salt were or were not then calcined in static air at 543 K for 3 or 4h and reduced in H2 flowing at 100 cm 3/min (at 523 K for lh). These preparative and pretreatment conditions were designed to give different Pt dispersions and to examine the role of the support characteristics, especially average pore size. The catalysts so prepared are given in Table 2 and are described in terms of the silica (A or B) the impregnation temperature, the calcination temperature and the reduction temperature, and are also defined in ref. 4. TABLE 2 Properties of 3% Pt/Si0 2 prepared from silicas in % pores Sample Temperature and time ~2 BET (nm) of 6-10 l) C K (hl (i. g 1* C R* (20-30) E F G
H I
J K
L M N 0
Table 1 in range Total pore v~lue r~8r2 (em ,g-l) - 0 (30-60) Hg N2 BET
273(5) 183 373(5) 180 523(1) 209 373(5) 523(1) 211 273(5) 205 373(5) 543(3) 212 373(5) 543(4)
122 222 124 103 200 67
74 85 83 80
273(5) 373(5) 523(1) 373(5) 373(5) 543(3) 273(5) 523(1)
37 44 110 123 26
(15)
(80)
0.78
(20) (14) (26)
(73) (72) (64)
0.94 0.97 0.93
*1 = impregnation C = calcination R = reduction
48 50 56 60 55
15 6 6 10
0.80 0.69 0.86 0.88
0.80 0.78 0.86 0.89
Sil i ea
A A A
A A A
B B B B B
33
Methods of Support and Catalyst Characterisation Mercury porosimetry data were obtained using a Carlo Erba 2000 series porosimeter which was capable of measuring pressures to 2000 Bar. The Washburn equation (ref. 5) was used to calculate pore radii down to 3.75 nm. All data were corrected for the compressibility of the mercury. Nitrogen adsorption data were obtained using a Carlo Erba Sorptomatic 1800 unit with the sample outgassed (298K, 16h, <53mPa). BET surface areas were calculated in the range 0.05 < plpo < 0.25 assuming that the molecular crosssection of nitrogen was 0.162 nm 2. Pore size distributions were calculated from isothermal desorption data using the Robert's method (ref. 6). Since this method is only applicable to pore radii 2 nm < r por e < 14 nm. it could only be applied to silica A; dvldr pore derived from N2 BET data was a maximum at r pore ~2 nm in silica B and this is not consistent with porosimetry in Table 1. The average size of supported Pt particles was estimated by selective hydrogen chemisorption on catalysts (and supports) determined in a conventional volumetric apparatus (ref. 4) at 298 K. The location of the platinum in catalysts was determined by scanning electron microscopy and microprobe analysis to cross-sections of the catalyst spheres. The rate and extent of adsorption of Pt onto 2.5 mm diameter silica spheres was determined by continuous atomic absorption spectroscopic analysis of the aqueous solution of hexach10rop1atinic acid (initial strength 50 ppm Pt) at constant temperature (273, 293 and 313 K) above a sample (lg) of silica spheres. The rate of penetration of aqueous solutions into the 2.5 mm diameter silica spheres was followed by direct video microscopy during the addition of 2 cm 3 of solution at constant temperature (273. 294 and 313 K) to 2g samples of the silica and image analysis at 1s intervals. The dry particles were opaque to transmitted optical illumination, but the wetted particles were translucent. Therefore measurement allowed the radii of the opaque centres of the silica spheres to be determined as a function of time. As a result average rates of water penetration were determined from image analysis of about 40 particles. RESULTS Support Porosity and Surface Area Fig. 2 shows that non-impregnated 2.5 mm spheres of silica have a N2 adsorption-desorption isotherm of type IV (ref. 7) at 77 K; Table 1 shows that the BET area calculated from these adsorption data is 211 m2.g-1, with a C value of 120. The hysteresis loop is typical of a mesoporous material and is designated type Hl (ref. 7). Total pore filling corresponding to the higher pressure hysteresis loop closure point in Fig. 2 is 608 cm3.g-1 at STP or a
34
Fig. 2. N2 adsorption isotherm at 77 K on untreated silica A with 2.5 mm diameter spheres (0 •• ) unreduced catalyst F (A •• ) and cata lyst G (0 •• ). Fi 11 ed symbols denote desorption points.
600
400
200
0.2
0.6
0.4
0.8
Fig. 3. Pore size distribution in untreated silica A (a). catalyst F (g) and catalyst G (c) derived from Hg porosimetry. pore volume (cm 3/g)
10
100
10
r
por e (nm)
100
10
100
35
liquid volume of 0.95 cm 3.g-1; the distribution in the range 7-11 nm of pore sizes in this support obtained from porosimetry is shown in Table 1 and Fig. 3a. In Fig. 4 silica 1.5 almost shows a type II N2 adsorption-desorrtion isotherm at 77 Kfrom which its BET surface area was deduced to be 56 m2.g- (see Table 1 which also shows that the C valuH wa~ very large). Unfortunately the pore sizes in this silica support were too large for accurate deduction from such data (although dv/dr pore was always a maximum ~1-2 nm); however, mercury porosimetry found 90% of its pores had radii at 20-60 nm (Fig. 5(a)). Catalyst Porosity and Surface Area Data in Table 2 and Figs. 3-5 (which can be compared with that before impregnation in Figs. 3-5 and Table 1) suggest that impregnation of both of the above silicas with hexachloroplatinic acid (giving samples E, F, K andL) leads to a lower surface area and pore volume than A and B in Table 1, although isotherm shapes (see Figs. 2 and 4) remain constant. This could result from blocking the pore structure by the impregnating salt or some dissolution-depolymerisation of the support at the pH of the impregnating solution. To show that the latter was not important both support silicas were treated with HCl solutions of the same pH as the impregnating solutions; the data for samples C and 0 in Table 1 are very similar to those for samples A and B respectively showing that this had no significant effect. It can also be seen from a comparison of data for samples E and F (and K and L) that the temperature of impregnation (273 or 373 K) had no measurable effect upon the pore volume and surface area of the final impregnated samples. It is possible that the difference between total pore volumes determined by BET and porosimetry methods for sample F in Table 2 is a reflection of the formation of 'ink-bottle' pores during the impregnation procedure. The results may also indicate a higher degree of pore blocking in samples prepared at 373 K over those prepared at 273 K; this may be related later to the location of the salt and Pt particles in the support pore structure. Calcination of impregnated materials (see results for samples I, J and N in Table 2) at 543 K appeared to raise the surface area and pore volume over that in samples F and L merely after impregnation; possibly the halide salt particles are converted to an oxidic phase with simultaneous decongestion of pores by particle migration. Reduction (see isotherms in Figs. 2 and 4 and porosity data in Figs. 3 and 5 and also Table 2 for samples G and M) at 523 K also increased the total surface area and pore volume of the samples, presumably by an equivalent mechanism. Such effects appeared larger in calcination at 523 K after impregnation at 273 K than impregnation at 373 K. Under none of the present conditions is there a general collapse of the silica support structure as previously postulated by Benesi et al. in ref. 1.
36
Fig. 4. N2 adsorption isotherm at 77 K on untreated silica B with 1.5 mm diameter spheres (0,0) and catalyst L (6,.6.). Filled symbol s denote desorption points.
-
J
• ;,
.'"
b
1•
~
I,
~~
~ •• ,.,
20<1-
6c:>
J.~
_/v-,_~." '_._6""._&:>,-6_."'O_&l'_ o
0.2
.._po.·._"_O_6.-6-o. .~O 0.4
0.6
0.8
p/po
Fig. 5. Pore size distribution in untreated silica B (a), catalyst L (b) and catalyst M (c) determined by Hg porosimetry. pore volume (cm3/ g)
0.750
rpore(nl'l)
37
TABLE 3 Pt characteristics in 3% Pt/Si0 2 determined by hydrogen chemisorption at ambient temperature Catalyst
Support
Dpt %
M
A A A B
0 Y
B
22 13 25 19 6
B
8
G
H X
2 -1 Spt (m .gpt) 59 36 69 52 17 21
apart (nm) 4.0 6.5 3.4 4.6 13.4 11.3
a par t average particle size measured by hydrogen chemisorption equals 2rpa r t • Samples X and Y were derived from catalysts I and N respectively in Table 2 by subsequent reduction in hydrogen. Platinum Dispersion and Location The dispersions of Pt on some catalysts prepared as described above measured by hydrogen chemisorption are shown in Table 3. Catalysts X and Y were obtained by reductions of samples I and N. As expected (see results of Dor1ing et a1. in 1967 and 1971 in ref. 1) higher metal dispersions were obtained when the higher surface area and smaller pore size silica (e.g. samples G, H and X which should be compared with M, 0 and Y respectively) had been used. However r pore > r part and it is not possible to say that r por e defines r par t obtained for the supported Pt particles (despite suggestions in Fig. Ib and ref. 3). Although the total surface area of catalysts G and H are very similar (209 and 211 m2.g- l), as are those of samples Mand 0, the micrographs in Fig.6 show that the samples impregnated at 278 K (e.g. H) have Pt located predominantly at the external edge of the support particles, while those prepared at 373 K (e.g. G) have the Pt particles located throughout the interior of the support. This may be a useful preparative parameter. DISCUSSION OF RELATIONSHIP BETWEEN PLATINUM DISPERSION-LOCATION AND SUPPORT POROSITY Rate of Adsorption and Solution Penetration Unfortunately, the fragmentation of the silica made direct AA analysis of the rate of adsorption of hexachlorop1atinic acid extremely difficult. Fig. 7 shows how the rates of penetration (dz/dt, where Z is the support particle radius, i.e. 0.75 or 1.25 mm here) of aqueous solutions into the pore structure of the silica samples with spheres 2.5mm diameter were readily
38
Fig. 6. Location of Pt in cross-sections of particles of catalysts G and H determined by microprobe analysis. The edge of the cross-section of the silica support is denoted ~ and the direction of liquid flow in impregnation by ~ . Impregnation at the lower temperature produces Pt preferentially at the outer edge of the support particle porous structure.
G
t
____- L--H
t
a
b •
•
• •
• • • • • -• •• • • • • •• • •
••
•
.•
••• •
• •
•
•
Fig. 7. Micrographs of silica A before (a)-and 15s after (b) addition to silica A of water at 293K to the point of incipient wetness. A few particles remained unwetted.
39
observed by the progress of an opaque-translucent interface into the particles as a function of time under isothermal conditions (273, 294 and 313 K). Rates of penetration decreased with time and penetration was initiated on different particles at different times; results were averaged over 40 particles. In all cases the rate of penetration decreased as temperature decreased (i.e. 0.237,0.071 and 0.010 mm.s- 1 at 273,294 and 313K at 15s). If the rate of adsorption over this narrow temperature range is largely constant then as the temperature decreases it would be expected that the Pt would be adsorbed to a greater extent at the external surface of the silica support structure. This appears to be the case as shown in Fig. 6. Certainly, silica is weakly acidic and its surface of silica is negatively charged (i .e. its PZC is at about pH 2) and adsorption of PtC1 62- is IJnlikely to be as extensive or as fast as on alumina (ref. 8). Assuming a cylindrical pore of radius r oor e exists from the surface of the silica sphere to its centre, then if the solution penetrates the pore in a equals streamline-laminar flow manner the driving force for flow TIr~ore'~P resistance to flow ~.dv/dr. 2TIr por e'z, where ~ is the dynamic viscosity of the solution, v is the velocity of the solution, z is the radial distance into the silica, and ~p is the pressure difference between pore entrance and termination. Integration gives the velocity v (or dz/dt) as r2.~p/4~z. However, ~p can be defined more precisely in terms of the pressure drop across the solution meniscus of radius of curvature R (i.e. ~p = 2y/R, where y is the surface tension of the liquid). Since the solution is in a pore of radius r por e then ~p (YCOs8/r pore' where 8 is the prevailing contact angle) is then the pressure for penetration of the pore during catalyst impregnation. dz/dt then equals rpore.y.COS8/4~z or apore.y.COS8/8~z on average during impregnation. The ratio y/~ is very temperature dependent. As temperature rises the time to penetrate a pore 8 nm radius (e.g. possibly one in silica A here) and one 48nm radius (e.g. possibly one in silica B here) decreases as the temperature rises and are 37.2 and 6.2 ~s respectively for A and B at 273 K and 157.8 and 9.5 ~s at 373 K. 1f the rate of PtC1 62- sorption is independent of temperature then rates of penetration could define the dispersion and location of the Pt in the catalyst. Naturally the model of impregnation is simplified to exclude the rate determining step being the rate of loss of air from the support pores; this may not be valid with pores open only at one end. Hence the observed rates of penetration during impregnation here are much slower than the maximum values predicted by the model as a result of this ~ffect.
40
Fig. 8. Tpr profiles of catalysts K, Land N and also E, F and I showing the higher impregnating temperatures produces a lower reduction temperature (as does sUbsequent calcination).
rate of reduction by H2 in tpr (arbitrary units)
~
:: "
"I,
"
II
II
I,
I I
I I
I'
,' I,
I I
'I
II
, , I
r
,
I
tI
\
\
\
N
~-
..._,-\ \
I
473
rv
373
I
,,I',,
......
I
,
\,
'-273
T(K)
41
CONCLUSIONS The silica surface area rather than the average pore size (or distribution) appears to affect the size and disperion of Pt produced by impregnation with hexachloroplatinic acid solutions. There is some pore blockage as a result of impregnation, but this is partly reversed by calcination and reduction and does not involve any significant depolymerisation-dissolution of the silica. despite the pH of the impregnating solution being as low as ~O.5; the solubility of silica rises with pH at these temperatures. It is likely that for other supports the acidity of this impregnating solution will be a more critical parameter. with much greater support modification. However. unexpectedly. the temperature of the impregnating solution does appear to be important in defining the location of the supported Pt particles. This may be a useful preparative parameter. since it does not modify or complicate the chemistry in a way that complexing agents do and may allow control of the degree of active phase penetration in catalysts (ref. 9). Fig. 8 indicates that impregnating at 373 K produces Pt in a state upon silica which is more easily reduced than that produced by impregnation at 273 K; calcination produces an even larger effect. Hence the temperature of impregnation can affect the chemical state and location of silica-supported platinum. ACKNOWLEDGEMENTS The authors gratefully acknowledge the provision of study leave for MAML by the Consejo (Madrid). REFERENCES 1 C.R. Adams. H.A. Benesi. R.M. Curtis and R.G. Meisenheimer. J. Catal .• 1 (1962) 336; T.A. Darling and R.L. Moss. J. Catal .• 5 (1966) 111; T.A. Dorling and R.L. Moss. J. Catal .• 7 (1967) 378; T.A. Dorling. C.J. Burlace and R.L. Moss. J. Catel .; 12 (1968) 207; P.C. Aben, J. Catal., 10 (1968) 224; H. Benesi. R.M. Curtis and P.S. Studen. J. Catal .• 10 (1968) 328; T.A. Darling. M.J. Eastlake and R.L. Moss. J. Catal., 14 (1969) 23; T.A. Dorling, B.W. Lynch and R.L. Moss. J. Catal .• 20 (1971) 190; J. Freel. J. Catal., 25 (1972) 139149; G.R. Wilson and W.K. Hall. J. Catal.. 24 (1972) 306; P.B. West, G.L. Haller and R.L. Burwell. J. Catal., 29 (1973) 486; D. Cormack, J. Pritchard and R.L. Moss, J. Catal •• 37 (1975) 548; G.C. Bond and B.D. Turnham. J. Catal .• 45 (1976) 128; S.R. Sashital. J.B. Cohen, R.L. Burwell and J.B. Butt. J. Catal .• 50 (1977) 479. 2 D. Pope. D.S. Walker. L. Whalley and R.L. Moss, J. Catal., 31 (1973) 335. 3 E. Ruckenstein and B. Pulvermacher, J. Catal .• 37 (1975) 416. 4 M.A. Martin Luengo, P.A. Sermon and A.T. Wurie. J. Chern. Soc. Far. Trans. (in press). 5 E.W. Washburn, Proc. Natl., Acad. Sci. USA. 7 (1921) 115. 6 B.F. Roberts. J. Coll. Int. Sci., 23 (1967) 266. 'Adsorption. Surface Area and Porosity'. Academic 7 S.J. Gregg and K.S.W. ~inq, Press (1982); K.S.W. Slng, D.H. Everett, R.A.W. Haul. L. Moscou, R.A. Pierotti, J. Rouquerol and T. Siemieniewska, Pure Appl. Chem.57 (1985)603. 8 E. Santacessaria, S. 'Carra and 1. Adami, Ind. Eng. Chern. Prod. Res. Dev. 16 (1977)41. 9 N.M. Ostrovskii, L.A. Karpova and V.K. Duplyakin, React. Kin. Catal. Lett. 26 (1984) 279; L.L. Hegedus, T.S. Chou, J.C. Summers and N.M. Potter,'Studies in Surface Science and Catalysis 3, p. 171, edited B. Delmon, P. Grange, P. Jacobs and G. Poncelet (1979).
42
DISCUSSION D.E. RESASCO : When you vary the impregnation temperature to study its influence on the dispersion and location of Pt, you also vary the temperature at which the drYin~ step occurs. The rate of drying at 273K or 373K will be much affected bY t e temperature. Do you think that this step, rather than the adsorption itself, could be the dominating factor in determining the dispersion and location of Pt? P. SERMON: It was the intention of the present work to show how the impregnation - drying temperature affected the dispersion and location of silicasupported Pt; this it appears to show quite well. However, you are right to highlight the fact that the locational-dispersional difference could arise from the rate of pore penetration or the rate of drying. The drying was carried out either at 373K or 273K in vacUO. Such drying steps were very fast and so we believe the pore penetratlon step is the more critical. J.A. SCHWARZ: Could the higher temperature during impregnation lead to a more complex speciation of the PtC12 6 anion and thus account for the differences in the TPR, spatial distribution and particle size? P. SERMON: Yes, it would be possible for the higher impregnation temperature to cause a greater exchange of Cl- ligands by surface -OH groups in the Pt anion. This would modify the TPR and metal dispersion-location observation if this species remains in a higher dispersed state. Alternatively, the greater rate of pore penetration at the higher temperature leads to the Pt anions being adsorbed more uniformly over the entire silica surface. As a result there is a greater chance of monodispersed Pt adsorbed species with greater involvement of OH-group ligands .. As a result the TPR will be modified from that obtained with lower rates of impregnation and poorer dispersions. These two are difficult to separate experimentally except we should possibly use spectroscopy to confirm the presence of PtCl 2that 6 (aql at 273K and 373K. Indeed, the reality is probably a mixture of the two. Perhaps what matters more is that we have tentatively shown here that a new preparative parameter exists - which is easy to use and which may be chemically simpler than the use of complexing agents. We hope that in the future this will prove useful in the preparation of heterogeneous catalysts; the precise causes of the effect remain to be resolved. D. REINALDA : To study the relationship between metal particle size and pore diameter 3% metal loading seems fairly low. Have you studied the effect of metal loading? P. SERMON: Commercial heterogeneous catalysts contain low transition metal loadings for economic and activity reasons. We believe the results obtained for these 3% Pt/Si02 catalysts do illustrate the effect of impregnation temperature well. Rowever you are right to suggest that measurements at higher Pt loadings might reveal the relationship between dpart and dpore more clearly. This work has not yet been completed. J. MARGITFALVI : What is the influence of the conditions of pre-wetting on the adsorption of the platinum compound? P. SERMON: If the silica is pre-wetted to the extent of leaving the pore filled with water and then the PtC12 6 solution is added, it is very difficult to follow adsorption to equilibrium since this is partly controlled by the slow ionic diffusion across this interface and within the pores. However, normally the impregnation step takes place with pores filled with air which has to be
43
displaced be solution entering. This has been used here. The silica here had surfaces which were fully hydroxylated; the degree of surface hydroxylation would have an effect on the extent and nature of PtC12 6 adsorption by this has not been studied here.
This page intentionally left blank
B. Delmon. P. Grange. P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV © 1987 Elsevier Science Publishers a.v., Amsterdam - Printed in The Netherlands
45
PREPARATION AND PROPERTIES OF THE CATALYSTS BY A CHEMICAL MIXING PROCEDURE F. MIZUKAMI, S. NIWA, M. TOBA, T. TSUCHIYA, K. SHIMIZU, S. IMAI and J. IMAMURA National Chemical Laboratory for Industry, Yatabe, Ibaraki 305 (Japan)
SUMMARY Various solid catalysts were prepared by a chemical mlxlng technique, composed of four steps (complexing, gelation, drying and activation) and characterized by the preparation of a homogeneous solution containing catalyst precursors and the uniform coagulation of the solution through hydrolysis. The mixed-oxides prepared by the technique were expected to be more homogeneous than those prepared by kneading and coprecipitation methods. The chemically mixed ruthenium catalysts were much more effective for the partial hydrogenation of benzene to cyclohexene without any assistance from poisons. Furthermore, the chemically mixed alumina-silicas showed much higher activities for the conversion of methanol to hydrocarbons than the kneaded a1umi na-si 1i ca. INTRODUCTION The pore size, surface area and homogeneity of solid catalysts are very important factors governing the heterogeneous catalytic reactions. Accordingly, the methods for controlling the pore size and surface area, and the procedures for preparing highly homogeneous catalysts have been groped because it is difficult to obtain the solid catalysts with controlled structure and uniform composition by traditional methods such as kneading, impregnation and coprecipitation. Recently, Ueno et al (refs. 1,2) and Nobe et al (refs. 2,4) have presented the interesting preparation techniques of catalysts, which involve the use of distillable organometalic intermediates and accompany chemical reactions. Such techniques seem to have the potential for giving catalysts with intimately mixed components. It has also been pointed out that natural gelation is effective for the production of mixed-oxide catalysts with homogeneous compositions. However, natural gelation exists only in a limited number of cases because the gelation operation requires that the mixed-oxide precursors formed in the solution must have a lyophilic nature (ref. 5). Intending to obtain the uniform solid catalysts, we have tried to combine the above-mentioned methods and to gelatinize smoothly without any precipitates in the solution by complexing the catalyst precursors with polar solvents having two or three functional groups (refs. 6-8). In this chemical mlxlng method, some organic materials seem to be enclosed in the colloidal polymers
46
of the catalyst precursors during gelation. Such organic material may be responsible for the formation of the surface structure of the solid catalysts (pore size and surface area). This paper will indicate that the improved method has the potential for providing the solid catalysts with homogeneous composition and controlled surface structre, and the catalysts prepared by the method have very interesting properties. EXPERIMENTAL Preparation of catalysts The chemical mixing method consists of complexing, gelation, drying and activation steps as shown in Fig. 1. The typical procedures are as follows.
step 2
step 1 1MXn _",.---'''--~
solution of t----"'--~--7I metal complexes
i acid or base
IE---tdry gel
'--_---I
step 6
step 5
mixed solution of metal complexes
1E:,------1 gelation
~--1coagulum
step 4
highly viscous solution step 3
Fig. 1. Preparation procedure of the catalysts by a chemical mixing technique (i) 2wt%Ru-O.2wt%Cu-Si0 2. One gram of ruthenium chloride hydrate and 0.065 g of copper chloride hydrate were dissolved in 20 cm 3 of ethanol at room temperature. To the sol uti on was added 50 g of ethyl ene glycol (EG), and the solution was mixed at 338 K for 0.5 h. Then, 63.3 g of tetraethoxysilane was added to the solution and stirred at the same temperature for 3 h. The tetraethoxysilane was immiscible with the ethylene glycol solution immediately after the addition, but it was gradually getting mixed and the homogeneous transparent solution was obtained at last. To this homogeneous solution was added 22 g of water. The solution was getting viscous and finally coagulated in agar. For the complete gelation, the coagulum was warmed at the same temperature for 0.5 h after the coagulation. The coagulum was dried under reduced pressure at 373 K for 24 h. The dry gel was finely powdered. Then, the powder was put in a quartz tube and was activated in a stream of hydrogen at 673 K for 8 h. (ii) 2wt%Ru-D.2wt%Cu-A1 Z03. Fifty grams of 2-methyl-2,4-pentanediol hexylene glycol, HG) was added to 10 cm 3 of ethanol solution containing
47
ruthenium chloride hydrate (1 g) and copper chloride hydrate (0.065 g) and stirred at 338 K for 0.5 h. Ninety-four grams of aluminium sec-butoxide was added to the solution and stirred at the same temperature for 3 h. To the homogeneous dark green solution obtained was slowly added 50 cm 3 of ethanol solution containing 6 g of water. After the solution was stirred for 1.5 h, 3 30 cm of ethanol solution containing 8 g of water was added to the solution. The solution was coagulated in jelly after a while. The coagulum was dried under reduced pressure at 373 K and finely powdered. The powder was activated in a stream of hydrogen at 673 K for 1 h. (iii) 10.2wt%A1 203-Si0 2. Eight grams of aluminium trichloride anhydride was dissolved in 15 cm 3 of ethanol. Seventy-nine grams of EG and 93.4 g of tetraethoxysilane were added to the solution and stirred at 353 K for 3 h. Thirty-six grams of water was added to the homogeneous solution. After a while the solution was coagulated in jelly. The coagulum was dried under reduced pressure at 373 K and finely powdered. The powder was calcined at 823 K for 12 h.
(iv) 30wt%Ti0 2-Si0 2. Fifty grams of l,2-cyclohexanediol (l,2-eHD) and 17.2 g of titanium tetraethoxide were dissolved in 50 cm 3 of ethanol. To the solution were added 48.7 g of tetraethoxysilane and 4.3 g of methanol containing 10wt% of hydrogen chloride, and the solution was stirred at 353 K for 3 h. Then, 11.5 g of water was added dropwise to the solution. The solution became viscous and finally coagulated in agar. The coagulum was dried under reduced pressure at 373 K and finely powdered. The powder was calcined at 823 K for 12 h. Measurement A Philips PW 1700 diffractometer equiped with a curved graphite monochrometer using Cu Ka-radiation was utilized to obtain X-ray diffraction XRD) over the range of 28=5-60°. RESULTS AND DISCUSSION Preparation of the catalysts The chemical mixing method was thought out as a technique for transfering the uniformity of homogeneous solutions containing various catalyst components to the solid state. In this method, thus, it is essential to prepare a homogeneous solution containing the catalyst precursors and to coagulate uniformly the solution. In order to satisfy the above requirements, the search for the appropriate catalyst precursors and solvents was carried out, and when the catalyst precursors were extremely different from one another in their hydrolysis rates, it was tried to level their reactivities by complexing with
48
polar solvents, or increasing their ligand exchange rates with addition of a acid or base. Generally, diols, ketoalcohols and aminoalcohols were suitable as solvents for the preparation of catalysts. Soluble metal salts such as nitrates, chlorides, acetates and alkoxides could be used as precursors for minor catalyst components in amount. For major catalyst components, metal alkoxides and some a-diketone complexes (for example, aluminium dialkoxy acetoacetic ester chelates), which can be easily gelatinize by hydrolysis, were suitable as precursors. However, aluminium, titanium and zirconium alkoxides very often formed insoluble precipitates with certain polar solvents as soon as the metal alkoxides were added into the solvents. Accordingly, the combinations of solvents and metal alkoxides to get homogeneous solution were examined as shown in Table 1. The table indicates that ketoalcohols, aminoalcohols and multibranched diols are suitable solvents when aluminium, titanium and zirconium alkoxides are used as catalyst precursors. TABLE 1 Combinations of metal alkoxides with solvents Solvent Ethylene glycol 1,2-flropanediol a 1,2-Butanediol a 2,3-Butanediol b Pinacol 1,2-Cyclohexanediolb l,3-Cyclohexanediol b 1,4-Gyclohexanediolb 1,3-Propanediol 1,S-Sutanediol a 2,4-Pentanediol b 3-Me-l,J-Butanediol Hexylene glycol a l,4-Butanediol Glycerine Diacetone alcohol Ethanolamine Propanolamine
Si (OEt)4 Al(Oi-Pr)3 Al(OsecrBu)3 Ti (Oi-flr)4 Zr(On-flr)4 no ppt. no ppt , no ppt. no ppt. no ppt. no ppt. no·ppt. no ppt. no ppt. no ppt. no ppt. no ppt. no ppt. no ppt. no ppt. no ppt. no ppt. no ppt.
ppt. ppt. ppt.
ppt. ppt.
ppt ,
ppt ,
no ppt. ppt.
no ppt. ppt. ppt. no ppt. ppt. ppt. ppt. no ppt. no ppt. ppt. ppt. no ppt. no ppt. no ppt.
ppt ,
no ppt. ppt. ppt. ppt. no ppt. no ppt. ppt. ppt. no ppt. no ppt. no ppt.
ppt ,
ppt. ppt. ppt. ppt. no ppt. no ppt. no ppt. no ppt. ppt. ppt. no ppt. ppt. no ppt. ppt. ppt. no ppt, no ppt. no ppt.
ppt. ppt. ppt. ppt. no ppt. no ppt., no ppt. no ppt. ppt. ppt. ppt. ppt. no ppt. ppt. ppt. no ppt. no ppt. no ppt.
ppt.: Combination which a precipitate occured, when a alkoxide was added into a solvent at room temperature. a Racemic isomers. b A mixture of racemic and geometrical isomers.
49
Homogeneity Generally, if all components in the solid composites are completely mixed to each other, it will be difficult for specific components to aggregate and crystsllize, because those components are surrounded by others. Accordingly, it may be possible to know whether the solid composites are homogeneous or not by inspecting the crystallinity of a given component in the composites. Figure 2 shows the X-ray diffraction patterns of seven metal oxides. Samples 1-4 were prepared by the chemical mixing technique, and samples 5 and 6 were prepared by coprecipitation and kneading methods, respectively. Sample 7 was prepared by precipitation method with ammonia as a standard of Ti02. Samples 1-4 do not show any clear diffraction peaks and are found to be amorphous. Samples 5-7 show the typical diffraction patterns of anatase Ti02, and samples
I~
(l)
(2)
,--
(4)
~
r--
(5) (6) I
(7)
-
1--
(3)
A50
- 1-"'-'lAo.
40
oJ
A
1.../
1
I
.;
20
30
500 cps
10
26
Fig. 2. X-Ray powder diffraction of mixed-oxides. (1), 3.3wt%A1 203-Si02 (raw materials, A1C1 3 and Si(OEt)4; solvent, 2,3-butanediol); (2), 30wt%Zr0 2-Si02 ( Zr(On-Pr)4 and Si(OEt)4; 1,2-CHD); (3), 3Owt%Ti02-Si02 (Ti(OEt)4 and Si(OEt)4; 1,2-CHD); (4), 50wt%Ti0 2-A1 203 (Al(Osea-Bu)3 and Ti(Oi-Pr)4; pinacol); (5), 50wt%Ti0 (NH precipitation); (6), 50wt%Ti0 2-A1 203 (kneading); (7). Ti02 2-A1 203 3 (NH 3 precipitation). Calcination: (1)-(4). 823 K, 12 h; (5)-(7). 823 K. 5 h.
50
5 and 6 are understood to have crystalline particles, that is, clear aggregates of Ti02. But. with the chemically mixed samples 2-4. any clear crystalline particles of Ti02 and Zr0 2 could not be observed by X-ray diffraction, although those samples were calcined enough at the temperature which Ti02 and Zr0 2 can easily crystallize. Especially. as samples 4-6 have the same composition. the above facts suggest that the samples prepared by the chemical mixing technique are more homogeneous than those prepared by the traditional methods of kneading and coprecipitation. The reason why the chemical mixing procedure produces relatively homogeneous solid composites will be considered as follows. At the complexing and gelation steps. the reaction (1)-(12) occur, and different components are uniformly incorporated with each other in colloidal polymers containing solvents, finally the coagulum with three-dimensional network (Fig. 3) is produced via twoM(OR)n + M' (OR)k
(l)
M(OR)n_x(OR'OH)x + xROH
xHOR'OH--~)
+ yHOR'OH
M' (OR)k -y (OR'OH) Y + yROH
~
I
J
M(OR)n_x(OR'OH)x + M'(OR)k_y(OR'OH)y
--~)
(2)
(L=OR. OR'O. I
M(OR)n_x(OR'OH)x + H20 ------~)
(3)
+ HZ
~~'-
Z=OR. OR'OH)
+ ROH (or HOR'OH)
~H
(4)
, M' (OR\_y(OR'OH)y + H20 --~) I
I
-M-L-M'- + H20 ------~) I I I
I
I
)
I
I
I
I
I
I
I
I
I
~'~H
+ I
I
I
I
--fj1-01'I
+ H20 I
(8)
I
) --M--O--M-L-M'- + H20 I I I
HD-M- + HO-M'M' (LH)
(7)
I
-M-OH + HO-M-L-M'I
(6)
I
I
I
I
I
-M-OH + M' (LH)
--~)
+ H
---rt-DH
I
I
(5)
HQ-M-l-M'- + ROH (or HOR'OH)
I
-M-l-M'- + H20 I I
-t;1'-{)H + ROH (or HOR'OH)
I
I
I
I
(9)
) HQ-M-O-M'- + ROH (or HORrOH) ---7)
I
--r!l'-l--h--o--A-oH + ROH (or HOR'OH) I
I
I
(10)
I
(11 )
I
I
I
I
I
I
-M '-l-M-O-M- + HO-M '-o-M- + 2H 20 ---.;..) -M '-l---M---O-M-O-M '---{)-M---{)H + I I I I I I I I I I H20
+
ROH (or HOR'OH)
(12)
51
I
I
OR
OR'OH
I
I
- M - OR'O - M - 0 - M - 0 - M - 0 -
oI
I
HX
0 ROH
I
I
- M - OR'O - M'-O -
I
o
HOR'OH
I
I
-M'- 0
I
I
OR OR'OH
-
6R"OM X
0
?R
I
I
I
I
M - OH
X-M'-OR'OH
I
I
0 ROH
I
M- 0 - M - 0 -
M -
I
J:
I
I
I
J:
I
I
o
H20 0
-
.0 HOR'OH 0 00=0 R' , 0 0 HOR OH
I MI
o I
M-O-M-O-M-O-M'-ORHO-M-
I
I
I
I X
I
Fig. 3. Probable structure of coagulum, dimensional polymers, On such polymerization, the homogeneity of composition in the liquid state seems to be maintained at the gel stage and finally reflected in the solid catalyst, Surface structure The dry gels before activation by the chemical mixing procedure (step 5 in Fig. 1) contain the solvents used for the catalyst preparation and the organic residues derived from the solvents and the alkoxy groups of starting materials. These solvents and organic residues, especially organic residues, not only affect the surface area of the final catalysts but contribute to forming pore. This detail has already stated in the previous reports (refs. 7, 8). Reactions Table 2 shows the highest cyclohexene yields and the concomitant benzene conversions observed in the hydrogenations of benzene with different ruthenium catalysts (ref. 7). The most important feature in the hydrogenations is that, in the formation of cyclohexene, the existence of corrosive additives in the reaction mixture is essential for the impregnated catalysts and is unnecessary for the chemically mixed catalysts. Table 3 shows the conversions of methanol with various alumina-silicas (ref. 8). Different from the kneaded and purchased alumina-silicas, the chemically mixed alumina-silicas are active for the conversion of methanol to hydrocarbons containing olefins in spite of their amorphism (see sample 1 in Fig. 2), although their activities are not so high as that of ZSM-5. Furthermore, it is found that the activities of the chemically mixed alumina-silicas for the formation of hydrocarbons depend on the solvents and starting materials used in the preparation of the catalysts.
52
TABLE 2 Partial hydrogenation of benzene with different ruthenium catalysts preparation No. Catalyst Solvent Procedure 1 2 3 4 5 6 7 8 9 10 11
EG Chem. mix. EG Chem. mix. Pinacol Chem. mix. Chem. mix. Pinacol Chem. mix. HG ~1e-l,3-RD Chem. mix. Chem. mix. DA EA Chem. mix. Impregnationa Impregnation a Impregnationa,b
Cyclohexene Catalyst composition yield (mol%) 2wt%Ru-Si0 2 2wt%Ru-O.2wt%Cu-Si0 2 2wt%Ru-A1 203 2wt%Ru-o.2wt%Cu-A1 203 2wt%Ru-O.2wt%Cu-A1 203 2wt%Ru-o.2wt%Cu-A1 203 2wt%Ru-o. 2wt%Cu-Al 203 2wt%Ru-O.2wt%Cu-A1 203 2wt%Ru-Si 02 2wt%Ru-A1 203 2wt%Ru-A1 203
27.0 31.4 14.2 24.3 24.5 23.7 17.4 12.2 8.6 5.3 27.2
Benzene conversion (%) 86.8 83.3 42.7 74.1 53.1 81.2 46.4 43.1 64.3 54.5 72.7
Chern. mix., chemical mixing procedure; Me-l,3-£D, 3-methyl-l,3-butanediol; DA, diacetone alcohol; EA, ethanolamine. Reaction coditions: benzene (160 cm 3), water (100 cm 3), catalyst (2g); reaction temperature, 453 K; reaction pressure, 7 MPa. a These catalysts were purchased from Nippon Engelhard Ltd. b Together with this catalysts, 5 g of CoS0 was added in a mixture of benezene 4 and water. TABLE 3 Methanol conversion with different alumina-silicas and ZSM-5 at 653 K No. 1 2 3 4 5 6 7 8
Catalyst preparation Solvent Raw material
Catalyst composition
Methanol Yield conv. (%) T.H.C. (C 2H4 + C 3H6) HG 3.3wt%A1 203-Si02 83.5 4.4 0.1 Si(OEt)4' A1C1 3 Pinacol 3.3wt%A1 11.5 2.7 Si(OEt)4' A1C1 3 203-Si02 82.9 13.8 2.6 Si(On-Pr)4' A1C1 3 1,2-4'0 3.3wt%A1 203-Si02 83.9 2,3-£0 3. 3wt%Al 203-Si 02 84.0 26.2 5.7 Si (OEt)4' A1C1 3 EG 27.6 6.1 Si(OEt)4' A1C1 3 3. 3wt%Al 203-5i 02 84.6 Kneadi ng 2.3 2.3 0.0 3. 3wt%Al 203-Si02 Purchased (Nikki N632HN) 0.1 28wt%A1 203-Si02 83.4 4.4 ZSM-5 Si02/A1 203=40 99.7 99.7 7.2
1,2-PO, 1,2-propanediol; T.H.C., total hydrocarbons. LHSV=2; Ar/CH 3OH, 50/50. Samples 1-5 were prepared by the chemical mixing procedure. REFERENCES A. Ueno, H. Suzuki and Y. Kotera, J. Chern. Soc., Faraday Trans. 1, 79 (1983) 127.
53
2 S. Takasaki, H. Suzuki, K. Takahashi, S. Tanabe, A. Veno and Y. Kotera, J. Chem. Soc., Faraday Trap 1, 80 (1984) 803. 3 I.M. Pearson, H. Ryu, w.e . ..ong and K. Nobe, Ind. Eng. Chem. Prod. Res. Dev., 22 (1983) 381. 4 W.C. Wong and K. Nobe, Ind. Eng. Chem. Prod. Res. Dev., 23 (1984) 564. 5 P. Courty and C. Marcilly, General synthesis methods for mixed oxide catalysts, in: P.A. Jacobs and G. Poncelet (Eds.), Preparation of Catalysts, Elsevier, Amsterdam, 1976, pp. 119--145. 6 S. Niwa, F. Mizukami, M. Kuno, K. Takeshita, H. Nakamura, T. Tsuchiya, K. Shimizu and J. Imamura, J. Mol. Catal., 34 (1986) 247. 7 S. Niwa, F. Mizukami, S. Isoyama, T. Tsuchiya, K. Shimizu, S. Imai and J. Imamura, J. Chem. Tech. Biotechnol., 36 (1986) 491. 8 F. Mizukami, Y. Kiyozumi, S. Niwa, S. Shin, J. Imamura, K. Noguchi and K. Matsuzaki, submitted for publication (J. Catal.).
54
DISCUSSION D. REINALDA : In the Ru/A1203 and Ru/SiOZ catalysts, is, after activation, all the metal exposed to the gas phase or is part of it retained in the matrix of the support ? and Ru/SiOZ prepared by a chemical mixing procedure F. MIZUKAMI : RU/A1ZO~ showed nigher activltles in the partial hydrogenation of benzene than the corresponding impregnated catalysts. So, we have thought that the amount of Ru in the support matrix is not so high, although a part of it is retained in the matrix. J. B.-NAGY: In TiOZ-SiOZ' you have supposed that the metal atoms are tetrahedrally coordinated. Do you have any evidence for the number of coordination of titanium atoms in the coagulum? F. MIZUKAMI : At present, we do not have any evidence which supports the existence of tetrahedrally coordinated titanium. Probable structure of coagulum only shows schematically one of the causes why homogeneous mixed oxides are formed. Rather, we think that the coordination number of titanium in the coagUlum is six because the coagulum contains many water, mono-alcohol and diol molecules which can easily coordinate to titanium ions. J. KIWI : At which temperature could you observe the disappearance of the alkoxide of Z% wt Ru - O.Z% wt CuSiOZ prepared catalyst in ethanol? What physical technique did you use for this observation? F. MIZUKAMI : The disappearance of the alkoxide of Z wt %Ru - O.Z wt %Cu-SiOZ prepared by a chemical mixing procedure was investigated by thermal gravimetric analysis (TGA) and it showed that the alkoxide is lost at about 20DoC (see J. Chern. Tech. Biotechnol. 36 (1986) 236). G.M. PAJDNK : If I understood correctly, your chemical mixing technique water plays the role of a reactant (and not of a solvent as usual). Your silica-alumina catalysts exhibit good properties in the MTG synthesis (methanol to gasoline) that are likely due to their acidities. Do you have information relative to this acidity (nature and strength) ? F. MIZUKAMI : As you have guessed. we have thought that the properties which the silica-aluminas prepared by a chemical mixing procedure had shown in MTG synthesis closely relate to their acidities. So, we are examining at present their acidities by means of NH3 TPD, titration and calorimetry. We will be able to report those data in the near future.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
55
© 1987 Elsevier Science PublishersB.V., Amsterdam - Printedin The Netherlands
Electrochemically Controlled Deposition-Precipitation. A New Method for the Production of Supported Catalysts. P.C.M. van Stiphout, H. Donker, C.R. Bayense and J.W. Geus Department of Inorganic Chemistry, University of Utrecht, Croesestraat 77A, 3522 AD Utrecht, The Netherlands. F. Versluis Harshaw chemie b.v. Strijkviertel 67, de Meern.
ABSTRACT A new method for the homogeneous deposition-precipitation of catalyst precursors has been developed. The precipitating agent is generated at the anode and/or the cathode in an electrochemical cell. Results of the preparation of nickel-on-silica are presented. Extensive characterization has been done and a comparison is made with a catalyst prepared by decomposition of urea. Catalysts made by the new method were found to be identical with those prepared by the conventional deposition-precipitation technique.
INTRODUCTION To produce catalysts of a high loading of the support with active
material
impregnation and drying is generally not suitable. Consequently coprecipitation of the support and the active precursor is mostly utilized to prepare catalysts exhibiting highly
an
loaded
elevated active surface area per unit volume, supports.
However,
catalysts is difficult to control and incomplete. This
Therefore
procedure allows
the
the
porous
structure
which calls of
for
coprecipitated
reduction to an active metal can
method of deposition-precipitation was
remain
developed.
achievement of a high loading of a separately
produced
support with a very uniformly distributed active precursor. Deposition-precipitation concentration to
be loaded.
precipitate bulk
is
carried
out
by
homogeneously
raising
of an active precursor dissolved in a suspension of the at
the
support
If the active precursor precipitates and if nucleation
of
the surface of the support proceeds more rapidly than
in
the
the active precursor is deposited exclusively onto
the
of the solution,
the
support provided the concentration is increased homogeneously. Two
different
precipitation reaction
from
general a
procedures
can
be
homogeneous solution onto
used a
to
effect
suspended
the
support,
of a dissolved compound to a precipitating agent and injection
above viz. of
a
56
precipitating procedure in
a
agent
solution
render the
into
the
suspension of
the
support.
With
the
the reacting compound is dissolved into a suspension of the of the active precursor at a temperature sufficiently
the rate of the reaction insignificant. suspension
has
been
homogenized
first support low
Raising the temperature
brings
about
an
increase
to
after
in
the
concentration of the species to be precipitated completely uniformly throughout the
solution.
below
the
precursor. stresses
The
second procedure uses injection of a
precipitating
surface of a suspension of the support in a solution of Since
agent
an
active
below the surface of the suspension sufficiently high
can be established,
shear
a rapid distribution of the precipitating
agent
can be affected owing to which the concentration increases homogeneously too. Both
of the above procedures have their limitations.
previously
dissolved
difficult
to
With reaction
compound a decrease of the pH value,
achieve.
for
Though the injection procedure can
be
of
instance, utilized
a is
more
universally, it must be carried out carefully and asks for a vigorous agitation of the suspension. When
the
deposition-precipitation
is carried out in
an
electrochemical
cell, the pH value can be increased or decreased and the valency of the ions to be
precipitated can be adapted to the precipitation process.
cathode anode
Thus around
of an electrochemical cell the pH value can be raised and
around
the electrochemical consumption of hydroxyl ions leads to a decrease
the the in
the pH value. Besides the very accurate control that an electrochemical process allows,
an
electrode
of
a large surface area can be
vigorous
agitation
around
species
generated
at the electrode throughout
stable
valency
of
precipitation,
used.
the electrode assures a rapid the
active precursors is often
electrochemical
reduction
suspension.
unfavourable
can
A sufficiently
dispersion
adjust
for
the
of
the
Since
the
deposition-
valency
for
a
subsequent deposition-precipitation onto the support. An important that
the
advantage of the utilization of an electrochemical
dissolution
of
a
precipitation onto a support. applied while onto
onto the
a
can
be
combined
with
its of
removal of hydrogen ions at the cathode brings about to
a
support.
cell
is
deposition-
Thus using one or more anodes of the metal to be
the support causes a gradual anodic dissolution
suspended
according
metal
Using electrodes
of
predetermined sequence affords the
different
the
precipitation
metals
production
metal,
of
switched catalysts
containing e.g. alloys. Discharge
of metal ions to be applied onto the support at the cathode
and
deposition
of the metal atoms thus formed onto the cathode can be prevented in
two ways,
viz. mounting a membrane permeable for hydrogen ions only around the
cathode, and covering the cathode by a chromium oxide layer. The new procedure will be illustrated by the deposition of nickel. For iron
57 on silica a similar procedure is applicable and precipitation by raising the pH was
also
utilized
to deposit zinc hydroxide onto
an
alumina
carrier.
The
precipitation of nickel(II) and copper(I) onto the same support was studied well.
Results
ourselves
of
these
experiments will be published.
Here we will
to the principles of the procedure and the deposition of
silica.
With
precipitated importance.
deposition-precipitation onto
We
the
surface
of the
the
transport
suspended
of
carrier
of
limit
nickel
species is
as on
to
be
paramount
therefore will consider especially the boundary layers
around
the electrodes. The
catalysts
absorption
were characterized by several
spectroscopy
temperature-programmed
to
confirm
reduction,
the
metal
magnetic
techniques. loading.
We We
measurements
did
atomic
further and
used
activity
measurements to compare the preparation techniques used. THEORETICAL The control of the anode and cathode reactions is based on Nernst's law for electrode reactions in
aqueous solution.
EO + (RT/nF)ln(aox/ared)
Eel
EO
where
(1)
is the standard reduction potential
Eel is the electrode potential F is the Faraday constant a ox is the activity of the oxidized species ared is the activity of the reduced species The potential between the anode and cathode is given by equation (2)
E
(2)
Ecathode - Eanode
Unless activity
activity coefficients are explicitely mentioned we will assume constants to be unity,
by concentrations.
the
which implies that activities can be replaced
If E is positive the reaction is spontaneous,
otherwise
a
potential has to be applied to accomplish the reaction. In figure 1 a schematic representation of an electrochemical cell is given. The positive electrode is the anode and the negative electrode the cathode. We have to consider the extra potential needed to overcome the (activation resistance
polarisation), of the system.
the
concentration
polarization
and
overvoltage the
The overvoltage is caused by the activation
ohmic energy
barrier of the electrode reactions, the concentration polarization is dependent on
the
concentration
of the active species near
the
electrodes.
Both
the
58
activation polarization and the concentration polarization are dependent on the current
density
[l,ZJ.
The
ohmic resistance can be minimized
by
adding
a
supporting electrolyte to the solution.
8
Figure 1.
A schematic representation of the electrochemical cell.
Now
we will
pay some attention to the precipitation
at
increasing
pH.
Metals which can be anodically oxidized are those that are more easily oxidized than the solvent i.e. on
water. The range of metals which can be oxidized depends
the values for the overvoltage for the oxygen evolution on the metal anodes
considered.
The
precipitation
of
metal compounds is often
limited
to
stability region of either the carrier or the precipitate to be deposited the
carrier.
As already mentioned this method provides
precipitating of vanadium,
the
possibility
ions of other valency and thus to control the stability molybdenum etc.
the onto of
regions
Here we use the nickel-water-silica system as an
example of an electrochemically prepared catalyst. The reactions at the electrodes are :
-->
Ni
Ni Z+ + Ze
Z HZO + Ze
-->
-0.Z3 V
Z OH- + HZ
o
E HZO/HZ
= -0.8Z8
(3)
V
(4)
The potential to be applied for the reaction is given by
E
= Eanode
- Ecathode
= (5)
The
problem arising when we oxidize nickel is the subsequent reduction
of
59
the
nickel
hydrogen
ions
at the cathode.
evolution
at
This reaction is more favourable
higher current densities and at
high
pH
than
prevent this reduction the cathode compartment can be separated from the compartment Another onto
by
means of a diaphragm,
the
cathode.
permeable
viz.
by a proton
is to use a chromium oxide layer
method
the
values.
To
anode
permeable
membrane.
electrochemically
deposited
The Nafion Nl17 [3,4] perfluoro sulfonic acid membrane
for protons and water,
metal ions of copper,
to a much less extent for hydroxyl ions
zinc and cadmium and impermeable for metal ions such
is and as
chromium(III), iron(III) and nickel(II). The of
Pourbaix diagram of nickel [5] helps to understand the
oxide
and hydroxide formation in aqueous solutions (see
thermodynamics
figure
2).
When
silica is present the solubility product of the nickel hydrosilicate limits the value
of
the
log-term in equation (5) and in the Pourbaix
diagram
the
Ni-
Ni(OH)2 equilibrium is replaced by the nickel-nickelhydrosilicate equilibrium. -2 2.2 EIV)2
a
-I
I
2
J
4
6
6
7
8
9
10
NiO~--?
II
13 14
12
15
,
1,6
,, , ,,
1.4
1,6 1,4
Nt
1,2
16 2.2 2
1,2
O,B
0.8
0.6 0,"
NiH
0,2
-e, -0,2
-----
-0,2
-- : --
-0,4 -0,6
-0,4 -0,6
- ........;.._ 2
-0,8
i
-I
.....
.
-
Ni
-1.2
f
-1,6 -1,8
°
I
2
J
4
5
6
7
8
S
10
-1,4 -1,6 11
12
13
14
-1,8 15 p H'6
The E-pH diagram of the system Ni-HZO.
Figure 2.
When we want to precipitate catalyst precursors, various
-I
-1,2
-1,4
-2 -!
-0,8
species
attention
are
of
great
to the static layer around the electrodes.
considered,
viz.
the concentrations of the
importance and therefore
we must
pay
Two effects have
some to
be
the electric double layer and the laminar layer dependent on
the degree of turbulence. We first consider the existence of an electrical double layer. a
schematic
cathode
there
representation of a diffusion layer is given. is
a region in wich the electrolyte
is
In figure 3
At both anode
concentrated.
If
and we
60
consider the Stern model [6] of a layer at a cathode there is a Helmholtz layer where
the
electrons
cations are concentrated and partially compensating the at the cathode.
excess
Outside this layer there is a diffuse double
of
layer
where the solution has an excess of cations. The outer Helmholtz plane (OHP) is the boundary between these layers.
¢m is the potential at the electrode, P2 is
the potential at the OHP.
Figure 3.
The Stern model and the corresponding potential profile.
The
thickness
of the Helmoltz layer is limited to a few nm
and
at
high
concentrations to only a few Angstrom [7]. Without knowledge of the value of;f2 we cannot estimate the exact concentration profile in the diffuse double layer. However
from
the literature [8] it is known that the effects of
the
diffuse
double layer are also confined to only a few nm. Besides the electrical double layer there is also a layer between electrode and
bulk
possible the
solution were there is no turbulence.
Convection is
and migration can only proceed by diffusion.
boundary
shear
(6)
Re is the Reynolds number;
,
(Jis the density of the solution; is the dynamic viscosity;
D is the diameter of a cylindrical electrode
or a wire diameter; Ux is the velocity;
not
Reynolds [9] postulated
to be where the potential energy equals the
ratio is the Reynolds number
where
therefore stress.
The
61
After
measuring Ux we can calculate the Reynolds number. We can measure the height of a waterco1umn in a small tube placed in the stirred solution. We use
a simplified equation of Bernouilli to find the velocity. (7)
a*h h is the height of the watercolumn in the tube a is 9.81 m/s 2 For
a
cylindrical
obstacle
of
lenght L and
diameter
D (L»D)
the
drag
coefficient (Cw) is 1.2 and (8)
where F is the force exerted on the obstacle; L*D is the largest cross-sectional area of the obstacle perpendicular to Ux ; with the relation
(9)
and
(10)
where
f is the thickness of the diffusion layer (11)
we can write
(12)
from which we can derive or If
bf gas
=
DI1/0.6Re
(13)
evolves at the electrode we also have to consider the change
in
mass
caused by the evolving gas.
Janssen and Barendrecht [10] found that the mass transfer coefficient for Fe(CN)6 3- to a hydrogen evolving nickel electrode did not change much with varying current density up to 1 kA/m 2• transfer
EXPERIMENTAL The nickel
preparation electrodes
technique is based on the oxidation of several (up to (initial diameter 6*10- 3 m and lenght 15*10- 2
evolution of hydrogen at a cylindrical platinum grid electrode 22*10-3 m 1enght 50*10-3 m). Several experiments are performed.
m) and (
5) the
diameter
62
First
the precipitation by means of the nafion membrane is
discussed.
In
figure 4 a schematic representation is given. r-----'
:
:potentiostate
!.
1
07 07
8
07 07 07
i' MEt-flRANE
Figure 4.
The precipitation vessel for the catalysts with the nafion membrane.
The
potential
potentiostate multimeter
(A).
over and The
the electrodes was regulated by a 30 the current was measured by
silica (Aerosil 380V,
Volt
Keithley
3
160
Ampere digital
Degussa) suspension in the
by a Heidolph R2R1 motor.
anode the
compartment was brought to 3.0 with hydrochloric acid and the pH in
the
was
stirred vigorously
a
The pH in
compartment anode
(V)
cathode compartment was brought to about 13 with a potassiumhydroxide solution. It the
was established that the diffusion of K+ and OH- through the exchange
experiment.
of
OH-
with
Cl- could be neglected during
the
membrane and time
of
the
To the solution in the anode compartment 50 ml of a nickelchloride
solution (97.3 gram nickel/liter) was added. This was necessary to decrease the ohmic resistance of the electrolyte. much lower initial current. useful
if
it
temperature settling
was
is
inert kept
at
An alternative was the precipitation at a
The addition of a supporting electrolyte and does 343 K.
not
migrate
To prevent the
through nickel
the
is
membrane.
hydrosilicate
onto the membrane surface some glass beads (diameter 2 to 5 *10- 3
only The from m)
63
were added. the
The voltage was limited to 2.5 Volt to prevent nickel reduction in
membrane.
It was observed that at a voltage of 5 the potential drop
over
the membrane exceeded 1.23 Volt which led to the establishment of two cells series. ions
in
In the 'anode compartment' nickel was oxidized at the anode and nickel
were reduced at and inside the membrane and in the 'cathode
compartment'
evolution of hydrogen at the cathode and oxygen at the membrane took place. The current was about 0.15 A. The pH curve was not measured continuously because of possible
damage
increased nickel
of
the pH electrode (glass beads) but we found that
from 3.0 to 5.9 and subsequently remained constant.
precipitated
onto
The
the
pH
amount
of
the silica was about 20 percent (based on
a
fully
reduced catalyst). The with
a
purpose
second
experiment was the precipitation of the
chromiumoxide
layer previously deposited on
nickel the
hydrosilicate
cathode
For
this
the platinum electrode was placed in a potassiumdichromate solution (1
gram/liter) and kept at a potential of -1.0 Volt compared to a nickel electrode (diameter 6*10- 3 m and lenght 15*10- 2 m). After 16 hours at a current of 1*10- 2 A a thin layer was deposited,
stable enough to remain intact at
densities used.
v
Figure 5. The precipitation vessel for the catalysts when the chromium oxide was used.
the
current
64 experiment was done in the 2 liter vessel as shown in figure 5.
The
suspension gram
of
3 to 6 gram silica (Aerosil 380V,
potassium
electrolyte.
chloride
was added to decrease the ohmic
One or two nickel electrodes were
two cylindrical platinum electrodes.
To
Degussa) in water about resistance
a
two
of
the
placed in the centre of one or
The platinum electrodes had
already been
covered with a chromium oxide layer. The distance between the anode and cathode was minimized and thus the ohmic resistance was substantially reduced. was
brought
temperature
to
a
was
value of 6.0 with
kept
at
a
potassium
hydroxide
solution.
363 K and with potentials of 7.2 and
currents of 0.08 and 0.70 A,
respectively,
The
12
pH The
Volt
and
a number of experiments were done.
pH and current were continuously registrated. We
also
prepared
some nickel hydrosilicate using urea
to
generate
the
hydroxyl ions. This method has been described earlier [11]. The temperature was kept
at 363 K and the nickel loading was about 20 percent.
vessel this
We used
the
same
We
used
as with the preparation by means of the chromium oxide layer. catalyst
precursor
catalyst precursors. lower
temperature,
to compare it with
the
electrochemically
prepared
To find out whether nickel hydrosilicate can be formed at we also prepared a 20 weight percent catalyst by means
of
the urea decomposition method at 345 K.
RESULTS AND DISCUSSION -Establishment of the precipitation region. As mentioned homogeneously in
the species to be precipitated
must
be
distributed
Therefore the region
the electrochemical cell must be established where the precipitation
place. in
above
over the surface of the suspended support.
Since convection provides a rapid transport of material,
the
turbulency
characterized
was
determined.
It turned out
that
the
takes
the variation turbulency
by the Reynolds number was rather uniform in the cell
with
as the
chromium oxide covered electrodes. We calculated the Reynolds number to be 6*10 4 with a corresponding thickness of the diffusion layer of about 5*10- 7 m at the anode and 2*10-8 at the cathode. The cell with the nafion membrane, on the other hand, diffusion
had a rather unfavourable turbulency. Near the membrane, was about 6*10 3 and the corresponding thickness of layer was calculated to be about 5.0*10-6 m. At 10- 2 m from
membrane,
the
Reynolds
number
Reynolds
number was substantially lower and too
small
to
the the the be
measured. We
also
wanted
to establish more accurately where the
reaction
nickel ions with the hydroxyl ions and the carrier takes place.
An
of
the
experiment
was done in which a nickel electrode was mounted besides a cylindrical platinum grid electrode (diameter 22*10-3m, lenght 50*10- 3 m) covered by a chromium
65
oxide
layer.
platinum closed
filter (diameter 30*10- 3 m) was
A soxhlet
placed
around
the
cylinder and the space within the platinum cylinder was filled with a glass
tube.
1
gram of silica was suspended into the
volume
of
the
solution inside the soxhlet filter and another gram of silica into the solution outside the filter.
Subsequently the experiment was started at a pH of 5.6 and
a current of 1 A. After 3 hours the green color of the silica in the suspension outside
the
reacted
with the hydroxyl ions and the silica,
soxh1et filter indicated the deposition of nickel ions
soxhlet filter had remained completely white. even
if
the
diffusion
that
had
whereas the silica inside
the
Therefore we must conclude that 4*10- 3 m, there is no
layer is expanded to about
precipitation in the cathodic diffusion layer.
0)
pH 17
3L.5K
U20 -
d1~17E-
363K
5
5
3 0
b)
" PH17 E05-N
12
8
1
16
000
5
00
"~7:-
5'
1
5
00 00 L.
Figure 6.
~)
"
-
L.
5
Fl~.
-
. Hl:~~---f
l(~)
3
3
'pHl'g-
00
-0
2
EII
2
3
I
401
123L.5 _
l(tnrs)
t(hours)
The pH-t and the I-t curves of the catalysts. * a second pair of electrodes, increasing the surface area was connected.
Experiments
with
the
nafion membrane cell and no dissolved
nickel
present at the beginning of the experiment indicated an accumulation of hydrosilicate non-ideal inside chromium
at the membrane if we did not use glass beads.
stirring,
the
In spite of
the nickel ions migrate to the membrane surface to
diffusion
oxide
ions nickel
layer.
coated cathodes
anodic diffusion layer.
This leads us to the suggestion that the nickel ions can also migrate out
Thus with the chromiumoxide experiments
the react
with
the
of
the
precipitation
must take place in the bulk of the suspension. Another
argument
for precipitation in the bulk of the suspension
can
be
66
found
in the pH overshoot.
supersaturation,
If the pH-versus-time curve displays an overshoot,
causing a large number of small particles to
been built up in the bulk of the suspension.
have
occurred
mainly
pH-electrode
nucleate,
must
If the reaction would
have
in the diffusion layer around the anode or the
cathode,
the
mounted in the bulk of the suspension would not have indicated
a
marked supersaturation. In
figure 6 the I-t and pH-t curves are given.
shows
no
ascribed
overshoot to
the
In figure 6b the pH
in contrast to figures 6a and
6c-f.
nucleation of hydrosilicate [11].
At
The low
curve
overshoot
was
temperatures
no
hydrosilicate is formed and no overshoot is exhibited. At 343 K no measurements of
homogeneous deposition-precipitation were available.
with nafion
experiments
nuclei, the
Separate measurements
urea indicated the presence of a nucleation barrier at 343
nafion
K.
With
the
followed by the formation of many small
did not occur near the pH electrode, but probably was established near
nafion
glass
supersaturation,
beads
membrane. had
membrane.
membrane
reacted
To prevent a non-uniform deposition onto
to be utilized to redisperse the carrier The
silica
in or near the difusion
with the nickel and hydroxyl ions;
settled
layer no
the at
carrier, onto
the
the
nafion
supersaturation
and
subsequent primary nucleation proceeded in the bulk of the suspension where the pH was measured. -The characterization of the catalysts. From temperature-programmed
reduction
i
hydrogen consumption la.u.J
profiles we can also
deduce
hydrogen consumption
(c.ul
U20
E20-N
E20-N2
EOS-N 500
Figure 7.
600
700
800 900 1000 temperature ( K ) -
500
600
700
800 900 1000 temperature ( K ) -
Temperature programmed reduction profiles of various catalysts.
that
67 nickel
hydrosilicate
decomposition
was
the
main
product
formed
method and the chromiumoxide cell,
with
both
the
urea
but not in experiments
with
the nafion cell. The broad reduction peaks in figure 7 are characteristic for the
reduction
of the nickel hydrosilicate. Some of the peaks end abruptly. This is due to the amounts of catalyst precursor used.
Small amounts give peak maxima to shift to
lower values and peak endings to be more abrupt. precursors prepared with the nafion membrane cell, show a sharp formation
E20-N, E20-N2 and E05-N, do
peak, due to the reduction of the nickel hydrosilicate.
Electronmicrographs metal
The reduction profiles of the
of the catalyst precursors clearly show hydrosilicate
in both types of experiments.
particles
The mean particle size of the
after a standard pretreatment is larger with the
nickel
E20-N
than
with the E20-Cr catalysts (see figure 8).
Figure 8.
Electronmicrographs of the catalysts. a) a E-Cr catalyst before reduction. b) a E20-Cr catalyst after reduction. c) a E20-N catalyst before reduction. d) a E20-N catalyst after reduction.
To estimate
the
particle size of the nickel particles after
a
standard
calcination in nitrogen during 7.2*10 4 sec (20 hours) at 723 K and a subsequent reduction in 10 percent hydrogen in nitrogen during 2.6*10 5 sec (72 hours) at 723
K,
vibrating pressed
we measured the magnetization of the nickel particles by means sample magnetometer [12].
of
a
For this purpose the catalyst powder
was
into tablets at 173*10 6 Pascal and fragmentated into particles of
0.6
68 to 0.85 mm.
From the data we calculated the particle size and the surface area
of the nickel particles [13,14J. The results are given in table 1. table
The metal loading,
1.
With the other catalysts the name denotes the weight percentage
catalysts.
* Based
nickel.
particle size and specific surface area of some
wt% Ni
catalyst
*
particle size (r 3)1/3 surface area(magn) (magnetic) m2/gr am metal (A)
E20-Cr
19.0
25
141
U20
19.8
23
179
The
of
on a fully reduced catalyst.
activities of the catalysts in the methanation reaction are
given
in
figure 9. 0
(kJ/mole) InkJ:lu)
link -1
98 129 95
-2
-3
.... -,
22·3 28·2 21·6
,
-4
E20-Cr·....
-5
U20
..........
-6 -7 -8
Figure 9.
E20-N
'·8
1·9
2·2
2·1
2·0
2·3
l0001T (K- 1) -
Activities of the 20 wt% catalysts.
CONCLUSIONS The electrochemically prepared nickel hydrosilicates from the chromium cell show a
good
resemblance
decomposition
method.
small
particles
nickel
with
the
hydrosilicates
prepared
by
the
urea
The magnetic measurements with the E20-Cr catalyst show as does
the
U20.
From the
temperature-programmed
reduction profiles we can see that in the nafion cell in most cases also nickel
69 hydroxide was formed due to non ideal stirring. If
we compare the activities of the catalysts in the methanation
we notice
that
the
E20-N catalyst is less active
Formation of nickel hydroxide, presumably
is
the
than
the
U20 catalyst.
which is reduced to larger metallic
cause of the relatively low activity.
For
prepared
in the chromium oxide cell hydroxide formation was
activity
equals the activity of the U20 catalyst.
reaction particles,
the
catalysts and
the
If the construction of
absent,
the
cell can be changed so as to provide a degree of turbulency as high as with the chromium oxide cell, the nafion method should give the same results. We
finally
conclude
deposition-precipitation precipitated
nickel
programmed reduction,
that indicate
pH-versus-time rather
over the silica.
curves
measured
accurately the
The consistent data
during
dispersion from
of
the the
temperature-
electron micrographs, magnetic measurements and activity
for the methanation reaction agree with the pH-versus-time curves.
LITERATURE 1. W.M. Latimer, oxidation potentials, Prentice Hall inc., Englewood Cliffs, N.J., sec. ed., (1964), 1 2. S.H. Maron and J.B. Lando, fundamentals of physical chemistry, Macmillan publ.co. New York, (1974), 554 3. Product bulletin, Du Pont Company, Wilmington D.E. U.S.A. 4. M. Seko, Ind. Eng. Chern. Prod. Res. Dev., 15 no 4, (1976),286 5. M. Pourbaix, Atlas d' equilibres electrochimique, Gauthier-Villars, Paris, 1963, 333 6. E. Gileadi, E. Kirowa Eisner and J. Penciner, interfacial electrochemistry, Addison-Wesley inc. London, (1975) 7. J. Albery, electrode kinetics, Clarendon press, Oxford, (1975) 8. W.A. Schultze, thesis, university of technology, Delft, (1970) 9. J.W. Geus, Physical transport fenomena, university of Utrecht, the Netherlands, (1983) 10. L.L.J. Janssen, E. Barendrecht, J. Applied Electrochemistry, 15 (1984) 549555. 11. J.W. Geus, Preparation of catalysts III, Elsevier science publishers Amsterdam, (1983),1 12. S. Foner, Rev. Sci. Instr., 30 no 7, (1959), 548 13. H. Dreyer, Z. anorg. allg. Chern., Band 362, (1968), 233 14. E.G.M. Kuijpers, thesis, University of Utrecht, (1982)
70
DISCUSSION L.H. STAAL: 1) Could you comment on the difference between catalysts prepared by the new procedure and catalysts prepared by "conventional" homogeneous precipitation like urea decomposition? 2) What is the influence of the difference in ionic strength? P.C.M. van STIPHOUT : An important part of this work was a comparison of nickel catalysts produced via the hydrolysis of urea and via electrochemical dissolution and subsequent precipitation of nickel. As dealt with in the paper no difference could be found between nickel-an-silica catalysts prepared in either way, provided a cathode covered by chromium oxide was used. As to your second question it is true that the ionic strength with the urea decomposition method on the one hand, and the electrochemical method with the chromium oxide covered cathode on the other other hand, might differ. However, the thickness of the electrostatic double layer will hardly be altered and the laminar double layer does not change. There is thus no difference in precipitation to be expected due to the different ionic strength of the suspension. E.K. POELS : 1) What is the maximum loading that can be obtained using this preparation method? 2) How long does a precipitation take to, say, a 20% Ni loading? P.C.M. van STIPHOUT : The metal loading that can be obtained varies from zero to at least the value corresponding with the complete reaction of the silica to nicke1hydrosi1icate (wt Ni/wt Si = 1.5). The metal loading we used varied from 5 to about 40 wt %nickel (in the reduced catalyst). However, higher loadings than the above Ni/Si weight ratio can also be applied with presumably good results. An advantage of the electrochemical procedure is that no blue soluble nickel ammine complexes are formed that render the production of highly loaded catalysts difficult. As can be seen in figure 6, we used a current of about 300 mA corresponding with a precipitation time of about 4.5 hour for a 20 wt %catalyst, if we have about 6 grams silica present. The precipitation has also been carried out with a current of 3 A, ten times faster. The parameter that determines the minimum precipitation time is the maximum current density. If we increase the surface area we can precipitate at a higher rate. XU XIAODING : Would you obtain a homogeneous deposit, if you work with two different ions? P.~.M. van STIPHOUT : To mention an instance we have sequentially precipitated Cu + and Ni 2+ and found with several characterization techniques that the copper and nickel were homogeneously distributed. Upon reduction we obtained small crystallites. The results were similar to the results obtained with the urea decomposition method. However, we did not yet measure the extent of alloying, the uniformity of the composition of the particles obtained by reduction. Before presenting definitive results we can say that, in the electron microscope, the reduced catalysts were exhibiting a uniform small particle size, which points to the formation of alloy particles of uniform composition. The composition of the alloy particles is presently established more accurately by magnetic measurements and by the catalytic activity in the CO hydrogenation.
B. Delmon. P. Grange. P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V.• Amsterdam - Printed in The Netherlands
71
PREPARATION OF HIGHLY DISPERSED, CARBON SUPPORTED, PLATINUM CATALYSTS
D. RICHARD and P. GALLEZOT Institut de Recherches sur la Catalyse, Laboratoire Propre du CNRS, conventionne a l'Universite Claude Bernard, Lyon I, 2 Avenue Albert Einstein, 69626 Villeurbanne cedex, France
SUMMARY Precise techniques of preparation are described to obtain highly dispersed platinum catalysts supported on high specific area graphite and on non porous carbon black (Vulcan 3). The essential steps are (I) the functionalization of the graphitic or pre-graphitic planes with strong oxidizing agents such as HN03, H202 and NaOC1, (II) the ion-exchange of platinum complex cations with these edge sites. The influence of the nature of the oxidizing agent on the number of functional groups and on the amount of exchanged platinum has been studied, as well as the influence of the counter-ion, concentration and exchange medium. INTRODUCTI ON The use of carbon as catalysts support has gained growing importance in the last two decades. However carbon have been less studied than other supports. and the method of preparation of carbon supported catalysts are described in a few number of works. In a recent review (ref.I). Ehrburger discussed the factors favouring the dispersion of metals on carbons. The importance of surface heterogeneities for the achievement of a high dispersion of metal was mentionned. Stress was 1aid on the role of the interaction of the metal precursor with the carbon atoms of the edge planes of graphites or grapitized carbons. This interaction could be connected with the presence of functional groups at these edges. It has been established that the burning of the support leads to an increase of these edges planes, the active specific area. relatively to the basal pl anes (refs. 2.3). Different oxidative treatments in liquid phase lead to a larger number of functionals groups (ref.4). However there has been very few attempts to take advantage of the increase of the active area and of the functionals groups on this area to prepare metal supported catalysts (refs.5-7). Among them. the exchange of palladium (ref.5) and silver (ref.6) aminocomplexes with the protons of carboxylic groups on an oxidized carbon support has to be mentionned.
72
These preparation techniques involving the functionalization of the carbon support and the ion-exchange method were used in the present work, in order to achieve a better dispersion of platinum particles on carbon supports. EXPERIMENTAL METHODS Starting materials One of the carbon materials used as support was a graphite (HSAG 12) from Lonza, with a high specific area (470 m2/g). The other support was a furnace black (Vulcan 3) from Cabot. This non porous carbon has a specific area of 74 m2/g. Functionalization of the supports These supports were first activated by partial combustion in flowing oxygen at 500·C. They were kept at this temperature for 3 hr. Under these conditions 60% of the graphite and 70% of the furnace black were burned off. Then, they were treated in liquid phase by different oxidizing agents: (I) Sodium hypochlorite; the suspension of carbon support in concentrated NaOCl (about 15% in Cl active) was stirred at room temperature for 24 hr. (II) Nitric acid; the concentrated HN03 (65%) oxidizing solution, was refluxed for 24 hr. (III) Hydrogen peroxide; the oxidation was performed at 40·C during 24 hr in concentrated H202 (aqueous solution 30%). Ion-exchange procedure The H+ ions of the functionalized supports were exchanged with IPt(Ntl3)4f+ions. Tetrammine platinous chloride monohydrate was supplied by Johnson Matthey. The chloride is converted to the hydroxide by passing a soluti.on of IPt(NH3)4IC12 through the hydroxyl form of an anion exchange resin (Amberlite IRN-78). The ion-exchanged catalysts were prepared by dropwise addition of tetrammine platinous hydroxide to a suspension of carbon support in N ammonia solution (ref.8). After the addition, the stirring of the suspension was maintained for 15 hr. The carbons were then filtered on Millipore MF type filters: RA with a pore size of 1.2 ~m for the graphite and VS with a pore size of 0.025 ~m for the furnace black. The catalysts were then washed thoroughly by water and dried at 100·C under flowing nitrogen. The dried catalysts were heated in flowing hydrogen for 2 hr at different temperatures from 200 to 500·C.
73
RF.SULTS AND DISCUSSION 1) Characterization of the functionalized supports The partial combustion treatment of graphite produces a decrease of the RET specific area from 470 m2/q to 310 m2/g. However this loss is probably due to the sticking of previously exfoliated graphite layers which reduces mainly the accessible area of the basal planes rather than that of the edge planes. Work is in progress to measure the relative fraction of edge planes at different steps of the treatment. The number of functional groups on the supports was estimated by NaOH titration, using the method described by Garten and al. (ref. 9). TABLE 1 Number of functional groups titrated by NaOH Number of functional groups (mmol/g) Support
Initial
After burning
HSAG 12 Vulcan 3
0,49 0,16
a 0,61
After oxydation by NaOCl HN03 H2 02 0,94 b
0,78 4,96
0,50 0,77
a not measured b the support was almost totally oxidized b.y HN03 no measurement was undertaken Considering the results given in Table 1, it is noteworthy that the nitric acid treatment leads to a 1arge number of functional groups on graphite, and even, almost totally oxidize the furnace black. The hypochlorite acts as milder oxidant and the hydrogen peroxide treatment leads to a smaller number of functional groups. It has to be pointed out that the increase in the number of functional groups following these oxidative treatments is more important on the Vulcan 3 than on the HSAG 12. That can be explained by the fact that in graphite, the functional groups are located at the edge of sheets, and so their number cannot be increased greatly if the area of these edges is not increased too. On the contrary, the furnace black which present many surface heterogeneities can accomodate a greater number of functional groups after oxidation. The creation or increase of the functional groups has also been checked by IR spectroscopy. In spite of the difficulty in recording IR spectrum using KBr pellets, because of the low transmission of carbon samples, a weak band at 1730 cm-1 has been observed on the Vulcan 3 support. It can be
74
attributed to carboxylic groups created by the oxidizing treatment. This has not been observed so obviously in the case of the graphite support because of the smaller number of these groups. 2. Study of the cationic exchange Platinum has been introduced on the supports by ion-exchange at different stage of their preparation. Table 2 shows that the amount of platinum fixed by the support is larger after oxidation, consistently this increase is more drastic in the case of Vulcan 3. TABLE 2 Percentage of Platinum fixed by differents supports after exchange with the same amount of Platinum in solution Support
Initial
After burning
HSAG 12 Vulcan 3
2,1 0,03
0,75
After oxydation by HN03 NaOCl H202 3,37
3,32
1,88
6,2
It should be noticed that after the partial burning, the graphite can fix less platinum than before. This fact can be explained by the temperature of the combustion (500·C), at which most of the surface functional groups should have left the graphite. Barton and Harrison (ref.10) mentionned a peak at 400·C in the thermodesorption of surface functional groups on graphite. Thus the burning step produces a decrease in the number of functional groups and therefore in the amount of platinum exchanged. It can be noticed that the increase in the amount of platinum exchanged after NaOCl oxidation is exactly proportional to the increase in the number of functional groups titrated by NaOH. The functional groups created by H202 and HN03 treatments are less effective to exchange platinum. Different parameters that may change the efficiency of the exchange have been investigated. (I) Amount of platinum added to the carbon supports. Figure 1 gives the platinum uptake as a function of the total amount of platinum introduced in the solution. It appears that the quantity of platinum fixed on the support reach a plateau when the amount of platinum added is increased. The limit corresponds probably to the exchange of the p1 atinum ions with all the sites available on the support.
75
Pt (wt%) (carbon supported)
3
2
Pt (wl%) (carbon in suspension)
10
Fi g. 1 solution
20
30
40
Pl atinum uptake on graphite vs total amount of pl atinum in
(II) Effect of the counter-ion of the platinum salt; by changing it from OW to Cl-, the percent of platinum fixed decreases from 3.3% to 2.7% for the graphite oxidated by NaOC1. The importance of the counter-ion on the exchange has already been noticed by Lowde and al. (ref. 10). (III) Exchange medium. The need for a basic solution during the cationic exchange, on silica and alumina has been mentionned by Benesi and al. (ref. 8). It was shown that the more basic the solution, the greater the capac ity of exchange of these supports. However the use of ammon i a solution could have another beneficial effect because this medium allows a competitive exchange between the NH4+ and IPt(NH3)412+ ions favouring the dispersion of platinum. This competitive effect has been shown by exchanging in water where the platinum uptake is 4.0% instead of 3.3% for the exchange in ammonia solution. 3. Characterization of the catalysts The dispersion of the metal was characterized by high resolution transmission electron microscopy. Figure 2 is a TEM view taken with a JEOL 100-C showing the platinum particles supported on HSAG 12 after reduction at 300·C. It should be noticed that the particle, in a size range 1-1.5 nm are mainly located along the edges of the graphite sheets. In addition they decorate the steps of graphite adlayers. This means that due to the interaction between the precursor and the functional groups at the edge plane of the graphite the
76
Fig. 2
TEM view of platinum particles on graphite support (HSAG 12) %
~-
20
~
t'l
Reduction at 300·C
-
---
10
o
J
~
--,......,. .-.-b
Reduction at 500·C
-
.--
'--
h-l
I
o Fig. 3
2
3
nm 4
Particle size distribution on graphite (HSAG 12) a) reduction at 300·C b) reduction at 500·C
77
pl atinum particles nucleates there, and since the mobility of metal particle is reduced there is little sintering. The increase of the reduction temperature to 500·C leads to slightly larger particle sizes, as shown on Figure 3 giving the histograms of particle sizes for reduction at 300·C and SOO·C. Figure 4 is a TEM view of platinum supported on Vulcan 3 after reduction at 300·C. The size distribution is again vey narrow and centered about 1.5 nm, but the knotty texture of the support prevents a precise localization of the particles.
Fig. 4 : TEM view of platinum particles on carbon black support (Vulcan 3) The catalysts were also studied with a field emission gun STEM (VG HBS01) which allows the recording of diffraction pattern of individual particles as small as 1 nm (ref. 12). In the case of graphite support, the nanodiffraction pattern of these particles shows that they are mutually oriented and in close epitaxy with the underlying plane. Figure Sa gives a typical pattern taken with a stationary beam on aI-loS nm particle and Figure 5b gives the pattern of graphite taken 2 nm apart from the particle. The diffraction spots and the background scattering are more intense on platinum graphite than on graphite
78
Figure 5 : Diffraction pattern taken with a stationnary beam in STEM a) on a 1.5 nm particle b) on the graphite support apart from the particle alone and even extra-spots due to platinum appear in this orientation. In other orientation the platinum particles are distorted and because of the close epitaxy no more extra-spot is observed. In the case of the carbon black support, the diffraction patterns show that the particles are not in epitaxy with the support (the support itself showing only diffraction rings but no spot) and not mutually oriented. CONCLUSION This work shows that by following a precise technique of preparation, it is possible to obtain well defined platinum catalyst supported on graphite or carbon. The preparation involves necessarily the following steps (I) a functionalization of the edges of graphitic or pre-graphitic basal planes. This creates exchangeable sites for the cations and anchoring sites for the catalyst precursor as well as for the final metal particle (II) the ion-exchange which introduces the metal precursor only on the anchoring sites. Under theses conditions, the final catalyst exhibit platinum particles homogeneously dispersed mostly in the size range 0.8-1.8 nm. In the case of Pt/graphite, these particles are located along the edges of the basal plane, decorating the steps of adlayers in epitaxial relationship with the underlying layer. Preliminary reaction data indicate that these catalysts have good selectivity in
79
hydrogenation reactions. ACKNOWLEDGEMENT ~e wish to thank Dr. P. Fouilloux for many valuable discussions and for suggesting to us the use of hypochlorite as oxidizing agent for carbons. REFERENCES 1. P. Ehrburger, Advances in Colloid and Interface Sciences, 21, (1984), 275-302 2. N.R. Laine, F.J. Vastola, P.L. Walker, J. Phys. Chem., 67, (1963) 2030-2034 3. P. Ehrburger, O.P. Mahajan, P.L. Walker, J. Catal., 43, (1976),61-67 4. J.B. Donnet, F. Hueber, C. Reitzer, J. Oddoux, G. Riess, Bull. Soc. Chim. Fr., (1962), 1728-1735 5. K. Morikawa, T. Shirasaki, M. Okada, Adv. Catal., 20, (1969), 97-133 6. P. Ehrburger, J. Dentzer, J. Lahaye, Proceedings 15 th Biennal Conference on Carbon, The American Carbon Society, Philadelphia, (1981), 254 7. A. Van Montfort, J.J.F. Scholten, Patent 2719006, (1976) 8. H.A. Benesi, R.M. Curtis, H.P. Studer, J. Catal., 10, (1968), 328-335 9. V.A. Garten, D.E. Weiss, J.B. Willis, Aust. J. Chem., 10, (1957), 295-308 10. S.S. Barton, B.H. Harrison, Carbon, 13, (1975), 283-288 11. D.R. Lowde, J.D. Williams, P.A. Attwood, R.J. Bird, B.D. McNicol, R.T. Short, J. Chern. Soc. Faraday Trans. I, 75, (1979), 2312-2324 12. P. Gallezot, C. Leclercq, I. Mutin, C. Nicot, D. Richard, J. Microsc. Spectrosc. Electron., 10, (1985), 479-484
80
DISCUSSION G.M. PAJONK : In order to determine the textural properties of a supported metal (for instance) it seems to me always important to double a physical mean (like TEM here) by a chemical one. 1) Did you chemisorb a probe like H2 to calculate the mean particle size? If so, did the results agree with the TEM values? 2) Can you give some information about the catalytic tests? D. RICHARD: 1) No chemisorption experiment has been yet undertaken on this catalyst, but such a verification of the dispersion is of course planned. 2) These catalysts have been tested in hydrogenation reactions: - competitive hydrogenation of benzene and toluene, which gives information about the electronic state of the metal, showing in this case a donor effect from the carbon support to the platinum, - hydrogenation of butadiene. A good selectivity in butenes is observed, - liquid phase hydrogenation of cinnamic aldehyde. Preliminary results indicate a high selectivity to cinnamic alcohol. G. JANNES : What is the decisive factor to obtain such highly dispersed metal crystallites: some special quality of the nucleation state, or some peculiar interaction between metal and support to prevent sintering? In the latter case, do you mean that the energy of this interaction well compares with the thermic activation at 500°C? D. RICHARD: We suppose that the interaction between metal and support is the key factor to achieve a high dispersion, because catalysts prepared by impregnation on similarly treated supports show bigger particles with a broader size distribution. J.A. SCHWARZ: 1) What was the ash content of the carbons? 2) Was there a change in surface area or porosity of the carbon after oxidative treatment? 3) How large ordered regions at oriented "basal plane" type structures are required in the STEM to give rise to the diffraction you observe? D. RICHARD: 1) The ash content of the carbon support was - for the HSAG 12, 0.13%, mainly ferric oxide - for the Vulcan 3, 0.03% before and 0.5% after hot HCl washing, the increase consisting of silica due to the glassware. 2) After the oxidative treatment, both supports, graphite and carbon black undergo a loss of surface area, but no porosimetry study was undertaken. 3) The probe size in the STEM is sufficient, even in the spot mode (fixed beam), to observe the diffraction pattern from the graphite basal planes. K. NOACK: How well are the Pt particles fixed on the carbon support? Are they lifted off during hydrogenations in organic solvents? D. RICHARD: The platinum particles are very tightly fixed on the carbon support, as can be seen on the TEM micrographs, since the preparation of grids for EM involves an ultrasonic dispersion of the solid in ethanol. After the hydrogenation of cinnamic aldehyde in 2-propanol, the particles remain on the support. D.E. RESASCO : It seems that your preparation method yields Pt particles strongly held to the carbon support. Could you speculate on the nature of the anchoring sites of the support during the adsorption, and for the reduction process, to account for such a strong interaction? D. RICHARD: On the carbon supports many different functional groups have been
81
found, for instance carboxilic acids, lactones, phenols and quinones in order of decreasing acidic strenth. We suppose that the platinum tetramine cations are exchanged with the more acidic of them, i.e. the carboxilic groups. However, it is known that the less acidic the functional groups, the more resistant they are toward thermal decomposition into CO or C02. Thus we can speculate that these weaker acidic groups playa major role during and after the reduction for the stabilisation of the metal particles.
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B. Delmon, P. Grange. P.A. Jacobs and G. Poncelet(Editors), Preparation of Catalysts IV
83
© 1987 Elsevier Science PublishersB.V., Amsterdam - PrintedinThe Netherlands
Application of Precursors of Catalytically Active Materials Q!! Preshaped Supports Qx. Impregnation with Solutions of Badly Crystallizing Compounds
G.R. Meima, B.G. Dekker, A.l. van Dillen and 1.W. Geus l 1.E. Bongaarts, F.R. van Buren, K. Delcour and 1.M. Wigman 2 IDepartment of Inorganic Chemistry, University of Utrecht, Croesestraat 77a, 3522 AD Utrecht, the Netherlands. 2Dow Chemical (Nederland) B.V., P.O. box 48, 4530 AA Terneuzen, the Netherlands.
ABSTRACT A new method is described for the preparation of catalysts by means of impregnation. The technique is especially suitable for obtaining a homogeneous distribution of the active material within preshaped support particles. The possibilities of the new procedure are lllustrated by impregnation of preshaped a-alumina pellets with various organometallic Sncomplexes. After calcination in air the size and distribution of the resulting Sn02 particles and the surface area and porous structure of the catalysts were investigated by several complementary techniques such as BET surface area measurements, mercury porosimetry, X-ray diffraction, electronmicroscopy and temperature-programmed reduction. It proved possible to control the mean particle size and the porosity within narrow limits.
INTRODUCTION It is common practice to apply catalytically active agents onto thermostable, highly porous supports to obtain the elevated active surface area per unit volume required in many catalytic reactions. Since the pressure drop over a fixed bed catalytic reactor must remain limited in technical operations, the bodies of the catalyst cannot have dimensions smaller than about 0.3 mm, It is therefore necessary to produce porous bodies of the support loaded with the catalytically active agentts) of a dimension of at least 0.3 mrn, To prevent the formation of smaller particles during loading of the catalytic reactor or during operation, the mechanical strength of the catalyst bodies has to be considerable. Application of a precursor of the catalytically active component onto a powdered support is relatively easy. However, the loaded powder subsequently has to be processed into mechanically strong bodies, often with a controlled pore volume and pore size
84
distribution. Usually this is difficult to achieve. It would therefore be favorable if a procedure could be provided for loading preshaped bodies of the desired support uniformly with catalytically active particles. With low loadings of the support impregnation of preshaped bodies with a solution of a precusor of the active component and subsequent drying may lead to excellent results. However, higher loadings are difficult to establish with a dense and uniform distribution of the active component over the generally extensive internal surface of the support. The elementary processes taking place during impregnation with a relatively concentrated solution and subsequent drying are complex and still poorly understood [I], During impregnation previous wetting with the pure solvent or evacuation of the support can strongly affect the resulting distribution of the active material. Likewise the rate of evaporation of the solvent can strongly affect the final distribution of the active component or its precursor. Especially with a large number of impregnated bodies of the support, it is difficult to establish a rate of evaporation of the solvent uniform for all porous bodies of the support. It has been found that raising the viscosity of the impregnating solution of the precursor of the catalytically active compound has a beneficial effect on the homogeneity of the distribution of the active component after drying. Accordingly Kotter and Riekert (2) arrived at a more homogeneous distribution of an active agent over the internal surface of a support by carrying out the impregnation with a solution of the precursor of an elevated viscosity. The authors obtained the best results by adding hydroxyethyl cellulose to the impregnating solution. They observed that after impregnation with a solution of a higher viscosity and drying of the impregnated support did not lead to an
inhomogeneous
distribution. However, a severe difficulty is that a solution of an elevated viscosity can hardly penetrate into the often narrow pores of a support. Also with the longer pores of larger bodies of the support, impregnation with a viscous solution leads to an inhomogeneous distribution of the active component. This paper deals with a special procedure in which bodies of the support are impregnated with a solution of the precursor of a catalytically active component of which the viscosity considerably rises during evaporation of the solvent. Besides the rise in the viscosity of the solution it is important that the precursor of the active component does not readily crystallize during evaporation of the solvent. The procedure is utilized to apply tin oxide onto silica and alumina carriers. In this paper we will report on the application of tin oxide onto a-AI 203• An important reason to investigate tin oxide is that besides the catalytic activity in combination with antimony oxide in selective oxidation reactions, tin oxide can excellently serve
as an "anchoring"
agent for catalytically active particles. Hence silver particles are thermally stabilized when they are applied onto alumina covered by tin oxide. These catalysts are especially suitable for the production of ethylene oxide. The "anchoring" of silver particles was also earlier demonstrated with Pt (3); however, uncovered Pt has a disastrous effect on the selectivity for ethylene oxide,
85
Though the above procedure allows to cover the surface of an alumina support completely by a uniform layer of tin oxide, the thermal stability of the catalyst obtained by subsequent application of silver onto an alumina support covered by a continuous layer of tin oxide is not substantially enhanced. A much higher therrnostability results when the continuous tin oxide layer is first decomposed into small tin oxide particles homogeneously covering the surface of the alumina carrier (4). Therefore a detailed study of the application of tin oxide and the sintering behaviour of these particles is necessary. It was established that an incipient wetness impregnation with organometallic tin complexes of pre-evacuated alumina-rings leads to the best results. The complexing agents used in this study are formic acid, citric acid and EDTA (ethylene diamino-tetraacetic acid). After drying the impregnated supports have been calcined for varying periods of time at different temperatures up to 1523 K. The influence of the loading has also been investigated. Assessment of the distribution of the tin oxide over the support and the further characterization has been carried out by (analytical) scanning and transmission electronmicroscopy, X-ray diffraction, mercury porosimetry, measurement of the BET surface area, both before and after loading of the support and by temperature-programmed reduction. EXPERIMENTAL Preparation of the catalysts All catalysts were prepared by means of the "incipient wetness" impregnation method. In this technique the impregnating solution is sucked into the previously evacuated porous support. The amount of solution is carefully controlled so as to just wet the support particles. In this study a-alumina pellets (Norton HPC 5525) were used as the support material. These pellets consist of cylindrical rings with a diameter of about 5 mm, a length of about 7 mm and a BET surface area of 0.27 ± 0.03 m 2/g. Mercury porosimetry showed that this support material mainly contains macropores. From the penetrated amount of mercury the pore volume was calculated to be 0.40 ml/g, To study also the influence of the amount of impregnating solution usually two varying amounts were employed: i) 0.30 - 0.35 ml/g support designated "dry". li] 0.40 - 0.45 ml/g support designated "wet". These typifications do not imply very large differences, they merely express that in the "dry" case no liquid film was formed over the external surface of the support pellets, After impregnation according to the "wet" method a thin liquid film could be noticed over the pellets. In both cases the total added amount of tin-complex remained the same. Three different complexes were studied, two carboxylates, formate and citrate and one chelate, EDTA (ethylene diaminotetra-acetic acid). The complexes were prepared in the following manner: tin(II)formate. Sn(II)oxalate was suspended into a 60 wt% formic acid solution the acidity of
86
which had previously been adjusted to a pH value of 3.9 with concentrated ammonia. (At this pH value the solution is completely clear; hydrolysis is observed at a pH level higher than about 7). tin(II)citrate and tin(II)EDTA. Sn(II)oxalate and the complexing agent to be used were, in an equimolar ratio, suspended into deionized water. The pH level was carefully raised to a value between 5.5 and 6.0 by slowly adding concentrated ammonia. After impregnation the catalysts were dried at 393 K for i6 to 17 hours, unless otherwise stated. Calcination at various temperatures was performed in air. The loading of the catalyst was calculated on the basis of:
* 100% wt% Sn(O) wt% Sn(O) + wt% AI 203 in order to avoid confusing debates on the form and state of the tin compound. During most experiments the loading of the catalysts prepared from the formate and EDTA-complex was 0.50 wt% and the loading of the catalyst originating from the citrate complex was 0.96 wt%. Catalysts with a different loading were also made using the latter complex. For comparison purposes especially for the TPR measurements, physical mixtures were also made. In this case Sn02 or SnO (Baker analysed reagent) and the ground a-alumina pellets were mechanically mixed. Apparatus and procedures The BET surface
~
of the catalysts both before and after calcination at various
temperatures was measured with a Quantasorb apparatus (Quantachrome corporation). With this apparatus the adsorption of nitrogen is dynamically measured at 77 K by means of a thermal conductivity detector. A sieve fraction between 0.5 and 1.0 mm and about 1 g of catalyst was used for the measurements. Mercury porosimetry to study the pore size and pore distribution was performed by means of a Carlo Erba Macropore Unit (pressure between 0 and 1 bar) and a Carlo Erba Porosimeter 2000 (pressure between 1 and 1990 bar). Catalyst samples of 0.7 to 1.0 g of the same sieve fraction as mentioned above were placed in the dilatometer, For the calculations a (more or less) arbitrary contact angle of 141.3° and a surface energy of 480 *10- 3 J/m2 was used. It was also assumed that the pores were non-Intersective and of
cylindrical shape. X-ray diffraction measurements to identify the phases present and to determine the average Sn02 particle size were performed with a Philips diffractometer placed on a Philips PW 1140 X-ray generator. Cu Ka radiation was used. Temperature-progammed reduction experiments were performed in a conventional atmospheric flow reactor !i.d.
8 mrn], Hydrogen consumption was
monitored by
measurement of the difference in thermal conductivity of the incoming and effluent gas. A detailed description is given elsewhere (5). Calibration of the amount of consumed hydrogen was performed by measuring several reduction profiles of exactly known amounts of CuO. It is well known that this compound gives rise to the formation of pure Cu during reduction.
87
The surface morphology and the main particle size of all samples was studied with a Philips scanning electronmicroscope equipped with a LaB6 cathode system and a Philips EM 420 transmission microscope. Electron diffraction measurements were also performed with the latter instrument. RESULTS AND DISCUSSION Impregnation All the impregnating solutions were visually clear at the pH levels mentioned. This is indicative of the formation of the above mentioned complexes as it is well known that in the absence of (strong) complexing agents tin(HHons are extensively hydrolyzed in aqueous solution. They then tend to form condensed basic ions such as SnOH+, Sn2(OH)22+ and predominantly Sn3(OH)l+ very rapidly, even at low pH levels (6,7). Comparative experiments in the absence of the complexing agents indeed showed the expected hydrolysis of the tlntlll-ions, However, it must be stated that when employing the citrate- and EDTA-complex the impregnation has to be carried out immediately after preparation of the complex. When these solutions are left to stand in air a precipitate is gradually formed. This is presumably due to a slow hydrolysis of the oxidized complex. A detailed study of the processes taking place during the preparation of the tin(II)citrate complex has been given by Kondruk et al (8). Simular studies have been reported (9,10) on the formation and hydrolysis of the Sn(II)EDTA complex. The latterstudies are confusing on the precise processes taking place during formation of the complex. However, it seems clear that the hydrolysis of the Sn(II)EDTA complex takes place at higher pH levels than 7 (I I, 12), presumably after oxidation of the tin(II)-species.
In the formic acid solutions the predominant species is the triformate stannate(II)-ion, Sn(HC0 2)3- (13). No hydrolysis was observed at the mentioned pH level, even after several weeks. Thermal stability; BET surface ~
and mercury porosimetry measurements
In figure 1 the measured BET surface area as a function of the calcination temperature (> 16 hrs) and impregnation procedure is shown for the various complexes. (With the citrate
complex only the "wet" method was employed). The results clearly show that after impregnation and drying at 393 K, the BET surface area does not differ from that of the support (represented by the dashed line in the figure). As will be argued later, decomposition of the complex and formation of Sn02 particles takes place after calcination at more elevated temperatures. This gives rise to an enhanced BET surface area. However, after an initial rise further calcination at higher temperatures leads to sintering of the Sn02 particles. A maximum is therefore observed in all the curves. With the EDTA- and formate- complex the maximum was attained at a lower temperature than with the citrate complex. This may be due to a thermodynamically more stable citrate complex and thus a
88
2
SnlIDEDTA
• dry o wet
01------+------1-----1 SnaIlcitrate
2
o wet
A (m7g1 1
01------+-----+------1 Sn(IIlformate
2
• dry o wet
~7L.3-------='':-::----=:-----:1:::773 773
1273
T (Kl
Fig. 1. 88ET as a function of calcination temperature impregnation with various 8n(II)complexes.
(t>16
hrs)
after
more elevated decomposition temperature in the latter case. Thermal analysis showed the citrate complex only slowly decomposed between 489 and 693 K, whereas it is known that the formate complex decomposes between 473 and 483 K (13,14). The "wet" impregnation method consequently gave rise to somewhat higher values in all cases. However, as can be seen, the sintering of the particles was also more rapid. The decomposition of the formate complex gave rise to a relatively high BET surface area. As also will be shown later this is caused by the formation of finely divided and very small 5n02 particles (more so than with the EDTA and citrate complex) thus enhancing the roughness of the initially very smooth ex-alumina surface. After calcination at temperatures equal or higher than 1273 K the 5n02 particles had slntered to such an extent, that no difference was observed between the different complexes. All catalysts then exhibited about the same surface area as the unloaded support. Figure 2 shows the influence of the calcination time on the BET surface area at a number of temperatures. Calcination at 773 K of the 5n(II)formate impregnate first leads to an initial rise In specific surface area followed by a decrease. After about 10 hours a constant specific surface area is obtained. Two calcination temperatures, 773 and 1273 K,
89 SnlIIlEDTA o "wet"
1.5
1.0
0.5
t::."wet" 1273 K
773 K
• "dry"
• "dry"
r
V
\t::._ _
."" ~
I
0.0
.
•
•
- -- ---
I
Sn(TI)formate
A (m~g)
1.S
• "dry" 773K
~
• •
•
•
1.0
0.5 0.0
o
I
5
10
15
I
I
20
25
t (hrs)-
Fig. 2. SBET as a function of calcination period at 773 K after with various Sn(II)complexes.
impregnation
were chosen to study the Sn(II)EDTA impregnate. Calcination at 773 K leads to a rapid rise in specific surface area followed by a more or less constant value after about 5 hours. No further decrease was measured as a function of time. Calcination at 1273 K gives rise to different results for the various preparation methods. With the "wet" impregnated sample no variation in surface area compared to the unsupported alumina was measured. With the "dry" impregnated sample the surface area initially rose and sharply dropped to the same value as the "wet" sample after about 3 hours. The above results are indicative of the following: I) calcination at constant temperatures mainly leads to changes In specific surface area during the first 10 hours. These changes are more rapid at higher temperatures. u) the "wet" Impregnated catalysts tend to sinter more rapid than the corresponding "dry" samples at elevated temperatures. In figure 3 a representative example Is given in which the specific surface area is plotted against the loading of the catalyst after a standard calcination procedure (923 K, 22 hrs), Clearly the specific surface area tends to rise with higher loadings. However, It must
be kept in mind that the exact shape of the curve depends on the calcination temperature. Mercury porosimetry measurements were performed with many samples. The relative pore size distribution was practically identical in all cases. Hardly any, or no variation
90
2.5 2.0
Sn(II)citrale "wet" 923K
1.5 A Im'Yg)
1.0
•
0,5 O.O!------L_..l.------.JL----L_--l.-------! o 2 4 6 loading wt% S n -
Fig. 3. SBET as a function of loading after impregnation with "wet" and calcination at 923 K for 22 hours.
Sn(II)citrate
0.50
r
0.25
cum. vol.
lml/gl
0.OOL---'--""TTTT1T""~---r-TT1"TTT1r-.,--J,J-r.I+Ir\M-J..Jl..1.J..f.u,uIJ+lrYrH..u::r:J:;i'l;.-r;rrnTI 1
10
100 1000 pore radius (nm) -
10000
100000
Fig. 4. A representative example of the pore size distribution of the catalysts as measured by mercury porosimetry.
91
between the various samples or even between the unloaded support was observed. A representative example of the pore size distribution is shown in figure 4. The pores with the most frequent occurrence have a radius between 1 and 10 micrometer. The variations in specific surface area as measured by this technique were only marginal. It is well known that mercury porosimetry can not measure pores with a smaller radius than about 4 nm at a pressure of about 2000 bar. Therefore it must be concluded that the rise in surface area as measured by the physisorptlon of nitrogen at 77 K is caused by the formation of "mesoand/or micropores" due to the formation of very small 5n02 particles. X-ray diffraction measurements Diffraction patterns of all catalysts have been recorded both before and after calcination in air. Depending on the specific complex and loading, only after calcination at temperatures more elevated than about 773 - 923 K peaks or bands other than those from the alumina support could be distinguished. Below these temperatures the 5n02 particles (if present) were amorphous, and only partly crystallized at more elevated temperatures. This was also confirmed by electron diffraction studies. At low calcination temperatures the diffraction patterns only consisted of diffuse halo's. At more elevated temperatures the
30 SnlIIlcitrate "wet.. 923 K
20
d [nm) 10 A12~
-
0 0
-A1 Z03
s-o,
6
Fig. 5. (left) Diffraction pattern obtained after calcination of the Sn(II)citrate wwetn sample at 1273 K.
(1101
5nOZ(101)
Fig. 6. (top) The mean particle size as a function of loading as estimated from ?
16
2 4 loading wt% 5n -
0
--eo
14
12
the X-ray line broadening. Catalyst prepared by impregnation with Sn(II)citrate wwetn followed by calcination at 923 K for 22 hours.
92
diffraction rings also contained spots, indicative of relatively large crystallites of tin oxide. All, except one of the observed peaks could be interpreted on the ASTM-JCPDS file values for Sn02 (IS) and Cl-AI 203• An additional peak was (often) observed at 8=14,08 (d=3.l7), especially after calcination of the samples at high temperatures. (A representative diffraction pattern is shown in figure 5). This peak could not be ascribed to other known tin oxide phases such as Sn304 or Sn203' The only tin compound listed in the ASTM file which could generate the observed peak is a rare modification of SnO, Le, "red" SnO (16). The pure compound is only stable in air up to about 543 K (I7). However, it was also mentioned that it is metastable at ordinary temperatures and that the true stability range of the material lies above the disproportionation temperature of "normal" SnO, viz. 658 K. It is also well known that alumina can stabilize SnO or retard its decomposition into metallic tin and Sn02 (I8 - 21). Choudhuryet al (22) and Muranaka et al (23) have also observed a slmular peak. The former ascribed it to an unknown tin oxide phase, whereas the latter ascribed it to the "red" SnO modification. However, it must be stated that the temperatures at which these studies were performed were much lower than in our case (maximum temperatures of about 723 K). In conclusion it can be said that two possibilities exist: i) the peak is caused by the "red" modification of 5nO, most presumably stabilized by the alumina support. Il) the peak is caused by an unknown tin oxide phase. We will further comment on these possibilities in the discussion of our TPR results. The mean Sn02 particle size was estimated from the broadening of the (1I0) and the (101) reflections using the well known Scherrer expression. Corrections for the instrumental broadening were performed according to the method as proposed by Whyte (24). Table gives the resulting diameters as a rounded- off average of the values obtained from both TABLE 1
Mean particle size as estimated from X-ray line broadening. Preparation proc.
Cel.c , Temp. (K)
d (tun)
Sn(II)citrate ·wet" 0.96 wt'% Sn
773 923 1023 1273 1523
9.5 8.5 15.5 76.5 128.0
"wet"
773 1273 773 1273 1523
11.0 67.5 14.0 47.0 129.5
1273 1523 1273 1523
55.0 128.0 76.0 128.0
Sn(II)BDTA
0.50 wt'% Sn "dry·
Sn(II)formate "dry" 0.50 w~.% Sn ·wet·
93
60
1 40 a (nml 20 Sn(IIlEDTA '~ry"
1273 K
oOl-----'-4--....la:----:''="2---:-:-t lhrs) -
Fig. 7. The mean particle size as a function of calcination time at 1273 K as estimated from the X-ray line broadening. Catalyst prepared by impregnation with Sn(II)EDTA ·dry·.
reflections after calcination of the samples for at least 16 hours at the listed temperatures. The results nicely correspond with the BET surface area measurements. Clearly the drop in specific surface area after calcination at elevated temperatures is caused by the sintering of the Sn02 particles. In figure 6 a representative example of the variation in the calculated mean particle size is plotted as a function of loading. All samples were prepared by means of the citrate complex and had been calcined in air for 22 hours at 923 K. (The corresponding BET measurements are shown in figure 3). As can be expected the mean particle size increases with higher loadings. The earlier mentioned influence of the calcination time on the mean particle size is illustrated in figure 7. After impregnation with the EDTA complex and drying at 393 K, the samples were calcined at 1273 K for varying
intervals. (The corresponding BET
measurements are shown In figure 2). The same measurements were performed after calcination at 773 K. Hardly any or no influence of the calcination time on the mean particle size was observed in this case. This was also observed for the Sn(II)EDTA "wet" sample; a mean particle size of about 57 nm was calculated, more or less Independant of the calcination time. Again a good agreement is obtained between the BET measurements and the observed X-ray line broadening. Electronmicroscopy Transmission as well as scanning electronmicroscopic study revealed that the catalysts made by impregnation with the above mentioned complexes lead to a complete and homogeneous coverage of the alumina support pellets. In all cases an initial film like structure of small deposited tin oxide particles was observed. Depending on the calcination temperature the deposited layer slowly broke up and larger Sn02 particles were formed. Usually discrete Sn02 particles were only present after calcination at temperatures higher
Fig. 8. Representative electron micrographs after calcination of the Sn(II)EDTA impregnate at 773 K for 21 hours (top) and 1273 K for 16 hours (bottom).
95
than about 923 K. Even then they were sometimes situated on top of a more or less amorphous layer of 5n02' Representative micrographs of the prepared catalysts are shown in figure 8. The mean particle size as measured from the micrographs agreed very well with those calculated from the X-ray diffraction patterns. However, exact measurements were sometimes obscured by the underlying film structure. Electron diffraction studies clearly showed that the crystallites became relatively large at calcination temperatures above about 923 K. Diffraction spots were observed in the diffuse halo's, indicative of crystalline particles. Line scan measurements as well as detailed studies of the distribution of the 5n02 particles within the pellet sometimes showed a somewhat higher loading on the outside of the pellet as compared to the fracture surface. However, usually this effect was only observed locally and limited to loadings higher than about one weight percent. Different results were obtained after calcination at 1523 K. The dense occupation with 5n02 particles had completely disappeared. In line with the other measurements it was established that severe sintering had taken place at this temperature. Temperature-programmed reduction Temperature programmed reduction experiments were performed to obtain information on the interaction between the 5n02 particles and the alumina carrier. Results from the literature (18-21) had clearly shown that reduction of 5n02 supported on v-alumina leads to stabilization of the 5n(II) species due to the strong interaction between 5n(II) and the alumina, even at temperatures well above the observed disproportionation temperature of (the thermodynamically unstable) pure 5nO, viz. 658 K (17). Therefore this technique could possibly discriminate between supported and unsupported 5n02' as it is well known that unsupported 5n02 is reduced to metallic Sn at elevated temperatures. Further information on the interaction of the 5n02 with the support and the extent of sintering at various temperatures can be obtained from factors such as the onset- and maximum reduction temperature, as well as the shape of the curve. Figures 9 and 10 show some typical TPR profiles as obtained after calcination at various temperatures for two different complexes. Figure 9 shows the results of the formate impregnate figure 10 of the EDTA impregnate. For comparison reasons the recorded TPR profiles of the physical mixtures of 5n02 and 5nO with the alumina support are also shown in both figures. The following features can be deducted from the spectra: i) calcination at higher temperatures leads to higher reduction temperatures, irrespective of
the complex used. The onset temperature is generally also shifted in the same direction. Ii) calcination at relatively low temperatures gives rise to a broad front-tail between 423
and 723 K. This band is not, or hardly observed with the physical mixture, nor after calcination at elevated temperatures (> 1023 K). Table 2 gives the relative peak areas as correlated to the peak area of the physical mixture of 5n02' The latter peak area exactly corresponded to the amount of consumed hydrogen
96
5
-
Sn02
--- Sno d
0
5 1523 K
c
'3 .5! c 0
:aE
0
5
:J III
c
1273 K
-£
b
0 0
0 SnlIntormate
5 573 K
a 573
773 T(KI -
973
1173
Fig. 9. TPR profiles for the Sn(II)formate "dry" impregnated sample. after calcination for 21 hours at a) 573 K b) 1273 K c) 1523 K d) the TPR profiles of the physical mixtures of SnO and Sn02' needed for the complete reduction of Sn02 to metallic Sn. This was calibrated by running several reduction profiles of known amounts of CuO. The catalysts made by impregnation and calcination had relative peak areas between 94 and 120%. It is therefore clear that the supported Sn02 is also completely reduced to metallic Sn, The spreading in the measured peak areas can be explained by factors such as base line drift and weighing faults. Some samples were run more than once and it proved that indeed a small, but reasonable spread could be expected. The area of the main peak was also calculated as to obtain information on the relative differences between the various samples. In alI cases the more or less arbitrary deconvolution of the peaks was performed in the same manner. In the first column of table 2 the percentage of the main peak area as compared to the total is given. The reduction to metallic Sn during TPR is surprising in view of the several reports (1821) which mention the stabilization of tin in the Sn(1I) state due to the interaction with the alumina support. A recent article (21) even reports the stabilization in this state after reduction of a mechanical mixture at 673 K. However, it must be stated that alI these
97
5 _
5n02
--- 5nO d Or-----t---~t=_--_+::.=-a"'--_____i
5 1523 K
c
::i
.!2
c 0 0
l5 ::J III C
0
773 K
IJ
b
N
:r
0 5
5n(II)EDTA 573 K
a
o 373
573
773
973
1173
T(K)-
Fig. 10. TPR profiles for the Sn(II)EDTA "dry" impregnated sample, after calcination for 21 hours at a) 573 K b) 773 K c) 1523 K d) the TPR profiles of the physical mixtures of SnO and Sn02' investigations were performed with y-A1 20 3 as a support material. Thus most presumably the stabilization of the Sn(II) compound is caused by the presence of acidic groups on the surface of the support. In this respect it is interesting that Dautzenberg et al (25) who used sodium-neutralized, non acidic alumina indeed only found metallic Sn after reduction. Another interesting feature in the TPR profiles is the front-tall after calcination at relatively low temperatures. This reduction at low temperatures was certainly not caused by a bi- or multi-disperse Sn02 particle size distribution; smaller particles giving rise to lower reduction temperatures. Electronmicroscopy proved that the particle size distribution was homogeneous throughout the entire pellet. The high percentage of the low temperature reduction peak, up to 54% (see table 2) certainly can not account for the relatively small variation in particle size. We believe that principally two factors can explain the reduction profiles: i) depending on the particle size, reduction proceeds via SnO, which is then further reduced to metallic Sn in a second step at more elevated temperatures, either by direct reduction or proceeded by the disproportionation of SnO to metallic Sn and Sn02'
98 ii) small Sn02 particles contain more defects which can enhance the rate of reduction.
After calcination at elevated temperatures larger sintered particles are formed and reduction is more difficult. This is also reflected in higher onset temperatures. To check the first hypothesis several TPR recordings were cut off at different temperatures and X-ray diffraction measurements were performed. Not once could detectable amounts of SnO be observed. Thus either a very small amount of (amorphous) SnO was formed, or the SnO which was produced at this temperature immediately disproportionated to give Sn02 and Sn, due to its thermodynamic instability. Conclusive: 2 Sn02 + 2H2 - - [2SnO + 2H20] - - Sn02 + Sn + 2H20 It is difficult to definitely rule out one of the above possibilities. It may well be that a combination of both is actually causing the reduction at lower temperatures. As stated earlier a peak possibly relating to the rare "red" modification of SnO was sometimes observed in the diffraction pattern, especially after calcination at elevated temperatures. It may well be that a small amount of this compound is stabilized by the alumina support, and is formed during pyrolysis of the complex. This may point to the possibility that the reduction to metallic tin is proceeded by the reduction of Sn02 to SnO.
TABLE 2 Peak areas of the TPR peaks of the catalysts in comparison the physical mixture of Sn02 and A1203. catalyst
peak area main peak(%) tota1(%)
Sn(II)formate "dry" 573 K 1273 K 1523 K Sn(II)EDTA 573 K 773K 1523 K Phys. mixture Sn02 SnO
column 1 column 2
52 100 94
114 120 94
46 83 100
66 78 114
105 116 114
63 67 100
100 23
100 23
100 100
with
*
100%
"dry"
FURTHER DISCUSSION AND CONCLUSIONS The results clearly show that the described
procedure of impregnation with
organometallic complexes leads to a homogeneous distribution of the active material over the preshaped support pellets. Electronmicroscopic study revealed that the Sn02 particles are at first deposited in a film-like structure. We believe that this is caused by a rise in viscosity during drying of the impregnation solution thus inhibiting the diffusion of the Sn
99 complex to the pore mouths. During evaporation experiments we have clearly witnessed this rise. Drying of the resulting treacly solution in some cases even caused a lower BET surface area than the pure support.
The
observed phenomena
of
combined
melting
and
decomposition (14) then gives rise to the observed homogeneous film structure. No crystallization took place during drying however, some hydrolysis of the EDTA and citrate complexes did take place during evaporation of the solvent. The decomposition of the complexes without an added support was rather spectacular. A foam was formed which overflowed the crucible. This was not observed after impregnation and subsequent drying and calcination, thus indicating that the decomposition of the complex only takes place on the support pellets. It proved possible to control the 5n02 particle size within narrow limits by choosing a suitable complex, loading, calcination temperature (and time). Especially the size of the complex is an important factor. The differences between the EDTA- and citratecomplexes are small compared with the formate complex. The rise in total BET surface area is much larger in the latter case due to the formation of very small 5n02 particles upon decomposition. The calcination temperature can also exhibit a large influence. Elevated temperatures lead to a break up of the film structure and to a sintering of the 5n02 particles. The Influence of the calcination time is somewhat smaller, especially at relatively low temperatures. The amount of impregnation solution is of minor influence however, it seems that a larger amount of solvent gives rise to a more rapid slntering, As the main conclusion it can be stated that the preparation procedure is a very suitable technique for obtaining a homogeneous distribution of the catalytically active material within preshaped support particles. Moreover, with 5n02 It proved possible to control the porosity within narrow limits. Especially with catalysts to be utilized in selective reactions (such as the epoxldation of ethylene) this is an important feature. This technique may also be utilized for the preparation of 5n02 films. These films are important mainly because of their industrial applications as transparent electrodes and as heat reflecting filters (26). Recently special attention has been directed to their potential use as gas sensors, since the resistivity of these films, especially those doped with Pd or Pt, is found to change appreciably upon exposure to various gases such as 02 and H2 (27). Methods now being used to prepare these films are reactive deposition, chemical vapour deposition and sputtering. The exact properties of these films are at least partly connected with the mean particle size (28). Therefore the preparation procedure as described in this report may also prove to be a valuable alternative for the preparation of these films. Acknowledgements The authors would like to express their sincere thanks to Mr's, A.Q.M. Boon, B.A. van Hassel, G.W. Koebrugge, M. van Leur and D.E. Stobbe for their part in the experimental work and to Mr. E.T.C. Vogt for his help In the preparation of the manuscript.
100
REFERENCES S.-Y. Lee and R. Aris. Catal. Rev. -Sci. Eng•• 27(1985)207 and references therein. 2 N. Kotter and L. Riekert. in B. Delmon. P. Grange. P. Jacobs and G. Poncelet (Eds.). Proc. 2nd Int. Symp. Preparation of Catalysts. Louvain-Ia-Neuve. Sept. 4-7. 1978. Elsevier. Amsterdam. 1979. p. 51. 3 K.P. de Jong and J.W. Geus. in G. Poncelet. P. Grange and P.A. Jacobs (Eds.). Proc. 3rd Int. Symp. Preparation of Catalysts. Louvain-1a-Neuve. Sept. 6-9. 1982. Elsevier. Amsterdam. 1983. p. 111. 4 G.R. Meima. A.J. van Dillen. J.W. Geus. J.E. Bongaarts. F.R. van Buren and K. Delcour. in preparation. 5 A.J.H.M. Kock. H.M. Fortuin and J.W. Geus. J. Catal •• 96(1985)261. 6 R.S. Tobias. Acta Chem. Scand •• 12(1958)198. Acta Odontol. Scand •• 16(1958)329. 7 P. Torell. E. Hals and T. M~rch. 8 E.1. Kondruk. G.V. Lavrova and V.A. Tsimmergakl. Russ. J. 1norg. Chem •• 15(1970)1667. 9 H.G. Langer. J. Inorg. Nucl. Chem •• 26(1964)59 and 767. H.G. Langer and R.F. Bogucki. J. Inorg. Nuc1. Chem •• 29(1967)495. 10 R.N. Lebedeva. E.M. Yakimets and E.F. Emlin. Russ. J. 1norg. Chem•• 12(1967)575 11 E. Bottari. A. Liberti and A. Rufo1o. J. 1norg. Nucl. Chem •• 30(1968)2173. 12 H.G. Langer and R.F. Bogucki. J. Chem. Soc. (A),(1967)1516. 13 J.D. Donaldson and J.F. Knifton. J. Chem. Soc •• (1964)4801. 14 J. Fenerty. p.G. Humphries and J. Pearce. Thermochim. Acta,61(1983)319. 15 Powder Diffraction File. Joint Committe on Powder Diffraction Standards. International Centre for Diffraction Data. Swarthmore. PAt 1976. card 211250. 16 Powder Diffraction File. see 15. card 13-111. 17 J.D. Donaldson. W. Moser and W.B. Simpson, J. Chem. Soc.,(1961)839. 18 A.C. Muller, P.A. Engelhard and J.E. Weisang, J. Catal •• 56(1979)65. 19 R. Burch. J. Catal •• 71(1981)348. 20 H. Lieske and J. V~lter. J. Catal.,90(1984)96. 21 R. Frety. M. Guenin. P. Bussiere and Y.L. Lam. in P. Barret and·L.-C. Dufour (Eds.), Proc. 10th Int. Symp. Reactivity of Solids, Dijon. Aug. 27-31. 1984, Elsevier. Amsterdam, 1985. p.1055. 22 N.S. Choudhury. R.P. Goehner. N. Lewis and R.W. Green. Thin Solid Films,122(1984)231. 23 S. Muranaka, Y. Bando and T. Takada. Thin Solid Films,86(1981)11. 24 T.E. Whyte jr•• Catal. Rev•• 8(1973)121. J.N. Helle. P. Biloen and W.M.H. Sachtler. J. 25 F.M. Dautzenberg, Catal •• 63(1980)119. 26 G. Frank. E. Kauer. H. Koestlin and F.-J. Schmitte in "Optical Coatings for Energy Effeciency and Solar Applications·. Proc. Soc. Photo-Opt. 1nstrum. Eng•• 324(1982)58. 27 T. Yamazaki, U. Mizutani and Y. Iwama, Jap. J. Appl. Phys.,21(1982)440. 28 H. Ogawa. A. Abe. M. Nishikawa and S. Heyekew«, J. Electrochem. Soc.,128(1981)685.
101
DISCUSSION A. VANNICE: How does O~ chemisorption vary on these Sn02!Ala03 systems as a function of calcination. How is 02 chemisorption on Ag dist1nguished from 02 chemisorption on the Sn02 (or reduced snO~) surface? Is 02 chemisorption a sensitive indicator of reduced Sn species. G.R. MEIMA : As yet we have not performed 02 chemisorption measurements on the pure Sn0 2!aA1 203 catalysts. However, we have obtained some preliminary results on the combined Ag-Sn02!aA1203 system. Reduction at elevated temperatures (400°C) leads to a large ennancement of the uptake amount of oxygen at 170°C, compared with reduction at 250°C. We believe that the larger uptake of oxygen is caused by the reduction of Sn02 at these temperatures, as the uptake amount of hydrogen after oxygen chemisorption is not enhanced. In accordance with your findings, we also have observed that it is possible to titrate the oxygen chemisorbed on silver with hydrogen. We therefore believe that it is only possible to titrate the oxygen chemisorbed on silver and not the oxygen on the reduced Sn-species. Whether or not 02 chemisorption is a sensitive technique for a reduced Sn species is still a matter of investigation. J.R.H. ROSS: Does the urea deposition method work for this type of a-A1 203based catalyst? G.R. MEIMA : When silica is used as the carrier material the urea method leads to excellent results. The deposition method, however, does not give rise to satisfying results when a-A1 203 is used as support. The precipitation of the Sn(hydr) oxide mainly occurs in the solution and not on the support. We believe this is caused by two factors: i) the high pH level of the zero point of charge of alumina (8-10) inhibits the adsorption of the positive Sn-species at low pH levels, and ii) the small number of OH-groups present on the surface of the support. K. NOACK: 1) What happens to the organic part of your precursor? completely burned off? 2) Did you use XPS to determine the oxidation state of Sn?
Is it
G.R. MEIMA : We have studied the burn-off of the organic precursor by means of a mass spectrometer. Depending on the specific complex used, the precursor is completely burned off at temperatures between 200 and 600°C. As yet we have not used XPS to determine the oxidation state of the Sn. However, X-ray diffraction and TPR measurements only indicate the presence of Sn02' M.V. TWIGG: What effect does the presence of tin oxide on the activity and selectivity of your catalysts in ethylene oxidation? G.R. MEIMA : As yet it is somewhat premature to comment on the exact influence of the Sn02 on the activity and selectivity of the Ag-Sn0 2!aA1 203 catalyst in the ethylene epoxidation reaction. We are still in the process of the evaluation of the catalyst with respect to the optimum particle size for both Sn02 and Ag. However, the results obtained so far are promising. The selectivity and activity of the catalysts is comparable with commercial catalysts of the same particle size if no additional promoter is added.
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B. Delmon, P. Grange. P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
103
PREPARATION AND CHARACTERIZATION OF SOL-GEL BASED CATALYSTS FOR THE SELECTIVE CATALYTIC REDUCTION OF NO WITH NH3 H.BARTEN, F. JANSSEN, F. V.D. KERKHOF and R. LEFERINK Arnhem
Institutions of the Dutch
Electricity Utilities.
N.V.KEMA, R&D Division,
Department of Chemical Research. P.O.Box 9035, 6800 ET Arnhem (The Netherlands). E.T.C. VOGT, A.J. van DILLEN and J.W. GEUS State University of Utrecht. Inorganic Chemistry Department, Croesestraat 77A. 3522 AD Utrecht (The Netherlands).
SUMMARY The preparation of sphere-shaped silica supported vanadium catalysts was carried out by using various techniques. The silica support was prepared by the sol-gel process, developed by KEMA. After the sphere formatIon, drying and sintering steps, the spheres of controlled diameters can be impregnated and coated by existing techniques, such as: wet impregnation, the socalled monolayer adsorption, which was developed at the Twente University of Technology (The Netherlands) and the homogeneous deposition precipitation, developed at the State University of Utrecht. The characterization of the catalysts was carried out by using various methods such as: activity measurements with a conventional fixed-bed reactor coupled with a mass spectrometer, temperature programmed reduction, scanning electron microscopy, X-ray fluorescence, BET-sorption and mercury-porosimetry,
INTRODUCTION The nitrogen OXides, present in stack and flue gases from fossile fueled burners are derived from two sources; one is the nitrogen fixation (called thermal NOx) and the other is the reaction of atmospheric and fuel oxygen with nitrogen present as nitrogen compounds in fossile fuels (fuel NOx)' A significant fraction of the thermal NOx can be reduced by combustion modification, I,e, staged combustion and flue gas recirculation. Also boiler and firing system geometries can influence the amount of NOx; low-NO x burners have been developed and are capable of reducing NOx by 40-65% compared with conventional burners. More powerful NOx removal techniques are the wet and dry processes, such as catalytic
104
and non-catalytic processes. In particular the selective catalytic reduction of nitric oxide with ammonia has been considered as a promising technique. Various proposed and tested catalysts have been described for the reduction of NOx by means of reducing agents NH3, CO, H2 and hydrocarbons. Supported vanadium oxide catalysts in particular show high activities for the reduction of nitric oxide with ammonia into nitrogen and water [1-4]. The main aim of the present project was to develope catalyst supports and catalysts, mainly vanadium, with the aid of a sol-gel technique. The sol-gel process makes it is possible to shape the support of the catalyst into a desired texture. Because of their uniform size sphere-shaped vanadia/silica catalysts are applicable in fluidized bed reactors and can be produced at low costs compared with titania supported catalysts. Moreover, vanadia/silica systems are stable in an acidic environment. Sol-gel processes were developed in relation to the production of nuclear energy at the KEMA laboratories since the fifties. Spheres of fissile material such as U0 2 and Th0 2 [5,6] have been prepared. The diameters of these dense spheres are 0.005 to 2.0 mm, often in small frequency distributions. During these studies, occasionally other compounds have been prepared such as oxides of other metals, mixed compounds or carbides. In the preparation processes the texture of the initially porous solids appeared very important in order to obtain maximum density on sintermg, Therefore, influences of process parameters on the texture of the products have been extensively studied. Recently the preparation of titania gel spheres was described [7]. The potential applications of the sol-gel technologies in the glass industry were reviewed [8]. Sol-gel processes have been used for the preparation of silica [9] and alumina [10] supports for catalysts for olefin hydration and reforming. This paper discusses in outline the general methods of preparation and characterisation of silica supports and of silica supported vanadia catalysts, and then discusses the activities and selectivities of some of the resultant materials.
EXPERIMENTAL Analar and reagent grade chemicals, from various sources, are used throughout.
Support preparation The sol is made by mixing sodium silicate and nitric acid (2.5 M) at 0 "C until a pH of 1.2 is obtained. Hexamethylenetetramine (0.5 M) is then added to the cooled solution till pH
= 4.3. This solution is introduced via a vibrating capillary into silicone oil or a
105 paraffin/perchloroethylene mixture at 80°C (Fig. I). The sol is converted into droplet form
2-
and
gelated within 5 to 10 seconds. The spheres
-
prepared
in
silicon
oil
and
in
paraffin/perchloroethylene will be designated in
..···.. ··
the text as spheres (sil) and spheres (par/per) respectively• The gelation liquid is removed by washing the
! 3
spheres with pure trichloroethylene. The spheres are washed with an ammonia solution (0.02 M), airdried during 18 hours and then dried at 120°C during 4 hours. The
..· ·.
spheres
are
characterized
with
various
techniques and treated further, as described in the results.
Preparation of the catalyst Three methods of preparation of silica supported catalysts are used In this study; wet impregnation 1::!;===~-6
with
an
aqueous
solution
of
ammonium metavanadate; monolayer adsorption of
Fig. 1
vanadyl acetylacetonate as described by Gellings
Apparatus for the manufacturing of silica spheres. 1. Storage tank with the cooled sol; 2. Vibrating capillary; 3. Silicone oil or paraffin/ perchloroethylene mixture; 4. Sieve for the removal of the spheres; 5. Heater; 6. Pump.
et
al.[11,12]
and
Homogeneous
Deposition
Precipitation (HDP) of V(III)oxyhydroxlde. The HDP method, which is described by Geus In general [13], is used because of the difficulties In the preparation of higher loaded vanadia/silica adlayered catalysts using the other methods.
Aqueous solutions of vanadium(III) sulphate are prepared by cathodic reduction of NH4V0 3 in sulphuric acid (pH = I) as will be described in detail elsewhere [14]. All catalysts are oven-dried at 120°C. The catalysts prepared by wet Impregnation and by monolayer adsorption are then calcined in air at 500 °C. The catalysts made by HDP are heated In a hydrogen/argon flow up to 500°C at a rate of 5 K/min. In Table 1 a survey is given of the catalysts used in this study. Characterization of the support and of the catalysts The supports and the catalysts are characterized by various methods and techniques. Temperature Programmed Reduction (TPR) is performed with a modified experimental setup, described elsewhere [IS]. The apparatus Is calibrated by injecting known amounts of hydrogen into the reactor. TPR is used to determine the vanadium content of the catalysts
106
and the method was checked with XRF (Philips PW 1410/20). Prior to TPR the catalysts are calcined in situ at 400°C for one hour in a stream of helium, containing oxygen (25 %) and subsequently cooled to 200°C. The BET surface area and the pore size distributions of the catalysts are determined by using a Carlo Erba Sorptomatic (type 1800) with N 2 as an adsorbate at 77 K and a mercury porosimeter (Carlo Erba 200). A steady state plug flow reactor is used to determine the activity of the catalysts; the influence of the oxygen concentration on the reduction reaction and concentration profiles of NO, NH3, N2' H20 and N20. The gases, 2200 ppm NO and 2000 ppm NH 3 and 20,000 ppm 02 in helium were purchased from Air Products and are used without further purification. A gas mixture of 1900 ppm 15NH3 in helium is made a by conventional gravimetric method. The gas mixtures are made-up with the aid of mass flow controllers (Matheson, USA). The reactor is coupled via an adjustable leak valve on a mass spectrometer. Further details of set-up and procedures are previously reported [I]. The various parameters used during this study are: inlet pressure 1 atm, the inletconcentratlons of NO and NH3 are 500 ppm. The balance is helium with 2% (vol.) oxygen. The whole system is processed by an Apple II data system. RESULTS The supports and the catalysts used and their 10'E
Ci
.s ~i'N" rn
(;;500
5
tual
~Q)
rn 400
:::l '" ~
l!!
300
0 5 c,. c: 01 Q)
E
200
main characteristics are listed in Table 1. The porosity of the spheres obtained after the drying step at 120°C is 50% approximately. The BET surface area is about 500 m 2/g and they have a small pore size distribution (pore radius 2.3 ± 0.5 nm), After heating the spheres (sil] in an oven
during one hour the relationship of the specific surface area and the mean pore radius as shown in Fig. 2 is obtained. Support A (Table 1) is prepared by heating the support in a helium stream at 300
100
°C. One part of the heated spheres are heated in
o 200
400
800
temperature (OC)
air at 650°C in order to enlarge the pore size; support B. In Fig. 3 the function of the pore radius of the spheres (par/per) and the heating time is given.
Fig. 2 Effect of the heat treatment of the spheres (sil) on the specific surface area and the mean pore radius during 1 hour in air.
'The .spheres are submerged in water in an autoclave during the heat treatment. The pore radius can be influenced by the temperature inside the
autoclave
and
the
presence
hexamethylenetetramine (hexa) in the sol. The
of
107
TABLE 1 The silica supported vanadia catalysts used in this study. Catalyst A is made by monolayer adsorption, B by wet impregnation, and E by HDP. CATALYST SUPPORT mean pore V205 pore volume content radius cm3/g 1)/2) wt% nm
Catalyst Support code prepared by Sol-gel Sol-gel Sol-gel Sol-gel Aerosil
A B C
D E 1)
2) 3) 4) 5)
3) 3) 4) 4) 5)
2
520 80 140 140 186
0.25 0.12 0.97 0.97
7
9 9
and C,
D
ppm N2/ mg cat. 300 DC
2.1/2.1 4.3/4.5 20/24/44/-
0.1 0.4 4.3 1.9 6.0
0 7
9 15 1
vanadia content determined by temperature programmed reduction after oxidation vanadia content determined by XRF Sol-gel silica manufactured by gelation in silicon oil Sol-gel silica manufactured by gelation in paraffin/perch1oroethy1ene Degussa, Aerosi1 200 V lower curve in Fig. 3 is obtained when no hexa is present in the sol and the temperature inside the autoclave is 150°C. The two upper curves are found when hex a is used and the temperature in
A the autoclave is 140 and 150 °C respectively.
Support C and D are the results of the treatment of the spheres (par/per) in the autoclave.
c Activity and selectivity of the catalysts Catalyst A of Table I starts to show some
o
2
4
6
8
heating time (hrs)
Fig. 3 Effect of the time of heating of the spheres (par/per) in water on the pore radius at different temperatures. The temperatures in the autoclave for the three curves A, Band Care 150 DC, 140°C and 150°C respectively. Curve C reflects the results of the sol-gel products without using hexamethylenetetramine. The data of curve A and Bare measured by mercury porosimetry and the data of curve Care determined by BET.
activity at 250 DC; the maximum
attainable
conversion is about 30% at 500 DC. However, at the temperature 350 DC the production of N20 starts and reaches a steady state level of 40 ppm at 500 DC. Catalyst B is prepared by wet impregnation and shows different behaviour in the reaction of NO, NH3 and 02' Fig. 4 presents the concentration levels of the reactants and products on 200 mg of catalyst
B,
at
different
temperatures
after
calcining the catalyst at 400 DC in a flow of helium containing 25% oxygen. Roughly two areas can be distinguished in Fig. 4. Up to 350 DC the following overall reactions take place
108
4NH3 + 302 ----) 4NO + 4NH3 + 02 ---->
2N2 + 6H20 4N2 + 6H20
(1) (2)
Reaction (1) dominates over reaction (2) at 350°C. Above 350 °C the selectivity of these reactions decreases and the ammonia becomes oxidized to nitric oxide; reaction (3)
4NH3 + 502 ----) 4NO + 6H20
(3)
Fig. 5 and Fig. 6 show the results of the
'E0- 8OO
reaction between NO, 15NH3 and 02 on the catalysts E and D (Table l) respectively.
S e
o
The behaviour of these catalysts is completely
i ..= 600
different from that of the catalyst described in
8c:
Fig. 4.
8
The catalyst is introduced in the reactor for
400
activity measurements after reducing the catalyst in the TPR apparatus (l hour at 50OOC). During the activity measurements the color of the catalysts
200
C, D and E changes from black to yellow-brown (V203
--> V205)·
As can be seen in the figures the conversion of NO temperature (0C)
to N2 is remarkable higher than in case of catalyst
B. The activity and the selectivity show a maximum
Fig. 4 Concentration
at 300°C. This activity remains constant even profiles of NO,
NH3' N2' H20 and N20.
[NOli = [NH3li = 500 ppm, [02]i = 2%, balance helium.
space velocity 30 l/g.h
after calcination of the catalysts during 20 hours. Above 300 °C 15NNO is produced from the reaction of NO and 15NH3 according to the overall reaction
4NO + 415NH3 + 302 -----) 415NNO + 6HZO
(4)
Above 3500 C both catalysts show a different behaviour with respect to the 15NH3 concentration. In Fig. 5 the concentration of 15NH3 decreases, while the concentration of 15NH3 in Fig. 6 increases. Moreover, at higher temperatures also the compounds 15NO, 15N2 and 15N20 are present at low concentrations (Table 2); the reaction products of the oxidation of 15NH3•
109
E
E
500
c:
400
~
0
0
~
E 300
2l
e 0
e
500
Q. Q.
Q.
.s,
~ 300 E
NO
C>
200 100 0
400
o e 200 0 o
15NH3
100
15NNO
15NH3
200
300
NO
0
400
400 temperature (0 C)
temperature (0 C)
Fig. 5
Fig. 6
The activity of the (44%)V205/SiOZ catalyst on the reaction of NO, 15NH3 and 02' The support is aerosil and the catalyst is prepared by deposition of V203 and afterwards calcining the catalyst at 400 ·C for 2 hours. [NO]i 500 ppm; [lSNH3]i 475 ppm; [OZ]i = 2%; balance helium; space velocity 120 l/g.h
The activity of the (24%)V205/Si02 catalyst on the reaction of NO, 15NH3 and 02' The support is prepared with the sol-gel technique and the catalyst is prepared by deposition of V203 and afterwards calcining the catalyst at 400 ·C in a stream of helium with 25% 02 for 20 hours. = [lSNH3]i = 500 ppm; [NO]i [OZ]i 2%; balance helium space velocity 60 l/g.h
TABLE 2
Some of the by-products formed as a result of the reaction of NO, at 400 ·C on three catalysts. Concentrations are given in ppm.
15NH3 and
02
Catalyst C
77
D E
108 110
63 26
47
41
28 25
5 10 8
Obviously a considerable amount of about 100-160 ppm of 15NH3 is oxidized by oxygen and about 70-110 ppm reacts with NO to form 15NNO at 400 ·C. However, at the temperature of maximum NO conversion only small amounts of N2 0 are found (Table I)
DISCUSSION Support The role of hexamethylenetetramine (hexa) is to produce ammonia, which is formed when heating the sol at 80°C [16]. As a result the pH of the sol increases. This stimulates the
110
polymerization of silanol groups in the sol. Two additional functions of hexa are to promote pore growing kinetics as can be seen in Fig. 3, and the unreacted hexa in the spheres leaves spaces behind when it is removed after washing and calcining. Activity It has been suggested [17] that the lack of interaction between V20S and the silica support during preparation results in badly dispersed V20S with a low thermal stability. This leads already at relatively low temperatures to a considerable sintering of the active component especially with higher loadings [18]. Those catalysts exhibit poor catalytic properties with respect to the conversion of NO with NH3 into N2 and H20 [19,20] In accordance with these results our catalysts prepared by impregnation show a negligIble activity for the reduction of NO. During the preparation by wet impregnation of the catalyst at pH values higher than about 2 the surface hydroxyl groups of the silica deprotonize, bringing about a negatively charged silica surface. which dis favours interaction with the negatively charged vanadium(V)oxy anions. Vanadium(I1I) ions in water as solvent exhibit a more basic behaviour and interaction of the silica support with its VO+ ion is more probable. The monolayer catalyst therefore, despite its low vanadia-content shows an interesting activity and selectivity. This indeed gives evidence for the supposition that interaction between the vanadia and the support is a prerequisite. Therefore the HDP catalysts can be expected to be comparable with catalysts like Fe203/Si02' V20S!y-AI203 and V20S/Ti02 [I].
formation Takagi-Kawai et al, observed N20 in the gas phase when NO was introduced at 300°C onto V20S/Si02, which was reduced by ammonia at 200°C [3]. ~2Q
This effect is also found on the freshly reduced catalysts C, D and E. When the catalyst is not calcined first N20 is produced in large amounts at low conversion of NO into N2• After calcining the catalyst in an oxygen/helium flow only N20 is formed which is coming from reaction (4). Thus for the N20 formation three overall reactions are responsible: the reduction of NO by the reduced surface, the oxidation of ammonia and reaction (4). CONCLUSIONS The interaction between vanadia and the silica support is a prerequisite for a good performance of the catalyst for the selective reduction of NO with NH3 in the presence of oxygen. The best method to prepare the vanadia/silica catalysts is the homogeneous deposition precipitation resulting in V 203/Si02, which can be oxidized to V20S/Si02•
111 The sol-gel technique is a suitable technique for support preparation of vanadia catalyst. The optimum temperature for the working catalyst is 300 o e. Above that temperature large amounts of by-products are found coming from the oxidation of NH3 by NO and 02'
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14 15 16 17 18 19 20
F.J.J.G. Janssen and F.M.G van den Kerkhof Kema Scientific & Technical Reports 3 (1985) 71-85. G.L.Bauerly, S.C. Wu and K. Nobe Ind. Eng. Chem •• Prod.Res.Dev. 17 (1978) 117-122. M. Takagi-Kawai. M. Soma. T.Onishi and K.Tamaru Can. J. Chern •• 58 (1980) 2132-2137. M. Inomete , A. Miyamoto. T. Ui , K.Kobayashi and Y. Murakami Ind. Eng. Chem •• Prod. Res. Dev., 21 (1982) 424-428. M.E.A.Hermans. Sci. Cer •• 5 (1970) 523-538 P.A.Haas. Nuc1. Technol •• 10 (1970) 283-292 S. Komarneni and R. Roy, Mat. Letters, 3 (1985) 165-167 J. Wenzel. J. Non-Crystall. Solids, 73 (1985) 693-699 T.G.Spek, European Patent 0067459, (1982) M.A.Day, European Patent 0115927. (1984) A.J. van Hengstum, J.G. van Ommen. H. Bosch and P.J.Gellings Applied Catal., 5 (1983) 207-217. J.G.van Ommen. K.Hoving. H. Bosch. A.J. van Hengstum and P.J.Gellings. Z. Phys. Chem •• N.F., 134 (1983) 99-106. J.W.Geus, In: Preparation of catalysts III, Studies in surface science and catalysis. Scientific bases for the preparation of heterogeneous catalysts. Vol. 16 (1983) 1-33. G. Poncelet. P. Grange and P.A.Jacobs (eds.). Elsevier. Amsterdam. E.T.C. Vogt, M. de Boer. A.J. van Dillen and J.W.Geus, to be published. H. Bosch. B.J. Kip. J.G. van Ommen and P.J. Gellings J. Chem. Soc., Faraday Trans., 1 (1984) 2479-2488 H. Barten. to be published. Y. Murakami, M. Inomata, K. Mori. T. Ui, K. Suzuki, A. Miyamoto and T. Hattori. In: Preparation of catalysts III. Vol. 16 (1983) 531-551 • See ref. 13. F. Roozeboom, M.C. Mittelmeyer-Hazeleger, J.A. Moulijn, J. Medema. V.H.J. de Beer and P.J. Gell ings , J. Pbys , Chern •• 84 (1980) 2783-2791 T. Shikada, K. Fujimoto. T. Kunugi, H. Tominaga, S. Kaneko and Y. Kubo, Ind. Eng. Chem. Prod. Res. Dev•• 20 (1981) 91-95 T. Shikada. K. Fujimoto. T. Kunugi and H. Tominaga J. Chem. Tech. Biotechnol. 33A (1983) 446-454
112
DISCUSSION E.B.M. DOESBURG Is the hydrothermal treatment the only way to control the texture of the s 1ica balls or are there other methods like, for instance, heating of the 1 quid in which the green balls are formed by gelling? H. BARTEN: Indeed, other methods can be used. In fact we heat the spheres in the gelation liquid, which enhances the mechanical strength. A partial drying step, prior to the hydrothermal treatment, influences the total pore volume. J. B.NAGY : How do surfactants or polymers influence the size of the sol-gel particles? H. BARTEN: These compounds influence the surface tension of the sol. In this way, the moment of breaking-up of the liquid system can be varied, e.g. the droplet size. L. HEINRICH: Do you have experience in resistance against HF and S02 (in flue gases)? H. BARTEN: So far, we have not obtained experience with SO~ and HF, but we expect the silica supported catalysts to be very stable agaTnst S02' G. CENTI : The use of Si0 2 as support instead of the most used Ti02 can be justified in order to prepare catalyst with better performances in fluid bed reactor. Have you data on comparison of the activity in NO x reduction with ammonia between the V-Si02 and V-Ti02 catalysts and data of applicability of V-Si02 to fluid bed reactor conditions? H. BARTEN: We have no experience with a fluid-bed application. The vanadia on titania is a more active catalyst than the vanadia on silica. Despite of lower activity (so far), the silica spheres of well control able texture may be preferred. Furthermore, the selectivity of our catalyst is 100% at 300°C. J.G. van OMMEN: - If V20S is a better catalyst for NO x removal, why do you use Si0 2 as a support? - Do you have an idea on the spreading of the V20S on Si0 2 compared with the spreading of V20S on Ti02? H. BARTEN: The texture of silica is much better control able than that of titania. Furthermore, besides the activity, the selectivity is very important. The dispersion of vanadium oxide is very good. More detailed information on this subject will be published shortly by the workers of the University of Utrecht. S. TAMHANKAR : What was the vanadium loading in your catalyst? Have you done any mechanistic study to understand the changes in the vanadium state that may take place during reaction? H. BARTEN: The vanadia loading of the catalysts was 2 to 24% (table 1). During the reaction, the color of the catalyst changes from yellow to black; this can indicate the formation of V(IV) on the catalyst surface.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
113
HIGH SELECTIVITY OF CVD Sn02/Si02 CATALYST FOR OXIDATIVE DEHYDROGENATION OF ETHYLBENZENE
T. HATTORI, S. ITOH, T. TAGAWA and Y. MURAKAMI Department of Synthetic Chemistry, Faculty of Engineering, Nagoya University, Furo-cho, Chikusa-ku, Nagoya 464, Japan.
SUMMARY CVD (chemical vapor deposition) method have been applied to prepare Sn02/ Si02 catalyst for further improvement of the sel ecti vity. The catalysts were prepared by the reaction of SnC14 vapor with Si0 2 surface followed by the hydrolysis of deposited Sn compound. CVD catalysts showed completely different behavior in the oxidative dehydrogenation of ethylbenzene from conventional impregnated (IMP) catalysts. Deposition of coke, which is unavoidable on IMP catalysts, was not observed, and CVD catalysts did not show apparent initial increase in the rate of styrene formation. HM, XRD, UV, IR, and acid-base titration indicated that CVD cata lysts had rather 1arge Sn02 particles and acid sites of moderate strength. High selectivity of CVD catalyst was attributed to moderate acid strength due to the controlled dispersion of Sn02' INTRODUCTION Chemical vapor deposition (CVD) may enable us to prepare a different type of supported catalysts from those prepared by conventional impregnation (IMP) method. In the CVDmethod,-it is possible to control the support surface before the deposition of catalyst component and to control the atmosphere during the deposition. These may lead to the different type of reaction of catalyst reagent with support surface from the reaction in excess sol vent. Or, in other words, catalyst component may be deposited on different surface sites of the support, resulting in the supported catalysts with the different structure and surface properties from IMP catalysts. We have conducted the screening of cata lyst for the oxidati ve dehydrogenation of ethyl benzene (ref. 1), and found that a Sn02-P205 catalyst showed high selectivity (ref. 2). Further, we have clarified that the acid sites of medium strength and the strong base sites are necessary for the selective formation of styrene (ref. 3), and, on the basis of the reaction mechanism, we have developed a Sn02/Si02 catalyst (ref. 4). Major by-products in the present reaction are carbon monoxide and carbon dioxide. On the selective catalysts such as Sn02-P205 and Sn02/Si02' the direct oxidation of ethylbenzene is not serious, but the consecutive reactions of styrene, such as the
114
formation of carbonaceous deposit, are significant. It has been suggested that carbon monoxide and carbon dioxide are formed by the oxidation of deposited material s (ref. 3). Therefore, for further improvement of sel ectivity, it is required to suppress the formation of carbonaceous materials. In the present study, CVO method(ref. 5) have been appl ied to prepare Sn02/Si02 catalyst for further improvement of catalysts, and the catalytic activity and selectivity in the oxidative dehydrogenation of ethyl benzene has been discussed in the light of the structure and the acid-base properties of catalysts. EXPERIMENTAL Preparation of Catalysts CVO catalysts were prepared by repeating the following CVO cycle at 453K: (1) the desiccation of si 1ica gel (Micro Beads Type 50, Fuji Davison Chem. Co., 240 m2/g) in a flow of dry N2, (2) the deposition of SnC14 vapor on sil ica gel surface, (3) the purge of excess SnC1 4 by flowing dry N2, and (4) the hydrolysis of deposited Sn compound by water vapor. The amount of deposited Sn was control led by the number of CVO cycles repeated. The number was represented by the numerals in catalyst symbol, e.g., CVO-3 catalyst was prepared by repeating the CVO cycle three times. The resulting material s were washed with 4N NH 40H five times to remove residual chloride ions and with water to remove NH 40H, and then dried and ca 1cined at 773K for 2h in flowing air. Impregnated (IMP) Sn02/Si02 catalysts were prepared by the impregnation of SnC1 2.2H20 from ethyl alcohol solution fol lowed by the treatment with NH 40H and water and by drying and calcination. The numeral s in catalyst symbol represent Sn02 content in wt%. Catalytic Property and Characterization The activity and selectivity in the oxidative dehydrogenation of ethylbenzene were measured by the conventional fl ow reaction as descri bed e 1sewhere (ref. 1-4). Standard reaction condition was as follows: catalyst weight was 19, total feed rate was 250 mmol/h, and partial pressures of ethyl benzene. oxygen and water vapor were 0.060. 0.061 and 0.485 atm. respectively. Acid and base properties were examined by the same procedures as those in the previ ous paper (ref. 3). TEM microphotographs were measured by using H 700H (Hitachi). The UV-visibl e diffuse refl ectance spectra were measured by using UV IDEC-505 (JASCO).
115
RE5ULT5 Physical Texture Table 1 summarizes physical texture of the catalysts. cycles resulted in the increase of 5n02 content. almost linearly in the initial three
The increase of CVD
5n02 content increased
cycles, but only less amount of 5n02
was deposited in the fourth and fifth cycles; in the initial three cycles 4-5 wt% of 5n02 was deposited in each cycle, but only a half of it was deposited in the 1ast two cyc 1es. In the X-ray diffraction patterns (XRD), only a very broad line due to 5i02 could be observed in the case of CVD-1 and CVD-2.
However, in the case
of CVD-3 and CVD-5, broad and weak 1 ines due to 5n02 also were observed. average part i c 1e di ameter was 5.6 nm for CVD- 3 and 8.3 nm for CVD-5.
The
In the
case of IMP catalysts, no diffraction lines due to 5n02 could be detected, though IMP-17.7 contains more 5n02 than CVD catalysts. Figure 1 shows transmission electron microscopic (TEM) photographs of CVD2 and IMP-U.5, respectively.
Dark spots represent 5n02 particles; diffrac-
tion pattern by selected area diffraction method agreed with that of 5n02' In the IMP-6.1 catalyst, 5n02 particles with a diameter of about 1 nm were dispersed uniformly over the whole primary particles of 5i0 2 support. In the cases of IMP-11.5 and IMP-I?? catalysts, essentially the same photographs were obtained except that the particle diameter was about 2 nm and the number of 5n02 particles increased with 5n02 content.
TABLE 1 Texture and Acid-Base Properties of 5n02/5i02 Catalysts Catalyst
5noi Con ent (wt%)
5n02 Particle 5ize (nm) TEM XRD
5i02
0.0
CVD-1 CVD-2 CVD-3 CVD-5
4.9 9.8 14.3 16.9
n.d. n.d. 5.6 8.3
IMP-6.1 IMP-II. 5 IMP-17.?
6.1 11.5 17.7
n.d. n. d. n.d .
5n02
40
ca. ca. 10 10 -
3 3 15 15
ca. 1 ca. 2 ca. 2
Ho max pKa max
Aciditya (mmol/g)
UV-Visible Absorptn max (nm)
+4.0
< +9.8
0.0
+1. 5 +1.5 +1.5 +1.5
(+17.2)b (+17.2) (+18.4) (+18.4)
0.053 0.010 0.037 0.030
25? 265 267 269
(+18.4) (+18.4) (+18.4)
0.095 0.081 0.115
255 256 258
(+15.0)
0.0
285
+1. 5 b (-5.6)b (-5.6) +1.5
a +1.5 > Ho > -5.6 b color of indicator changed a little.
116
In the case of CVD catalysts, the results were astonishing.
As shown in
Fig. I-a, Sn02 particles were deposited on only a part of Si0 2 primary particles, forming clusters. But some of Si02 particles remained vacant. Simil ar photographs were obtained in the other CVD catalysts. particle diameter of Sn02 was about 3 nm.
In CVD-1,
In CVD-2, particle diameter
increased a 1 ittle, and the number of Sn02 particles in each cluster decreased.
In CVD-3 and CVD-5, the di ameter of 5n02 parti c les
increased up
to 10-15nm, and new clusters consisting of small Sn02 particles were formed. It shaul d be noted that 5n02 was deposi ted only a few primary parti cl es of 5i02' but not on the other Si0 2 primary particl es. UV-visible spectra of Sn02 were essentially the same as that reported by Sala and Trifro (ref. 6).
It had an absorption maximum at 285 nm.
As
shown in Table I, the absorption maximum was shifted to lower wavelength. The shift was more remarkable in IMP catalysts than CVD catalysts, suggesting stronger interaction of Sn02 with the support in IMP catalysts. Acid and Base Properties The maximum acid and base strength (Ho max and pKa max) of catalysts were shown in Tabl e 1.
Si0 2 had on 1y very weak aci d sites of Ho >+3.3, and 5n02 had weak acid sites of Ho > -3.0. The maximum acid strength of CVD catalysts
was identical to that of Sn02' but IMP catalysts with high Sn02 content had stronger acid sites.
Although Si02 and Sn02 did not have strong base sites,
both CVD and IMP catalysts had such strong base sites of pKa
Fig. 1
> 17.2.
TEM photographs of CVD-I (a) and IMP-11.5 (b) catalysts.
117
It has been shown in the previous paper (ref. 8), that the acid sites of 1.5
> Ho > -5.6 are necessary to adsorb and activate ethyl benzene. The
acidi ties of concern are shown in Tab le 1. sites of Ho
< -5.6
It sho ul d be noted that acid
were not present on CVO catalysts, and were very few on
IMP catalysts.
Thus, Table 1 shows the total amount of acid sites stronger
than Ho = 1.5.
The acidity of IMP catalysts increased with Sn02 content.
However, CVO catalysts again gave very astonishing results.
The acidity was
the largest on CVO-l, took the minimum on CVO-2, and then increased on CVO-3 and CVO-5.
Thi s resul ts sugges t that the aci d sites on CVO-l may be diffe-
rent in nature from those on the other CVO catalysts. IR spectra of adsorbed pyridine also were measured. On Si0 2, only bands due to hydrogen bonded pyridine (HPy) were observed at 1447 and 1597 cm-l, but no bands due to pyridinium ion (BPy) and coordinately held pyridine (LPy) could not be observed.
On CVO-l, LPy bands at 1453 cm-l and 1613 cm-l as
well as HPy bands were observed.
The intensity of LPy bands decreased on
CVO-2, and increased again on CVO-3 and CVO-5.
In the case of IMP catalysts
strong bands of LPy as we1 1 as HPy bands were observed. Catalytic Properties Figures 2 and 3 show the time course of oxidative dehydrogenation of ethyl benzene. The results on CVO catalysts were again astonishing.
It has
been reported that the rate of styrene formation increased with time on stream in the initial several hours on many catalysts, such as Si02-A1203 (ref. 3), zirconium phosphate (ref. 7 and 8) and A1203 (ref. 9) and the initial increase in the styrene formation has been ascribed to the condensa4
..-....
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~
."
n:l
s, 0
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n:l
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<E---
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Q)
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~
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.
0::
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.' ,.
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...,'"
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" .... " styrene
.,,"'. "
~
2 I-
.... .....-
A- -.-~--.-:t
100
I
°
2
°
(h)
Fig. 2 Time course of oxidative dehydrogenation of ethyl benzene on IMP-ll.5 at 747K.
118
0.50
-l:i.-.f:..-/i
- t:..-t:..-t:..-~-l:i.-IJ ~
K
"<::
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o
r-,
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...,OJ
0.25
ttl
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ttl
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-
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o
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+ CO2 -[J-[J-[J-[J""""'iD-[J 2 3 4 ~CO
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o
o
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Time on stream Fig. 3 7I3K.
-0-0-°
ttl
<=
OJ 0:::
100
~
( h)
°
-n-.n-[J 2
o
Time course of oxidative dehydrogenation of ethyl benzene on CVD-5 at
tion of produced styrene to form carbonaceous material s. On IMP-17.7, as shown in Fig. 3, the rate of styrene formation increased remarkably with time on stream, and the color of the catalyst turned from white to black. And the calcination of used catalyst in flowing oxygen decreased the rate of styrene formation to the initial value. These results indicate the deposi-
1. 50
Sn02 content (wt%) 15 10
5
.---- ----.."......-. ",,' ,""
c<,
.;,
......
'0 1.0
It:..;'
!
II
...,
ttl
s,
......., 0
/ '
I ,.
<=
'
,~
90~
~
80 .;; ....
..., u
OJ
Qj
70 Vl
I
~/~O
u
ttl
/
OJ 0:::
0
"
/ ,.styren~O-
I
0.5
100
t:..-._ ,-
I:i.
~
..c:::
17
a
/
co, CO,
60
~ ----..-------
-- -
4 3 2 Number of eva cycle
5
Fig. 4 Effect of CVD cycle on activity and selectivity of CVD catalysts (open marks) and IMP catalysts (closed marks) at 723K.
119
tion of carbonaceous materials. On CVD catalysts, however, gradual decrease in the rate of styrene formation was observed rather than the increase, as Further, the color of the catalyst remained white after shown in Fig. 2. the reaction, indicating the absence of the deposition of carbonaceous materials. Figure 4 shows the activity and selectivity of catalysts as a function of Sn02 content. On both CVD and IMP catalysts, the activity increased with Sn02 content, but IMP catalyst had higher activity than CVD catalysts. The se l ectiv i ty al so increased with Sn02 content, and it appeared to reach constant value which depend on the type of catalyst. The selectivity was higher on CVD catalysts than IMP catalysts, even at the same activity level. CVD-5 had almost the same activity as that of IMP-ll.5, but the selectivity of CVD-5 was higher than I MP-ll.5. DISCUSSION Structure and Acid-Base Properties CVD method was app1 ied to prepare Sn02/Si02 catalyst for the oxidative dehydrogenation of ethyl benzene. In CVD method, SnC1 4 has been supposed to react with surface OH group to form surface chloride compound which is, in turn, hydrolyzed to form oxide layer. Repeating CVD cycles is supposed to resul t in the increase of thickness of Sn02 1ayer. Sn02 content increased almost linearly in the initial three CVD cycles, but it did not increase so much in the fourth and fifth cyc 1es. It s hou 1d be noted that the amount of Sn atoms deposited in the initial three CVD cycles, 2.72 nm- 2, was very close to the amount of Sn atoms for the monolayer coverage, 2.79 nm- 2, calculated from the diameter of SnC1 4 molecule and to a half of the amount of DH group, 5.38 nm- 2, calculated from the ignition loss of pre-desiccated Si0 2 support. This result strongly suggests that the surface OH group plays important role in the deposition of SnC1 4. As expected, SnC14 may react with the surface OH group. However, TEM and XRD gave striking results. High and uniform dispersion of Sn02 was obtained by the IMP method rather than the CVD method, as shown by the schematic models in Fig. 5 which summarizes the results of the characterization. In the CVD catalysts, Sn02 was deposited on only a part of primary particles of Si02 to form clusters, remaining some of Si02 primary particl es vacant. Sn compound deposited woul d aggregates in any of steps in catalyst preparation. NH40H treatment appeared to promote the aggregation, because XRD 1ine intensities were weaker in CVD-Sn02/Si02 prepared without NH 40H treatment. The acid properties, shown in Table 1, indicate that new acid sites are
120
CVD-l
y Si02
""",1ff!It!I!!!a
v
CVD-2
L-acid
A -v /
CVD-3
v Fig. 5
\ IMP
5chematic model of CVD and IMP catalysts.
formed by supporting 5n02 on 5i0 2. The strong interaction of 5n02 with 5i0 2 a 1so was observed in UV- vi s ib 1e spectra summari zed in Table 1. The absorption maximum shifted to lower wavelength in 5n02/5i02 catalysts, and the shift in CVD-l and IMP catalysts was 1 arger than those of CVD-3 and CVO-5. The rel ation between the aci dity of 1.5
> Ho > -5.6 with CVO eye 1es suggests
that the acid sites on CVD-l are different in nature from those on CVD-3 and CVO-5.
Thus, high acidities of CVO-l and IMP catalysts can be ascribed to
the strong interaction of 5n02 with 5i02 which was due to the high dispersion or thin oxide layer.
In the cases of CVO-3 and CVD-5, 5n02 particles are too
thick for the acid sites to be formed on 5n02 particles. The acid sites may be formed around the interface of 5n02 and 5i02'
The moderate acid strength
of CVO catalysts with high 5n02 content al so may be ascribed to the weak interaction or, in other words, to the large 5n02 particles. Activity and Acid-Base Properties The fol lowing reaction mechanism has been proposed for the present reaction in the previous paper (ref. 3). adsorbed on acid sites of 1.5
> Ho > -5.6,
Ethylbenzene is dissociatively and the catalytic activity per
acid sites of concern, i.e., turnover frequency, is determined by the base property of catalysts.
Figure 6 shows the effect of CVO cycles on the amount
of adsorbed ethylbenzene in the pulse reaction.
The trend is quite similar
to the acidity of concern shown by the closed marks, indicating that ethyl-
121
5n02 content 8
0
trl
;:l
s;
0
4f-
VI
-0 to
co
L.LJ
I
6 ....
15
17
120
I
0
5
2 n.
I
_----e
~
trl
"0
~
,.."
E
,
I
Fig.6 Effect of > -5.6 (b).
eVD
u
eVD
40
-c
-y
I
0
""'-
17
C/
i-
>,
u
'"
;:l
0CI.J
4
i-
I.L..
/
s,
> 2 0
!-
IMP
~--------
cf
'"s,
::l
i-
0
-.
0
s;
CI.J
0
0
1
Number of
3
eVD
4
5
(wt%)
d
b
CI.J ....,
17
.~
VI
-0
u
c(
0.4
<,
'"
CI.J trl
e-,
x
0
/C»-C5--
0.2
CI.J
>
----. IMP
...., u
,~
c(
2
5n02 content TO 5 15
0.6
I
a 6
3 1 2 4 5 Numver of eVD cycles
(wt%)
8 VI
0
cycle on ethyl benzene adsorbed (a) and acidity of 1.5
5n02 content 5 10 15
I
-- --
",,-- IMP
j0'v0
~ .~
-0
3 1 2 4 5 Number of eVD cycles
~
17
I
CVD
~
,
--.
15
•, •
, ,,
80
;:l
(0\0-9 Q
10
b
,..• ...IM;
-0 CI.J
.0
10
I
(wt%)
5n02 content
a
'<,
'0 E
5
(wt%)
0
cycles
0
4 5 Number of eVD cycles 1
2
3
Fig. 7 Effect of eVD cycle on turnover frequency (a) and amount of active oxygen per acid site (b). benzene is adsorbed on the acid sites of 1.5
> Ho > -5.6.
Figure 7 shows the initial turnover frequencies and the amount of active oxygen.
The latter was measured instead of the basicity, because the basici-
ty cou1 d not be determi ned due to obscure change of the co lor of i ndi cators.
122
The amount of oxygen was calcul ated from the total yields of styrene and carbon oxides obtained by repeating ethyl benzene pul se until any oxidation products could not be detected. The trends agreed well with turnover frequency, indicating that the lattice oxygen is responsible for the activity. Selectivity and Acid-Base Properties In the case of IMP catalysts, the rate of styrene formation increased with time on stream, and the reoxidation of used catalysts resulted in the same initial activity and the same time course of acti vity decl ine, as shown in Fig. 2. At the initial stage, styrene formed polymerizes to form coke, resul ting in the decrease of apparent rate of styrene formation. Coke thus formed blocks active sites for polymerization and suppresses the polymerization, which results in the apparent increase of the rate of styrene formation. It has been said that the coke deposit is oxidized to form CO and CO 2 (ref. 3 and 8). Thus, it may be expected that the suppresion of coke formation wi 11 increase the sel ecti vity. On CVD catalysts, coke deposition was not observed; the rate of styrene formation on CVD catalysts did not increase with time on stream, as shown in Fig. 3, and the color of the CVD catalysts remained white after the reaction. And, as expected, CVD catalysts had high selectivity, as shown in Fig. 4. Since styrene is polymerized on strong acid sites (ref. 3), the moderate acid strength of CVD catalysts may suppress the coke deposition. Thus, the absence of strong acid sites, which is due to the controlled dispersion of Sn02' resulted in high selectivity of CVD-Sn02/Si02 catalysts. ACKNOWLEDGMENT We would like to thank Dr. Masatoshi Yamada (JGC Co.) for TEM measurement. This work was partially supported by a Grant-in-Aid for Scientific Research (No. 59550546) from the Ministry of Education, Science and Culture, Japan. REFERENCES 1 Y. Murakami, K. Iwayama, H. Uchida, T. Hattori and T. Tagawa, App1. Cata1., 2 (1982) 67. 2 Y. Murakami, K. Iwayama, H. Uchida, T. Hattori and T. Tagawa, J. Cata1., 71 (1981) 257. 3 T.Tagawa,T.Hattori and Y. Murakami, J. Cata1., 75 (1982) 56, 66. T. Tagawa, K. Iwayama, T. Hattori and Y. Murakami, J. Catal., 79 (1983) 47. 4 T.Tagawa,S. Kataoka, T. Hattori and Y. Murakami, App1. Cata1., 4 (1982) 1. 5 T. Tagawa, S. Itoh, T. Hattori and Y. Murakami, Reac t. Ki net. Cata 1. Lett., 19 (1982) 135. 6 F. Sala and F. Trifro, J. Catal., 34 (1974) 68. 7 T. Hattori, H. Hanai and Y. Murakami, J. Catal., 56 (1979) 294. 8 G. Emig and °H. Hofmann, J. Catal., 84 (1983) 15. 9 R. Fiedorow, W. Przystajko, M. Sopa and I.G. Dalla Lana, J. Catal., 68 (1981) 33.
123
DISCUSSION ZHAO JIUSHENG : From the data shown in Table 1, it seems that you have measured the acidity by a titration method. Did you measure the acidity of the catalyst using TPD method and what was the result? T. HATTORI: The acidic properties of mixed Sn02-Si02 catalysts, prepared from SnC1 2 and Si(OC 2H5)4 ' were examined by TPD of NA~. Two desorption peaks were observed at 345K and 475K. After the reaction for several hours, the peak area of the high temperature peak was reduced at about one half of the initial value, and the peak temperature also decreased to 455K, indicating that the strong acid sites were blocked by carbonaceous materials but the medium acid sites remained. J.L. MARGITFALVI : 1) How can you explain such a significant difference in the coke formation between CVD and IMP catalysts? 2) How will the amount of surface OH groups on the silica influence the amount of tin introduced? 3) How can you avoid the surface reaction between the tin precursor (SnC14) and residual water? T. HATTORI: The difference in the coke formation is due to the difference in the acid properties. IMP catalysts contain stronger and more acid sites than CVD catalysts, as shown in Table 1. Styrene polymerizes on the strong acid sites, resulting in the significant coke formation on IMP catalysts. The controlled dispersion of SnO particles on CVD catalysts results in the weaker acid strength, which suppres~es the coke formation, leading to high selectivity. 2. It may be expected that the high concentration of surface OH groups will result to a large extent from the tin compound deposited, although the effect of the amount of surface OH groups has not been examined experimentally. 3. Silica gel was dried in a flow of dry nitrogen at 453K for 3 hours prior to the introduction of SnC1 4 vapor. It was confirmed by TGA that the adsorbed water can be removed by such a pretreatment; further weight loss was not observed in TGA after in-situ pretreatment under the above-mentioned condition. R. CAHEN : What is the conversion? What is the catalyst cycle or life time? T. HATTORI: 1) The conversion level in the present paper was 10 to 20%. As for the IMP catalyst prepared by using another silica gel as a support, the selectivity remained high even at a conversion of 60%, as reported previously (4) •
2) The activity of CVD catalyst decreased only in the initial several hours and further decrease in the activity was not observed, although prolonged life time test has not been conducted. M.V. TWIGG: What is the reactivity of conventional iron-based dehydrogenation catalyst under the conditions of your activity experiments? T. HATTORI: Iron-based catalysts gave significant conversion in the presence of oxygen, but they produced mainly carbon monoxide and carbon dioxide, resulting in low selectivity.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
125
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
THE PREPARATION AND CHARACTERIZATION OF VANADIA SUPPORTED RHODIUM CATALYSTS Y.-J. Lin l, R. J. Fenoglio 2, D. E. Resasco 2 and G. L. Haller l (1) YALE UNIVERSITY, Department of Chemical Engineering P.O.Box 2159, New Haven, CT 06520 (USA) (2) INTEMA (Institute of Materials Science and Technology) Universidad Nacional de Mar del Plata - CONICET Juan B. Justo 4302, Mar del Plata (ARGENI'INA) SlMiARY
The effects of the preparation variables on the extent of the metal-suwort interactions taking place in Rh/V.,o3 catalysts have been investigated. Under reduction conditions the catalysEs undergo morphology changes which may play a crucial role in the metal-oxide interaction. These morphological changes are strongly dependent on the catalyst preparation variables. The pH of the impregnating solution appears to have a more significant influence than prereduction of V~3 with regard to the resUlting interaction. Dissolution of the suwort followed by deposition over the rhodium chloride particles during the inpregnation step, as well as incorporation of rhodium chloride into a surface gel formed on the vanadia, may be responsible of a more intimate contact between the metal and the suwort.
INTROOUCTION
Metal catalysts supported on reducible oxides have been the subject of detailed studies during the last few years [11. The effects of high temperature reduction on the catalytic activity of these systems have been extensively investigated. It is widely accepted that in titania-suworted metal catalysts the metal particles become partially covered by suhoxide species migrating from the support during reduction at 5000C [21. In the case of vanadia-supported metal catalyst, very little is known about the nature of the metal-support interaction. Probably, the first workers who studied this interaction were Tauster and Fung [3], who reported losses in hydrogen and CO chemisorption capacities after reduction at 5000C in a series of reducible oxide supported catalysts. More recent works on vanadia-supported catalysts have reported increases in selectivity toward olefins for the CO hydrogenation reaction [41 and small binding-energy shifts to lower values with increasing reduction [5]. OUr own kinetic data for the n-butane reactions over Rh/V203 [61 have shown unusual behavior for these catalysts after BTR ,i. e., high selectivities toward isomerization and dehydrogenation, significant changes in activation energy,
126
etc., which were not exhibited by Rh/Ti0 2 catalysts under similar conditions. Furthermore, we have observed that the extent of alteration of kinetic parameters strongly depends on the preparation variables. It is possible that the initial interaction between the support and the precursor metal complex may play a role in the subsequent metal-support interaction taking place during the reduction step. Therefore, we have initiated an investigation of the effects of the preparation variables on the catalytic properties of Rh/V203 catalysts. In particular, we analyze here the effects of varying the thermal pre-treatment and the pH of the inpregnating solution using two different vanadia supports. The influence of these variables on the resulting catalysts have been characterized by hydrogen and carbon monoxide chemisorption. EXPERIMENTAL
To study the processes taking place in the aqueous solution during the inpregnation step we followed the evolution of pH after placing 0.5 g of vanadia support in contact with 30 em3 of a 6xlO~ M rhodium chloride solution in a well stirred vessel. The pH was continuously monitored by a digital Corning pH-meter. In order to determine whether a dissolution of the support takes place during the inpregnation procedure, we measured vanadium ion concentrations in solution as a function of contact time. For each measurement, aliquots of the solution were withdrawn from the stirred vessel and immediately centrifuged, filtered and analyzed. Quantitative analysis of vanadium in solution was performed by a spectrophotanetric method and by atanic absorption. To analyze the effects of the initial pH of the inpregnating solution and the thermal pre-treatment of the support on the catalytic properties we prepared three different sanples by an incipient wetness iIli'regnation method varying these two parameters. Two different VfJ3 supports were used. One was sUWlied by Alfa Products (10U03268l) and the other by Aldrich Chemical Co. (10U34l2LL). The rhodium content in every catalyst was 1.0% by weight. The volume of liquid used per gram of catalyst during the incipient wetness inpregnation was 0.75 em3 for the Alfa support and 1.3 em3 for the Aldrich support. Table 1 summarizes the main features of the catalysts investigated. As indicated, for two catalysts the support was treated in hydrogen at 5000c for 2h before the impregnation step. In one case, the support was impregnated as received. Mter the inpregnation, the catalysts were dried in air for 1-3 days at room temperature. Then, they were further dried in oven at 120 0C for 4h. Subsequently, they were reduced in hydrogen at 500°C for 2h, then oxidized in pure oxygen at 4000c and finally stored. Before any adsorption measurement, the catalysts were reduced "in situ· for 2h at either 2500c (LTR) or 5000c (HTR).
127
X-Ray diffraction measurements were performed in a GE XRD - 5 (Cu Kelt) diffractometer. Static H2 and CO uptakes were obtained in a conventional volumetric system. '!hermal programmed desorption (TPD) spectra were obtained in a gas chranatograph system using purified Ar as a carrier gas and a heating rate of lOoC/min.
TABLE 1 Characteristics of the catalysts investigated. CATALYST
SOURCE
OF
THE SUPPORT
U4.5 Pl.5 P4.5
Alfa Alfa Aldrich
PRETREA'IMENl' OF
THE
SUPPORT
none 500OC(H2) 2h 500OC(H2) 2h
INITIAL pH OF THE IMPREX;NATING SOLtJI'ION 4.5 1.5 4.5
RESULTS
SOlubility of the s\1PP.Qrt in the !!lP.r~ting solution We have observed that when the vanadia support is placed in contact with the rhodium impregnating solution an attack of the support takes place causing a partial dissolution of the vanadium oxide. Figure 1 shows vanadium concentrations in the rhodium solution (initial pH = 2.90) as a function of time after addition of un-reduced vtJ3 (curve u) and pre-reduced vtJ3 (curve p). Both curves are similar. Apparently, the hydrogen pre-treatment slightly enhances the dissolution of the support during the i.Ilpregnation step. We have observed that the solubility of the support in aqueous BCl solutions also increases as the PI of the solution is increased. This trend would suggest some acidity of the vanadilDll oxide support.
Evolution of pH during the impregnation SUbstantial differences between the pre-reduced and the un-reduced supports were observed in the pH evolution measurements. Figure 2 shows pH evolution curves obtained after adding the vanadia support to aqueous solutions. Curves u(4.5), u(2.9) and u(1.5) correspond to the addition of un-pretreated V203 to RhC1 3 solutions of the same concentration (6 x 10-4 M) but varying their initial PI values. '!hese should be capred to curves p(4.5), p(2.9) and p(l.S) obtained
128
after addition of pre-treated V203 to the respective RhC1 3 solutions. TWo clear trends are immediately obvious by inspection of these pH evolution curves. When the initial pH of the impregnating solution was below 3-4 the addition of the vanadia support caused an increase in pH. In contrast, when the initial pH value was higher than 3-4 it dropped after the support was added. Even though these trends were independent of the suppport treatment, the extent of pH change was significantly affected by the hydrogen pre-treatment. ~
100
Il>
-+'
0--
InItIal pH c
o
(pi
.b
-+'
2 -+'
50
o
E :::l
c
.sq , •
,...,
..ff
:::l
____
o
-
C1)
(j)
O .,..--.
.•.....•.•' . - -
-+'
If.'/' '"'
c
o o c o
/
= 2.9
o p14.51 t\ 0_0-0-0-01\ 0·0·0' u (2.91 • 0-
L -+'
~
"
1~.o·O·o
.O'O'O'O'6.~
..•........
af..
-.-.
2_._9_1_ _
0-
_._.-.-
u14.5J
4-
o
./
•........-.---.---.-
~.
u
"""0
o c o >
100 TIME Irn 1 nJ
Figure 1. solubility of the support in RhC13 solution. Vanadium concentration as a function of time. Un-reduced V-:P3 (u). Pre-reduced V-:P3 (P)
(1. 5)
o
200 TIME (sec)
Figure 2. pH evolution after the addition of V203 to the solution.
X-ray diffraction We have studied the morphological changes occurring on the vanadia support during the thermal treatment and the impregnation steps by X-ray diffraction. '!be two supports investigated exhibited essentially the same diffraction pattern. For the vanadia support supplied by Alfa, small differences in the relative intensities were observed after the hydrogen pre-treatment and after the inpregnation with the rhodium solution followed by reduction and oxidation. But in every case, the V203 corundum structure was maintained. However, when the Aldrich vanadia support was used, the X-ray diffraction pattern was consistent with the existence of a VtJS structure after the catalyst thermal treatments. In addition, the solid presented a metallic luster and when it was placed in acid it partially dissolved giving a deep blue solution. These characteristics would
129
indicate the presence of VO and V02 phases [7]. On the other hand, differences between catalysts PI.5 (Alfa support prereduced) and U4.5 (Alfa support un-reduced) were also evident. Then, even though the pre-treatment of the support does not cause noticeable structural changes, it may promote more profound changes during the subsequent catalyst preparation. The two vanadia samples (Alfa and Aldrich) had almost the same crystal structure. But, they presented substantial differences in surface areas and densities. Therefore, the drastic differences in structure observed after impregnation with the rhodium solution and the thermal treatment might be attributed to the differences in texture. No diffraction peak corresponding to Rh crystals was observed in any of the catalysts investigated. This ilK3icates that the Rh particles were not larger than 40-50 ft. .!'2 and 00 chemisorption we have investigated the effect of the initial pH of the impregnating solution and the hydrogen pre-treatment of the vanadia support on the chemisorptive properties of the various Rh/V 203 catalysts by thermal desorption and static adsorption techniques. The H2 adsorption isotherms obtained after both LTR and BTR for catalysts U4.5 and PI.5 are shown in Fig. 3 and 4 respectively. The hydrogen adsorption isotherms obtained for catalyst P4.5 are not shown. '!bey were very similar to those obtained for catalyst U4.5. In general, the three catalysts exhibited typical strong metal-support interaction (SMSI) behavior, i. e., significant losses in the H2 chemisorption capacities after HTR. However,
catalyst U4 ..;; ___ ••• 0
.J::. 0:::
LTR
0---
__ 0
__
.J::. 0:::
____ 0
<,
catalyst P1.5
o
____
<, ......... 0
u 0.2
-
8
L
0--LTR
0'---
o~
.i->' 0 _ _ 0 - •
0.2
............ 0
L
o
o
.J::. 0:::
.J::. 0:::
<,
<, I
-HTR
.:.::;:--
I
.. . -;,,~
• ___
•
..
I' ...
HTR
-----
-
. .-::::::::=: _e-
0.00
PRESSURE
(tor-r-)
200
0.0 0
200
PRESSURE
(tor-r-)
Figures 3 and 4. 82 (.) and CO (.) adsorption isotherms for catalysts U4.S and PI.S obtained after reduction at 2S0<>C (L'lR) and SOO<>C (m'R).
130
the metal-support interaction in catalyst Pl.S apgears to have different consequences than on the other two catalysts. The LTR isotherms obtained for catalysts U4.S and Pl.S are almost identical and indicate a similar percentage of Rh exposure on both catalysts. But, the HTR isotherms are clearly different. For catalyst U4.S the H2 uptake after HTR drops by more than an order of magnitude while the slope of the straight portion of the isotherm is about the same as that of the LTR isotherm. By way of contrast, catalyst Pl.S exhibits an increased slope after HTR. Accordingly, at high enough pressures, the H2 uptakes obtained after HTR were conparable to those obtained after LTR. A noticeable difference was also observed in the rate of hydrogen adsorption on the high temperature reduced catalyst Pl.S (HTR) compared to that on the low temperature reduced one, Pl.S(LTR). This is clearly evidenced by the hydrogen pressure evolution curves obtained during the adsorption and depicted in Fig. S.
1.0
\
-0
e
°
\
N
----1
0 E
-,
L
0
r::
'-'0.8 ,... ~
1... f/) f/) ~
1...
0c
~O. 0
0
•
\
:J
6
L
-0 c-,
:::I:
0
P1.5
°
•
-.
"~"~
LTR
--. --. --.--
0 __
"'-.
5 TIME l rn I n l
Figure S. Evolution of pressure during hydrogen adsorption on catalyst Pl.S after LTR and HTR.
For the Pl.S(LTR) catalyst about 40% of the equilibrium amount was taken up during the first minute, while for the same period of time only 10% of the equilibrium H2 uptake was reached on the Pl.S(HTR) catalyst. '!his slow hydrogen uptake was only observed for the Pl.S catalyst after HTR. Figures 3 and 4 also show the CO adsorption isotherms obtained on catalysts
131
U4.5 and PI.5 after both LTR and HTR. For the two catalysts the HTR caused a significant decay in the CO uptake being more pronounced on catalyst U4.5. No change in the slope of the isotherms nor in the rate of adsorption was observed for this case. In contrast to the very similar CO and H2 adsorption isotherms observed for catalysts Pl.5 and U4.5 after LTR, the TPD spectra of hydrogen evidenced clear differences even after LTR (see Fig. 6 and 7). For catalyst Pl.5(LTR) two desorption maxima are clearly evident. A large peak appears at 283°C and a smaller one at 128°C. After HTR, the first noticeable effect is a drastic decrease in the amount of hydrogen desorbed. In addition, new adsorption states appear to be formed. The peak at 283°C, which was dominant after LTR, still appears after HTR but it is almost as small as those appearing at 100°C, 340 0C and 400OC. These spectra would indicate that after HTR most of the metal surface is not available for adsorption while new sites with lower (IOOoC) and higher (400OC) adsorption heats are generated.
U4.5
x 1
~ 20
0
P1.5
x 1
~ 20
0
LTR
x
LTR
x
J
J
--<
--<
u,
u,
c
c
&10 0
&10 0
\...
\...
-0 », I
-0 s-, I
o0
500 Temper-otur-e (C)
TPD spectra of hydrogen Figure 6. fran catalyst U4.5 after LTR and HTR
o0
500 Temper-otur-e (Cl
Figure 7. TPD spectra of hydrogen fran catalyst Pl.5 after LTR and HTR
For catalyst U4.5 several desorption peaks appear after LTR. The main peak lies at the same temperature as that for catalyst PI.S (LTR), Le, 283°C. New maxima appear at 85°C, ISOoC and 222°C. The TPD spectrum obtained after HTR . shows a drastic decrease in the amount of hydrogen desorbed, similar to that observed for catalyst Pl.5. Table 2 summarizes the effects of HTR on H2 and CO uptakes on the catalyst
132
series. It is interesting to compare the relative decrease in H2 uptake after HTR as measured by TPD and volumetric adsorption respectively. For catalyst U4.5, rather good agreement between the TPD data and the static measurements was obtained. l\ccording to the TPD peak areas, after HTR the amount of H2 desorbing between 200 C and 500 0 C is about 30 times lower than after LTR. Likewise, the adsorption isotherms would indicate that the hydrogen uptake after HTR is about 20 times lower than after LTR. However, for catalyst Pl.5, while the TPD peak area after HTR is 17 times lower than that after LTR, the adsorption isotherms show only a drop of 60% in the hydrogen uptake after HTR. It would then appear that for this catalyst the losses in H2 chemisorption capacities after HTR can be detected by TPD but are not evident by static adsorption measurements. TABLE 2
Chemisorption capacities after low and high tenperature reductions. CATALYST H/Rh
U4.5 P1.5 P4.5
LTR
HTR
LTR
HTR
Hads
TP%
0.19 0.19 0.35
0.01 0.07 0.02
0.24 0.20
0.02 0.08
19 2.7 17
12 2.5
30 17
DISCUSSION we believe that the different effects observed on this catalyst series as a function of the initial pH of the iJlpregnating solution can be at least partially explained by the dissolution of the vanadium oxide taking place when the support is immersed in the rhodium chloride solution. This process may take different paths, depending on the pH of the solution. The pH evolutions shown by curves u(4.5), p(4.5) and p(1.5) would approximately describe the variation of pH in the impregnating solutions during the preparation of catalysts U4.5, P4.5 and Pl.5 respectively. Therefore, we might use those observed trends to help ourselves in the interpretation of the iJlpregnation process. Vanadium (III) oxide is basic and dissolves in acid solutions giving V(III) aqua ions [7]. However, when exposed to air it readily chemisorbs oxygen, causing an electron transfer from vanadium to oxygen, which can be easily detected by UPS [8]. The surface of V-P3 might then be considered to be covered by a few layers of V205' This oxide is more acidic and less soluble than V203 [7]. Thus, the solubility of the support in the acidic iJrpregnation solutions would be significantly lowered. Therefore, the concentration of vanadium in solution would be low. In this case, for pH values below 3.5 the main species present in soIution is vo2+ [9]. This species may result fran the reaction:
133
VtJS + 2 H+ ---> 2 V02+ + H20 (1) The net effect of this reaction would be the dissolution of the vanadium oxide accompanied by a H+ consumption, which would make the pH increase. This is in fact observed for curves u(l.S), p(l.S) and p(2.9). On the other hand, if the initial pH is above 3.S the dominant species is H2V04-, which can be formed by the follCMing dissolution reaction: VtJS + 3 Hi> -> 2 H2V04+ 2 H+ (2) In this reaction, protons are released making the pH of the solution decrease as the support is added. This is observed for curves u(4.S) and p(4.5). During the drying process following the impregnation step, the concentration of these species in solution will rise. Then, reaction with other ions and precipitation may occur. Therefore, vanadium species may be deposited over the rhodium precursors during this step [10]. Accordingly, we might especulate that V02+ and H2V04- will yield different eatp)unds in intimate contact with the Rh precursors. Thus, catalysts prepared at different pH will behave differently during the subsequent reduction. The low H/Rh ratios measured by hydrogen chemisorption after reduction at 2500C contrast with the small rhodium particle sizes indicated by the abscence of X-ray diffraction peaks. One possible explanation could be that the rhodium particles are partially covered by vanadium oxide even after L~ As explained above, this can occur by the process of vanadia dissolution followed by deposition over the metal precursor. However, this encapsulation may also result fran the insertion of rhodium into a gel structure formed on the vanadia surface when it is placed in contact with aqueous solutions [11]. Vanadium (III) oxide has the corundum structure. By reduction, oxygen vacancies can be easily created in the lattice. However, the corundum structure can be retained up to a concentration of oxygen vacancies of 10% [9]. This property of V203 might explain why the pre-reduced and un-reduced supports have similar X-ray diffraction patterns. These patterns do differ following catalyst preparation, suggesting a reacti vi ty of the oxide dependent on vacancy concentration. We can expect that the extent of surface reduction would have an effect on the extent of support dissolution, on the type of species released to the solution and subsequently on the reaction and re-deposition. However, the influence of the pre-reduction on the ultimate properties of the catalyst is secondary carpared to the effect of the pH of the inpregnating solution. The much greater losses in hydrogen and CO adsorption capacities observed in the static volumetric system after HTR for catalyst U4.5 would indicate that the extent of metal-support interaction is much greater on this catalyst than on Pl.5. However, the small amounts of hydrogen desorbing in the temperature range 20OC-500oC evidenced by TPO after HTR on catalyst PI.5 would indicate that most
134
of the hydrogen taken up dur Inq-che static measurements is not adsorbed on the rhodium surface. This surface hydrogen would be either too weakly or too strongly adsorbed to be detected by TPD. Keeping in mind the slow hydrogen adsorption rate and the increase in uptake with pressure observed for this catalyst after HTR, we propose a spillover process from the metal to the vanadium oxide followed by hydrogen bronze formation. A similar phenomenum has been previously reported to occur on pt/V205 catalysts [12]. Likewise, the relatively high co uptakes obtained on this catalyst after HTR might be due to adsorption on reduced vanadia species. At the present, we cannot differentiate whether those species are associated with the rhodium particles or with the bulk oxide. Our results demonstrate that dissolution and re-deposition of the support during impregnation can profoundly affect catalyst properties. We believe that the pH of the impregnating solution is a partiCUlarly important variable because it affects not only the solubility of the support but may also change the nature of the species re-deposited on the metal precursor. While these conclusions are specific to vanadium oxide and may have only a modest effect on some oxides, e. g. silica, they may generally apply for most oxide supports. Acm:mLEDGEMENl'S We gratefully acknowledge financial support from the National Science Foundation, USA (grant number INT-8504l84) and from CONICET, Argentina, (Leg.13974; Res.ll09/85). REFERENCES
exanple nMetal-Support and Metal-Additive Effects in catalysis n (ed. B. Imelik et al.) Stud.Surf.SCi.catal. 11, (1982) Elsevier Sci.Pub. 2 J. santos, J. Phillips and J. Dumesic, J.cata1.8l (1983) 147; D. Resasco and G. Haller, J.catal. 82 (1983) 279. 3 S. Tauster and S. Fung, J.catal., 54 (1978) 29. 4 E. Kikuchi, H. Nomura and M. MatslmlOto, Appl.catalysis 7 (1983) 1. 5 B. sexton, A. Hughes and K. Feger, J.catal. 77 (1982) 85. 6 Y.J. Lin, D. Resasco and G. Haller, in preparation. 7 F. COtton and G. Wilkinson in nJldvanced Inorganic Chemistry" 4th edn, J. Wiley, New York, 1978 8 R. Kurtz and V. Henrich, Phys. Rev. B 28 (1983) 6699. 9 M. Pqle and B. Dale, ().lart. Rev., 22 (1968) 527. 10 V. Ponec, personal cOlmlUJ1ication. 11 J. Legendre and J. Livage, J.C011.Interf.Sci. 94 (1983) 75. 12 J. Marcq, G. Poncelet and J. Fripiat, J.catal. 87 (1984) 339. 1
see for
135
DISCUSSION P.G. MENON: How sensitive is your catalyst in its SMSI state toward oxygen at ambient temperature? We have observed (P.G. Menon and G.F. Froment, Applied Cata1. 1 (1981) 31) that in the case of Pt/Ti0 2 catalysts subjected to a high temperature reduction, much of the so-called SMSI effect can be neutralized by a few 02-H2 cycles at room temperature. Many others have also reported the effect of 02-H cycles at 25°C in neutralizing the SMSI effect (suppression of and hydrogeno1ysis activity). Have you observed any such both H-chem,so~ption effect of oxygen at 25°C on your catalysts subjected to HTR? D.E. RESASCO : Compared to the TiO -supported catalysts, these vanadiasupported catalysts exhibit a rema~kab1y higher stability in their SMSI state. As you mention, the Ti02-supported catalysts could be easily reversed from their SMSI state, i.e. restore the original activity and chemisorption capacity by room temperature 02-H2 cycles or by the presence of traces of 02 or H20 in the feed. However, on these V203-supported catalysts the effects of the high temperature reduction cannot be reversed by mild oxidation treatments but only after oxidation at high temperatures (e.g., 400°C) followed by low temperature reduction. G.C. BOND : In the case of the SMSI effect brought about by the presence of titania, it is commonly believed that hydrogr?Ispillover is a necessary prerequisite, with the consequent formation of Ti species which migrate onto the metal. My question is : is hydrogen spillover also implicated in the SMSI effect with V203 and what is the oxidation state of the vanadium when it is on the metal particles? D.E. RESASCO : As discussed in the paper, depending on the preparation variables used, H2 spillover may be involved in the SMSI process. However, we cannot, at this time, call it a prerequisite nor can we conclude whether the vanadium species transported over the metal have oxidation states lower than III. G.R. LESTER: I believe that vanadium trioxide has a rather low melting point and might be expected to have a significant vapor pressure at the conditions corresponding to the high temperature reduction (S60°C in H2). Could this be the mechanism of the migration of some vanadium species observed here? D.E. RESASCO : Actually, the form of vanadium oxide which has a relatively low melting point is V20S (about 660°C). The melting point of V203 is rather high (about 1960°C) (R.J.H. Clark, in "Comprehensive Inorganic Chemlstry", Pergamon Press, N.Y., 1973). It is then interesting to note that the surface migration that we propose takes place under conditions at which less mobile species are present. L. GUCZI : Hydrogenolysis of n-butane is an excellent tool to study changes at the surface. Four-fold selectivities (multiple, statistical, middle bond and terminal splitting) as well as a change of the state of reaction vs. hydrogen pressure makes it possible to study the nature of metallic site as a function of SMSI as well as the depletion of surface hydrogen (a shift in rate maximum). I wonder to which extent were these parameters studied in your work, because the change in activity alone would not give too much information on SMSI effect; it simply indicates its existence. D.E. RESASCO : We agree that the study of these kinetic parameters provides important information. In fact, we have used them in the study of highly dispersed Rh/Ti0 2 catalysts (D.E. Resasco and G.L. Haller, J. Phys. Chem. 88, 4552, 1984). In that case, we observed that, as the reduction temperature-was increased, the change in selectivity toward n-butane isomerization and dehydrogenation was accompanied by a drastic change in the bond splitting
136
selectivity, i.e. the terminal/middle bond splitting ratio increased, suggesting a change in the reaction mechanism under SMSI conditions. In contrast, the V?O -supported catalysts showed no change in the terminal/middle bond splitting raiio after HTR, although they exhibit a significant enhancement in the selectivity toward isomerization and dehydrogenation.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV
137
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
PREPARATION OF HIGHLY SELECTIVE AND ABRASION-RESISTANT THICK SHELL CATALYSTS FOR HETEROGENEOUSLY CATALYZED EXOTHERMIC OXIDATION REACTIONS D. ARNTZ and G. PRESCHER DEGUSSA AG, Hanau (Federal Republic of Germanyl
SUMMARY Shell catalysts were prepared by coating a catalytically active material onto a superficially rough inert support in a specific coating device. Even for thick shells with up to 10 X of the weight related to the final catalyst, high abrasion resistance was achieved. Catalyst properties are highly dependent on production parameters such as moisture control during coating, mechanical energy input through drum rotation, and adjustments for differing thermal expansion coefficients of the active material and the support. Selectivity and kinetic stUdies of this catalyst show, that pore diffusion was suppressed by using expanding agents during the shell formation.
INTRODUCTION The rate of diffusion limitation in heterogeneous oxidation reactions is widely discussed in literature
[1
-
4). While some
oxidation catalyst usually need only a low content of active catalyst phase. i.e. V-P-TiOx-catalysts for phthalic acid or maleic anhydrid, most bismuth-molybdate catalysts are bUlk catalysts because of their lower activities. To avoid these diffusion limitations of such catalysts, in a large number of pUblications. various proposals were made to get a better bismuth molybdate catalyst [4 - 1]. Shell catalysts have promising properties: Pore diffusion rates are low. because of the short diffusion paths for gaseous reactants. Local overheating is avoided by the temperature equalizing effects of the support material and the adaption of catalytic activity to the technically achievable heat transfer. The pressure drop in the reactor can be reduced by minimizing variations in pellet diameters. The quantity of expensive catalytically active material can be minimized to the requirements of the reaction kinetic. In the production of thick shell catalysts - catalysts with a shell weight of more than 20 X of the total catalyst - the usual
138
conventional coating processes result in a poor abrasion resistance of the shell. not satisfactory for use in fixed bed reactors, and in a too large variation in the pellet diameters. To overcome these insufficiencies an improved process was developed.
EXPERIMENTAL The new process was developed by using the advantages offered by devices designed for filmcoating of pellets, particulary favorable is the Driacoater [9.10J. A flow sheet of the process is given in fig. 1.
The catalytic material in form of an oxidic prestage of the
catalytically active phase. prepared by standard methods, is coated onto an inert support core. In the first step a suspension of the powdered prestage is prepared by dispersing the powder in water. Sufficient water content of the suspension should allow a reliable continuous pumping and spraying operation. At the same time it needs to be minimized to reduce
the energy consumption for evaporation and to shorten the
run time of the coating step. A typical water content is 40 - 50 weight percent relative to the catalyst. Usually 2 -5 percent of binder. i.e. glucose. urea etc., is added to the suspension to improve the abrasion resistance of the final catalyst. For the preparation of thick shell catalysts an additional expanding material may be added to the suspension to increase the macroporosity of the final catalyst. The coating procedure begins by loading the coater drum with the inert support. For oxidation reactions the support should have a low surface area and an outer surface with a high degree of roughness to guarantee a firmely anchored shell to the support. ChemicallY the support consist i.e. of a-A1203, Mg-silicates or Al-silicates. Next the support is fluidized mechanically by rotating the drum and simultaneously loosened by drying air. The air is injected exclusively from the bottom of the fluidized bed through hollow ribs attached to the inside wall of the drum. The moisture-loaden exhaust air is drawn off above the material through the hollow receiving lug of the rotation axis of the drum. The suspension is sprayed onto the
139
catalyst precursor powder
binder support
"driacoater ..
shell
catalysts Fig. 1. Preparation of she11 cata1ysts f1uidized bed by two component nozz1es with a suspension pressure upstream of the nozz1es of about
1
to 3 bar generated by a
suspension pump. Therefore the suspension is being sprayed in counter f10w to the drying air. whi1e the f1uidized partic1e bed is moving across the f10w of both streams. During the coating process the f10w rate and the temperature of the drying air is norma11y he1d constant at about 15-30 Nm
3/h
per 1iter support and 80-100 0 e .
Sma11er air throughputs resu1t in distinctly slower drying rates, 1ess uniform f10w through the entire f1uidized bed due to bypass at the drum wa11 and in substantially 10nger preparation times. Higher air throughputs cause the suspension to dry out too rapid1y on its way from the nozzle to the charge surface, causing a discharge of dried prestage powder with the exhaust air. In order to achieve an abrasion resistant and rigid1y anchored she11 the precise contr01 of the she11 humidity during coating is necessary. The moisture of the she11 surface has to be he1d constant during the coating process. This can easily be
contr011ed by the
amount of suspension sprayed per unit of time. 80th the air temperature above the charge and the humidity of the exhaust air
permit a
sensitive contr01 of the drying process. They can both be used as the measured variab1es to contro1 the spraying process.
140
These process parameters also allow a fully automatic spraying with the help of a suitable control algorithm
that has been developed
for this application. This spraying operation is followed by a consolidation phase of 5 minutes while the drum continues to rotate. and a drying phase of 20 minutes without continuous drum rotation. After air-drying over night the catalyst is activated by conventional treatment at elevated temperature.
RESULTS AND DISCUSSION OF PREPARATION CONDITIONS The
developed coating process showed to be severly sensitive to
variations in the process parameters: - drying air: temperature, flow rate and moisture content - suspension: flow rate, moisture content - pan revolution and coating time - properties of the support and of the prestage of the active catalyst powder. Moistyre Control While the drying
air
conditions are held constant. the
process parameters during the coating process are controlled by the suspension flow rate. When the shell is too moist during coating. several particles agglomerate with one another causing the final diameter distribution to deteriorate. On the other hand, the desired anchoring of the shell to the support and the consolidation of the shell itself cannot be obtained,if the shell gets to dry. Typical temperature profiles of the gas temperature above the fluidized bed during preparation are shown in fig.
2.
Initially the fresh air
shows a sharp temperature drop. Later in the production cycle there is a slight temperature decrease and the moisture content of the effluent exhaust increases at the same time. While the suspension feed rate during the whole preparation periode stays constant. the moisture increase is caused by the removal of water from the inner shell. Therefore the moisture content at the shell surface is different from the average moisture content of the shell.
141
~. E
2
~
"0
::'" s
.\
:::>
45
i
27 29
1.0 1.0
7.0 ;0
35
2.0 l.O
106
fresh air
abrasion [W'/,]
2.1
---35--------35---------27_ 2" -=3P!9 1-._, '_
~
E .S!
I Vol'/, Hz 0 ]
37
~ .~
~
up. No.
w ·-27 ---17--29
a
af
..'" :::>
.c ""
,.,.------
~.as---
---
--29---'::::::---27=' ~,~
3,00
5:
37-
..------35-----.2 9 - ' - 7 -----
7
"il"------~
Vi
\..-.-27-'-'
~ ~
0
31:
10
20
40
30
-
50
time (min 1- - -
Fig.
2. Temperature and moisture profi1es during coating process
This average moisture content was determined at different times during the coating process by taking samp1es which then were dried unti1 their weights stayed constant. A s1ight decrease in the she11 moisture was observed. which is necessary to obtain high1y abrasion
resistant cata1yst she11s. The typica1 moisture content
of the shell for the used multicomponent bismuth-molybdate catalyst powder shows about 30 weight percent of the she11 materia1 short1y after beginning of the coating process and decreases to about 20 weight percent of the production cyc1e. Fig. 2 a1so shows the narrow 1imits. the air temperature profi1es above the f1uidized bed need to be contr011ed
during the
coating process. These profi1es are examp1es of measurements taken inside the deve10pmental
machine. the Driacoater 500. with a 10ad 3/min and 800C. They
of 6 kg support and a drying air stream of 2 Nm
are direct1y transferab1e to production machines up to 300 - 500 kg capacity for one charge.
142
t,\,
,\.\'
\.
'---
\.\.
.. .."
on
~~
.c x
4.0 Vol '10 HID
---__
~---~-~-
-.. . . . .
--+-----
----._-
~__
2.0Vol '10 HID
--..I.
45
0.5 Vol'" HID
10
.20
30
-
50
40
coating time [minI -
Fig. 3. Optimal coating temperature profiles at different H20 contents of drying air The pan temperature during the preparation of the two samples no. 27 and no. 29 differs less than 1 degree when the moisture content of the fresh drying air was held at 1 Vol percent H20. Nevertheless in experiment no. 27 the shell moisture was found to be too low
• caused by sufficient slurry flow rate during preparation.
resulting abrasion resistance in example no.
The
27 is very different
from that in example no. 29 with 7.0 vs. 2.0 weight percent*. On the other hand in trial no. 37 the shell
is found to be too moisty
during coating and also the abrasion resistance is insufficient compared to trial no. 35. This shows that the temperature profile needs to be adjusted, when the moisture content at the drying air varies. While the profile of no. 29 is optimal with respect to a moisture content of 1 percent of the drying air, it leads to a too.moisty shell for drying air having a moisture content of 2 Vol-percent (trial no. 371. *The abrasion resistance of the samples was tested after the final activation in the La Roche friabilator (20 rotations per minute, running time of 7 minutesl.
It determines the loss in weight of the
catalyst after the test relative to the weight of the fresh catalyst. For technical use the abrasion resistance value should be less than 3 weight percent.
143
As these examples show the exhaust air temperatures and moisture contents are very sensitive parameters in controlling the coating process. 80th can be used in the automatization of the coating process: The amount of supension sprayed is preferably adjusted to obtain the optimum temperature profile depending on the moisture content of the fresh drying air. Typical temperature profiles for 3 different drying air qualities are shown in fig.
3.
At higher
moisture contents smaller suspension feed rates could be utilized resulting in an increased coating time. For the preparation of a catalyst with 50 percent active material this time will increase from 50 minutes for 0.5 Vol X H20 in the drying air up to about 85 minutes for a 4 Vol X H20-content. Mechanical Energic Incyt If the preparation time exceeds specific values depending on powder properties and if the moisture content of the drying air is too high. the mechanical energy input may cause a too high of a consolidation of the shell. The shell forms very fine cracks. which enlarge with an increase of production time. This will cause a break of the shell during further processing of the catalyst and will lower its abrasion resistance. In batch 8 the coating duration was increased compared to batch 1 by reducing the drying air stream. This results in a lower abrasion resistance (tab. 1). The same occurs. if the spraying process is followed by too long of a consolidation period batch
with continued drum rotation (batch 8 compared to
C).
TA8LE 1 Mechanical energy inpyt dyring creparation Batch No.
coa ting Time (min)
Consolidation Time (minI
Abrasion Resistance (weight X)
A
65
5
1.3
B
105
5
7.6
C
65
15
4.5
144 Thermal Expansion Coefficient During the process development it was found that the thermal expansion coefficient of the powdered active material and the support material should be approximately equal.
If these coeffi-
cients differ by a larger amount (more than 10 - 20 7.l. the shell will crack in the final tempering step of the catalyst activation. These cracks substantially reduce the abrasion resistance of the final catalyst: If the thermal expansion coefficients are very different from each other, the shell
chips off in flakes.
The
increase in abrasion resistance with respect to the difference in expansion coefficients for three examples is given in table 2. TABLE 2 Influence of the coefficient of thermal expansion on abrasion resistance of final catalyst.
Linear Expansion Coefficient Batch No.
Calcin. Temp. [OC]
Spec. Surface [m 2/g]
(50 to 500°C) Catalyst Powder [10- 7/ oC]
Abrasion
Support £l0-7/o C]
Resistance [weight 7.]
0
330
52
n.m.
71
35.2
E
450
35
41
71
8.7
F
520
23
77.3
71
0.6
The catalyst powder used for the sample preparations was a multi component bismuth molybdate with the composition
M012
Bi, Fe,
Ni,0 Coo.! P, Ox which was prepared by standard methods [8]. The support was the SA 5218 of Norton Compo While suitable supports 7/ fall in the narrow range of about 70 - 90 X 10- oC, it was found that the thermal expansion coefficients of prestage powder could be adjusted by temperature treatments at 250 to 600°C. For a catalyst powder calcinated at a temperature lower than 350°C (batch 0) the linear expansion coefficient changes widely during heating from 500C up to 600°C. This coefficient could not be determined with standard equipment.
145
The catalyst shells prepared from such a powder cracked
during the
activation. As the expansion coefficient difference is decreased by changing the calcination temperature,
the abrasion resistance
increases sharply (table 2). Care must be taken that this treatment is being carried out not only for a specific temperature,
but for
the entire temperature range of the final activation operation. Expanding Agents A characteristic property of oxidation processes is pore diffusion limitation. which results in undesired combustion products such as carbon oxides.
In the oxidation of propene to acrolein for
example. it is observed that when with shell catalyst pore diffusion occurs, its influence increases with increasing shell thickness. The acrolein selectivity of a shell catalyst consisting of 50 percent active component, a common value for catalyst in industrial application. is about 4 percent less than one with 38 percent (fig. 5). Pore diffusion may be sharply reduced by incorporating an expanding agent into the shell during production and by removing it later during the activation step. These agents can be easily incorporated when they are added to the suspension and thus are coated together with the catalytic powder. For this methode they should nearly be insoluble in the dispersing agent and be removable from the formed shell by thermolysis or oxidation at temperatures below the activation temperature. Suitable materials are pentaerythritol, polymethacrylates, polystyrene. Fig. 5 shows a shell catalyst with 50 weight percent of active material prepared with 5 weight percent expanding pentaerythritol exhibiting thesame acrolein selectivities as a catalyst with 38 percent active material without expanding agent.
This result is explained by a re-
markable decrease of pore diffusion limitation by using expanding agents. CATALYST PROPERTIES particle Size pistribytion The progress in the production of shell catalysts by using the new process is shown also by the narrow particle size distribution obtained. This size distribution sharply affects the pressure drop in a commercially used narrow tube reactor for oxidation reactions. and so affects directly the energy costs of the processes.
146
without ,xpanding ogents
c J• ., '/, active component
.~
o sew'/, active co_pantnt ..Ih .."'iIlg
.,.ls
............04'----0
•
.50.'/, active cDmpon.nt /
~/----/,/"
'"
//
//
PE , prop•••
10
I.
rtaclor Itt.
10
IS
prop••• ·ca.vtrSlo. [ 10~~OI
Pf]--
Fig. 4. Diameter distribution of
Fig. 5. Pore diffusion limitation
shell catalysts prepared in
of shell catalysts for the
different devices
propene oxidation to acrolein
Figure 4 shows the diameter distribution of samples made in a Driacoater and in a conventional rotating drum, where the drying air is not passing through the fluidized bed. Whereas the average diameter of the particles is about 5.3 mm for both samples. the mean deviation of the conventional preparation is s driacoater preparation it is s
=
=
0.7 mm,
for the
0.3.
Kinetic Properties Detailed studies of the evaluation of the reaction kinetics were carried out [11]. The reaction scheme is depicted in fig. 6 and the obtained effective kinetic parameters are shown in table 3 for catalysts with 30 and 50 weight percent active material respectively. Both catalyst were prepared with pentaerythritol as an expanding agent. whereas all other preparation parameters were held constant.
147
~Acro~ein~
.t
Propene
~CO'
• AcrYlic Acid C02
Fig. 6. Reaction Scheme TABLE 3
Results for Effective Kinetic Parameters
I
I
Active Phase i f the Catalyst i
30 weight r. A.
50 weight r. A.
1
Kmole/m 3 .s.Pascal C[ j
Ei
nil
ni 2
)
n i j
3 (1: n ) Kmole/m .s.Pascal j ij
,
.i!Kmole
i
I
I
1
16.7 x 10- 6
2
1.3 x 10- 6
3
1.28x 10- 3
4
77 .1 x 10- 3
It was surprising
ni 3
1
6 30.4 x 106 2.26x 103 2.03x 103 272.5 x 10-
47.4xl0 6
0.44
0.93
42.8xl0 6
0
0.54
0.54
0
52.8x10 6
0.66
0
0
93.2x10 6
0
0
1
that for the three parallel reactions origina-
ting from propene the activation energies Eidid not change, whereas the preexponential factors Ai increased by a conversion factor which was nearly equivalent to the increase of active material in the catalyst. REFERENCES Satterfield, C.N., Mass. Transfer in Heterogeneous Catalysis, MIT Press Cambridge/Mass. (1970) [21 Smith. J.M. Chemical Engineering Kinetics, 3rd edition Mc Graw Hill (1981) [31 Eigenberger, G., Ruppel. W., Chem. Ing. Techn. 57 (1985) 181 [41 Krabets. R., Chem. Ing. Techn. 46 (1974) 1029 [51 SOHIO, DE-PS 23 51 151 /19721 [61 BASF. EP-15 565 /19791 [71 Nippon Shokubai Kaga Kogyo, DE-OS 33 00 04 (1983) [81 DEGUSSA DE-PS 20 49 583 119701, DOS 31 25 061 119811 [9 J DRIAM, DE-OS 28 05 801 119781 [10J Bohne, L., Paul, M., Pharm. Ind. 40 /19781, 12, 1366 [ 11J Arntz, D., Prescher, G. et al, ACS Symposium Series 196 1198213 [lJ
1~
DISCUSSION B. DELMON : Achieving the adequate porosity in the active phase deposited "shell" probably depends on the size and particle size distribution of the starting active phase powder. Ceramic science tells us that those parameters also influence mechanical (abrasion) resistance. As a whole, do you generally prefer a starting powder with a very narrow size distribution, or, inversely, a mixture of relatively large agglomeration (e.g. spherical) and smaller particle, which might act as a "cement"? D. ARNTZ : Because of a given powder production process as base of the process development described (and any problem involving this given particle size during the development), these dependences were not studied in details. We agree with your statement of the principle influence of particle size. The use of large particles in narrow distribution (without small particles) should give better porosity even without expanding agents. But we found indeed a lower abrasion resistance for such catalysts; therefore larger particle size should be preferred. L. HEINRICH : The optimum of temperature treatment was 520°C, for the Bi/Mosystem. Other catalytic systems should need other temperatures. Can you give us an idea of the chemical background of this temperature influence (the thermal expansion being only a secondary effect)? D. ARNTZ: Usually, the precursor of the catalytically active material is prepared by precipitation methods of salt mixture i.e. nitrates. These have to be transformed to the active oxidic stage. This transformation step can be done before or after the formulation step. So, there is a free variable, if the transfer will be partly conducted before the formulation step. It is correct that the terminal pretreatment is dependent on the active phase used and it has to be adjusted for every catalytic system. R. KRABETZ : Has the particle diameter of the powder an influence not only on the abrasion resistance but also on the selectivity of the catalyst? D. ARNTZ: In the range examined (10-l50~m particle diameter) no influence of particle size on abrasion resistance or selectivity was observed. N. PERNICONE : You have reported data obtained with ceramic spheres as supports. Do you think that similar behaviour is to be expected when one uses ring-shaped ceramic supports, as for instance in the Montedison catalyst for maleic anhydride? D. ARNTZ: This would be possible only for low load of active materials up to 5 to 10 weight %. At higher loads preferably the hole of the ring will be filled with the active material. J.B. MOFFAT: You indicated that for optimum abrasion resistance, the linear expansion coefficients of catalyst and support should be as similar as possible. Although this apepars intuitively reasonable, it is not clear from the oral presentation how such similarity was achieved. D. ARNTZ: As shown in Table 2, the adjustment of the terminal expansion coefficient was reached by a thermal pretreatment of the powder of the active material. A. HOLT: Is your method of preparation suitable for the production of 1 mm attrition resistant spheres of catalytic materials?
149
D. ARNTZ: Depending on the density of the support particles, 1 mm diameter is within or out of the range of our method. If the density is too low, the consolidation of the shell by mechanical energic input is insufficient. S. TAMHANKAR : My concern in the fixed-bed pellet reactor operation is the carryover of fine particles in the gas phase. Are your results on abrasionresistance applicable to such a case? D. ARNTZ: The process developed is characterized by the simultaneous consolidation of the shell by mechanical energy input and high drying rates at the same time. Without the mechanically induced consolidation, abrasion resistance necessary for technical use is not obtained; therefore the process has also a limitation for support particle smaller than 1-2 mm diameter. E. NEWSON : Is the purity of the inert support an important parameter; for example, have you observed impurities from the support diffusing into the active layer during calcination? D. ARNTZ: Used supported are checked for chemical inactivity under reaction conditions. No change of the active phase has been observed for the support components neither during calcination nor technical use of the catalyst.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
151
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
THE PREPARATION AND CHARACTERISATION OF IRON AND VANADIUM OXIDE MONOLAYER CATALYSTS ON Ti0 2 AND Zr0 2 J.G. VAN OMMEN, H. BOSCH, P.J. GELLINGS, J.R.H. ROSS Department of Chemical Technology, Twente University of Technology, P.O. Box 217, 7500 AE Enschede, The Netherlands
SUMMARY Iron oxide and vanadium oxide monolayer catalysts have been made by reacting toluene solutions of the corresponding acetylacetonates with the surface OH groups of Ti0 or Zr0 • The product of the reaction, acetylacetone, competes 2 2 with the acetonates for the adsorption sites. Of the four adsorption methods attempted, continuous adsorption and high-temperature adsorption gave approximately complete monolayers; the latter is the more promising method as it is more rapid and consumes fewer chemicals. TPR and CO-oxidation measurements for the Fe 203-Ti02 samples show that the monolayer is well-spread and that there is a sign~ficant interaction with the support. The iron cannot be reduced beyond the Fell state, possibly due the formation of a surface ilmenite. The Fe 0 monolayers are less active for CO 2 3 oxidation than is bulk Fe 03' Impregnat~on of Ti0 by Fe(N0 ) gives rise to 2 small crystallites of Fe 2 which have higher act~vities for3 eO oxidation than do the monolayer materials.3
6
INTRODUCTION Oxidation catalysts containing transition metal oxides play an important role in industrial processess. However, the choice of the catalyst for these processes is mostly empirical [1-2] as correlations between the structure of such catalysts and those properties which determine their activities and selectivities are not well established. In order to attempt to establish such correlations, we have prepared and studied well-dispersed oxides supported on various high-area oxidic materials using a general preparative method involving adsorption of Me(AcAc)n (AcAc = acetylacetone); as reported previously [3-4], this method probably gives monolayer spreading. In particular, we have described the preparation of catalysts containing Fe and V on various supports. 203 205 Monolayer materials have the advantage that the active metal oxide is more effectively used than if it is present as crystallites on the support, as is often the case when wet or dry impregnation are used. Moreover, the mechanical and thermal stabilities are also increased, larger surface areas (depending on the support used) can be obtained and the influence of the support is at its maximum. In previous publications, we have shown how the structure of the support in such catalysts influences their activity and selectivity for the oxidation of methanol, toluene or o-xylene [5]. The preparation technique used can
152
have great influence on the spreading of the oxide on the support. We have now studied in greater detail the method of preparation of the monolayer materials, paying particular attention to the way in which the Me(AcAc)n complexes are reacted with surface hydroxyl groups. The aim is thus to study the factors that influence the adsorption of Me(AcAc)n on oxidic supports in the preparation of monolayer catalysts and to determine the differences between monolayer and bulk oxides. In the present paper, we report on the preparation of iron and vanadium oxide materials supported on titanium and zirconium oxides, on the use of temperature programmed reduction to characterise the catalysts and on the use of CO oxidation as a test reaction. EXPERIMENTAL Materials 2 -1 2 -1 (80m g ) and Ti0 (68% anatase and 32% rutile; Sam g ) were 2 203 2g-1) obtained from Degussa, Zr0 (monoclinic; 11.Sm from Viking Chemicals, 2 and Fe(AcAc)3' VO(AcAc)2 and toluene from Merck (Analytical Grade).
Al
Methods Four methods are used to prepare the supported oxidic catalysts. (i) Continuous adsorption. A solution of M(AcAc)n in toluene is passed slowly (50 cm 3h- 1) through a bed of support particles. The adsorption is assumed to be complete when the concentrations at the outlet and inlet are the same. (ii) Batch adsorption. A solution of M(AcAc)n in toluene is added to the support particles and the slurry is stirred or shaken for a period between 20 and 500 h. (iii) Batch adsorption at high temperature. A solution of M(AcAc)n in of toluene is added to the support particles and is heated with stirring at 130°C for a period between 1 and 3 h. (iv) Wet impregnation. Water is slowly evaporated from a stirred slurry of the support particles suspended in an aqueous solution of Fe(N03)3' the temper0C ature being maintained at 95 (2 to 2.5 h). After the preparation, the resultant catalyst material is dried for 1h at 1000C and is then calcined in air for 2h at 450°C. The test reaction used here, CO oxidation, is carried out in a continuous flow reactor with a fixed bed of 0,5g catalyst (particle size 0,3-0,6mm); a gas flow of 50mlmin- 1 (1 atm., 20°C) consisting of 1 - 1,S% CO, 1 - 1,5% 02 and 97 98% He is used. The composition of the product gas mixture is measured by gas chromatography. TPR (temperature programmed reduction) experiments are performed with an apparatus described elsewhere [6] using a heating rate of l6°Cmin- 1 and a flow of 1 10mlmin- of 6% H2 in Ar. The amount of catalyst used was chosen so that the
153
amount of vanadium or iron in the reactor in every experiment was about 1.5 mg. Analysis of the Catalysts Metal contents of the catalysts were determined by X-ray fluorescence; BET surface areas and X-ray powder diffraction patterns were obtained as described previously. RESULTS AND DISCUSSION Catalyst Preparation Preliminary experiments showed that the Fe(AcAc)3 could be brought onto a Ti0
surface either from from the gas phase or from the liquid phase. Sublima2 t i on of Fe(AcAc)3 onto the Ti0 resulted in a catalyst which was visibly inhom2 ogeneous, consisting of light and dark grains; analysis of these showed them to have Fe contents of 1.07 and 2.7 wt.% Fe respectively. This method was therefore not used further. Adsorption of Fe(AcAc)3 from toluene on Ti0
was found to be a very slow 2 process as can be seen from the results of Fig. 1. Only after some 20 h was
02t= T: __~ ~ ~ ~ ~ . ,.
Nads mmolFe
o
10 - - hours
30
Fig. 1. Depe~qence on time of the amount of Fe(AcAc)3 adsorbed on Ti0 at 20 oC; 2 Co = 4 mmol 1 ,volume toluene = 300 ml, 1 g Ti0 Fe). (O.lmmol = 0.56 wt% 2• adsorption apparently almost complete. However, a series of experiments was carried out with adsorption times of 14 days and it was found that the extent of adsorption was now approximately 30% higher. Fig. 2 shows the dependence of the amount of Fe adsorbed on Ti0
as a function of the concentration of the impreg2 nating solution for adsorption times of 28 hand 14 days; in both cases, the influence of the concentration was found to be negligible above 1.5 mmol 1-1
It
is probable that the limiting uptakes shown in Figs. 1 and 2 correspond to approximately a monolayer of Fe(AcAc)3 adsorbed on the Ti0 If one takes the 2• 2 area occupied by one Fe(AcAc)3 molecule to be 50 A [7], a quantity of 0.16 mmole Fe(AcAc)3 (0.9 wt% Fe) is calculated as the monolayer capacity of the Ti0 2 2g-1) and this is the limiting value reached after 28 h adsorp- tion (area 50 m (Figs. 1 and 2). The level of adsorption achieved after 14 days must therefore correspond to the adsorption of more than a monolayer of Fe(AcAc)3' From
154 03 t" 14 days
-e.._------
t· 26 hrs
i
0·1
o
1
2 3 Ce mmol.lo'
_
) 0-+--.-o
4
Fig. 2. Dependence on concentration of the amount of Fe(AcAc)3 adsorbed on Ti0 2 for two different times at 20 0Cj volume of toluene = 300 ml, 1 g Ti0 2•
02
---r,--,'--,----T
··r·
1
2
-
3
5
ce (mmol/l)
Fig. 3. Dependence on concentration of amount of VO(AcAc)2 adsorbed at 200C on Ti0 2 for two different volumes and two diff:rent amounts of T~02; 2ooml, Zg T:OZ; ZOOml, 19 T10 x 100ml; Z; wt%V) 2g T10Z' (0.1 mmol V = 0.5
°
°
thermogravimetric experiments carried out after adsorption [3], it appears that the adsorbed Fe(AcAc)3 looses AcAc-groups by adsorption with time. The same type of behaviour as described above was found (see Fig. 3) for the adsorption of VO(AcAc)2 on the same Ti0 2, it also being found that the extent of adsorption depends strongly on the total volume of the solution from which adsorption occurs; the amount adsorbed is much higher if a large volume of solution is used relative to the weight of TiO We believe that this is a result Z' of an increased desorption of AcAc groups (as acetyl acetone HAcAc [3]) when the volume of the solution is higher. To show that the reaction product HAcAc competes with the metal complex for sites on the support surface, the results shown in Fig.4 for the adsorption of
0-47! continuous ads. equilibrium
continuous ads. no equilibrium
02 mmol Fe
,,
I
gr. Ti0 2
\
01
\ \
q, 0-_ -0-
_
Of----.,-----F:=.<=.:::.r==-==-==r-===="( 4
____ vol °/0 AcAc
Fig. 4. Dependence on concentration of HAcAc of the amount of Fe(AcAc)3 adsorbed on Ti0 at ZOoC; volume of 2 toluene = 100mI, 2g TiOZ'
155
Fe(AcAc)3 on Ti0
were obtained; different amounts of HAcAc were added to a 2 Fe(AcAc)3 adsorption solution in a constant volume (100 ml) of toluene to which
a constant weight (2g) of Ti0 2 was added. In order to keep the concentration of the HAcAc formed during the adsorption as low as possible, continuous adsorption experiments were carried out by leading a fresh solution of Me(AcAc)n over the support; this allows the product HAcAc to be removed together with the unused reactant. A typical example is given in Fig. 5, where the concentration of the effluent from the adsorption
r=..
••
c (mmel/I)
t 2
4
6
~t(hrs)
Fig. 5. Breakthrough curve for the continuous adsorption of Fe(AcAc)3 on Ti0 2: results for ~12.l wt% Fe catal: yst; Co = 24mmol 1 ,flow = 62 ml h l vessel is given as a function of time. In this case, the amount of Fe adsorbed on Ti0 was found to be 2.1 wt%. The results for continuous adsorption experi2 ments with both the iron and vanadium complexes on Ti0 and Zr0 are summarised 2 2 in Table 1 where they are compared with the results of comparable batch adsorption experiments for the Ti0
support and for the V complex on Zr0 a 2; 2 limited number of these results were also given in reference [3]. Also included
in the table are the percentages of the theoretical coverages achieved, these being based on figures of 2.8 wt% Fe or 2.5 wt% V for the maximum coverage calculated for a Ti02 support with an area of 50m2g-1 [3J; it should be noted that these theoretical monolayer coverages are based now on adsorbed oxide species, the assumption being that all the AcAc groups of the adsorbing complexes can be removed during the adsorption process. For the Ti0
support, it 2 is clear that a higher coverage is achieved in the continuous adsorption
experiments than in the batch experiments. This appears not to be the case for the Zr0 support. A surface concentration of 2.64 wt% Fe (96% monolayer) can be 2 reached on Ti0 by the continuous adsorption method after adsorption times 2 longer than 30 h. In both batch and continuous adsorption experiments, it takes a long time to
156
TABLE 1 Comparison of batch and continuous adsorption of Fe(AcAc)3 and VO(AcAc)2 on Ti0 2 2 -1 2 -1 0 (50 m g ) and Zr0 (11.5 m g ) at 20 C. 2 Support
Weight
Me(acac)n
Ig Ti0 2 Ti0 2 Ti0 2 Ti0 2 Zr0 2 Zr0 2 Zr0 2
*
1.0 2.0 2.0 5.1 2.5 1.0 1.9
Fe Fe V V Fe V V
CO
Time Immoll- 1 Ih 11.10 24.00 9.08 4.40 24.40 6.04 3.73
70 31* 72 11
1** 70 24
volume
%Me
%monolayer
Im1
~~~t+
~E~: ~~~t+
55 94 49 96 59 64 64
1.55 2.64 1.22 2.47 0.38 0.37 0.37
Presented in reference [3] erroneously with adsorption time 1 h. presented in reference [3]; gi~rn here for comp1etness. Continuous adsorption: flow 50-60 m1 h •
t* Already
TABLE 2 High temperature adsorption (130 oC) of Fe(AcAc)3 on Ti0 (50 m2g-1 ). 2 Adsorption method
Concentration
Batch* Batch Batch Continuous** Batch two*** Batch three*** 3B+C+
I gl -1 2.75 2.75 8.25 7.5 7.3 7.3 7.5
Fe I%wt 1.11 1.61 2.35 2.85 4.4 5.5 6.35
k k500 SBET 500 I m2g-1 110- 8msec-1 I 10-8 m3g-1 sec -1 45 45 42 43 39 32 35
oC; * Volume oC: = 80ml; T = 130 time = 1-3? ** T = 80 time = ISh; flow = 15ml h- • Repeated adsorption (e~ther two or three *** catalyst for 2h at 450 C.
+
1.5 1.3 1.2 1.6 1.2
63 56 47 51 35
times) after calcination of the
Batch adsorption conditions (three adsorptions) as aboye. then one continuous adsorption at BOoC for 10 h; flow = 15m! h- •
157
reach almost the theoretical monolayer coverages. In order to try to reduce the adsorption time, high temperature adsorption at 110 - 1300C was tried using 2 -1 both methods on Ti0 2 (area = 50 m g ) for both VO(AcAc)2 and Fe(AcAc)3; the results for the Fe complex are shown in Table 2. In both cases, monolayer or near-monolayer coverages were achieved within 3 h. Short adsorption times at low concentrations (2.75 g 1-1 is 1.07 mmol.l- 1) of Fe(acac)3 yielded submonolayer catalysts, while higher adsorption times at higher concentrations (8.25 g 1-1 is 32.23 mmol.l- 1) yielded reproducibly 2.35 wt% Fe on Ti0 (average 2 of three measurements), i.e. 84% of the theoretical monolayer. Continuous adsorption at 80
°c for
15 hours yielded 102 % of the theoretical monolayer. By
successive adsorption experiments each after calcination at 450
°c for
two
hours, the higher amounts given in Table 2 were obtained. From these results, it is clear that an extra amount is adsorbed after calcination but that this is much less than the first amount adsorbed, apparently because less Fe(AcAc)3 can be adsorbed on Fe
than on the fresh support. At the higher percentages, 203 the BET surface area also decreases, an observation which was not made for mono-
layer or submonolayer catalysts. No Fe
203
diffraction lines could be detected
for any of the catalysts examined. Catalyst Characterisation TPR measurements showed a striking difference between catalysts with 2.85 wt% Fe or higher compared with those of a lower Fe content, as is shown in the results of Fig. 6; at and above 2.85 wt% Fe, a second reduction peak is found, its magnitude increasing with the percentage iron. The temperature of reduction corresponding to the first peak also increases with Fe content (from 360 to 400 °C)
Fe (wt ...)
-- 6-05
I
N
100
300
~ T<°Cl
700
Fig. 6. TPR profiles for Fe on Ti0 2 203 for different wt% Fe.
158
TPR Fe
4
(wl%)
Fig. 7. Dependence of the H consumption during TPR for the data2 of Fig.7.
o
2
4
6
- F e (wtOfo)
while maximum of the second peak remains almost constant. The area of the first peak increases linearly with Fe content, as shown in Fig. 7, until the iron content is 2.5 wt% Fe, i.e. approximately the monolayer coverage for this Ti0 2 2g-l). Beyond the monolayer coverage, the size of the first sample (area 50 m peak increases less rapidly and the area of the second peak begins simultaneously to increase. The amounts of hydrogen consumed in the reduction process agree with the amount of iron in the catalyst only if we assume that the well-spread Fe 0 giving the first peak is reduced only to FeO while the Fe 0 giving 2 3 2 3 rise to the second peak is reduced fully to Fe; the lack of complete reduction in the first stage is probably caused by a strong interaction between TiO and
Z
the FeO formed on reduction. We therefore suggest that there are two Fe Z03 phases present in the calcined catalyst, the oxidic species corresponding to the monolayer and a second phase perhaps consisting of micro-crystallites. The TPR profiles of the reduction of samples of bulk Fe 0
Z 3
prepared from Fe(AcAc)3 and
Fe(N0 are shown in Fig. 8; these show a small peak at 385-395 0 C and a split 3)3 0 peak with maxima at 600 - 605 C and 654 or 718 °c respectively. From the consumption of HZ per gram of Fe
Z03, it is calculated that the first peak
Fig. 8. TPR profiles for unsupported Fe 0 3 prepared from Fe(AcAc)3 and 2 Fe tN03)3·
II II
\1
\1
\\
\\
100
300
500
-T(OC)
700
900
159
accounts for the reduction of Fe for the reduction of Fe
to Fe and the second and the third peaks Z03 304 to Fe, possibly in the two steps:
304 1/3 Fe -> FeO -> Fe 304 Thus, if the second peak in the TPR profile of catalysts with more than a monolayer (Fig. 6) corresponds with the reduction from Fe
to Fe, than part of 304 the first peak for these samples should belong to the reduction of Fe to Z03 Fe This is not, however, consistent with iron percentages determined by 304• X-ray fluorescence. We shall return to this problem in a later section.
CO Oxidation Experiments CO oxidation experiments with these catalysts given in Fig. 9 are consistent
4 k!lOO
(10:;')
t
Fig. 9. Dependence on wt% Fe of the rate constant for CO oxidation at sOOK.
2
o
o
o
t-rrr-
4 -
70
Fe(wt'l~
with the TPR results. Up to Z.8 wt% Fe, there is an lineair increase in activity (Give in legend: part of the results marked
° are from our earlier publication
(3». But at higher percentages, the activity stays constant as long as the lower total surface area is taken into account; if this is not done, then the Z activity decreases slightly. The activity per m of unsupported Fe is much 203 higher than that of the monolayer catalysts. It would thus appear that TiO has Z a negative influence on the activity of the monolayers of Fe in CO oxidation Z03 and one would expect that Fe crystals (i.e. more than a monolayer) on Ti0 2 203 would be less influenced. The catalysts shown in Table 3 were therefore prepared TABLE 3 Results for CO oxidation for FeZ0 3/TiO catalysts prepared by impregnation of Ti0 with e(N03)3' 2
t
Fe %wt 6.5 6.7 9.1 12.3
ZO.S Fe
203
SBET 2 -1 mg 44 45 43 38 42 10
k
sOO
1O-8msec-1 4.1 4.4 9.6 11.6 14.5 4.6
k
sOO
10-8m3g-1 sec-1 180
ZOO 410 440 610 47
160
by impregnation of TiO in Fe(N0 followed by calcination, X-ray diffraction Z 3)3 showed the presence of lines due to small crystals of Fe The activity in CO Z03, oxidation for catalysts with high percentages of these small crystals is larger than for pure Fe
Z03,
This argument also holds if the comparism is made per unit
of surface area. Model for the Catalysts This work shows that it is possible to prepare well spread oxides on TiOZ or ZrO by both batch and continuous adsorption of Fe(AcAc)3 or VO(AcAc)Z' Earlier Z measurements [3] demonstrated that other oxides such as Al Z03 or CeO Z can also be used as supports, The product of the reaction of the metal complex with the surface, acetylacetone (HAcAc) competes with the complex for the available adsorption sites (Fig, 4) and even at low concentrations inhibits the further adsorption of Me(AcAc)n' As a consequence, use of a continuous adsorption method gives rise to the most complete monolayer material, but the method has the draw-back that it is very time- and chemical-consuming. High temperature adsorption appears to favour the adsorption of the Me(AcAc)n and gives rise in a rather short time to monolayer or near monolayer coverages, All these results are consistent with the mechanism proposed earlier [3,4] for the adsorption of acetylacetonates on TiO According to Boehm [8,9], there are two hydroxyle Z' groups per TiO entity at the surface of Degussa P-Z5 and we suggest that these Z can react with two AcAc groups, binding the acetylacetonate to the surface:
......'OH Ti
+ Fe(acac)3 ~
'OH
/o\. Ti
te(acac) + Z Hacac
'0/
The process produces two molecules of HAcAc which can apparently also adsorb on inhibiting further adsorption of Me(AcAc)n' How the acetylacetone is TiO Z' bonded to the TiO surface is not yet clear; it may possibly occur by a coordinZ ative bond, thereby forming a surface titanium oxy-acetylacetonate, The TPR results (Figs, 6 and 7) strongly suggest that until the theoretical monolayer is reached, an iron oxide structure exists on the TiO is different from bulk Fe
surface which Z This structure can be reduced at lower tempera-
Z03, tures than can bulk (i,e, unsupported) Fe
but it can be reduced no further Z03 than FeO and this occurs in one step, This contrasts with the two-step reduction
of bulk Fe
which occurs via Fe • A similar difference in behaviour was Z03 304 found for V on TiO and unsupported VZOS [5,6,10], the reduction of the Z05 Z supported V stopping at V this was explained by suggesting that there Z05 Z03; exists a strong interaction between V and TiO which prevents further Z03 Z reduction [5], Such a strong interaction is also possible in the case of Fe Z03
161
on Ti0 2; the formation of FeTi0 (ilmenite) could take place on reduction [11]. 3 If such a well-defined structure formed, this would explain the lower reduction temperature observed and the fact that it is impossible to achieve further 2 reduction of the associated Fe +; the surface ilmenite is unlikely to be reduced at these temperatures. FeO is unstable below 570 easily be reduced [lZ].
°c if
unsupported, so it should
It has been shown for other supports such as A1
203 [13], MgO [14], or SiO [15] that part of the Fe cannot be reduced beyond the Z Z03 Fe 2+ state. The TPR profiles for samples with iron contents above 2.85 wt% are more complex, there now being an additional pair of poorly-resolved peaks (Fig. 6). It can be deduced that a second Fe 3+ phase, different from that discussed in the last paragraph, had been formed (see Fig. 7) and that this probably comprises of Fe
crystallites which are too small to be detected by XRD but which can be Z03 reduced to metallic Fe in one or two steps. A comparison of the data of Fig. 6
with those of Fig. 8 for the reduction of a-Fe Fe(AcAc)3 or Fe(N0
prepared by decomposition of Z03 shows that there is an extra peak for the bulk oxide at a
3)3 temperature of about 390 0C, probably corresponding to the reduction from Fe
203 to Fe304• We cannot fully exclude the possibility that this process also occurs for the crystallites present in the supported catalysts, as it would occur at the same temperature as the peak attributed to the reduction of the monolayer; however, calculations based on the hydrogen consumption and catalyst composition lead us to believe that this is not so. If such a process does occur, the resultant will still be small in comparison to the first reductions peaks of Fig. 6: the area will be approximately an eighth of the total of the highertemperature peaks if the reduction occurs first from Fe
Z03
to Fe
304
and
subsequently to Fe. Fig. 7 shows the behaviour of these peaks as a function of Fe content. If all the iron beyond a monolayer is present as Fe 203-like crystallites, line (1) would be expected to rise to a steady value and then stop. This is clearly not the case. We therefore conclude that a second layer of Fe 3+ species builds up gradually on top of the monolayer in addition to the additional Fe
crystalZ03 lites formed and that, as with the monolayer and because of the interaction with the support, this additional layer can only be reduced to Fe 2+; as the thickness
of the layer builds up, the interaction with the support decreases slightly, as evidenced by the movement of the lowest reduction peak to higher temperatures. It should be noted that this model was used in the construction of Fig. 7 from the data of Fig. 6: no other possible model was found to fit the data. The
co
oxidation results (Fig. 9) are, within experimental uncertainty, in
good agreement with this model. Until the theoretical monolayer of Fe
on Ti0 2 303 has been achieved, there is a linear increase of activity with wt% Fe. At higher
Fe contents, the activity per unit surface area is more or less constant. The
162
activity of a-Fe itself is much higher. Catalysts prepared by impregnation of Z03 TiOZ with Fe(N0 also show a higher activity than the monolayer materials and 3)3 also than pure Fe (Table 4). The relatively bad performence of the monolayer 203 catalysts in CO oxidation must be connected with the fact that it can only be reducted to FeO and the fact that a relatively inactive surface ilmenite structure may be formed. ACKNOWLEDGEMENT We wish to thank K. Hoving, B.J. Kip, J.A. Roos, E. Schasfoort and H.G.J. Lansink Rotgerink for performing part of the experimental work, H. Kruidhof for the X-ray fluorescence results and J. Boeijsma for the X-ray diffraction measurements, G.Altena for technical assistance, A.J. van Hengstum for helpful suggestions. REFERENCES 1 A. Bielanski and J. Haber, Catal. Rev. Sci. Eng. ,12 (1979) 1. 2 D.J. Hucknall, in Selective Oxidation of Hydrocarbons, Academic Press, London, (1974). 3 J.G. van Ommen, K. Hoving, H. Bosch, A.J. van Hengstum and P.J. Gellings, Z. Phys. Chern. N.F., 134 (1983) 99. van Ommen, H.Bosch and P.J. Gellings, Appl. Catal., 5 4 A.J. van Hengstum,~G. (1983) 207. 5 A.J. van Hengstum, J.G. van Ommen, H. Bosch and P.J. Gellings, in Proc. 8th. Int. Congr. Catal., Berlin, Verlag Chern. Weinheim (1984), Vol.4, p.Z97. 6 H. Bosch, B.J. Kip, J.G. van Ommen, P.J. Gellings, J. Chern. Soc., Faraday Trans. I, 80 (1984) 2479-2488. 7 R.B. Roof jr., Acta. Cryst., 9 (1956) 781 8 H.P. Boehm, Z. Anorg. Allg. Chern., 352 (1967) 781. 9 H.P. Boehm, Z. Anorg. Allg. Chern., 368 (1969) 156. 10 G.C. Bond, J.P. Zurita, S.Flamerz, P.J. Gellings, H. Bosch, J.G. van Ommen, B.J. Kip, Appl. Catal., 22 (1986) 361 11 P.J. Gellings, K.A. de Jonge, G.M.H. van de Velde, Chemistry and Industry, (1971) 1433. 12 A.J.H.M. Kock, Thesis, University of Utrecht, The Netherlands (1985). 13 R.L. Garter, D.F. Ollis, J. Catal., 35 (1974) 232. 14 M. Boudart, A.Delbouille, J.A. Dumesic, S. Khammonma, H. Topsoe, J. Catal., 37 (1975) 486. 15 M.C. Hobson, A.D. Cambell jr., J. Catal., 8 (1967) 294.
163
DISCUSSION G. CENTI : You indicate a parallel behaviour between V- and Fe-Ti monolayer catalysts. On V-Ti oxide, the V interacting with TiD? surface was suggested to be present in a different coordination situation (tetrahedral vs. octahedral). Have you any evidence (for example Mossbauer data) that a similar situation may be present also on Fe-Ti02 samples? J.G. van OMMEN: Our results with Fe203 monolayer catalysts indicate that the first layer of Fe?03 is in strong interaction with the TiO?, in the same way as we proposed with Bond for V20S mono1ayers. This is pos~ib1y connected with a different coordination situation. However, the difference in coordination was suggested for low coverages of V20S on A1 203; on Ti02, the situation could be different. From Mossbauer spectroscopy, we only have evidence that the Fe203 present in the monolayer is very well dispersed even after a long measurement time of about one week; no splitting of the iron signal has occurred. L. HEINRICH: The Degussa P2S is a very active type of Ti02. The interactions between this Ti02 and the iron and vanadium oxides should be observed by IR measurements : t~e characteristic IR absorption of the components will be disturbed or disappear. Did you observe an upper temperature limit, where the monolayer changes to agglomerates of iron or vanadium oxide? This effect must be observed by the appearance of typical IR bands of the support and catalytic components. J.G. van OMMEN: For the Fe203 monolayer, we did not observe an upper temperature limit of stability because we did not heat it higher than SOO°C. Until this temperature, the monolayer is For a V20S monolayer on Ti02 it was found that V20S dissolves as VIVtab1e. into the Ti02 lattice at temperatures above SSO°C. From IR measurements, we have until now no proof that either V205 or Fe203 on Ti02 is different from bulk V20S or Fe203; because of the deep colOur of these sampTes, IR spectra are very d,fficu1t fo measure. We can however observe that the OH groups of Ti02 disappear if a monolayer of Fe203 or V20S is adsorbed on this support. R.I. BICKLEY: Did you prepare your catalysts in the absence of daylight, or artificial light, since it is well known that organometallic reagents can easily be photocata1ytica11y degraded at room temperature? J.G. van OMMEN: We were aware of the photocatalytic degradation of Fe(AcAc)3 in toluene solution, especially for the long measurement times; the samples were therefore stored in the dark and a reference solution of Fe(AcAc)3 was also measured to check that such degradation did not occur in our case. L. GUCZI : As far as the reducibility is concerned. it is known from the literature that. at low concentration; the iron fills up the ~~ordinatively unsaturated sites and thus one cannot reduce it beyond Fe . This is in line with the TPR experiments. My question is how stable is the disordered iron oxide layer? Is there any evidence. e.g. XRD. for the stability of the noncrystall ine layer during CO oxidation? J.G. van OMMEN: As far as stability of the monolayer of Fe203 is concerned, we can only state that there is no decrease in activity durlng CO oxidation for at least one week after calcination at 4S0°C for 2 hours. Further calcination at 4S0°C for 17 hours did not decrease the activity of the catalyst.
1M
M.J. LEDOUX: We have found, on alumina, that the important points to obtain a monolayer (checked by EXAFS, see paper F6) were: 1/ Control of the OH concentration on the support before impregnation; 2/ Washing of the catalyst before calcination to extract the non-reacted AcAc complex; 3/ Impregnation at 120°C in order to displace the equilibrium by the evaporation of HAcAc (boiling point 118°C). Did you make the same observations ? J.G. van OMMEN: We observed the same phenomena as you mention. However, our high temperature adsorption was performed at 110°C; the increased adsorption of the AcAc complex can also be caused by differences in adsorption between the complex and HAcAc at this temperature, giving rise to displacement of HAcAc by the complex. The boiling point of HAcAc is 139°C (at 746 mm Hg). So, it could be that its evaporation also plays an important role, as you suggested. F. CAVANI : Have you experimental evidences, beyond the theoretical ones, based on geometrical considerations, about the formation of a complete monolayer of iron oxide over the TiO ? Do you think the "monolayer" is constituted by isolated unities of i~on oxide, or are these in some way ineracting with one another to form a uniform bidimensional structure? J.G. van OMMEN: We only have indirect evidence to support our suggestion: no XRD pattern or Laser Raman spectrum of Fe 0 ; the BET surface area remains the same after adsorption, and the amount ads6rtled during continuous adsorption stops at the theoretical amount. From the results presented in the paper, it can be seen that Fe?03 formed up until the theoretical monolayer percentage is reduced at a lower temperature than that above the monolayer and the monolayer and submonolayer material can only be reduced to FeO. We think that the Fe 0 monolayer forms a bidimensional structure on top of TiO in a way simila~ to that presented in our publication with Bond over V20S ~Onolayers. J.B. MOFFAT: You presented indirect evidence for the existence of a monolayer. Do you have any direct evidence? What experiments are you contemplating to generate such direct evidence ? J.G. van OMMEN: We have no· other evidence as that mentioned above for a wellspread Fe20 3 monolayer on TiO? We can possibly generate more direct evidence from EXAFs or ESCA measurements. E.V.W. GRITZ: 1/ The support has a strong influence on the success of Fe monolayer formation according to our own research. But is there evidence that the Fe monolayer survives the catalytic reaction? Did you characterize your spent catalyst with TPR ? 2/ You presented a model of active Fe surface formation for TiO -support. High metal loading on Zr02 was not possible according to your ex~erimental data. Is your model of monolayer formation only limited to the Ti02-support ? J.G. van OMMEN: 1/ We did not characterize our spent catalyst with TPR but as answered in relation to the question on the acitivity of the catalyst in CO oxidation, the catalyst did not change its activity during CO oxidation, nor at the start of the measurement nor after seven days of CO oxidation. 2/ We think that our model is not limited to Ti0 monolayers can also be obtained on A1 201 , for instance,. We do not know2;why the loading on Zr02 is not as good as on TiO. It is possibly caused by a lower surface coverage of OH groups. The forma~ion of the monolayer of Fe on Ti0 is caused by the surface reaction of Fe(AcAc)3 with the available 20O~ groups.2 If, for instance, TiO? is dehydrated and lost part of its OH groups, less Fe(AcAc)3 can be adsorbed.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV 1987 Elsevier Science Publishers B. V., Amsterdam - Printed in The Netherlands
165
©
INTERLAYER ACCESSIBILITY IN LAYERED DOUBLE-METAL HYDROXIDES
F.A.P. CAVALCANTI. A. SCHUTZ and P. BILOEN Chemical and Petroleum Engineering Department. University of Pittsburgh Pittsburgh. PA 15261. U.S.A.
ABSTRACT The accessibility of the inter1ayer space in layered double-metal hydroxides is being reflected by thei r sorption of nitrogen and hydrocarbons. The results i ndicab; that the i nter1 aYffr space 30f dehydrated materials intercalated with [Fe I (CN)6J4- and [Fe (CN)6J - anions is accessible to molecules as large as methy1cyc1ohexane whereas even nitrogen is exc1uded from the i nter1 ayer space in the materi ali nterca 1ated with N03anions. These findings confirm that their microporosity can be widely changed by varying the size and charge of the intercalating anions.
INTRODUCTI ON Layered double-metal hydroxides. hereafter called hydrota1cite. HT.-like compounds have the general formula: [AYH OJ [MIl MIl I (OH) JX+ I-x x 2 layer x/y· n 2 interlayer where MIl is a divalent metal cation such as Mg 2\ Ni 2+. etc. MIII is a tri valent metal cat i on such as A1 3+, Cr3+, etc, and AY- is an exchangeable NOj. [Fell(CN)6J4-, [Felll(CN)6J3-, etc. anion such as CO~-, HT-1ike materials are easily synthesized and can be generally obtained with X ranging from 0.15 to 0.35 (ref. 1). Their thermal stability varies with the nature of the metals, correspondi ng approximately to the dehydroxylation temperature of MI I(OH)2' with limits between 250 and 450°C. HT-1ike compounds have found use as ani on-exchangers (ref. 2) and they are catalytically active for aldol condensation after calcination at 400 0C (ref. 3). The intercalation of iso and heteropoly anions in HT-like compounds has been reported in the patent 1iterature (ref. 4). The use of these materials in the separation of the components of a mixture of H2 and air has
166
Another possible application of HT-like also been reported (ref. 5). The interlayer water can be removed at compounds is as desiccants. temperatures ranging from 120 to 250 0C leaving dehydrated products which show two-dimensional microporosity and which can sorb water in amounts from 10 to 25% of their weight. In this paper we report results concerning the sorption of nitrogen and hydrocarbons in several HT-like compounds. Our purpose is to study the accessibility of the interlayer space to molecules.
EXPERIMENTAL Materials The basic material, [M91-x Al x (OH)2] [(N03)x • n H20], was obtained by coprecipitation, at ambient temperature and at pH 8.5, of a solution of magnesium and aluminum nitrates in the appropriate ratio with NaOH. After washing, the material was subjected to a hydrothermal treatment at 210 0C for 18 hours, a procedure which leads to crystallites in the um size range. Materials where X = 0.20 and 0.25 were prepared and these correspond to anion exchange capacities of 2.8 ad 3.5 meq/g, respectively. Materials intercalated with the [FeII(CN)6]4- and [FeII I(CN)6]3- anions were obtained by anion exchange of the hydrothermally treated nitrate material This usually involved 2 or 3 until total exchange was accomplished. successive exchanges of 30 minutes each at 600C using the stoichiometric amount of the desired anion. Before the sorption experiments were carried out, the materials were subjected to a pretreatment consisting of heating under vacuum for 16 hours. This was done to remove the interlayer water and leave only anions in the interlayer space.
X-ray Powder Diffraction X-ray powder diffractograms were obtained on a GE XRD-700 unit using CuKa radiation and a graphite monochromator. Samples were prepared by the drying at 75 0C of a suspension of the desired material which had been deposited on a glass slide. This yields oriented specimens which show only the 001 reflections.
167
Sorption Isotherms Sorption isotherms were measured in a volumetric apparatus under static conditions. The measurements were carried out at 240C for the hydrocarbons sorption and at liquid nitrogen temperature for nitrogen sorption.
RESULTS AND DISCUSSION Structure The general formula of the dehydrated materials used in this study is gi ven by [M9 1- x Al x
(OH)2]~;yer
[A~iY]interlayer
where X = 0.25 and 0.20, and AY- = N03' [FeI I(CN)6]4- and [FeII I(CN)6]3-, and which from now on we will abbreviate as HTx-A. These materials crystallize with the C6 (CdI2) structure which consists of the hexagonal close-packing of OH- ions in two sheets between whi ch Mg 2+ and A13+ ions occupy a11 the octahedral sites. This unit bears a net positive charge which is balanced by the presence of ani ons the i nterl ayer space. Under appropri ate conditi ons, HT-like materials crystallize as hexagonal plates of approximately 1 and O.l~m diameter and thickness, respectively, as shown in Figure 1.
Fig. 1 TEM image of crystallites of HT O. 25-nitrate hydrothermally treated at 210°C for 18 hours.
168
The basal spacing of these materials varies with the size and orientation of the inter1ayer anions.
For HT o•25-nitrate, d003 =
s.s
A.
ions are oriented at an angle to the hydroxide layers (ref. 2). sma11 charge and therefore 1arge abundancy they occupy interlayer space.
The planar N0 3 Due to their
almost a 11 of the
In the case of HTO.25-ferro and HTO.25-ferri d003=1l.O A.
This corresponds to an interlayer spacing of 6.2 A which suggests that both anions are oriented with two of the octahedron faces parallel to the hydroxide layers (ref. 6). In order to estimate the dimensions and amount of void region in the interlayer space for the HTO.2b-ferro material we have built the model shown in Figure 2.
With x=O.25 it is possible to draw an ordered distribution of
the A1 3+ ions (ref.
7).
This enables us to build a hexagonal supercell with a = 12.24 A which contains four net positive charges from four A1 3+ ions (Figure 2A). One [Fe II(CN)6]4- anion is associated with this supercell.
Based on this model we obtain the following values for the void region in the inter1ayer space:
inter1ayer spacing- 6.2 A; minimum and maximum distance
between anions- 3.5 and 4.5
A. respectively.
LA)
L a
(8)
a Fig. 2 Schematic model of HTO 25-ferro: (B) Projection along the b-axl~. 3 eMg 2+; e -A1 +;o - OH- ; e - CN-
(A) Projection along the c-axis.
169
In the cases of x
= 0.20
or [FeIII(CN)6J3- it is not possible to build
simple models. However, when comparing HT O•20·ferro with HT O•2S-ferro one can expect an increase in all dimensions of the void region due to the lower layer charge of the former. When comparing HT O•2S-ferri with HTO.2~-ferro decrease is to be expected due to the lower anion charge of the former.
a
Surface Area The external
surface area of the materials studied can be calculated
using the crystallite size information obtained from TEM images such as shown in Figure 1. A theoretical value of 12±2 m2/g is obtained. Nitrogen sorption isotherms at nOK for HTo.2S-nitrate,-ferro and -ferri are shown in Figure 3.
All three samples were dehydrated under vacuum at lS0 0C for 16 hours.
Both the HT0.2S-ferro and -ferri materi a1s show Type I (Langmui r) isotherms whereas HTO.2S-nitrate sorbs only a relatively small amount. The measured surface areas are: HTO.2S-nitrate- 13.9 m2/g; HTO.2S-ferro- 246 m2/g; HT O•25ferri- 35S m2/g. Ni trogen sorpti on by HTO.25-nitrate is confi ned to the external surface of the crysta 11ites, as i ndi cated by the shape of the isotherm and by the agreement between the theoret i ca 1 and measured surface areas.
Thi s confi rms
the assertion that due to thei r low charge and to thei r geometric size and configuration,
the
NOj
ions
render
the
interlayer
space
essentially
inaccessible to molecules as small as nitrogen (with a kinetic diameter of about 3.6S A). The results for N2 sorption on HTO.2S-ferro and HTO.25-ferri were not expected. Since both the [Fe I I(CN)6]4- and [Fe I I I(CN)6]3- anions have similar sizes the voi d regi on is essent i ally determi ned by the dens i ty of ani ons in the inter1ayer space, which depends on their respective charges.
Therefore.
one should expect a larger interlayer void HT0.2S-ferri.
region in HT O•25-ferro than in The dimensi ons of the channels in the i nterl ayer space all ow
for its accessibility by nitrogen molecules and the final result should be a higher nitrogen sorption and higher surface area in HT O•2S-ferro than in HTO.2S-ferri •
The experi ments cons i sten ly showed the oppos i te to be true.
This led us to look into the details of the changes that occur during the synthesis and conditioning of these two materials. Both HT O•2S-ferro and HT O•2S-ferri are synthesized by anion-exchange from HT O•2S-nitrate. In an aqueous suspension the HT o•25-ferro solid is white, but it turns brownish-yellow upon drying in air at lOoC. HT O• 2S-ferri is
170
100
-r--------------------------.. . .
90
_.r- -------------------
80 70
60
50 40
,,+-+-t-+-t-+-+-+
30
20 10
.-.--._---.-.---.---.
0+----.----r------,;----~-----r---_1
o
0.2
0.4
0.6
RELATNE PRESSURE
Fig. 3 N sorption isotherms
o~
2 + -HTO• 2S-ferro; • -HTO. 2S-ferrl.
HT-like materials: _ -HTO. 2S-ferri;
brownish-yellow in suspension and remains the same color upon drying. Furthermore. upon vacuum dehydration between 100 and 150 0C HT O• 25-ferro turns dark green whereas HT O• 25-ferri retains its color. These changes in color seem to i ndicate that the or; gi na1 HT0.25-ferro exchange product undergoes changes when exposed to ai r and/or heat i ng. We have carried out some I.R. work on the materials dried at 700C which has yielded the following qualitative observations; (i) both HT O• 2S-ferri and solid K3[Fe II I(CN)6] show a sharp band at 2117 cm- I; (i 1) there is a sharp contrast between the l.R. spectra of HTO• 2S-ferro and solid K4[Fe l l(CN)6] in the region between 2120 and 2020 cm- I• which involves shifts in the relative intensity and in the position and the disapperance of some bands in HTO.25-ferro. and notably the appearance of a band centered at 2115 cm- I which is completely absent in solid K4[Fe l l (CN)6]; (iii) a broad band centered at 1375 cm- I• characteristic of carbonate. is present in HTO.2S-ferro. Based on the above observations we propose that the unexpected difference in nitrogen sorption and surface area between HTO.2S-ferro and HT o•25-ferri is
171
due to changes in the former upon exposure to air and/or heating. The color changes and I.R. results seem to indicate that for the HT O• 2S-ferro material. dried in air at 70oC. an oxidation process is taking place which transforms at 1east some of the i nterca 1ated ferro ani ons to ferri. Since charge balance between the positive hydroxide layers and the anions must be kept some CO 2 from ai r is taken up by the oxidi zed HTO.2S-ferro materiaI Ieadi ng to the formation of carbonate anions in the interlayer space as indicated by the I.R. results. Contamination by C02 from air leading to carbonate is a very common problem with HT-like materials (ref. 8) and this additional number of anions in the i nterl ayer space of the oxidi zed HTO.2S-ferro materiaI cauld account for its lower nitrogen sorption and surface area compared to HT o•2S-ferri.
Hydrocarbons Sorption In order to evaluate the accessibility of the interlayer space to larger molecules we have carried out the sorption of various hydrocarbons in HT O• 2Snitrate. -ferro. and -ferri, and in HT O• 20-ferro. Figure 4 shows the sorption 110 100 90
~ E g
80
1XI
60
It:
_/
/
70
~
50
!Z :::I
40
~
30
0
+-+----- +---+
+,...-+-+-
0
III
--
----- -------------
20
.-+-+-----------+
10
o
0.2
0.4
0.6
RELATIVE PRESSURE
Fig. 4 i-CsH sorption isotherms (24°C) on HT-like materials (dehydrated at IS00C for 1I)2 hours):. -HTO. 2S-ferri; + -HTO. 2S-ferro; • -HTO. 2S-nitrate. of isopentane, i-CSH12, in HTO.2S-nitrate. -ferro and -ferri. The results obtained are consistent with the surface area measurements for these materials and indicate that the sorption of i-CSHI2 is confined to the external surface
172 of HT O•2S-nitrate whereas the interlayer space in HTO.2S-ferro and HT O.2Sferri is accessible to this hydrocarbon. Once again we observe the unexpected result of higher sorption of i-CSHI2 in HTO.2S-ferri than in HT O•2S-ferro. This provides further evidence to the assertion that the HT O•2S-ferro material is undergoing changes, possibly by oxidation of the anion, which result in a smaller void region in its interlayer space. Figure isopentane,
S
shows
i-C SH I2'
HT O•2S-ferro.
the
results
cyclohexane.
for
the
sorption
of
pentane,
CSH I 2• C6H12' and methyl-cyclohexane, C7HlS' in
Within the accuracy of our measurements this material sorbs
identical amounts of CSH 12 and i-CSHI2 indicating that the difference in their structures is not large enough to cause the exclusion of i-CSHI2 from the interlayer space. small.
The sorption of C6H12 is lower, whereas C7H12 sorption is whereas CSH12 and i-C SH I 2 sorption follow a Type I
Furthermore,
isotherm and is almost instantaneous, this is not the case for C6H 12 and C7H1S• The latter two hydrocarbons showed pronounced kinetic effects during their sorption as indicated by a very slow approach to equilibrium. It seems that both C6H12 and C7HlS can still penetrate in the interlayer space of HT O•2S-ferro but only with considerable difficulty, especially in the case of
70
---+-+ +---_..--. .-
60
+-+
Cl
<, Cl
50
E
fj III It: 0
+-/ +................
-.-
---.
_. •
40
fIl
~
•
30
~
::J 0
~
20
10
0
v·
o
0.2
»->:': 0.<4-
0.6
0.8
RELATIVE PRESSURE
Fig. S Hydrocarbons sorption isotherms (24 0C) on HT O.2S-ferro (dehydrated at 150 0C fat' 16 hours) : +-i-C SH I 2; - -C SH I 2; • -C 6H I 2; ... -C 7H I 5.
173
C7H1S' It is noteworthy to point out that C6H12 has a kinetic diameter of 6.2 A which is approximately the size of the interlayer distance in HT O• 2S-ferro. A possible way of tuning the sorption properties of HT-like materials is by changing the charge density of the hydroxide layers by varying the amount of A1 3+ cations in them, i.e., X. A lower X results in less charge balancing anions and therefore in a larger void region in the interlayer space. This trend is shown in Figure 6 where the results for the sorption of i-CSH12 in HTO.20 and HTO.2S-ferro are shown. As expected, HT O• 20-ferro was able to sorb
6 0 . , . . . - - - - - - - - - - - - - - - - - - - - - - - -......
------- ----+----+ +-+ --+ --+ ~--_.-
50
.... 40 -,,~
30
+'
~
_.------,
+.....
20
10 -
O+--'"T"'---r----r-----,r---r---.,.....---r---~
o
0.2
0.4
0.6
0.8
RELATIVE PRESSURE 0C)
Fig. 6 i-CsH 12 sorption isotherms (24 on HT-like materials (dehydrated at 100 0C for 15 hours):. -HTO. 20-ferro; + -HT O. 2S-ferro. more i-CSH12 than HTO.2S-ferro. This result is further substantiated by the measured surface areas of 224 and 177 m2/g, respectively.
CONCLUSIONS (i) The interlayer space of HT-like material is accessible to organic Specifically, for the case of HT O• 2S-ferro there is no molecules. molecular sieving effect between CSH12 and i-CSH12• Both C6H12 and C7H15 can still access, although with difficulty, the interlayer space.
174
(ii)
The nature of the interlayer anion determines whether or not interlayer sorption will occur. Specifically, sorption of both N2 and i-CSH12 is confined to the external surface in HTO• 25-nitrate whereas it takes place in the interlayer space of HT O• 25-ferro and HT O• 25-ferri.
(iii) It is possible to tune the sorption capacity of HT-like materials by varying the charge density of the hydroxide layers. (iv)
The material HTO.25-ferro seems to undergo changes (possibly involving the oxidation of some of the ferro anions to ferri with simultaneous uptake of CO 2 from air leading to carbonate anions for charge compensation) which result in a lower void region in the interlayer space than predicted from a theoretical structural model.
ACKNOWLEDGMENTS The authors thank the U.S. Department of Energy, Office of Fossil Energy Program (Contract DE-AC22-84PC70020) and the Exxon Education Foundation for financial support of this work.
REFERENCES 1 2 3 4 5 6 7 8
W. Feitknecht, and M. Gerber, Helvetica Chim. Acta, 25 (1942) 131-137. S. Miyata, Clays and Clay Minerals, 31 (1983) 30~-311. W.R. Reichle, J. Catal., 94 (1985) 547-557. G.M. Wo 1termann, U.S. Patent 4, 454, 244 (to Ashl and Oil, Inc., June 12, 1984). S. Miyata and T. Kumura, Chem. Lett., (1973) 843-848. S. Kikkawa and M. Klizumi, Mat. Res. Bull., 17 (1982) 191-198. G.W. Brindley and S. Kikkawa, Am. Mineral., 64 (1979) 836-843. G. Mascolo and O. Marino, Mineral Mag., 43 (1980) 619-621.
B. Delman, P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
175
MOLYBDENUM CATALYZED OXYGEN-TRANSFER REACTIONS.HETEROGENIZATION OF HOMOGENEOUS CATALYSTS BY USING NEW DIHYDROXYBORYL-SUBSTITUTED RESINS
E.TEMPESTI,L.GIUFFRE' ,C.MAZZOCCHIA,F.DI RENZO and P.GRONCHI politecnico di Milano,Dipartimento di Chimica Industriale e Ingegneria Chimica,P.L.daVinci 32 - 20133 Milano, Italia
ABSTRACT Conventional molybdenum catalysts have been heterogenized by using new dihydroxyboryl-substituted resins. The specific activity of the grafted catalyst increases with increasing protonation of the grafted complex. The optimum molar ratio of molybdenum to boronic groups of the support is equal to one. No degradation of the catalyst is observed. INTRODUCTION Propylene oxide is commercially made by the Halcon (ref. 1) process in which isobutane or ethylbenzene is first oxidized to the corresponding hydroperoxide which then reacts with propylene yielding propylene oxide and an alcohol byproduct. The catalyst is a molybdenum complex such as a naphthenate or a stearate which is soluble in hydrocarbons although a major disadvantage is represen-ted by its irreversible loss at the end of the reaction. Many attempts have been made in order to heterogenize typical molybdenum catalysts but the results obtained are very often complicated by catalyst degradation which is more or less evident. More specifically it seems that when a chemical combination of molybdenum (VI) with the support is achieved through definite acid functions such as carboxylic (ref. 2) or phosphoric (ref. 3) acid groups no substantial degradation of the catalyst is found. In other istances when this combination does not occur (ref. 4) molybdenum-leaching phenomena become evident on a macroscopic scale. In order to test this assumption we have decided to heterogenize conventional molybdenum catalysts by using a suitably functionali-zed polymeric support such as the following dihydroxyboryl-substituted resin which is commercially available (Aldrich) and which has already been used for column chromatographic separation of mixtures of sugars and nucleic acids which form complexes of
176
varying stability
yH 3
-CH -C2 I
O=C-NH~
~B(OH)2 n with boronic acid groups. METHODOLOGICAL APPROACH The classical impregnation method followed by calcination for catalyst activation has been employed extensively in the prepara-tion of supported catalysts. For a molybdenum-impregnated catalyst inorganic acids or their salts like ammonium para-molybdate in an aqueous solution are generally used. Oxymolybde-nic species in aqueous solution are in equilibrium between 2n- depending upon the pH, temperature and (M00 4) and (MOxO y) concentration of the solution (ref. 5). Consequently a non-uniform distribution of active sites and heterogeneous properties of the catalyst may ensue. As a result the mechanism of the surface reaction may become ambiguous and conflicting results may sometimes be drawn. Rather than by impregnation we have found that it is possible to fix the molybdenum by adequate interaction with the surface boronic acid moieties. This interaction has been postulated to occur on the basis of an acid-base interaction accompanied by elimination of water as follows:
o
}~-OH OH
+
II
HO-Mo-OH II
o
}
~
B-O-Mo-OH I
OH
+
II
0
As a result the fixed (ref. 6) catalyst has been synthesized molecularly by taking advantage of the ready reaction between molybdic acid and the surface boronic acid groups. By governing the Mo fixation on the polymeric resin through the population of B(OH)2 groups at the support surface it is evident that the distribution of active sites can be controlled by varying the number and topography of the surface acid groups.
177 RESULTS AND DISCUSSION We have found that the optimum conditions for the fixation of molybdenum (VI) by the dihydroxyborylsubstituted resin lay in the pH region between 1.0 and 3.0. In a more acidic medium the fixation of molybdic acid is prevented by protonation of the boronic acid groups of the support while in a more alkaline medium the fraction of anionic oxyrnolybdenic species increases and the interaction with the surface acid groups may occur as follows:
1-~-OH
+ (anionic oxyrnolybdenic
speCieS)~}BorOmOlYbdeniC
acid + nOH-
OH In the case of grafted boronic groups within the cited pH region the condition of electroneutrality of the surface induces us to assume that the grafted acid group must be either fully protonated or present in salt form:
r
o11"1-OH 2 B-O-Mo-OH I
OH
II ~
0 OH
2
t
o
OH
11'1-
~
B-O-Mo-O Na I
OH
+
II"
0 OH
2
On the basis of the results obtained in the epoxidation of cyclohexene with t-butylhydroperoxide catalyzed by different samples of dispersed molybdenum-containing dihydroxyboryl-substituted resins it is possible to discriminate between these two possibilities. In fact in terms of selectivities (S, moles of epoxide/ moles of consumed hydroperoxide) and yields (Y, moles of epoxide/ moles of initial hydroperoxidel it has been found that the catalytic activity observed with different catalysts prepared by treating the resin with varying solutions of Na
2Mo04 at different pH increases with decreasing the pH of the solution
used for fixing molybdenum (e.g., at pH=1.0 S=85% and Y=56% after one hour of reaction). Thus it may be safely assumed that the protonated form has a higher catalytic activity in epoxidation than the sodium form. It has also been found that the specific catalytic activity of the catalysts grafted at constant pH reaches a maximum when the molybdenum content on the resin (evaluated in g.atom/g by atomic
178
absorption through a preliminary chemical attack and solubiliza-tion of the grafted sample) equals the number of boronic acid groups of the support (expressed in g.atom of boron/g). When the ratio Mo/B is further increased by increasing at constant pH the amount of Na
initially present in the solution, 2Mo04 the specific catalytic activity of the grafted catalyst remains
constant or decreases. This can be due to the formation of grafted complexes containing two or more molybdenum atoms for every boronic acid group. In the case of repeated use of the same weight of catalyst prepared under optimum experimental conditions at controlled pH the catalytic activity observed remains practically constant. Indeed with the more active catalysts the relative deviations of activity are less than 5% while with the less active catalysts the deviations do not exceed 15%. The observed effects may thus be attributed to the grafted catalyst and not to minor amounts of solubilized molybdenum. The substantial absence of leaching phenomena is confirmed by the very low conversions of hydro-peroxide observed in the case of repeated heating of spent solutions after separation from the grafted catalyst. HETEROGENEOUS versus HOMOGENEOUS CATALYSIS From a synthetic point of view and relative to new covalent boron (III)-molybdenum(VI) mixed oxo derivatives such as
o
@:C " I
0
0
II B-O-Mo-X
0""'"
o"
which have already been tested as homogeneous catalysts for oxygen-transfer reactions to olefins (ref. 7), the following distinctive features have been observed: 1- when boronic rather than borinic acid group moieties are involved in the formation of the B-O-Mo chemical bonding the preparation requires more drastic experimental conditions. This may be due to the change of hybridization of boron which is implied in the formation of a five-membered ring boron compound since it is known that this effect is mainly responsible for the enhanced acidity which allows for the direct titration in
179 10) aqueous solutions of the weak acid (boric acid: K a=5.8X10when 1,2-diols (which readily form dioxaborolane rings) are present. 2- by comparing the characteristic Mo0 infrared absorptions 2 observed with both type of catalysts as prepared, it is evident that with the heterogenized catalysts the structure and environment of the transition element affects to a lesser degree the electronic properties or degree of covalency (ref. 8) of the Mo-O bond as the active site prior to oxygen-transfer step. 3- with the heterogeneous catalysts obtained by Mo-fixation and relative to the homogeneous catalysts reported above no chemical degradation of the catalyst is observed. These differences of chemical stability necessarily imply that the electronic requirements for hydroperoxide activation through its coordination to the metal centre and subsequent oxygen transfer to the olefin are less affected in the former case by the proximity of a stable B-O covalent bond. These differences should also be dealt with satisfactorily in terms of mechanism of oxygen transfer.
REFERENCES
2 3 4 5 6
7
8
U.S.Pat. 3350422 (1967); U.S.Pat. 3351635 (1967) both to J.Kollar. S.Ivanov,R.Boeva and S.Tanielyan, J.Cat.,56 (1979) 150. A.P.Filippov and O.A.Polishchuk, Kinetika i Kataliz, 25(6) (1984) 1341. M.B.Ward,K.Mizuno andJ.H.Lunsford, J.Mol.Cat.,27 (1984) 1 Y.lwasawa,Y.Nakano and S.Ogasawara, J.Chem.Soc.,Faraday Trans., 74 (1978) 2968. The term "fixed catalyst" is proposed to indicate a well-defined supported catalyst with chemical bonding between the transition element and the support in order to distinguish this type of catalyst from conventionally impregnated catalysts. E.Tempesti,L.Giuffre,C.Mazzocchia and F.Di Renzo, New Covalent Boron (III)-Molybdenum(VI) Mixed Oxo Model Compounds As Eligible Hetero Bimetallic Catalysts For Propylene Epoxidation, in:B. Imelik et al. (Ed.), Catalysis by Acids and Bases, Elsevier, Amsterdam, 1985, 177; E.Tempesti,L.Giuffre,C.Mazzocchia,G.~1odica and E.Montoneri, Covalent Arylboron(III)-Molybdenum(VI) Mixed oxo Derivatives As New Model Compounds For Catalytic Oxygen-Transfer Reactions To Olefins, presented to the Fourth International Symposium on Homogeneous Catalysis,Leningrad, USSR, September 24-28, 1984. G.K.Boreskov, Proc. 5th Int.Congr.Catalysis, 2 (1972) 981.
180
DISCUSSION E. NEWSON : Have you tri ed heterogeneisation procedures wi th other supports such as alumina, silica, carbon, etc ... and with what results? F. 01 RENZO : Some years ago, Forzatti and Trifiro studying the epoxidation reaction reported the results obtained by supporting molybdenum over silica, alumina and magnesia (React. Kinet. Catal. Lett. 1 (1974) 367). They observed considerable leaching of the catalyst; accordingly, they considered the system of no industrial usefulness. We think that the boronic group is quite specific in fixing molybdenum to the support. Further work is in progress in order to introduce the boronic group into different resin networks.
B. Oelrnon. P. Grange. P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V .• Amsterdam - Printed in The Netherlands
181
PREPARATION OF VO(HP04)·O.5HZO AND ITS TRANSFORMATION TO THE MALEIC ANHYDRIDE CATALYST (VO)ZP207 Jack W. JOHNSON, D. C. JOHNSTON, AND A. J. JACOBSON Corporate Research-Science Laboratories, Company, Annandale, NJ 08801
Exxon
Research
and
Engineering
SUMMARY The blue vanadium(IV) compound VO(HPO g)·O.5H 20 has been identified as precursor to the maleic anhydride catalyst tVO)2P207. Structural similarities between the two compounds suggest a topotactic lransformation from hydrogen phosphate to pyrophosphate as the mechanism of catalyst activation. In both compounds, measurement of the magnetic susceptibility as a function of temperature reveals that the major bulk fraction of the vanadium spins are antiferromagnetically coupled, while a variable minor fraction of the defect spins are nearly isolated from the other spins. Two new phases of VOP0 4 are formed upon oxidation of VO(HP0 4)·O.5H 20. INTRODUCTION Maleic anhydride is an important industrial chemical with a 1980 worldwide capacity in excess of 700,000 metric tons per year. The bulk of maleic anhydride production goes into polyester resins, although it also finds uses in lube additives, agricultural chemicals, specialty chemicals. and other products (ref. 1). Since 1933. maleic anhydride has been produced commercially by the oxidation of benzene over V205-Mo03 catalysts. More recently, since the 1970'S, it has become more economical to replace benzene with C4 feedstocks such as butane in the production of maleic anhydride by partial oxidation (ref. 1-3). One of the developments that has made the C4 route to maleic anhydride more attractive is the discovery that vanadium phosphorus oxide catalysts are well suited for the oxidation of C4 feeds, particularly butane. The 1iterature deal ing with the scientific. rather than economic, aspects of maleic anhydride production from butane over vanadium phosphorus oxide catalysts has recently been comprehensively reviewed (ref. 4). Trifiro and coworkers (ref. 5-8) have extensively studied the effect of preparation conditions and P/V ratio on the performance of maleic anhydride. catalysts. Many variations of the preparative conditions lead to a common precursor that was identified as (VO)2P207·ZHZO. We have prepared VO(HP0 4)·0.5H20 and find it identical to the previously synthesized maleic anhydride catalyst precursor by comparison of the reported X-ray powder and infrared spectral data (ref. 9). In this paper, we describe procedures for synthesizing VO(HP04) ·0.5H20 and the mechanism by which it transforms to (VO)2P207.
182
SYNTHESIS OF VO(HP04)·O.5HZO AND ITS TOPOTACTIC CONVERSION TO (VO)ZPZ07 Vanadyl hydrogen phosphate hemihydrate, VO(HP0 4)·0.SH20, can be synthesized by reaction of the vanadium(V) compound VOP0 4·2H20 with organic compounds. Alcohols are oxidized by VOP04·2H 20 to give the corresponding ketones or aldehydes.
The
reduced
vanadium(IV)
product
of
these
reactions
is
VO(HP04)·0.SH20.
For example, when VOP0 4·2H20 is refluxed in 2-butanol, highly crystallized VO(HP04)·O.SH20 is produced. The well-formed plate-like
crystallites have diameters on the order of 3 um with thicknesses of O.2um or less.
The same compound can also be prepared more conveniently by simply
ref1uxing a suspension of V20S in an ethanolic solution of phosphoric acid. The ethanol reduces the vanadium(V) and as the reduction progresses, the light blue-green,
insoluble VO(HP04)·O.SH20 precipitates.
When prepared by this
method, the crystallites are also plate-like but of smaller dimension,
~.5
~
across and only 0.08 um thick or less. When aqueous solutions of vanadium (IV) are treated with phosphoric acid and concentrated, the tetrahydrate VO(HP04)·4H20 slowly crystallizes at room temperature, while at sligh1y higher temperatures VO(HP04)·0.5H20 is deposited along with a glassy phase upon slow evaporation. from aqueous solution are much larger (_10 5 ~3) solution,
and as
a
result
possible. (ref. 10). vanadium
octahedra
VO(HP04)·O.SH20,
single crystal
The crystall ites deposited than those from organic
X-ray diffraction studies were
The structure of VO(HP04)·4H20 consists of chains of and
these
phosphate chains
tetrahedra.
are
condensed
(Figure 1) that contain the vanadium atoms
into
In
the
structure
two-dimensional
of
layers
in triply-bridged pairs.
When
VO(HP0 4)·0.5H20 is heated in an inert atmosphere above 380°C, it loses water and
transforms
smoothly
into
vanadyl
pyrophosphate.
The
structure
of
(VO)2P207 has been reported in the literature (ref. 11) and is illustrated in Figure 2.
The transformation of VO(HP04)·O.SH20 proceeds topotactically, as
has is suggested by comparison of the lattice constants of the two compounds (Table 1). TABLE 1.
Both compounds are orthorhombic.
The layer spacing decreases from
Relationship of Lattice Constants. (VO)2 P207
VO(HP04)·0.5H20 a
= 7.420(1)
a =
0= 9.609(2)
I
= 5.6931(7) +
+
+
+
a - - + b/2 b ....... c +
c
+
• a/2
7.725(3)
b
= 16.576(4)
c
9.573(3)
=
183
Figure 1. Structure of VO(HP04)·O.5H20 in the ab plane, represent vanadium atoms.
shaded circles
zl!: y Figure 2. Structure of (VO)2P207 in the!£ plane. vanadium atoms.
Shaded circles represent
5.6911 (£) in the hemihydrate to 3.86A (Y2) in the pyrophosphate. This is consistent with removing the water molecule shared by the vanadium pairs, and filling the resulting vacant vanadium coordination site with the oxygen atoms
184
of vanadyl
groups from the layer above.
The P-OH groups of the hydrogen
phosphate condense across the interlayer space to form pyrophosphate groups. The in-plane lattice constant that is shown vertically in the figures remains essentially constant, while the other in-plane lattice constant expands from 7.42A ~) in the hydrogen phosphate to 8.23J.1 (Ef2) as a result of changing from triply-bridged vanadium pairs or face-sharing pairs of vanadium octahedra to double bridged vanadium pairs or edge-sharing pairs of octahedra along this direction, represented as horizontal
in the figures.
It is clear from the
figures that the transformation from VO(HP04)·O.5H20 to (VO)2P207 can occur while maintaining V-P-O connectivity between the layers.
The crystallites of
(VO)2P207' when observed by scanning electron microscopy, look identical to the crystall ites of topotactic
nature
the VO(HP04)·O.5H20 from which they are formed. of
the
transformation
has
also
been
confirmed
The by
transmission electron microscopy and electron diffraction (ref. 12). Although the conversion of the hemihydrate to the pyrophosphate can take place without breaking any V-O-P bonds, closer inspection of Figures 1 and 2 show that some adjustments must be made along the direction perpendicular to the 1ayers.
In Figure I, the terminal vanadyl oxygen atoms in the face-sha red
octahedral pairs of VO(HP04)·O.5H20 have a ~ arrangement, while in (VO)2P207 the vanadyl oxygen atoms in the edge-shared pairs are in anti positions. Also, the phosphate hydroxyl groups alternate regularly down and up along the
.! axis in the hemihydrate structure, while in the pyrophosphate, the P207 groups go in sequence of two down, two up along the corresponding causes the doubling of the unit cell in this direction). in
l
axis (this
These rearrangements
the layer stacking direction, along with the 11% expansion along the
VO(HP04)·O.5H20 .! axis,
result
in
the
initial
formation
of
(VO)2P207
crystallites with many defects.
EFFECTS OF ANNEALING ON (VO)2PZ07 Maleic anhydride catalysts are usually activated by heating the precursor phase to a temperature of 380°C or above (ref. 4).
This is the temperature at
which VO(HP04)·O.5H20 transforms to (VO)2P207 (ref. 9).
However, the nature
of the (VO)2P207 formed depends strongly on the activation conditions. vanadyl
pyrophosphate
is
formed
by
heating
vanadyl
hydrogen
When
phosphate
hemihydrate in flowing hel ium for 19h at 420°C, the X-ray powder diffraction pattern is not very well defined (Sample 1 in Figure 3). the 200 reflection at 2e
=
The sharpest peak is
22.9°, which corresponds to the layer separation.
The width of this reflection is significantly less than in the diffraction pattern reported recently (ref. 13) for a (VO)2P207 sample prepared by calcining a precursor prepared in an organic medium at 420°C first in air and
185
~r I
~
-- -1
~ ~
)
\..---.J,
A
2
--loA
3
A
i
A A
1
L_~
o
10
A.
I
30
20
-
A
I
40
50
60
28 (CuKa)
Figure then
3. X-ray diffraction patterns for (VO)ZPZ07 successively higher temperature (see text).
in
severity improves.
1% butane/air, of
the
samples
annealed
wh i ch showed a broadened ZOO refl ect i on.
anneal ing
conditions
is
increased,
the
at
As the
crystall inity
Sample Z of Figure 3 was heated in flowing hel ium for 19 .lnours at
790°C, and sample 3 was prepared by heating sample Z in a sealed, evacuated tube for 49 days remained
at
relatively
775°C. constant
The BET surface areas of the three samples at
6.Z,
6.7,
and
4.9
mZ/g,
respectively.
Examination of the samples by SEM showed no discernible difference between Samples 1 and Z, but the longer term heat treatment for Sample 3 did alter the crystallite morphology, as shown in Figure 4.
Figure 4.
SEM photographs (6000x) of (VO)ZPZ07 sample Z (right) and 3 (left).
186
In both VO(HP04)·O.5H20 and (VO)2P207' the vanadium atoms are in the +4 oxidation state, possessing a single d electron. In the structure of the hemihydrate, these atoms occur in
isolated pairs.
Hence, we expected and
found a temperature dependent magnetic susceptibility x(T) for the bulk spins characteristic of
In the structure of (VO)2P207' the vanadium pairs are connected in
(ref. 9). the
~
isolated antiferromagnetically coupled pairs of spin 1/2
direction to form an infinite ladder though the crystal perpendicular to
the layers.
Again, antiferromagnetic coupling between the spins is observed.
but the form of x{t) for the bulk spins is more complex (ref. 14 and 15).
The
important finding from our magnetic studies of (VO)2P207 catalysts relevant here is that the V4+ defects have a different x(T) behavior than that of the bulk spins and that the defect densities could therefore be measured.
Sample
1 had 25 moll of its vanadium as defects, Sample 2 had 71 and Sample 3 had 2.31.
It
is
well
known
that
heating
a maleic
anhydride
catalyst
to
temperatures above 450°C will decrease the total yield of maleic anhydride after the temperature excursion. We postulate that this behavior is associated with heal ing of the magnet ic defects in the (VO)2P207 structure. Indeed. preliminary catalytic tests on {VO)2P207 samples annealed at 450 to
soooe
show that the rate of conversion of butane to maleic anhydride can be
correlated with the number of magnetically detected defects (ref. 16). The exact role of the defects in the catalytic reaction is not clear. In view of the low surface area and relatively high number of magnetic defects. they must
be located primarily
in
the bulk.
A recent study of
butane
oxidation over {VO)2P207 showed through isotopic labe11 ing techniques that only the surface layer of oxygen atoms are involved in the catalytic reaction (ref. 17).
Close examination of the crystal structure of (VO)2P207 (ref. 11)
reveals that the longest (V-O) bonds in the plane of the layer are those that hold the vanadium pairs together.
Shear planes breaking these bonds could
isolate the vanadium atoms and destroy the antiferromagnetic coupling.
The
propagation of these defects to the surface could then result in the sites that are able to activate butane.
TRANSFORMATION OF VO(HP04)·0.5H20 TO VOP04 When vanadyl hydrogen phosphate hemihydrate is heated in the presence of oxygen. the vanadium is oxidized from the +4 to the +5 oxidation state during and after the loss of water. A yellow compound of stoichiometry VOP04 is obtained
with
a
structure
that
depends
on
the
reaction
temperature.
Experiments were performed in flowing 02 at 400°C or in air at 750°C.
The
x-
ray powder diffraction patterns and infrared spectra show that two different phases were obtained.
The infrared spectra (Figure 5) are different from each
187
wavenumbers Figure 5.
Infrared spectra of
(right) and y-VOP04 (left). other, part tcul arly in the region around 600 cm- 1• Vanadium and phosphorus analysis of both compounds confirmed the stoichiometry VOP04. The X-ray diffraction pattern of the low temperature product, ~-VOP04' is broad, but the peaks cannot be assigned to a known VOP04 phase. When ~-VOP04 is heated to 700 0 e in air, the high temperature phase y-VOP04 is produced. This phase also results when well-crystallized (VO)2P207 is heated in air at 600-700 0 e. The powder pattern of y-VOP04 is quite sharp. The transformation y-VOP04 ( ) (VO)2P207 has been studied by Bordes, et al ,; and a structure ~-VOP04
TABLE 2. X-ray Powder Diffraction Pattern of 6-VOP04 d
d( A)
4.54 4.03 3.68 3.21 3.13 2.96 2.73
M
VS M
W S W VW
2.58 2.15 2.03 1.898 1.781 1.563 1.477 1.427
I W
W, Br W M M
W W
W
188
TABLE 3. X-ray Powder Diffraction Pattern of y-VOP04 d(A)
d(A) 6.157 4.886 4.346 4.258 4.151 3.910 3.838 3.501 3.442 3.214 3.089 3.056 2.897 2.774 2.726 2.698 2.613
7.3 66.6 39.9 18.2 56.1 100.0 61.7 62.0 11.0 45.2 42.8 56.8 4.6 7.8 12.2 16.2 4.3
2.487 2.448 2.415 2.338 2.163 2.055 2.005 1.991 1.961 1.951 1.901 1.859 1.824 1.788 1.770 1.747 1.723
10.0 24.1 24.6 2.7 8.6 5.2 8.9 6.7 10.1 9.4 8.9 6.0 13.0 4.5 9.4 5.0 5.3
for the former compound has been proposed (ref. 18). At the temperature that maleic anhydride catalysts typically operate, the important phase in oxidized regions of the catalyst bed is most likely o-VOP04• The phase relationships within the V-P-O system have recently been examined, and Bordes and Courtine postulate that (VO)2P207 prepared from VO(HP0 4)·0.5H20, which gives y-VOP0 4 upon air oxidation at 700°C, is a distinct phase, different from the (VO)2P207 phase prepared by classical sol id state methods that give S-VOP0 4 upon air oxidation (ref. 19).
ACKNOWLEDGEMENT We thank J. F. Brody, M. E. Leonowicz, and D. P. Goshorn for experimental assistance and P. Courtine for helpful discussions.
REFERENCES 1. M. Malow, Hydrocarbon Processing, (Nov. 1980) 149-153. 2. R. L. Varma and D. N. Saraf, Ind. Eng. Prod. Res. Dev., 18 (1979) 7-13. 3. K. Wohlfart and G. Emig, Hydrocarbon Processing, (June, 1980) 83-90. 4. B. K. Hodnett, Catal. Rev. Sci. Eng., 27 (1985) 373-424. 5. G. Poli, I. Resta, O. Ruggeri, and F. Trifiro, Appl. Catal., 1 (1979) 395-404. 6. F. Cavani, G. Centi, I. Manenti, A. Riva, and F. Trifiro, Ind. Eng. Chem. Prod. Res. Dev., 22 (1983) 565-570. 7. G. Centi, I. Manenti, A. Riva, and F. Trifiro, Appl. Catal., 9 (1984) 177-190. 8. F. Cavani, G. Centi, and F. Trifiro, Appl. Catal. 9 (1984) 191-202. 9. J. W. Johnson, D. C. Johnston, A. J. Jacobson, and J. F. Brody, J. Am. Chem. Soc., 106 (1984) 8123-8128.
189
10. 11.
M. Eo Leonowicz, J. W. Johnson, J. F. Brody, H. F. Shannon, and J. M Newsam, J. Sol id State Chem., 56 (1985) 370-378. Yu. E. Gorbunova and S. A. Linde, Ookl. Akad. Nauk SSSR, 245 (1979) 584-
588. 12.
E. Bordes, P. Courtine, and J. W. Johnson, J. Solid State Chern., 55
13.
F. Cavani, G. Centi, and F. Trifiro, J. C. S. Chem. Comm., (l985) 492-
14.
D. C. Johnston, J. W. Johnson, O. P. Goshorn, and A. J. Jacobson, unpublished results, 1986. D. C. Johnston and J. W. Johnson, J. C. S. Chem. Comm., (1985) 1720-1722. J. W. Johnson, D. C. Johnston, J. S. Buchanan, and S. Sundaresan, unpublished results, 1986. M. A. Pepara, J. L. Callahan, M. J. Desmond, E. C. Millberger, P. R. Blum, and N. J. Bremer, J. Am. Chem. Soc., 107 (1985) 4883-4892. E. Bordes, J. W. Johnson, A. Raminosona, and P. Courtine, Mater. Sci. Monographs, 288 (1985) 887-892. E. Bordes and P. Courtine, J. C. S. Chern. Comrn., (1985) 294-296.
(1984) 270-279.
494. 15. 16. 17. 18. 19.
190
DISCUSSION F. CAVAN I : It has been reported that the transformation of VO(HP0 4).0.5 Y20 to (VO)2P201 u~ua11y occurs with a s~rong increase of the surface area (from 5 -:- 10 m /g to about 40 -:- 50 m /s). The low values of surface area of (VO) P you obtain suggest that the samples are not prepared from VO(HP04j.6.~ H?O transformation. Do you believe the correlation you find between catalytic activity and defective properties is valid also for the high surface area catalysts ?
°
°
J.W. JOHNSON: Our samples of (VO) P were prepared by the topotactic dehydration of well-crystallized sa~pfel of VO(HP0 4).0.5 H20 that contain no residual alcohol. We see no increase in surface atea, and indeed would expect none due to the topotactic nature of the transformation. We have not measured the magnetic susceptibility of high surface area samples, but in the absence of evidence for the contrary, I would expect the correlation of activity with the presence of paramagnetic defects to apply to these samples as well.
L. HEINRICH: During the calcination process the crystalline VO(HP04).0.5 H20 changes to the crystalline system (VO)2P207 through an amorphous state. Can you give an idea of the importance of the precursor's structure for the formation of the active crystalline (VO)2P207 and of the amorphous phase? J.W. JOHNSON: When (VO)2P?07 is formed by dehydration of VO(HP0 H?O at ~ 400°C in helium, as oar samples were, an amorphous phase is 4).0.5 not observed if the reaction is allowed to go to completion as monitored by weight loss. The X-ray powder pattern is broadened, but the major lines of (VO) P are always observed. The structure of the precursor is important in t~e2fb110wing respects. Due to the crystal structure of VO(HPO ).0.5 H 0, the active catalyst (VO)?P 20 7 can be formed by a topotactic ~eaction~ When the precursor is synthesizea ln an organic solvent, thin plate-like crystallites result. This plate-like morphology is preserved in the active catalyst, due to the topotactic transformation. The predominant crystallite face exposed is the one with maximum selectivity in butane oxidation.
°
G. CENTI : You mentioned a relationship between the amount of defects in (VO) P and activity in n-butane oxidation. Do you think that some alcohol mo1egufel can remain trapped inside the crystalline structure of VO(HP0 4). 0.5 H20 (F. Busca, F. Cavani, G. Centi and F. Trifiro, J. Cata1., Aug. 1986), and tnen can create defects during the transformation to (VO)2P207 (for instance reducing some vanadium sites) ?
°
J.W. JOHNSON: In both methods we used to synthesize VO(HP0 4).0.5 H?O, adequate water was present to conform to the stoichiometry observed 1n the product. In the first method, we used ethanol that contained 5% water, and in the second, we used VOP04.2H 20 as starting material. In contrast, samples formed under strictly anhydtous conditions may be able to incorporate alcohol molecules between the layers to complete the vanadium coordination sphere, resulting in desorder in the stacking direction. Since the transformation to (VO)2P207 is topotactic, one would expect defects or disorder in the precursor to result in defects or disorder in the product. LV.W. GRITZ: At elevated temperatures the catalyst deactivates and the percentage of vanadyl-ions in defect places declines. From plant reactor runs, it is known that the synthesis reaction forms a hot-spot with a temperature excess up to + 100°C. Why is the lifetime of a MSA catalyst more than a year in a commercial reactor? J.W. JOHNSON : In our experiments annealang temperatures in excess of 500°C were required to heal the paramagnetic V + defects in (VO) PO. Even with the hot spot, temperatures much higher than this are unlik~lY to be encountered in stabile commercial reactors.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
191
EFFECT OF ULTRASONIC TREATMENT ON THE PHYSICOCHEMICAL PROPERTIES OF Cr-Mo-O CATALYSTS FOR METHANOL OXIDATION 2 1 and J.PESHEVA 2 T.S.POPOV1 ,D.G.KLISSURSKI ,K.I.IVANOV
IDepartment of Chemistry,Higher Institute of Agriculture "V.Kolarov" ,Plovdiv 2Institute of General & Inorganic Chemistry,Bulgarian Academy of Sciences,Sofia I040,Bulgaria
SUMMARY
A systematic study has been carried out on the influence of ultrasonic treatment on the physicochemical properties of chromium-molybdenum oxide catalysts for methanol oxidation to formaldehyde.Three catalyst series are prepared by precipitation from (J.IIH,4.),4.Mo 70 24.4H,O and Cr(NO')3.9H20 solutio!t5 at a final pH=5-5,5: a)unt:tea1:el1 cati!lysts(series A);o)catalysts subjected to ultrasonic,treatment after filtering of the precipitate(Series B);c) catalysts subjected to ultrasonic treatment during the precipitation and once more after washing the precipitate(Series C). It is established that ultrasonic treatment leads to a gradual decrease in specific surface area of the catalysts and has practically no effect on their specific catalytic activity and selectivity with respect to the oxidation of methanol to formaldehyde.
INTRODUCTION. Chromium-molybdenum oxide catalysts have a higher thermal stability and better mechanical properties than the iron-molybdenum oxide catalysts applied in the industrial production of formaldehyde,selectivity being practically the same in both cases.(I,2). In order to achieve further improvement of the properties of this type of catalysts,it was of interest to study the effect of ultrasonic treatment at different stages of their preparation.
EXPERIMENTAL The catalysts were prepared by a conventional procedure via precipitation from aqueous solutions of(NH4)4M07024.4H20(20 wt %) and Cr(N03)3.9H20(I5 wt.%). at a final pH of 5-5,5 (1,2). The ultrasonic treatment was carried out using a GU 1-06 type
192
generator with an exit power of 0,63 kW and a frequency of 20,4 -23,6 kHz.After washing the pr-e cLpd t a t ed catalysts were dried and calcined finally for 4 hours ,at 550 0C according to a procedure described previously(3).A final Ko:Cr atomic ratio of 2 was choosen and confirmed by chemical and X-ray fluorescence analysis The X-ray analysis of the calcined catalysts was performed with a "Dron-3" apparatus.The infrared spectra were recorded by "Specord-75" spectrometer.A "Jeol" apparatus was used for electronmicroscope observations. The catalytic activity and selectivity of the samples were evaluated using a standard flow apparatus(4).Their specific surface areas were determined by a chromatographic method via argon desorption.
RESULTS AND DISCUSSION As can be seen in Fig.l,the X-ray analysis confirmed the formation in all cases of two phases:Cr 2(fuo0 and Mo0 con4)3 3.This clusion was also shown by DTA curves(Fig.2).Only two endothermic peaks, corresponding to the melting of molybdena and to the decomposition of Cr2(fuo0 were observed. 4)3
x
Cr/Mo0I,)3
o Mo0
3
2
5
6
o
d)A
Fig.I.X-ray diffraction patterns for series B(IO min. treatment) In Fig.3.,a characteristic infrared spectrum of catalysts obtained by the conventional method(Series A) is shownv Phe absorbtion bands in the regions 450-600 cm- I can be attributed to the vibrations characteristic for Mo0 tetrahedra.A more detailed 4
193
200
1.00
600
BOO
7000 t,°e
Fig.2.Characteristic DTA curve(Series B,IO min. treatment) analysis of the spectrum has also shown that in the calcined catalysts a stoichiometric Cr2(Mo04)3 with its characteristic absorbtion bands at 820-860 cm-I and 960 cm- I and M00 600 cm~I 860 3(450cm-I and 1005 em-I) are present. The infrared spectra which are also identical for all preparations,show the existence of well crystallized chromium(III) molybdate. This 1200 conclusion is 7000 800 600 -1 400 em supported by the low symmetry of Fig.3.Infrared spectrum for Series B(60 min. the Mo0 tetra4 treatment) hedra.The stretching modes for Mo0 underr,oes a splitting to two components( 4 830 cm- I and 860 em_I) and the band corresponding to the stretching modes at 960 cm- I appears.
194
A typical distribution of the pore volume as a function of the pore radii for catalysts of series A and B is shown in Fig.4.
100 -
~ 0
\ \
-::
\
~u
\
80
~
r
1.
\2.
lu ~
:::> -..J
60
0 ~
-..J
s
~O
~
lu
20
~
Q:
~
5360
2010
755
28~
PORE RADIUS
A
Fig.4.Pore distribution;curve I-series B (10 min.treatment);curve 2-Series A ~
~%
£100 -.;;.
~
~ ~
~ ~
~50 U) ~
t:: 'q:
~
~
(..)
2570 1230 591 28' 136
A
65
31
PORE RADIUS. Fig.5.Surface area distribution according to pore radius;curve I-series B(IO min. treatment;curve 2-Series A
195
For catalysts of series C and B-subjected to an ultrasonic treatment,weak tendency to increase in volume and surface area corresponding to pores with small radii was observed(J!'ig.5). Data illustratine the tempera5,% ture dependence 100 of the de['ree of conversion of 60 methanol to formaldehyde and the 50 selectivity of conversion with 1.0 50 respect to the same reaction 30 are shown in Fig.
20
6. It has been established that 10 there is no substantial diffe200 225 250 275 300 325 rence between the t,Oc values of these Fig.6~Temperature dependence of the parameters for degree of conversion of methanol to formaldehyde and of the selectivity for all series of series A and for series B (. )min. catalysts(Series treatment A,B and C). Ac.c·ording to some previous studies (5,6) , the ultrasonic treatment ,'during the preparation of precipitated catalysts may have a significant favourahle effect on their dispersity.lt has been assumed(5) that such a treatment can lead to an increase in number of defects in the crystals,which can act as new active sites with respect to some catalytic reactions.Fornlation of an atomic phase was also supposed.A decrease of the mean pore radius of some oxide materials and an increase of the specific surface area upon subjecting the starting hydroxides to ultrasonic treatment have been also reported( 7 ). On the contrary our results have sho\~(Fig~7) that the specific surface area of the catalysts subjected to ultrasonic treatment for I to &.0 min. gradually decreases.Characteristic data for series B and series C catalysts are shown in Fig.7
196
zo 6.0
5.0
o
10
20
30
40
50
60
1: min
Fig.7.Dependence of the specific surface areas of the catalysts on the ultrasonic treatment duration;Series B (0); Series C(e) This can be attributed to two contradictory effects of the ultrasonic treatment. The primary effect should be dispergation of the crystal phase,which is formed. The secondary effect of the acoustic carttation should be an increase of the number of collisions between the nuclei formed and their agregation.1n the case of precipitation from concentrated solutions the second effect seems to be predominating.
REI!'ERENCES 1.T.S.Popov,B.1.Popov,V.N.Bibin,G.M.Bliznakov,and G.K.Boreskov, Catalytic properties of chromium-molybdenum oxide catalysts in methanol oXidation,React.Kinet.Catal.Lett.,1(2)(1975) 169-175. 2.~.S.Popov,V.N.Bibin,B.1.Popov,G.K.Bore8kov,G.M.Bliznakov,L.G.
Karakchieva,G.N.Kustova,L.M.Plyasova,Oxidation of methanol to formaldehyde on cromium-molybdeUum oxide catalysts,Proc.3-rd 1nt.Conf.Heterogen.Catalysis,Varna,1975,Ed.BAN,Sofia,1975,p.I02 3.T.S.Popov,Ph.n~Thesis,University of Plovdiv,Plovdiv 1973 4.G.Bliznakov,M.Marinov,D.Klissurski,V.Ko zhuharov,J.pJsheva,Oxidation ofrmethanol to formaldehyde on v2o - Te0 2 catalysts, Commun.Chemistry,BAN,IS(3)(I982),261-265. S 5.A.N.Mal'tsev,Activation of Heterogeneous Catalysts and Heterogeneous catalytic reactions by ultrasonic treatment,Zhurn.Fiz. khim.(Russ.),50(7)(1976) 1641-1652.
197
6.1.Y.Solov'eva,A.N.Mal'tsev,Zhurn.Fiz.Khim.,46(1970),2970-2978. 7.T.Y.Parijchak,A.N.Mal'tsev,N.1.Kobozev,Zhurn.Fiz.Khim.(Russ.), 41(1967) 1206-1212.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
199
MODIFICATION OF THE TEXTURAL AND ACID PROPERTIES OF A1P0 4 WITH SULFATE ANIONS J. M. CAMPELO, A. GARCIA, D. LUNA and J. M. MARINAS Department of Organi c Chemi stry, Faculty of Sciences, Uni vers ity of Cordoba, E-14004 CORDOBA, Spain.
ABSTRACT The effect of SO-2 on the surface acid properties of A1PO catalysts, based in cyclohexene skeletal isomerizati~n (CSI) has been on its catalytic ac~ivity studi ed by usi ng a mi crocata1yti c pu1 se reactor and temperatures in the range 523-673 K. The kinetic rate constants, activation energies and selectivities at 1-methy1cyc10pentene have been calculated and compared to those_~btained using pure A1PO as the catalyst. It has been found that the A1PO -SO catalysts are more activ\ for the isomerization process than the A1P0 4 cataly~ts, indicating an increased strong acidity. The increase in acidity 1S due to the inductive effects of sulfate on neighbouring hydroxyls which enhance their protonic mobility. The catalytic activity of A1PO -treated with sulfate ion (1-5 wt%l showed a maximun at about 2.5 - 3 wt% and ~aried with further temperature treatment, being highest at a temperature of 573 K. The texture of A1P0 4 is practically unchanged for sulfate contents lower than 4 wt%. INTRODUCTION The presence of acid-base sites on A1P0 4 surface (where both their number and strength are important) has been demonstrated through the use of two mai n experimental procedures. The first one involves the quantitative determination of surface acid-base sites by a spectrophotometric method from, respectively, cyclohexane solutions of organic bases or acids as titrants. The second one deals with the study of catalytic reactions promoted by these sites, such as dehydration, isomerization, alkylation, cracking and retroaldolisation reactions (ref. 1-6). The present paper is part of a systematic study that we have recently undertaken to shed more light on the role of additives in the properties of the A1P0 4 catalysts, especially in the variations of the acid-base properties of modified A1P0 4 in connection with its catalytic properties. Our recent works (ref. 7-11) concerning the promotion of A1P0 4 have provided information on the acidity of the samples and its effect on catalytic activity and selectivity for CSI, a reaction that requires the presence of acid sites of high strength. A clear decrease in acidity, observed after the introduction of alkali ions on A1P0 4 (ref. 7) and an increase in acidity after the introduction of fluoride ions (ref. 8) were found. Also we have established that the course of the change in catalytic activity for the CSI is parallel to that of change in acidity (ref. 9-10). Thus,
200
we found that the CSI can be used to evaluate the stronger acid sites of A1P0 4 catalysts since the catalytic activity and selectivity are a linear function of strong surface aci dity measured vs. anil i ne usi ng a spectrophotometri c method (ref. 5). In the present investigation, we have focused on the textural properties and surface acid sites of a series of A1P0 4 impregnated with different sulfate contents (1-5 wt% S04 2) upon calcination at successively higher temperatures (573773 K) in order to give additional information about the potential use of A1P~~ as acid catalyst and catalyst support. The evolution of acidity in A1P0 4-S04 samples was followed by their catalytic activities in the CSI in order to examine whether it would be possible to generate strong surface acid sites within the A1P0 4 catalysts by incorporating sulfate anions without a great variation in textural properties. Thus, it is shown that conversion was highest when the A1P0 4-S0 42 was calcined at a low temperature, and when the A1P0 4 contained low sulfate loading. So, A1P0 4 has been found to be a solid whose surface acidity is greatly enhanced by the addition of small amounts of sulfate ion. EXPERIMENTAL Material s A1P0 4 (Al/P = 1) was prepared from aqueous solutions of A1C1 3.6H20 and H3P0 4 (85 wt%) by precipitation with propylene oxide, followed by washing with 2-propanol and drying at 393 K for 24 h and then calcining at 923 K for 3 h (ref. 7). Sulfate anions were introduced onto the A1P0 4 surface by impregnation until incipient wetness with aqueous solutions of (NH4)2S04 of varying concentration to obtain catalysts differing in the amount of S04 2 (1-5 wt%). The impregnated A1P0 4 were dried at 393 K for 24 h, then calcined for 3 h, in an electric muffle furnace at temperatures in the range 573-773 K and stored in a desiccator. Reference to the sampl es thus prepared wi 11 be made by gi vi ng A1P0 followed by 4 their composition as wt% S04 2 and the temperature of calcination after the sulfate impregnation. Hence, A1P0 4-1S-573 denotes a sample containing 1 wt% S04 2 calcined at 573 K. Textural properties Textural properties were volumetrically determined for each of the A1P0 4-S0 42 materials using the adsorption isotherms of nitrogen at its liquid temperature. The surface area was obtained by the BET method and pore size analysis was made usi ng the "corrected mode lless" method developed for the analysi s of mesopores (ref. 12). X-ray diffraction analysis X-ray diffraction powder analysis were performed with Fe-filtered Co Ka ra-
201
di ati on (>. = 1. 79026 A) with a standard Phi 1i ps X-ray diffractometer operati ng with a scanning speed of 2Q min- l between 2 e = 15 - 80Q • Catalytic activity The CSI process was carried out in a microcatalytic pulse reactor. The schematic details of the reaction system were described in a previous work (ref. 5). The reactor was packed with catalyst particles (20 - 100 mg, > 400 mesh) he1ded by small plugs of pyrex glass wool. The catalyst in the reactor was first standardized under N2 flow of 40 ml min- l at a temperature of 573 K for 2 h. Then, the reaction was studied under the following conditions: hydrocarbon pulse, 0.2 - 1.6 ~l; temperature, 523 - 673 K. The products were analyzed with a Hew1 ett-Packard gas chromatograph equipped with a flame ionization detector and a Varian 4270 data processor using two columns in series packed with, respectively, 5% polypheny1eter and 5% squa1ane on Chromosorb G AW-DMCS 80/100. The GC was operated in a oven temperature of 323 K. The response factors were equal for all a1 kenes obtai ned: 1-, 3-, and 4methylcyc10pentenes and cyc1ohexene. RESULTS AND DISCUSSION Textural properties Full adsorption-desorption isotherms were determined up to the saturation vapor pressure of nitrogen at 77 Kon each of the A1P0 4-S04-2 samples. A representative isotherm for A1P0 4-S0 2 (1 - 3 wt%) samples is shown in Figure 1. The isotherms, which correspond to type IV of the BDDT classification (ref. 13) and exhibit type A hysteresis loops of the de Boer classification (ref. 14), are characteristic of adsorption on mesoporous solids. That this is the case is confirmed by the state of the n-plots (ref. 15) which are all linear over the initial range of p/po and deviate upwards from linearity with the onset of capillary condensation. This capillary condensation starts at relative pressures simil ar to the corresponding hysteresi sloops. These facts i ndicate that in the texture analysis we are dealing essentially with a cylindrical pore system. Besides, none of the isotherms show saturation in the vicinity of p/po = 1, where complete mesopore filling may be safely assumed. This is indicative of a continuous passage from the meso- to macroporosity. For A1P0 4-S0 2 (4 - 5 wt%) samples, type II isotherms and linear n-plots in all p/po range were found as a result of unrestricted monolayer-multilayer adsorption on the non porous solids. The agreement between the values of SBn (at p/po < 0.3) and St (from the slope of the linear part on the n-plots) in Table 1 for all solids is noteworthy, thus confirming the validity of the standard isotherm used (de Boer's t curve) on the basis of the CBET constant (ref. 15).
4
4
202
Fig. 1. Adsorption-desorption isotherm of nitrogen at 77 K of the AP-1S-673 sample, Pore size distributions have been determined by application of the "corrected modelless" method developed for the analysis of mesopores (ref. 12) following cylindrical idealization applied to the adsorption branch of the isotherm. The pore volume, vp' the cumulative pore volume, ~vp' and cumulative surface area, ~p' for the samples are also shown in Table 1. The cumulative pore volume and surface area correspond quite well to Vp and SBET' Table 2 shows the pore size distributions of studied samples. From the data co11 ected in Tables 1 and 2 some differences in textural properties of samples are patent, especially for A1P0 4 containing 4 - 5 wt% S04 2, which strongly decreases the surface area and porosity due to the crystallization of A1P0 4 , For A1P0 4-S0 42 (1 - 3 wt%) samples, impregnation with sulfate simultaneously decreases surface area and pore volume (almost independently of the sulfate content) but the pore radii spectrum is displaced towards even increasing values (Table 2). These facts are independent of the calcination tem-
203
TABLE 1 Textural properties of A1P0 4 impregnated with sulfate anions and calcined at different temperatures. CATALYST
SBET (m2 g -1)
St (m2 g-l)
Vp (ml 9-1 )
:c.V a p (ml g-l)
:c.S a p (m2 9-1)
AP AP-1S-573 AP-2S-573 AP-2.5S-573 AP-3S-573 AP-4S-573 AP-5S-573 AP-1S-673 AP-2S-673 AP-2.5S-673 AP-3S-673 AP-4S-673 AP-5S-673 AP-1S-773 AP-2S-773 AP-2.5S-773 AP-3S-773 AP-4S-773 AP-5S-773
79 60 46 45 40 18 16 61 54 47 46 15 15 60 53 50 50 10 10
80 61 45 45 40
0.503 0.324 0.272 0.247 0.227 0.036 0.036 0.332 0.296 0.296 0.298 0.029 0.029 0.377 0.364 0.346 0.327 0.029 0.029
0.509 0.331 0.276 0.251 0.232 0.038 0.038 0.335 0.301 0.301 0.300 0.033 0.033 0.381 0.370 0.354 0.332 0.031 0.031
72 58 43 43 40 17 17 54 50 45 48 16 16 60 50 53 52 10 10
61 56 49 48 61 52 53 53
aFr om the adsorption branch of the isotherms using the mode11ess method and cylindrical idealization.
TABLE 2 Pore size distributions of A1P0 4 impregnated with sulfate anions and calcined at different temperatures DISTR IBUnON. Vol %
CATALYST
AP AP-1S-573 AP-2S-573 AP-2.5S-573 AP-3S-573 AP-1S-673 AP-2S-673 AP-2.5S-673 AP-3S-673 AP-1S-773 AP-2S-773 AP-2.5S-773 AP-3S-773
>300 A
200-300
14.8 18.7 22.4 15.2 26.0 21. 9 17.5 18.8 31.6 24.8 29.7 30.4 34.9
18.2 21.6 8.3 18.1 27.0 19.3 19.6 22.7 16.8 32.2 19.3 24.6 24.2
100-200 52.9 43.3 39.9 40.5 30.2 40.8 43.7 40.7 36.2 29.6 38.8 28.8 26.5
50-100
20-50
<20
12.2 11.0 15.2 21.7 10.1 12.9 14.7 13.7 10.9 10.4 10.1 12.8 11. 1
0.7 1.4 2.5 2.9 2.6 2.8 2.2 2.2 1.4 0.4 0.2 0.7 1.0
1.2 4.0 1.7 1.6 4.1 2.2 2.3 1.9 3.1 2.6 2.0 2.6 2.4
204
perature, 573 - 773 K. Besides, for the same sulfate loading (1 - 3 wt%), an increase in the further treatment temperature (573 -773 K) leads to a slight increase in surface area and pore volume. The overall effect is a slight variation in the textural properties of A1P0
catalysts for sulfate loadings between 4 1 - 3 wt%, with a large increase in surface acid properties, as is indicated below.
X-ray diffraction analysis X-ray diffraction effects on the sampl es are gi yen in Fi gures 2 and 3 both 2 and
before and after impregnati on with ammoni um sul fate up to 1 - 5 wt%
S04
calcination between 573-773 K.
AP-4S AP-5S
A-3S
40
30
20 2 9
Fig. 2. X-ray diffraction profiles for samples calcined at 673 K (AP-1S, AP-2S, AP-3S, AP-4S and AP-5S) showing sulfate effects on crystallinity.
205
773K
673 K
573 K
50
40
30
20
2
e
Fig. 3. X-ray diffraction profiles for AP-2.55 samples calcined at 573, 673 and 773 K.
2 XRD spectra of sulfate containing A1P0 catalyts show that as the 50 con4 tent increases, sharper lines are obtained indicating a development of crystallinity which is almost complete at 4 - 5 wt% 50 2 (Fig. 2). This effect is
4
4
independent of the treatment temperature back to impregnation although, for a given sulfate content, the increase in the calcination temperature does not affect the degree of crysta11i nity, si nce the intensity of the XRD bands is unaltered (Fig. 3). The structure corresponds to the pseudohexagonalone of the tri dymi te a1though the relative peak intensities differ slightly from those reported by Florke (ref. 16). The band at d = 4.31 A (2 e = 24.0 2 ) is related to the presence of structural defects. The X-ray diffraction patterns given for A1P0
4
samples containing 1 - 3 wt%
206
S04 2 are practically identical to pure A1P0 4, thus maintaining a low degree of crystallinity in the samples and hence, their porous texture, as has been discussed above. This low sulfate content lets us assume that the sulfate anions are incorporated as surface groups. The fully developed crystallinity for 4 - 5 wt% S04 2 samples explains the extremely low surface area of the samples. Catalytic activity The operation variables were so chosen as to eliminate diffusion control on the reaction rate. The CSI process yields a mixture of 1-, 3- and 4-methylcyclopentenes (1-, 3- and 4-MCP) although the fractional conversion to 4-MCP is almost negligible due to the short contact time in the reaction process. Side reactions such as dehydrogenation to benzene and cyclohexadienes or cracking were not observed. Thermal isomerization is also negligible. The gas-chromatographic method permits a simple quantitative study of first order reactions only, where the fractional conversion of reactant to products is independent of pressure. Besides, the rate of adsorption must be fast relative to the rate of the surface reaction, the latter being the rate-controlling step. Under our experimental conditions, these facts were confirmed and thus the fractional conversion of a pulse of reactant is given by an equation analogous to that for conversion in a steady-state flow reactor (under similar conditions). So, the isomerization activity (expressed by the total conversion to 1-, 3and 4-MCP, which is only exact for conversions lower than 20 wt% where the equilibrium reaction can be neglected) is fitted to the Bassett-Habgood (ref. 17) equation: ln __1__
t-x
= R T kK
~ Q
where X is the molar fractionary conversion, R the gas constant, T the reaction temperature, k the rate constant of the surface process, K the adsorption constant of cyclohexene on the catalyst, Wthe catalyst weight and Q the flow rate of the carrier gas. Besides, on all catalysts, conversion was found to be insensitive to pulse size between 0.1 - 1 pl. This behaviour significantly ensured linear chromatography in the pulse mode, i.e. ensuring equilibrium chromatography. Measurements of catalytic activity and determination of selectivity parameters have been performed at reaction temperatures between 523 - 673 K at 50 K intervals on A1P0 4-S042 (1 - 5 wt%) samples calcined at 573, 673 and 773 K for 3 h. The catalytic activity and selectivity were found to be affected by such factors as these involved in the preparation of the catalyst. Apparent rate constants, kK, at 673 K reaction temperature from the 1i near
207
plots In(l/(l-X)) vs. W, activation energies, Ea, and preexponential factors of the Arrhenius equation, ln A, for all catalysts studied are compiled in Table 3 together with the kinetic selectivity factors (a) to l-MCP. All values were reproducible to within about 7%. Table 3 also includes the same values, selected for comparison, previously obtained for A1P0 4-F (1 - 3 wt%) catalysts (ref. 10). From Table 3 it is evident that the catalytic activity and selectivity of A1P0 4S04 2 catalysts, and hence their strong surface acidity (since the CSI is known to be a probe reaction for the strong acid sites on catalyst surfaces (ref. 18)) are higher than that of A1P0 4-F catalysts (ref. 8, 10) when containing the same amount of addit i ve and subjected to the same heat treatment. Thi s shows that the specific effect of anions with an increase in the surface acidity of A1P0 4 is markedly stronger for S04 2 than for F- as occurs for other metal oxide catalysts such as A1 203 (ref. 19-21l, Ti02 (ref. 22-25), Zr02 (ref. 25-29), Fe203 (ref. 30), Sn0 2 (ref. 31, 32) and Hf0 2 (ref. 33) or mi xed oxides 1i ke Zr02-Sn02 (ref. 32). TABLE 3 Apparent rate consta nts at 673 K, kK, acti vati on energi es, Ea, preexponenti al factors of the Arrhenius equation, ln A, and selectivity factors to l-MCP, ~ for cyclohexene skeletal isomerization on pure A1P0 4 and A1P0 4-S042 catalysts CATALYST
kK 106 (mol atm-l g-l s-1)
AP-1S-573 AP-2S-573 AP-2.5S-573 AP-3S-573 AP-4S-573 AP-5S-573 AP-1F-573 AP-2F-573 AP-3F-573 AP-1S-673 AP-2S-673 AP-2.5S-673 AP-3S-673 AP-4S-673 AP-5S-673 AP-l S-773 AP-2S-773 AP-2.5S-773 AP-3S-773 AP-4S-773 AP-5S-773
14.30 75.02 123.82 125.00 1.60 1. 50 22.25 24.79 25.82 12.36 57.34 67.80 66.72 1.63 1.59 9.50 35.82 45.58 37.80 1.55 2.70
Ea (KJ mol-l)
ln a
60.9 59.9 61. 7 59.8 61.7 61.0 63.8 59.1 60.7 60.2 59.7 63.2 61.0 58.9 64.8 63.9 58.8 68.8 62.8 62.5 64.5
8.27 9.26 10.23 9.71 7.40 7.30 9.46 8.66 8.54 8.27 9.25 9.89 9.60 6.51 7.44 9.09 9.18 10.98 9.65 7.36 8.09
aA is expressed in mol atm-l g-l s-l bRatio of the fractional conversion (X l/X3)
a
b
--2.51 3.56 3.77 4.08 2.00 1. 69 3.07 3.14 3.25 2.33 3.55 3.60 3.56 1.08 1.11 2.18 3.09 2.28 3.19 1.10 1.18
208
The changes in the apparent rate constants, kK, as a function of sulfate content are shown in Figure 4 for different calcination temperatures. This Figure, which also includes the results previously obtained for pure A1P0 4 catalyst (ref. 10), shows the drastic effect that the sulfate content and the calcination temperature have on CSI activity. Figure 4 indicates that catalytic activity, and hence strong surface acid sites, increase with an increase in sulfate content, pass through a maximun at about 2.5 - 3 wt% and then strongly decrease to below the initial value (without added S04 2) , independently of the calcination temperature. However, the calcination temperature of the catalyst back to impregnation also influences the activity and thus, an appreciable decline was found by calcining at temperatures above 573 K. Among the samples heated at 773 K only those containing 2 - 3 wt%
100
50
o
3
4 wt % SO-2
5
4
Fig. 4. Variation of the apparent first order rate constants with S04 2 loading: (0) AP-S-573; (e) AP-S-673; (0) AP-S-773.
209
are catalytically active. Deactivation at 4 - 5 wt% 5°4 2 may be ascribed to the considerable decrease in the surface area of the catalyst due to the crystallization of A1P0 4, which equally implies a great decrease in surface acid properties (ref. 7). Presumably the activity maximun at low sulfate content results from two opposing factors: as the sulfate content increases, the acidity of the remaining surface hydroxyls must increase, but the number of such hydroxyls will decrease as they are replaced by sulfate. Thus, the strength of acid sites is increased by the inductive effects of sulfate anion on neighbouring hydroxy1s which enhance their protonic mobility by weakening the O-H bond. Besides, the sulfate attached to Al or P polarises the A1-0-P bonds increasing the Lewis acid strength of A1P0 4. The whole effect is a decrease in the number of OH groups and an increase in the acid strength of those remaining. The i.r. spectra (ref. 34) confirm this asumption. In either case, the apparent activation energies for CSI remained practically unchanged after different sul fate and temperature treatments. Thus, the transition state is of the same type for all catalysts, since we are dealing with first order kinetics, and differences in catalytic activities can be explained through the more or less developed carbocationic character of the reaction intermediates which is a function of the acid strength of the catalyst. Selectivity For all catalysts studied, the concentration of l-MCP was higher than that of 3-MCP. The Optimun Performance Envelope (OPE) curves, which describe the selectivity behaviour of products, were obtained by plotting the fractional conversion (X) of a particular product against the total conversion (X T) for different weight ratios of catalyst to introduced CH, such as has been described by Wojciechowski et al , (ref. 35). By applying this method to 1- and 3-MCP (Figure 5) it can be inferred that both are stable primary reaction products coming from CH, since the plots of Xl (or X3) vs. XT show straight lines. A primary product is defined as that whi ch is produced from the reactant no matter how many surface i ntermedi ates are involved in its formation. The application of the Wheeler criterion (ref. 36) on the kinetic selectivity factor (a), by plotting Xl vs. X3 (Figure 6), lets us assume that both isomerization products are really competitors through a parallel process with a first order kinetics. As Figures 5 and 6 contai ned points at all reacti on temperatures, sel ectivity is independent from the extent of reaction and apparent activation energies for 1- and 3-MCP formation are identical on each catalyst. So, the different a values are due to the different A l/A3 ratios according to:
210
10
o Fig. 5. OPE selectivity curves for CH isomerization on the AP-35-573 [( 0) Xl; (e) X1 ) ,AP-45-673 [(0) Xl; (-) X1) and AP-2.55-773 [(6) Xl; (.) X 3) cata1ys~s. (Data at all reaction temperatares and catalyst weights).
a =
~ k
3
=
A1 exp (-Ea/RT) A exp (-Ea/RT) 3
=
~ A 3
As the preexponentia1 factor is a function of the number of high strength acid sites and these are different on each catalyst, the selectivities on A1P0 42 5°4 are different. As can be seen from Table 3 the selectivity to 1-MCP follows the same behaviour as
the catalytic activity,
although the increase in
i.e., it increases by sulfate impregnation
a value is less important. Besides, a is higher for
sulfate-containing A1P0 than pure A1P0 or f1uorided-A1P0 showing that the 4 4 4, formati on of 1-MCP is enhanced with respect to 3-MCP formati on on more aci di c A1P0 4· 50, A1P0
the carbocati oni c reacti on pathway, previ ous 1y proposed for A1P0 and 4 2 catalysts (ref. 5, 10), encounters application for A1P0 catalysts 4-F 4-504
211
o
2
4
Fig. 6. Selectivity factor X vs. X on AP-3S-573 (0), AP-4S-673 (t::.) and AP-2.5S-773 (0 ) catalysts.l(Data a1: all reaction temperatures and catalyst weights). through the action of intermediates with a more developed carbocationic character. This fact in turn increases the conversion with a rise in the selectivity factor to l-MCP due to the higher stability of their tertiary intermediate carbenium ion in relation to the secondary carbenium ion that gives to 3-MCP. The increase in stronger acid sites also explained the slight formation of 4-MCP in relation to pure A1P0 4 catalysts. CONCLUSIONS In this work, the catalytic activities of A1P0 4-S0 2 were studied with an emphasis on anion loading and the further temperature treatment. The results presented demonstrated the role of 5°4 2 in the modification of acidity in A1P0 4 catalysts. Thus, the impregnation of A1P04 with sulfate anions up to 2.5 - 3
4
212
wt%, followed by drying and calcination at a low temperature (573 K) leads to highly active and selective catalysts for the CS1 process which brings us to the conc1usi on that, in practice, they bear strong aci d centres on thei r surface. Thus, it was found that the acti vity increases si gnifi cant ly with i ncreasing sulfate content up to 3 wt% independently of the studied treatment temperature, although the activity was highest at the lower temperature (573 K). These resul ts are in accord with the previously reported by Tada et a1. (ref. 37, 38) on the dehydration of l-butanol on A1P0 4 both unpromoted and promoted with sulfate ions. These authors find that so~2 ion produced an enhanced conversion at lower concentrat ion, but a poisoning effect in 1arger amounts, and the selectivity is shifted from ether to l-butene. Besides, as far as the A1P0 4 are concerned, the bulk structure remained practically unaltered up to 2.5 - 3 wt% SO~2. This fact is important, since the catalytic properties of A1P0 4 are strongly dependent on thei r textural and surface aci d properti es and thus sulfate impregnation is able to vary the latter while the textural properties of the parent structure remain. Sulfate impregnation also provides a way to control the surface properties of metal supported catalysts as well as the possible SMSI through the decrease in the point of zero charge in the final A1P0 4 used as the catalyst support. ACKNOWLEDGEMENTS The authors gratefully acknowledge the financial support for this work from the Comi si on Asesora de Investi gacion Ci entifi ca y Tecnica (CAICYT, Project 0257/84), Ministerio de Educacion y Ciencia, Espana. REFERENCES 1 J.M. Marinas, C. Jimenez, J.M. Campelo, M.A. Aramendia, V. Borau and D. Luna, Proc. 7th Iberoamerican Symposium on Catalysis, La Plata, Argentina, 1980, p, 79. 2 J.M. Campelo, A. Garcia, J.M. Gutierrez, D. Luna and J.M. Marinas, Colloids Surf., 8 (1984) 353. 3 J.M. Campelo, A. Garcia, D. Luna and J.M. Marinas, Can. J. Chern., 62 (1984) 638. 4 M.A. Aramendia, J.M. Campelo, S. Esteban, C. Jimenez, J.M. Marinas and J.V. Sinisterra, Rev. Inst. Mex. Petrol., 12 (1980) 61. 5 J.M. Campelo, A. Garcia, J.M. Gutierrez, D. Luna and J.M. Marinas, Can. J. Chern., 61 (1983) 2567. 6 J.M. Campelo and J.M. Marinas, Afinidad, 38 (1981) 398. 7 J.M. Campelo, A. Garcia, J.M. Gutierrez, D. Luna and J.M. Marinas, J. Colloid Interface Sci., 95 (1983) 544. 8 J.M. Campelo, A. Garcia, J.M. Gutierrez, D. Luna and J.M. Marinas, J. Colloid Interface Sci., 102 (1984) 107. 9 J.M. Campelo, A. Garcia, D. Luna and J.M. Marinas, React. Kinet. Catal. Lett., 30 (1986) 165. 10 J.M. Campelo, A. Garcia, D. Luna and J.M. Marinas, J. Catal., submitted for publication. 11 J.M. Campelo, A. Garcia, D. Luna and J.M. Marinas, J. Catal., in press.
213
12 S. Brunauer, R.Sh. Mikhail and E. Bodor, J. Colloid Interface Sci., 24 (1967) 541. 13 S. Brunauer, L.S. Deming, W.S. Deming and E. Teller, J. Am. Chern. Soc., 62 (1940) 1723. 14 J.H. de Boer, in D.H. Everett and F.S. Stone (Eds . l , The Structure and Properties of Porous Materials, Academic Press, New York, 1958. 15 A. Lecloux and J. Pirard, J. Colloid Interface Sci., 70 (1979) 265. 16 O.W. Florke, Zeit. Krist., 125 (1967) 184. 17 D. Bassett and H.W. Habgood, J. Phys. Chern., 69 (1960) 220. 18 H. Pines, J. Catal., 78 (1982) 1. 19 K. Jiratova and L. Beranek, Appl. Catal., 2 (1982) 125. 20 LA. Paukshti s , P.1. Soltanov, LN. Yurchenko and K. Jiratova, Collect. Czech. Chern. Commun., 47 (1984) 2044. 21 W. Przystajko, R. Fiedorov and I.G. Dalla-Lana, App1. Catal., 15 (1985) 265. 22 K. Tanabe, M. Itoh, M. Morishigue and H. Hattori, in B. Delmon, P.A. Jacobs and G. Poncelet (Eds.), Preparation of Catalysts, Elsevier, Amsterdam, 1976 p. 65. 23 M. Hino and K. Arata, J. Chern. Soc. Chern. Commun., (1979) 1148. 24 G. Busca, H. Saussey, O. Saur, J.C. Lavalley and V. Lorenzelli, Appl. Catal., 14 (1985) 245. 25 M. Hino and K. Arata, Appl. Catal., 18 (1985) 401. 26 M. Hino and K. Arata, J. Chern. Soc. Chern. Commun., (1980) 851. 27 M. Hino, S. Kobayashi and K. Arata, J. Am. Chern. Soc., 101 (1979) 6439. 28 M. Hino and K. Arata, Chern. Lett., (1981) 1671. 29 K. Tanabe, Mater. Chern. Phys., 13 (1985) 347. 30 K. Tanabe, A. Kayo and T. Yamaguchi, J. Chern. Soc. Chern. Commun., (1981) 602. 31 G. Wang, H. Hattori and K. Tanabe, Chern. Lett., (1983) 277. 32 G. Wang, H. Hattori and K. Tanabe, Chern. Lett., (1983) 959. 33 K. Arata and M. Hino, React. Kinet. Catal. Lett., 25 (1984) 143. 34 J.M. Campel0, A. Garcia, D. Luna and J.M. Marinas, unpublished results. 35 D. Best and H.W. Wojciechowski, J. Catal., 47 (1977) 11. 36 A. Wheeler, Adv. Cata1., 3 (1951) 250. 37 H. Itoh and A. Tada, Nippon Kagaku Kaishi, (1976) 698. 38 A. Tada, M. Yoshida and M. Hirai, Nippon Kagaku Kaishi, (1973) 1379.
214
DISCUSSION B. NOTARI : It is known that A1203 catalysts suitably prepared have acidic properties and can indeed perform the isomerization of cyclohexene to methylcyclopentene. A1Z03 catalysts are very easily obtained, very convenient to employ, and are used in many industrial plants where they operate for years. Could you explain why one should employ a more complex material and certainly more expensive as SO~ modified A1P04 to perform this reaction? Complex and expensive catalysts are required where simple and cheap ones do not give satisfactory results. J.M. CAMPELO PEREZ: At the present time, there is an increasing interest in obtaining catalysts with high acidity to be employed in Fluid Catalytic Cracking plants in order to obtain better gasolines. In this connection, it is justified to search for new materials as S04 modified A1P04, exhibiting stronger acidity than A1Z03-SiOZ commercial catalysts. J.B. MOFFAT: You have shown that the addition of sulphate ion to aluminum phosphate alters the pore size distribution. However you attribute the change in selectivity to a change in the strength of the acidic sites. Do you have information on the latter? J.M. CAMPELO PEREZ: From the present results we have not obtained any correlation between the pore size distribution and the selectivity while the latter exhibits the same evolution as the apparent rate constant with respect to the sulfate content, as shown in Table 3. Here it is also possible to see a certain correlation between selectivity factors and catalytic activity which is directly associated to the number of strong acid sites spectrophotometrically measured, using aniline as titrant agent (pK a = 4.6) (ref. 3 and 5). G. JANNES : Concerning the figures in Table 3 and also your selectivity measurements, are your pre-exponential factor values and your activation energy values based on actual rate constants, or on apparent rate constants? In this latter case, are your effects mixed activity-adsorption effects? J.M. CAMPELO PEREZ : Values of Ea and 1n A were obtained from apparent rate constants values, consequently they contain activity and adsorption effects, while the selectivity factors values, a, are obtained experimentally (by gas chromatography). J.L. DALLONS : In order to get information on the acid strength of the surface, do you have performed TPD measurements of NH3 or IR spectroscopy of pyridine ? These methods will give you a good understanding of the repartition of acid sites of your modified catalysts. J.M. CAMPELO PEREZ: We agree with you on the necessity to get more information on the nature of the acid sites exhibiting catalytic properties. At the present time, our information is limited to know that the number of these active sites in the CSI can be spectrophotometrically measured by using aniline (pK a = 4.6) as titrating agent. The number of acid sites determined by n-butylamine (pK a = 10.6) or pyridine (pKa = 5.25) do not exhibit good correlations with the catalytic activity. XU Xiaoding : What is the thermal stability of the acidity of your catalysts? To what temperature can they still retain their acidity? J.M. CAMPELO PEREZ: The thermal stability of the acidity may be visualized in Fig. 4, which shows how the catalytic activity (which is a linear function of the strong surface acidity) strongly decreases upon calcining from 573 to 773 K. The upper limit of temperature at which the acidity is retained is determined by the stability of the A1P04' which, upon increasing the calcination
215
temperature, loses its amorphous character. Thus at 1073 K it is practically a crystalline material with a very small surface area and, consequently, no acidic properties. However, when the A1P04 catalyst is obtained by precipitation with ethylene oxide, the surface acidity is kept at calcination temperatures of 1073 K. Further treatment at 1273 K for 3h destroyes the acidity. G. JANNES : You have published elsewhere papers on the used of your A1P04 materials as catalyst supports. Do you believe that the special catalyt1c properties of these supported catalysts may be explained by a simple bifunctional picture, or do you believe in the influence of some metal-support specific interaction ? J.M. CAMPELO PEREZ: We believe that the support effects have to be attributed to a strong metal support interaction (SMSI) because most of our results obtained using A1P04 as support of several metals are related to liquid-phase hydrogenation process at room temperature. So, it is very difficult to imagine participation of the active sites of the support. Besides, the phenomenon of SMSI actually is not only restricted to A1P04 or Ti02 but it is described in most usual supports, like A1203 and although its nature is not well established. The participation of the support throughout a transfer of electrons between the metal atoms and the acidic (oxidizing) or basic (reducing) sites on the surface of the supports is now becoming apparent.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
217
PROCESSES OF THE FORMATION OF THE ACTIVE STRUCTURES OF THE V-Mo-O CATALYSTS FOR SELECTIVE BENZENE OXIDATION M. Najbar, A. Biela.D.ski, J. Oamra , E. Biela.D.ska Department of Chemistry, Jagiellonian University, Cracow, Poland W.Wal, J. Chrz~szcz, W. Ormaniec Institute of Industrial Chemistry, Warsaw, Poland SUMMARY Phase composition of an industrial tr,ype promoted V 0 -MoO catalyst obtained by the evaporation on a support of ~n5apprd priate solution containing oxalic acid was studied using i.r. spectroscopy, X-ray and electronographic phase analysis. It has been shown that the catalytically active mass in the fresh and used catalyst contained relatively large crystallites of -Na-V-Mo bronze as the predominant phase. Besides this fresh catalyst contained also V 0 -MoO so11d solution and V Mo 0 X-r~y but detectable elec~roBo~aphi phase both not detectable2b~ cally. Electronographic analysis has shown in the used catalyst additionally also MoO and some lower molybdenum oxides. The bronze beiog catalytically ooo-active sodium vanadia-molybd~na the observed high activitr,y and selectiv1t;X of the catalyst had to be ascribed to very small (below 100 I) crystallites of V20R-MoO~ solid solutioo and/or VqMo h040 dispersed on the -nronze crystallites. It is presumed'~hat the presence of -bronze stabilizes the dispersion of catalytically active phases. INTRODUCTION Athough the V205-MOO~ system is the basis of most of the industrial catalysts for the oxidation of benzene to maleic anhydride they usually contain some additives the presence of which frequently results in the formation of oxydic phases different then those which correspond to V205-MoO~ equilibrium diagram: solid solution of MOO~ in V205, the intermediate compound V M0 6040 and molybdenum trioxide (refs. 1-2). The conditions at 9 which the catalytically aotive mass is prepared are also strongly influencing its phase composition. 1nglot and Wenda (ref.~) e.g. have shown that the presence of oxalic acid favoured the formatioD of j3-broDze phase while this was not observed if the active mass was obtaiDed by the fusioD of the oxides or by the evaporation of a solution containing hydrochloric ac1d. Hence the forma tion of the vanadia-mo lybde De ca talys ts of the
218
industrial type deserves special studies a~m~ng to determine the phase composition of catalytically active mass, its evolution in the course of the prolonged use in the catalytic reactor 88 well as to characterize the role of particular components in determing activity and selectivity of the catalyst. Such study of a promoted catalyst prepared by the evaporation from the solution containing oxalic acid was undertaken. EXPERIMENTAL The catalyst containing V20 and M00 in molar ratio 0.65 : 3 5 0.35 was obtained by the evaporation on a sintered corundum support of an aqueous solution of ammonium vanadate, molybdic acid, oxalic acid as well as Na, Ni and Ta promoters introduced in the form of chlorides. Dry catalyst was calcined at 320 3400 0 in the atmosphere of air. The catalyst contained also some P205 as promoter. The properties of this catalyst were described in (ref. 4). For some experiments also the samples of ammonium-vanadium bronzes were used. They were obtained by the thermal decomposition of the product obtained by the evaporation of the solution of NH4 V0 and oxalic acid. 3 The catalytic experiments were carried out in a tubular flow reactor (length 3 m, diameter 21 Im1 ) immersed in a heating bath of temperature about 385°0. In the upper section of the reactor where the intensity of the oxidation process was the highest the tempera tare exceeded tha t of the hea ting ba th by abou t 1000 • The morphology of the catalysts was studied using JEOL 50 A Electron Scanning Microscope. X-ray phase analysis was performed using DRON X-Ray Diffractograph with a copper lamp as the radiation source (CuK~ = 1.538 R). A Ni Filter was used in order to eliminate CU-Ks radiation. The X-ray analysis was carried out after detaching the active mass from the support. I.r. spectra were registered in KBr pellets using a PerkinElmer-580 IR Spectrometer. RESULTS Fig. 1 shows the examples of the scanning electron microscope photographs (SEI) of the surface of the catalyst. It is seen that the catalytic mass is contained in the form of globUlar aggregates (2.5 - 3.5~m) of much smaller crystallites.
219
Fig. 1. SEl photograph of the surface of a fresh catalyst
111 i 13
{ 211
100
h k I 002
a
Do 'Of. 111
202 104 (X)'.
b
I
,
12
10
Fig. 2. X-ray diffraction patterns of the fresh catalyst (a) and of the catalyst used in the reactor for 16 months (b) Fig_ 2 shows the X-ray diffraction patterns of the fresh catal;yst (a) and the same catal;yst after the use in the reactor
220
for 16 months. Essentially the same reflections are present in both cases although the diffraction lines of the used catalyst are somewhat sharper than those of the non-used one. All these reflections correspond to the sodium or ammonium vanadia (5-bronze Me20.V204·5V205 (refs. 5-6). However the same diffraction pattern appeared also in the case of fused Na 20-Mo0 3-v20 5 00 mole % M00 ) samples and i t was concluded that at the pre3 sence of Mo0 sodium vanadia-molybdena bronze is present (ref.?). 3 It is hence very much probable that the phase of fi-bronze detected in the present investigation is in fact vanadia-molybdena bronze. The broad reflection at 2 & = 10.8° which is present only in the diffraction pattern of the non-used catalyst may be due to the non-reacted (InI4)2V12029 (ref. 8). The fact that the diffraction lines of p-bronze, which is a predominant phase in the catalyst samples, become sharper after prolonged use of the catalyst indicate that its crystallinity has augmented during this treatment. Some increase in the specific surface area of the catalyst (from 1 to 3 m2/g) indicate that simultaneously some morphological changes occurred. Fig. 3 shows the i.r. spectra of the catalyst freshly prepared (b) and also of that used for 16 months (a). Besides the bands at 990, 960 and 940 cm-1 characteristic for p-bronze phase the band at 1010 cm-1 is present in both spectra which can be ascribed to the M00 solid solution. This band is present in 3-V205
r Q)
u
c
o E III
C
o
-L.
o
co
Ln
'() 6 x102 cm-1 Fig. 3. I.r. spectra of the fresh catalyst (ba) and of the catalyst used in the reactor for 16 months ( )
221
pure V20 at 1025 cm-1 and the increasing amount of molybdena S in the solid solution shifts its frequency to the lower values (ref. 9). The i.r. investigations indicated, therefore, that besides ~-bronze phase also some amounts of, perhaps highly dispersed, M00 - V20 solid solution phase was present. In order to test 3 5 this conclusion electron diffraction patterns were obtained using the samples which were very thoroughly ground~ Suspension of small amount of such material in acetone was introduced on the carbon film deposited on the copper grid. The examples of the obtained diffraction patterns are shown in Fig. 4.
Fig. 4. Examples of selected area diffraction patterns of the crystals present in the fresh catalyst They enabled to identify the phases present in particular very small crystallites (below 100 j thick). The Table gives the results of the phase identification of 10 statistically chosen crystallites from the unused catalyst and 43 ones from the catalyst after 16 months of working in the reactor. The indexing of the electronograms W8S performed using the X-ray diffraction data. The data concerning the phase composition of the non-used ca talyst clearly indica te tha t besides the (J -bronze phase detectable by X-ray analysis elso V20 (solid solution of M00 in S 3 V20 ' and V M0 6040 phases were present. Evidently the two latter 9 S ones were not detected by the X-rays owing to very small size of their crystallites. Electronographic investigations confirm therefore - the results of i.r. studies on the basis of which the presence of V20 phese was postulated. It shOUld be observed 5 here that V 6040 phase apparently was not detected by i.r. 9Mo bacause of the fact that its characteristic band et 985 cm-1
222
overlaps the 990 cm-1 band of ~-bronze. TABLE: Phases identified in statistically chosen crystals of the active mass of the fresh catalyst Oa talyst Identified phase
fresh
V20 5 V 6040 9Mo (wi th the s tructure of V MOOe (ref.11» 2
a -bronze (Na20.v204·5V205)
used for 16 months in reactor
Number of crystals of identified phase
References Orystallographic system to X-ray diffraction data
.3
(ref. 10)
orthorhombic
3
(ref. 12)
monoclinic
4
(ref. 5)
monoclinic
8
(ref. 10)
orthorhombic
V20 5 V9'!06040 (11)
12
(ref. 12) monoclinic
a -bronze
10
(ref. 5)
monoclinic
Mo17047 MOO - h 3
4
(ref. 13)
orthorhombic
4
(ref. 14) hexagonal
V6013 Mo 4011
5
(ref. 15) monoclinic
2
(ref. 16)
orthorhombic
Mo 5014
1
(ref. 17)
te tragona 1
The phase composition of the catalyst which has been used for a prolonged time for benzene oxidation is somewhat more complicated. Besides three initial phases present already in the fresh catalyst (~-bronze, V20 phase and V 6040) MOO,; and three 9M0 5 lower molybdenum oxides as well as the V6013 phase were detected. This indicates that at the conditions of the catalytic reaction demixing phenomena occurred resulting in the segregation of molybdenum oxides similarly as it was observed in the previous investigations carried out in our laboratory (refs. 18-23). Parallel to the changes in the phase composition of the working catalyst a certain decrease in the selectivity of maleic anhydride production and an increase in the total activity was observed.
223
The nature of ~-bronze phase needs some further discussion. The X-ray diffraction pattern of sodium and ammonium vanadiamolybdena !3-bronzes are almost identical (refs. 5-6). Principally both can be present in the samples of the catalysts because of the presence of NH4 + and Na+ ions in the substr?tes. The discrimination between ammonium snd sodium bronzes could be done by investigating their thermal stabili~. Fig. 5 shows i.r. spectra of the products of calcination of the sample obtained by the evaporation of the solution of NH4V0 in oxalic acid. The 3 calcination was carried out for 3 h in air st 320°0 (a), 420°0 (b), 500°0 (c). It shoul~ be here observed that the presence of strongly reducing oxalic acid enables the reduction of the sample necessary for the formation of bronze phase ~ref. 3). In all cases the i.r. bands at 990, 960 and 940 cm- characteristic of the -bronze phase were present indicating that ammoniumVBnadia bronze is stable even at the temperature as high as 50000 However the results shown in Fig. 6 show that the sample obtained from oxalic acid solution and NH4V0 exhibiting still at 420°0 3
P
1100 1000 cm1
Fig. 5. I.r. spectra of the sample obtained by the evaporation of the solution of NH~VO~ in oxalic acid ana oalcination in air a~ tem~era tures:oa) 320 0, b) 420°0, c) 500 0
am'6&) em1-1 Fig. 6. I.r. spectra of a) the sample Obta~nea oy the evaporat10n o~ ~ne solution of NH4VO; in axslic acid and calcination a~ 420 0 b) the sample obtained by the evaporation of the mixture of V 0 and oxalicoacid and then calci6a~ion at 320 c) the sample "a" used as a catalyst for 8 h
224
t;ypical i.r. bands of (!J-bronze did significantly changed after being used for 8 h at 340-400 oC as the catalyst. Its i.r. spectrum bacame now similar to that of vanadia sample obtained b;y the evaporation of the mixture of V20 and oxalic acid and cal5 cination at 320 oC, thus indicating the decomposition of ammonium fi-bronze at the catalytic conditions. Fig. 7b shows that ammonium ,B-bronze is still less stable at the presence of both vanadium and mol;ybdenum. It can ba therefore concluded that practically onl;y the sodium bronze is stable at the catalytic conditions. The experiments carried out with stoichiometric sodium f3 -bronze have shown tha t this phase is comple tely inactive in the catalytic oxidation of benzene. This is in good accordance with the previous results obtained with fused Na 20-Mo0 20 3-V 5 samples (9.27 mole % Na 20) containing also j3-bronze phase (ref. 7). However, the catalysts used by us obtained from the solution
'000
abo
'em·'
Fig. 7. l.r. spectra of: a) commercial V 0 , b) the sample obtained by the evaporation of the solutfod 8f NH4VO~, molybdic acid and oxalic acid, and calcination at 320 0, CJ tne sample obtained by the evaporationoof the solution of NH4V0 1n oxalic 3 acid and calcination at 320 O.
225
of oxalic acid exhibited fully satisfactory properties concerning both the selectivity to maleic anhydride and the activity. They also have shown quite appreciable stability of their properties during their prolonged use. The X-ray detectable f> bronze, presumably the predominant phase, being catalytically non-active the observed catalytical activity and selectivity had to be therefore - ascribed to the V20 (solid solution) and/or V 6040 9M0 5 present in a state of high dispersion. It can be suggested that in this fairly complicated catalytic system the non-active p-bronze phase forming relatively large crystallites stabilizes the dispersion and hence also the activity of catalytically active phases. REFERENCES 1 A. Bielanski, K. Dyrek, J. Pozniczek, E. Wenda, Bull. Pol. Ac. :Chem., 19 (1971) 507 and A. Bielanski, M. Najbar, Pol. J. Chem., 52 (1978) 883-884. 2 T.Ekstrem and M. Nygren, Acta. Chem. Scand., 26 (1972) 1827. 3 A. Inglot, E. Wenda, Bull. Pol. Ac.: Chem., 28 (1982) 815-822. 4 J. Chrz~szcz, J. Ob16j, W. Ormaniec and W. Wal, in B. Grzybowska-Swierkosz and J. Haber (Eds.), Vanadia Catalysts for Processes of Oxidation of Aromatic Hydrocarbons, Polish Scientific PUblishers, Warsaw-Cracow, 1984, ~. 11. 5 N.S. Butt, A. Fish, F.Z. Saleeb, J. Catal' 5 (1966) 508. 6 A. Inglot, Bull. Pol. Ac.: Chem., 29 (1982)t 293. 7 A. Bielanski, J. Pozniczek, E. Wenda, Bull. Pol. Ac.: Chern., 24 (1976) 493. 8 J. TUdo, Compt. rend. Acad. Sci., 269 C (1969) 112. 9 A. Bielanski, K. Dyrek, A. Kozlowska-Rog, Bull. Pol. Ac.: Chem., 20 (1972) 1055. 10 H.G. Bachmann, F.R. Ahmed and W.H. Barnes, i. Krist., 115 (1961) 110. 11 R.H. Jarman, P.G. Dickens, A.J. Jacobson, Mat. Res. BUll., 17 (1982) 325. 12 H. Eick, L. Kihlborg, Acta Chem. Scand., 20 (1966) 1658. 13 L. Kihlborg, Acta Chem. Scand., 14 (1960) 1612, 17 (1963)1485. 14 Powder Diffraction Fille (Joint Committee on Powder Diffraction Standards, Philadelphia, Pensylvania) 21569. 15 K.A. Wilhelmi, K. Waltersson and L. Kihlborg, Acta Chem. Scand., 25 (1971) 2675. 16 L. Kihlborg, Ark. Kemi., 21 (1963) 471. 17 L. Kihlborg, Acta Chem. scand., 13 (1959) 954. 18 M. Najbar, S. Niziol, Solid State Chem., 26 (1978) 339. 19 A. Bielanski, J. Camra, M.Najbar, J. Catal., 57 (1979) 326. 20 M. Najbar, E. Bielanska, in K. Dyrek, J. Haber (Eds.) Proc. 9th Int. Symp. React. Solids, Cracow, September 1-6, 1980, Elsevier, Amsterdam, 1981, p, 657. 21 M. Najbar, K. Stadnioka, J. Chem. Soo., FaradaY Trans., 79 (1983) 27. 22 M. Najbar, Proe. 8th Int Congr. Catal., Berlin(West), JUly 2-6, 1984, DECHEMA, Frankfurt am Main, 1984, voL, V, p, 323. 23 M. Najbar, J. Chern. Soc., Faraday Trans., 82 (1986) - in print.
226
DISCUSSION E. NEWSON: Was the molybdenum content of your industrially used catalyst the same as fresh catalyst since molybdenum sublimation around the hot-spot is a known deactivation mechanism for these catalysts and systems? W. WAL : We observed a decrease of the molybdenum content in industrially used catalysts in comparison with the content of the fresh catalyst. In the freshly prepared catalysts, the atomic ratio Mo/V = 0.45, while in the catalyst used for two years in the industry this ratio decreases to about 0.30, especially in the region of the thermal peak. The sublimation of molybdenum is one of the most important reasons for the catalyst deactivation, or more strictly, for the decrease of selectivity and yield of maleic anhydride in the process of soft benzene oxidation. The problem has been described in details elsewhere (A. Bielanski, N. Najbar, J. Chrzaszcz and W. Wal, Proc. Intern. Symp. Catalyst Deactivation, Antwerp 1980, Elsevier, Amsterdam 1980, p. 127). D.+ARNTZ : Would you please comment on the quantity of the promoters beside Na used and on their influence on the catalyst structure. W. WAL : The vanadia - molybdena catalysts contain the following promoters (in wt %) : P : 0.04 - Na : 0.35 - Ni : 0.15 - Ta : 0.02. The phosphorous promoter improves the selectivity of vanadia - molybdena catalyst used in the process of benzene oxidation. However, the catalysts with high phosphorous content show low activity. On the contrary, the tantalum promoter improves the activity of the catalysts but the increase in tantalum concentration causes a decrease of the selectivity. Nickel and other metals of the iron group have no essential influence upon the catalytic properties of vanadia - molybdena systems. Phosphorous promoter present in the catalyst in the mentioned quantity does not form any new phase with the vanadium molybdenum oxides, detectable by X-ray diffraction analysis. However, phosphorus strongly influences the oxidation level of vanadium. Parallel to of the phosphorus content phase is an increase of the concentrathe increa~e tion of V ions, as well in the freshly prepared as in the used catalysts. We have not looked at the influence of nickel and tantalum promoters on the structure of vanadia - molybdena catalysts.
B. Delmon, P. Grange. P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
227
© 1987 Elsevier Science Publishers B.V .• Amsterdam - Printed in The Netherlands
PREPARATION CHEMISTRY OF V-Ti-O MIXED OXIDES. COMPARISON OF COPRECIPITATION, GRAFTING AND IMPREGNATION METHODS F. CAVANI, G. CENT I , F. PARRINELLO and F. TRIFIRO' Istituto Tecnologie Chimiche Speciali, Viale Risorgimento 4, 40136 Bologna, Italy.
SUMMARY Vanadium-titanium oxide catalysts prepared by different methods (flash drying, coprecipitation, wet impregnation, grafting) were investigated. Results indicate that three vanadium species are present in addition to crystalline V20S (found at coverages higher than the theoretical monolayer) : two vanadium species (VIV and VV) strongly interacting with the titanium surface and a weakly interacting VV species. All three of these species have an infrared vV=O band different from crystalline V205' Preparation procedure strongly influences the relative ratio of these three vanadium species, preparations in which an oxo-hydrated titanium gel is involved instead of preformed Ti02 give rise to the formation of a greater amount of strongly interacting vanadium species. A comparison of catalytic properties in o-xylene selective oxidation suggests that all these species are involved in the mechanism of oxidation, and in particular VIV, in the stage of activation and vV, in the selective step. INTRODUCTION Interest in V-Ti-O catalysts stems from the high selectivity these materials display in several reactions such as i) the oxidation of o-xylene, ii) the ammoxidation of aromatic derivatives, and iii) the selective reduction of nitrous oxides with ammonia. On the other hand, from a fundamental point of view, interest lies also in the special physico-chemical features of the "vanadium oxide monolayer" which forms at the interface with Ti0 supporting the vanadium oxide and 2 that, in agreement with almost all authors (1-5), is assumed to be the active phase. Although some disagreement is present in the literature on the concept of monolayer, a general good reference is to assume the "theoretical monolayer" as the value of V205 corresponding to the formation of a bidimensional V0 sheet 2. S on the entire Ti0 surface. 2 The aim of this work was to examine more deeply the chemistry involved in the preparation of V-Ti oxides by different techniques in order to obtain a more complete and generalized picture of the nature of "vanadium oxide monolayer" and, by comparison with catalytic data, of the nature of the active phase(s) present in "monolayer" catalysts.
228
EXPERIMENTAL Preparation of catalysts. IV V was reduced by means of oxalic acid until a deep blue V 205 solution was obtained. TiC1 was dissolved in HC1/H (pH lower than 0.5) and 20 4 the two solutions were mixed in the relative amounts necessary to obtain the
Flash drying.
desidered Ti/V ratio. The solution was then dropped on a hot ceramic plate kept at 190 C ; immediate solvent evaporation occurred, leading to a powder which was then calcined at 380 C for 3 h. IV IV Coprecipitation. The V and Ti solutions were prepared and mixed in the same way as for the flash drying method. To this solution, an ammoniacal solution was rapidly added under vigorous stirring. The amount of NH added corresponded to 3 that necessary to have a final pH after neutralization of about 7. The hydrogel obtained was filtred, washed, dried at 120 C and calcined at 380 C for 3 h. Wet impregnation. utilizing
The impregnation of preformed Ti0 was carried out 2(anatase) i) a vIV/oxalic acid/water solution and evaporating the solvent in
rotavapor at 50
c,
ii) a boiling NH
solution and evaporating at 4V03/water around 110 C. After the impregnation step, the catalysts were dried overnight at
150 C and calcined at 380 C for 3 h. The preformed Ti0 were prepared by neutra2 IV/HC1/H lization of a Ti solution with NH , washing, drying and calcin20 3/H20 ation at different temperatures. In the case of the impregnated catalyst on oxo- hydrate
of titanium (Catalyst C of Table I), the same initial procedure of
wet impregnation was adopted, but using the titanium hydroxygel obtained in the neutralization step only dried at 120 C. Grafting. Grafted catalysts were prepared on preformed Ti0 synthetized as before, 2 using a specific reaction in anhydrous benzene (1,6) between the Ti0 hydroxylated 2 surface (by dropping a small amount of water on hot Ti0 and VOC1 The VOC1 2) 3" 3 added corresponded to that necessary in each case to have 1.5 times the amount necessary to form the complete theoretical monolayer. After filtration of the slurry and drying of the solid at 150 C, the chlorine was removed by hydrolysis, adding small amounts of water to hot solid. After further drying, the catalyst was calcined at 380 C for 3 h. Characterization of catalysts. Weighed amounts of calcined samples were treated with an ammoniacal solution (pH around 10). After filtration, the solid was dissolved in a concentred H 2S04/ /H solution. The two solutions, the first containing the extracted vanadium, 20
229
and the second the dissolved catalyst after the extraction, were then analyzed IV V (7) to determine the absolute amount of V and V present, X-ray diffraction (XRD) analyses were performed using a Philips computer-controlled diffractometer with Ni-fi1tered CuK a radiation, Infrared (IR) analyses were performed in air using a Jasco AZOZ spectrophotometer and the KBr disk technique. Spectra were digitalized and the TiO z contribution was subtracted using weighed amounts of TiO in KBr. Z Catalytic tests were carried out in a tubular flow reactor at atmospheric pressure, using 0.38 g of catalyst, a flow rate of 84 cc/min and the following reagent composition: 0.78% o-xy1ene. 13.5% 0Z' 87.3% N Reagent and reaction Z' products were analyzed using two on-line gas-chromatographs; further details on the method of analysis and on the reactor are reported elsewhere (8). RESULTS AND DISCUSSION Chemical characterization. Coprecipitation. Washing V-Ti-O calcined catalysts with a basic medium showed that at least three vanadium species are present. The vanadium species which ~
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FIG. 1 Chemical analyses of coprecipitated catalysts prepared by flash drying (A) and at pH = 7.0 (B). (0.): vanadium not extracted in a basic medium, VIV and vV respectively; (D): vanadium extracted in a basic medium. (-): value of theoretical monolayer. can be dissolved is all vV, while the other two species which remain bound to Ti0 surface, are present both as VIV and vV, respectively. Reported in Figure 2 are the amounts of the different species vs. the nominal amount of V used in Z05 the preparation, for two techniques of coprecipitation. All three vanadium species are expressed in terms of corresponding %wt of VZO in order to have comparable S' data. The first technique (Fig. 1 A) consisted in flash drying of the solution IV IV containing Ti and v , by dropping it on a hot ceramic plate ( T around 190 C).
230
TABLE I Comparison of different V-Ti-O samples calcined in air
* % wt (V Catalyst Surface area 20 5)/g(Ti0 2) IV Ti0 m2/g VV Theor. mono- nominal V vanadium V 2, layer deposited extracted interacting interacting A B C
A B C D E
58 64 56
D
17
E
18
7.4 8.2 7.2 2.2 2.3
11.0 11.6 20.0
n.o
10.0
0.2 8.5 7.5 2.3 8.5
5.5 1.0 7.2 5.3 1.3
5.3 2.1 5.3 3.4 0.2
Coprecipitated catalyst (flash drying) calcined at 380 C. Wet impregnated catalyst on preformed TiD calcined at 380 C. Wet impregnated catalyst on oxo-hydrated €itanium calcined at 380 C. Coprecipitated catalyst (flash drying) calcined at 490 C for 3 h. Wet impregnated catalyst on preformed Ti0 2•
The second one (Fig. 1 B) was classic coprecipitation by rapid mixing of V-Ti solution (final pH around 7). Using the first method, 3/H20 catalysts with final surface areas after calcination at 380 C of about 60-70 m2/g solution with an NH
are obtained, whereas the second method allows catalysts with higher surface areas (about 100-110 m2/g) to be prepared. In all cases only the anatase form of Ti0 is present. The following considerations derive from these results 2 1. The amount of V species strongly interacting with Ti0 is higher than the 2 amount of dissolved species. 2. In both preparations the vIV/vV ratio in the interacting species decreases as the nominal amount of vanadium oxide increases. 3. In catalysts prepared by flash drying (Fig. 1 A), as the nominal vanadium IV content increases, the V interacting species tends to reach constant values near to that of the theoretical monolayer, whereas the VV interacting species increases almost linearly. 4. In the two preparations the absolute amount of vanadium bound to Ti0 is sim2 ilar, but due to the higher surface areas of catalysts prepared by second method (Fig. 1 B), the interacting vanadium accounts a lower fraction of the theoretical monolayer. Wet Impregnations. Reported in Figure 2 is the effect of Ti0 surface area on 2 the amounts of VIV and VV which remain bound to Ti0 after washing, for catalysts 2 prepared by impregnation. Similar results are obtained using VIV/oxalic acid or NH 4V03 aqueous solutions. In all cases the amount of nominal V deposited is 205 11.6 %wt • corresponding to theoretical monolayer of a catalyst with 90 m2/g.
231 10
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FIG. 2 Chemical analyses of impregn- FIG. 3 Chemical analyses of grafted ated catalysts. Nominal vanadium, 11.6 catalysts. (.): fraction of deposited vanadium with respect to theor. monolayer. %wt (V 20S)' Legend. Vanadium not extracted in basic medium, expressed as %wt (V : 20S) ( 0) VIY , (.) yY • Fraction of not-extracted vanadium with respect to theoretical mono1ayer , %wt (V20S) : (6.). Amount of V deposed by grafti ng : (D). The most important considerations that emerge are the following: IV V 1. Just as for the coprecipitated catalysts, both V and V are present in the insoluble fraction; however the absolute amount of interacting Y is much lower than in the case of coprecipitated catalysts, as shown by the comparison between coprecipitated and impregnated catalysts characterized by similar surface areas and nominal vanadium oxide contents (Table I, compare catalysts A,S and D,E). Table I reports also the results of chemical analysis of a sample prepared by impregnation on oxo-hydrated titanium (sample C). The strong analogy with the coprecipitated catalyst (sample A) indicates that the main difference between coprecipitation and impregnation methods lies in the different hydroxilate nature of the titanium surface on which V reacts, an oxo-hydrated titanium in the coprecipitation method and preformed Ti0 in 2 the impregnation method. 2. As the surface area of the Ti0 support increases, the interacting species 2 account for a lower fraction of the theoretical monolayer. In particular, the IV y remains approximately constant whereas the amount of interacting yV increases linearly. Grafting. The results for grafted catalysts on Ti0 anatase with different surfa2 ce areas are reported in Figure 3. In all cases the amount of initial VOC1 for 3 the grafting procedure corresponds to 1.S times the amount necessary to have
232
complete monolayer coverage on t h e s p e c i f i c T i 0 which remains on t h e T i 0
used. The amount o f vanadium 2 s u r f a c e i n c r e a s e s as t h e s u r f a c e area o f t h e s u p p o r t
2 i n c r e a s e s , up t o a c o n s t a n t v a l u e ; c o r r e s p o n d i n g l y , t h e f r a c t i o n o f monolayer
coverage, which i s p r a c t i c a l l y u n i t a r y a t low surface areas, decreases ( F i g . 3 ) .
As i n p r e v i o u s cases, p a r t o f t h e vanadium d e p o s i t e d i s s o l u b l e i n a b a s i c medium. F i g u r e 3 shows t h a t t h e t o t a l amount o f i n t e r a c t i n g V i s v e r y s i m i l a r t o t h a t o f impregnated c a t a l y s t s ; however, t h e r e l a t i v e amounts of V I V and V v i n t h i s i n t e r a c t i n g p a r t , change as compared t o impregnated c a t a l y s t s . I n p a r t i c u l a r , i n a l l V samples, t h e amount o f t h e V I V s p e c i e s i s g r e a t e r t h a n t h a t o f t h e V species, IV v a l t h o u g h t h e r e l a t i v e r a t i o o f V / V decreases as t h e s u r f a c e area i n c r e a s e s . Physi co-chemi c a l c h a r a c t e r i z a t i o n . I n a l l c a t a l y s t s examined, XRD and I R analyses show t h a t c r y s t a l l i n e V 0 i s 2 5 p r e s e n t o n l y i n impregnated and g r a f t e d c a t a l y s t s when t h e amount o f nominal vanadium s l i g h t l y exceeds t h a t o f t h e o r e t i c a l monolayer. F o r example, r e f e r r i n g 2 t o F i g u r e 2, i n t h e c a t a l y s t w i t h a s u r f a c e a r e a of 63 m /g, c r y s t a l l i n e V 0 was 2 5 d e t e c t e d even though t h e nominal vanadium i s o n l y 1.3 t i m e s h i g h e r t h a n t h e t h e o 2 r e t i c a l monolayer coverage, whereas i n t h e sample w i t h 110 m /g, no c r y s t a l l i n e V205 was d e t e c t e d ( i n t h i s case nominal vanadium i s 0.8 t i m e s t h a t o f t h e t h e o r e t i c a l m o n o l a y e r ) . I n c o p r e c i p i t a t e d c a t a l y s t s , c r y s t a l l i n e V205 i s d e t e c t e d o n l y when t h e nominal V l a r g e l y exceeds t h a t necessary t o f o r m t h e t h e o r e t i c a l monolayer ( a b o u t t h r e e t i m e s h i g h e r ) . F o r example, i n t h e f l a s h d r i e d c a t a l y s t w i t h a nominal V o f 25 w t % ( F i g . 1 A ) no c r y s t a l l i n e V 0 was d e t e c t e d even though 2 5 t h e nominal V i s 2.7 t i m e s h i g h e r t h a n t h a t of t h e t h e o r e t i c a l monolayer. Washing t h e c a t a l y s t removes some o f t h e n o n - i n t e r a c t i n g vanadium,
leading t o a catalyst
i n which t h e amount o f vanadium p r e s e n t i s s t i l l h i g h e r t h a n t h a t o f t h e t h e o r e t i c a l monolayer (2.3 t i m e s h i g h e r ) . The m o d i f i c a t i o n s o c c u r r i n g i n t h e XRD p a t t e r n s a f t e r washing t h e c a t a l y s t c o n s i s t o n l y i n t h e disappearance o f a weak background due t o an amorphous phase. Reported i n F i g u r e 4 a r e t h e I R s p e c t r a i n t h e vV=O r e g i o n ( a f t e r T i 0
2
con-
t r i b u t i o n s have been s u b t r a c t e d ) o f some V-Ti-0 c a l c i n e d samples prepared by d i f f e r e n t methods, b e f o r e and a f t e r washing, t o g e t h e r w i t h chemical d a t a ( T a b l e 1 1 ) . The s o l u b l e vanadium i s r e p o r t e d as two f r a c t i o n s ; t h e f i r s t one r e p r e s e n t s t h e amount o f vanadium which s h o u l d be added t o t h e i n t e r a c t i n g vanadium i n o r d e r t o complete t h e monolayer coverage.
I t i s shown t h a t when t h e amount o f nominal
vanadium i s h i g h e r t h a n t h a t necessary t o f o r m t h e t h e o r e t i c a l monolayer, a w e l l
233
<: ~
TABLE II
d
980
960
~ ~
Chemical analyses of catalysts of Fig.4
Cat. Surf.area m2/g
~ e~ 960
94 0
980 960
940
1020
FIG. 4 Infrared spectra of some V-Ti-O catalysts in the v V=O regi on.
a b c d e f g h
34 64 110 110 110 110 60 58 58
vanadium monolayer V~V V~(*) V~ 1.1 1.1 1.1 1.1 4.7 4.7 2.9 5.0 7.4
1.1 2.1 3.4 3.4 4.1 4.1 1.6 4.8 2.3
2.1 5.0 9.6
V 205{*) 7.3 3.4
3.6
a. Strong-interacting vanadium species b. Weak-interacting vanadium species V (*) The weak-interacting V species has been assumed as equal to the amount of soluble vV species necessary in addition to strong-interacting vanadium, to complete the theoretical monolayer j the part of dissolved vanadium exceeding this amount has been assumed as crystalline V 205•
defined band centred at 1020 cm -1 , typical of vV=O in V is present (spectra 205 a and b of Fig. 4). At lower nominal vanadium concentrations with respect to the theoretical monolayer coverage, only a broad band centred around 940-980 cm- l is found, in agreement with other authors (2-5) who indicated that this broad absorption is characteristic of vanadium monolayer species (spectra c-.i). However, more detailed analysis of this broad band, taking into account the corresponding chemical analyses of the samples (Table II) and the effect of the washing, suggests the following considerations 1. The broad band probably consists of three components, at around 940 , 960 and 980 cm- 1. 2. The interacting part of vanadium, which is isolated by the washing procedure, is characterized by the bands at 940 and 960 cm- 1, while the non-interacting -1
part which completes the monolayer, by the band at 980 cm . These considerations derive from the following points -1
1. The spectrum of impregnated catalyst (c), originally centred at 980 cm , but presenting a shoulder at lower frequencies, after washing shows only the weak component at lower frequencies (960 and 940 cm- l) (spectrum d). 2. The spectrum of coprecipitated catalyst (e), in which the strong interacting
234
vanadium species is present in higher amounts than that of impregnated catalyst (c), presents the principal band at 960 cm- l, with shoulders at 980 and 940 cm -1 . The washing procedure (spectrum f) removes only the higher frequency shoulder. 3. The spectrum of catalyst (g), where the interacting part of vanadium is con-1 IV, stituted principally by v shows the principal band at 940 cm . Catalyst (h), on the contrary, according to chemical analysis, shows both bands at -1 940 and 960 cm ; the latter disappears when the sample is reduced (spectrum
i).
Comparison of methods of preparation. Both chemical and physico-chemical analyses show that the nature of the monolayer species is more complex than the only one VV species usually considered in the literature (l-S). At least three species can be evidenced by chemical and infrared analyses, two species (V IV and VV) which strongly interact with Ti02, and a weak interacting vV species, soluble in a basic medium, but not present as crystalline V 20S• The relative distribution of these species is strongly influenced by the specific modality of preparation, and in particular two classes of preparations can be evidenced according to starting titanium material: i) Ti0 (impregnated 2 and grafted catalysts) or ii) oxo-hydrated titanium (coprecipitated catalysts and impregnated catalysts on oxo-hydrated titanium gel). The main difference lies in the amount of interacting vanadium species, which is about three times higher in the latter preparations. This difference can be rationalized considering the specific reaction that can occur between vanadium and free OH groups of the titanium surface. In the case of oxo-hydrated titanium the higher number of active sites for this specific reaction allows an higher surface reactivity toward vanadium, and the comparison between different preparations suggests that this reactions leads to the formation of the strongly-bonded vanadium species, insoluble in a basic medium. The evidence that grafted and impregnated catalysts show the same absolute amount of interacting vanadium supports this hypothesis of the presence of a limited number of OH groups on preformed Ti0 surfaces. 2 Accordingly, a recent paper (9) indicated, on the basis of IR and TPD measurements, that only about one third of the Ti0 anatase surface shows free OH groups, 2 even after hydroxylation treatments with water. The remaining part of the Ti0 2 surface is covered by strongly adsorbed water molecules.
235
In coprecipitated catalysts using both methods, the amount of interacting vIV reaches a plateau,near to the value for theoretical monolayer in the samples with lower surface areas, while interacting vV increases linearly. It should also be noted that i) chemical analysis is carried out on calcined samples and thus refers to species having a different oxidation stability, and ii) infrared analysis on grafted catalysts at low vanadium surface coverage (6) shows that at low calcination temperatures (up to about 300 C) the interacting vanadium has free OH groups (V-OH). All these data indicate that the different reactivity of interacting vanadium can be interpreted as the direct formation of a Ti-O-V bond leading to a stable IV v species (the first monolayer) ; as vanadium content increases a further reaction can occurs on this first vanadium layer leading to the strong interacting
v
.
V spec1es. SURFACE AREA:
HIGH Strong· Interacting
coprec.:.
~
11III v'v
~
IE]
VV
weak·lnteractlng
_
Impreg.:~
f/
TI02
~
VV
crystalline v20S
FIG. 5 Schematic representation of the surface situations of vanadium-titanium oxides as a function of preparation procedures. This model also seems to be valid for the impregnated and grafted catalysts, although some further considerations are necessary. Grafted and impregnated catalysts possess similar amounts of interacting vanadium, but different relative IV ratios of the V and VV species. This could be related to the differences in preparation procedures. However, it also is possible that vanadium interacting with the titanium surface also can be present as VV after calcination, probably related to the presence of different types of sites on the TiO oxide surface as
z
compared with the oxo-hydrated surface. This also can explain the observed effect IV of TiO surface area on the relative distribution of interacting V and VV Z species. Furthermore, the IR and XRD analyses show that in the case of preformed TiO the surface coverage is completed by an amorphous vanadium species different Z from bulk V and with spectroscopic characteristics very similar to interacting Z05
236
vanadium. The presence of this phase in a structural organization (for example, bidimensional) different from bulk V205 can be justified by considering the presence of an interaction with Ti02 (for example, crystallographic fit). The solution of this phase by washing shows that the interaction is weak, however sufficient to contrast the variation in surface free energy in respect to the formation of bulk V (10). 205 These three different vanadium-monolayer species also can be evidenced in the infrared spectra
( vv=o region), as previously discussed. In particular, these
three vanadium species give rise to three components of the broad vV=O band found by many authors using both infrared and Raman spectroscopy and assigned to vanadium-monolayer species (2,4,5).
It should be noted that the strong effect
of preparation on the relative distribution of these species also explains the disagreements found in the literature on the position of the maximum of vV=O. A schematic representation of all these possible vanadium phases present in titanium-vanadium oxides is given in Figure 5. Nature of active species. Reported in Figure 6 are the results of o-xylene oxidation to pthalic anhydride for impregnated catalyst (before and after washing), for a grafted catalyst and for a coprecipitated catalysts,sinterized (however, still with Ti0 in the 2 anatase form) in order to have catalysts with similar surface areas and comparable with literature results (1-17). ~ °.
U
50
.!! >Q)
40
U
...
U
>-30 s:
c
co .~
20
"§.
10
'?300
conversion, %
340
380
temperature, C
FIG. 6 Catalytic tests (o-xylene oxidation) on impregnated catalysts before (0) and after (0) washing, on grafted (6) catalyst and on coprecipitated (.) catalyst. In Table III are reported the chemical analyses of these catalysts after these catalytic tests.
237
TABLE III Chemical analyses of catalysts of Fig. 6 after catalytic tests. Catalyst Impregnat. before wash. after wash. Grafted Coprecipitated
Surface area m2/ g
vanadium monolayer VIV VV VV (*)
18 18 18
17
a
1.5 1.3 1.5 5.3
a.Strong-interacting vanadium species (*) See note in Table II
a
b
0.1
0.7
7.7
0.2 0.1
0.6
3.4 ~.Weak-interacting
2.3
vanadium species.
Four main suggestions derive from these data 1. The three catalysts prepared with different methods and presenting different amounts of vanadium, shown a similar diagram selectivity-conversion, except for the washed catalyst not selective at high conversion. 2. The activity of the catalyst depends on the amount of strongly interacting vIV, as shown by the comparison between coprecipitated and impregnated or grafted catalysts . 3. The vanadium exceeding monolayer (crystalline V does not play any signifi205) cant catalytic role (comparison between grafted catalyst and impregnated catalyst before washing). 4. The weak-interacting VV species (vanadium necessary to complete the monolayer in addition to strongly-interacting V species) seems to playa role in the selective stage, as evidenced by the strong decrease in selectivity at high conversion for impregnated catalysts before and after washing. In conclusion, all three vanadium-monolayer species seems to be involved in the general mechanism of o-xy1ene oxidation, although in different steps, VIV in the activation step and vV in the consecutive steps to form the selective product. The method of preparation influences the relative distribution of these species, and thus the catalytic behavior. In particular, it was shown that the coprecipitation method can enhances the amount of strongly interacting VIV species increasing the activity of the catalyst. ACKNOWLEDGEMENTS Financial support of "Ministero Pubb1ica Istruzione - ItalY" is gratefully ackow1edged.
238
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15 16
17
G.C. Bond, K. BrUckmann, Faraday Discuss., 72 (1981) 235 ; G.C. Bond, P. Konig, J. Catal., 77 (1982) 309. I.E. Wacks, R.Y. Saleh, S.S. Chan, C. Chersich, Appl. Catal., 15 (1985) 339; R.Y. Saleh, I.E. Wachs, S.S. Chan, C.C. Chersich, J. Catal., 98 (1986) 102. A.J. van Hengstum, J.G. van Ommen, H. Bosh, P.J. Gellings, Appl. Catal., 8 (1983) 369 ; ibidem, 5 (1983) 207. Y.Nakagawa, T. Ono, H. Miyato, Y. Kubokawa, J. Chem. Soc. Faraday Trans. I, 79 (1983) 2929. F. Rozeboon, C. Mittelmejer-Hazeleger, A. Moulijin, J. Medema, V.H.J. de Beer, P.J. Gellings, J. Phys. Chem., 84 (1980) 2783. G. Busca, L. Marchetti, G. Centi, F. Trifiro, J. Chem. Soc. Faraday Trans. I, 81 (1985) 1003. G. Centi, G. Fornasari, F. Trifiro, J. Catal., 89 (1984) 44. P. Cavalli, F. Cavani, M. El Sawi, 1. Manenti, F. Trifiro, Ind. Eng. Chem. Fundam., submitt. K. Morishige, F. Kanno, S. Ogawara, S. Sasaki, J. Phys. Chem., 89 (1985) 4404. G. Busca, G. Centi, L. Marchetti, F. Trifiro, Langmuir, submitt. M. Inomata, K. Mori, A. Miyamoto, T. Ui, Y. Murakami, J. Phys. Chem., 83 (1983) 754. M. Gasior, I. Gasior, B. Grzybowska, Appl. Catal., 19 (1985) 13 ; M. Gasior, T. Machej, J. Catal., 83 (1983) 472. J. Haber, Pure Appl. Chem., 56 (1984) 1663. A. Anderson, S.L.T. Anderson, in "Solid State Chemistry in Catalysis" , R.K. Grasselli, J.F. Brazdil Eds., ACS Symp. Series 279, ACS Pub. 1983, 121; A. Anderson, J. Catal., 76 (1982) 144 ; A. Anderson, J-O. Bovin, P. Walter, J. Catal., 98 (1986) 204. P. Kozlowski, R.F. Pettifer, J.M. Thomas, J. Phys. Chem., 87 (1983) 5176. A. Vejux, P. Courtine, J. Solid State Chem., 23 (1978) 83 ; A. Vejux, P. Courtine, C.R. Acad. Sci., 286C (1978) 135 ; J. Papachryssanthou, E. Bordes, A. Vejux, P. Courtine, R. Marchand, M. Tournoux, in " Preprints , European Workshop Meeting - New Developments in Selective Oxidation ", Louvain-laNeuve March 1986. P. Cavalli, F. Cavani, I. Manenti, F. Trifiro, in "Preprints, European Workshop Meeting - New Developments in Selective Oxidation It, Louvain-laNeuve , March 1986 ; G.C. Bond, S. Flamerz, L. van Wijk, ibidem.
239
DISCUSSION J. SCHEVE: It is well known that the ratio of anatase/rutile determines the oxidation state of the vanadium cations. Are you sure that this ratio is the same for all your samples? G. CENTI : All data concerning this communication refer to catalysts where the Ti02 is present in the anatase form, free of impurities of the rutile modTfication. J.G. ~~n OMMEN: There is a4~orrelation between activity in o-xylene oxydation and V present. Is this V measured before or after the r~~ction with o-xylene. What do you think about the concept of creating V during o-xylene oxidation depending on the ratio of oxygen/o-xylene ? G. CENTI : The amount of V4+ in the catalyst utilized for the correlation with the catalytic activity (measured after reaching steady-state conditions) is determined after the catalytic tests. In fact after the catalytic tests, we observed a slight increase in the V4+ amount in the case of o-xylene oxidation and a much higher increase after the catalytic tests in presence of ammonia. A redistribution of the different vanadia species occurs during reaction, ~eing the insoluble VS+ (not present as crystalline VZOS) partially reduced to V + and partially segregated forming crystalline VZOS (F. Cavani, E. Foresti, F. Trifiro and G. Busca, submitted to J. Catal.). Furthermore. we observed that the amount of V4+ tends to reach a value close to that of the theoretical monolayer, at least in the catalysts prepared by coprecipitation and flash drying techniques, whereas in the catalysts prepared by grafting and wet impregnation these phenomena are much less evident (G. Centi, F. Trifiro, G. Busca. L. Marchetti, Langmuir, in press). We did not analyze the influence of the o-xylene/OZ ratio on the amount of V4+ in the catalyst, but all our observations suggest that this value can be influenced by this factor. However, in our opinion the strongly interacting V4+ forms during catalytic tests by specific reaction with the water vapor formed as by-product. The amount of water vapour is thus a second factor which must be considered in any correlation of the amount of V4+ with the reaction conditions. J.G. van OMMEN: What is your concept of monolayer of V20S on Ti02' because we think that the amount needed for a monolayer is less than the amount you indicate (for Ti02 with S8 m2/g, your amount is 7.8% VZOS; we merely think that with 1% per 10 m2, you would get S.8%). G. CENTI : We choose as value for theoretical monolayer the one indicated by Roozeboom et al. (J. Phys. Chern., 84 (1980) 2783) and Bond et al. (Faraday Disc. 72 (1981) 235). This value was calculated on the basis of geometrical consideration of the unit cell of VZOS' The empirical value you suggest in the last paper in collaboration with Bond group (Appl. Catal. 22 (1986) 361) does not seem to us a correct value for reference. Moreover, our data effectively have shown that the maximum v~ue of interacting V4+ is very near to the one calculated on the basis of geometrical considerations (Po Cavalli, F. Cavani. I. Manenti, F. Trifiro. lEC Product Res. Dev., accepted). J.W. GEUS : Usually it is difficult to reduce V(V) compounds, especially the oxide V20S' to compounds where the vanadium exhibits a lower valency. With maleic anhydride catalysts, for instance, reduction by HCl or oxalic acid is required to reduce the vanadium to V(lV), which is stabilized by phosphate. For your results I get the impression that the V20S/Ti02 oxides, reduction with e.g. organic compounds is not resuired to obtain lower valence vanadium oxide. Hence, calcination in air suffices to get vanadium species of a valency lower than (V). Is that true? Is calcination sufficient to release oxygen and reduce vanadium(V) ?
240
G. CENTI : For the preparation of our samples, generally a V(IV)-oxalic acid solution was employed. In some cases, we utilized a V(V) compound, but it should be noted that during calcination there is the formation of NH3 (from NH4V03) or organic contaminants (VOC13 in benzene solution) are present. Therefore, in all cases we have the presence of a possible reducing agent for vanadium. When V4+ is formed, the interaction with Ti02 stabilizes this covalency state and avoids its oxidation even during caTcination in air up to 500°C. In conclusion, we don't have indications that the reaction of V(V) with Ti02 surfa~e is sU~ficient for the redu~tio~ of vanadium, but only on the fact that this lnteractl0n prevents the reoxldatlon of V(IV) to V(V). A. LYCOURGHIOTIS : Can you estimate the fraction of vanadium deposited by adsorption vs. the one deposited by evaporation of the solvent? G. CENTI : Generally, we utilize the technique of wet impregnation by solvent evaporation. Some tests were made by equilibrium adsorption, but the amount of V found on the catalyst using this technique was less than about 1% V20S (wt). This value corresponds to about 10% of the theoretical monolayer and for this reason we did no longer utilize this method of preparation.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV © 1987 Elsevier Science Publishers BV., Amsterdam - Printed in The Netherlands
241
CATION EFFECTS IN THE PREPARATION OF AND CATALYSIS BY HETEROPOLY OXOMETALLATES
J.B. MOFFAT Department of Chemi stry and Guelph-Waterloo Centre for Graduate Work in Chemistry, University of Waterloo, Waterloo, Ontario, Canada N2L 3Gl
SUMMARY The cation is shown to have a profound influence on the structure, porosity, adsorptive properties, thermal stability, and catalytic properties of heteropoly oxometa 1ates. Such salts are frequent ly found to be nonstoi chi 0metric, with the cations being accompanied by residual quantities of protons. The effect of the cations and residual protons on the aforementioned properties is evaluated by the application of a variety of techniques. INTRODUCTION Heteropoly Oxometalates are ionic solids with large, cagelike, high molecular weight anions (1). The anions of interest in the present work may be represented by the stoichiometry [XM120~O]qand are often referred to as having Keggi n structure. The heteropoly anions contai n a tetrahedron X0 4 (X=P, Si, As, for example) which is surrounded by twelve octahedra (MO G) with oxygen atoms at the vert ices and a metal atom M (W, Mo, V, for examp 1e) at their centres (Fig. 1). The twelve octahedra can be visualized as composed of four groups of three edge-shared octahedra (M 30 1 3 ) which are linked by shared corners to each other and to the central X04 tetrahedron (1). Although at least two isomers of the Keggin structure have been identified, the aforementioned form is labelled as the a isomer. Heteropoly anions of Keggin structure have been prepared from at least sixty-five elements as the central atom, but the peripheral metal elements appear to be restricted to those with a Figure 1: Anion of SiMo12040-~ particular combination of ionic radius with Keggin structure. and charge as well as the abil ity to
242
form dn - pIT M-O bonds (1). In the present work phosphorus, silicon, and arsenic have served as central metal atoms while tungsten, molybdenum, and vanadium have been employed as peripheral metal atoms. STRUCTURE Three types of oxygen are found in the heteropo ly ani ons of Keggi n structure, one bridging the central atom and the peripheral metal atoms, a second bri dgi ng each two of the 1atter atoms, and a thi rd bonded to the peripheral metal atoms and protruding from the anion. It is convenient to refer to the latter oxygen atoms as terminal or peripheral. While there is relatively little ambiguity concerning the semi-quantitative aspects of the anion or Keggin structure, there is considerably less certainty with respect to the secondary structures, that is, the arrangement of cati ons and ani ons , Fortunately, Brown and co-workers, with the aid of X-ray and neutron diffraction, have provided valuable information on the structure of 12-tungstophosphoric acid (2). The anion of this heteropoly oxometalate possesses the Keggin structure previously described. Each of the protons is surrounded by four water molecules only two of which, as a result of a twofold thermal disorder, are hydrogen-bonded to the proton (Fig. 2). The hydrogen atoms of the four water molecules are themselves hydrogen-bonded to the terminal oxygen atoms of four of the anions (Fi g. 3).
,,
, ,,
....
~,;.,.,. -VO(4)
Figure 2:
Proton surrounded by water molecules and four terminal oxygen atoms of anion (2).
Figure 3:
Arrangement of anions, protons and water in 12-tungstophosphoric acid (2).
243
Information on the effect of the central atom and peripheral metal atoms of the anion on the charges held on the terminal oxygen atoms and the M-O bond strengths has been obtained from semiempirical extended Huckel (EXH) calculations on simulated heteropoly anions (1-7). The latter were composed of the trimetall ic group with the central tetrahedron attached to yield XM 30 1 6-n (Fig.4). The use of this fragment of the heteropoly anion allowed not only the effect of changes in the metallic or peripheral atom component to be evaluated, but also variations in the nature of the central atom to be s imulated. For convenience the fragments will be referred to as XM,where X and M (0) are the central and peripheral elements, respectively. It should be noted that, of course, the absolute values of parameters calcul ated for the fragments cannot be expected to equal those of the Figure 4: (a) Fragment XM g0 1 6 - of complete anion. However it is expected heteropoly anion, (b) heteropoly anion that the changes in such quantities showing position of fragwith variation in X or Mwill provide a ment. reasonable reflection of those for the complete anion (8). Although the EXH technique will permit the calculation of values for a variety of parameters, only two will be considered here. These are the bond partitioned energy (BE) of the M-O (outer) bonds, which may be employed to represent the bond energy. and the atomic charge (AC) on the outer oxygen atoms. The results of the calculations show two significant features (Fig.S) A change in the central atom from phosphorus to silicon, for example, produced relatively little change in either the BE or AC. while SUbstituting tungsten by molybdenum as the peripheral metal element decreases the BE and increases the magnitude of the AC significantly. As will be apparent in later sections the decrease in BE when tungsten is substituted by molybdenum appears to reflect the increased lability of the terminal oxygen atoms in the molybdenum-containing anion, which is consistent with the observed superior catalytic activity of the latter species in oxidation processes. The observat i on that the magni tude of the AC of the termi na1 oxygen atoms is smaller with tungsten as the peripheral metal element than with molybdenum can be interpreted as reflecting the greater proton mobility (and hence Bronsted
244
acidity) in the former than in the latter case. As will be evident in a later section this is consistent with the selectivities of catalysts containing these anions when employed in the methanol conversion process. 0.0
~
___________--------- b
J:
U
~ ~
Q
a
~ !2.,
6
I=!l!
1.0
----------------- w -=~:::::::::-:--:-;.~--
il
4
~
-------------------- 0
Q
2
~
~
0
10 t-
l;
ffi ffi [l
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\
5
i
%
u \
,,
~-1.0
,, ,
................._-------------::::11- 0
~ -20
0
a
b
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~
'-
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___ b
..:::::::------
-~::::::::=
c
,, ,
~
Z
i
---------'\
a
b
c
d
c
Q1
Figure 5: The effect on the net atomic charge of the terminal oxygen atoms and the M-O bond partitioned energy of changing central atom or peripheral metal atom. (a) PW S01 6 - 9 , (b) SiWS016-10, (c) PMO SOI 6 - 9 (b, bridging; 0, outer.)
Figure 6: The effect on the net atomic charge of the terminal oxygen atoms and the M-O bond partitioned energy resulting from stepwise substitution of Mo for Win PWs 01 6 - 9 , (a) PW SO I 6 - 9 , (b) PW zMoO I 6 - 9 , (c) PWMO Z0 1 6 - 9 , (d)PMo s01 6 - ;
The effect of stepwise substitution of tungsten by molybdenum may be seen in Figure 6. The BE for the W-O bonds and the AC for the outer and bridged oxygen atoms are shown for PW at the 1eft of Figure 6. The outer W-O bonds possess a larger BE than that shown by the bridged W-O linkages, but two of the latter must presumably be severed to release an oxygen atom. On substitution of one molybdenum for one tungsten atom, a splitting of the bridging W-O BE results, since one of the bridging oxygen atoms becomes coordinated with both a molybdenum and a tungsten atom. Two Mo-O PE are also evident with that for the outer oxygen atom being the 1arger. In genera1 the BE for the W-O bonds are significantly larger than those found for the Mo-O bonds. While the value for atomic charge also splits, the charge on the outer
245
oxygen atom attached to the molybdenum atom is considerably more negative than that bound to the tungsten atom. Further sUbstitution of tungsten by molybdenum produces results consistent with the aforementioned general conclusions, although it is to be noted that the changes are not monotonic. For example the magnitude of charge on the outer oxygen atoms is predicted to be a maximum when only one molybdenum for tungsten substitution has occurred. CATALYTIC PROPERTIES Earlier work from this laboratory has shown that 12-tungstophosphoric acid (H3PWl2040) and 12-tungstosilicic acid (H4SiWl2040) are active and selective catalysts in the conversion of methanol to hydrocarbons (9,10). It is important to note however that unlike the ZSM series of catalysts (11) no significant quantities of cyclic hydrocarbons were found in the products. With HPW at reaction temperatures between 325 and 400·C, greater than 90% of the methanol is consumed. However at small residence times (W/F) dimethyl ether is the predominant product. With increase in residence time the amounts of Cl) hydrocarbons (primarily olefinic) increase, pass through a maximum and, with larger residence times, decrease. At 400·C the principal product for higher residence times is methane. The product distribution was also shown to be dependent on the catalyst pretreatment conditi ons , Pretreatment in helium at 350·C was optimal, but that in hydrogen at 400-450·C was also appropriate. However, pretreatment in air at 400·C almost eliminated the (> Cl) hydrocarbons from the methanol conversion products. Salts of HPW are also effective catalysts for the conversion of methanol to hydrocarbons (10,11). However the product distributions are dependent on the cation. For example, the yield of C4 hydrocarbons was found to increase with decreasing magnitude of charge on the terminal anionic ollYgen atoms, although the yield with the parent acid itself fell well below that expected on this basis. It was tentatively suggested that the strength of acid sites in the solid acid is sufficiently high to promote irreversible chemisorption
(>
leading to the formation of coke precursors. The results obtained for the conversion of methanol on the ammonium salt of 12-tungstophosphoric acid were of particular interest (10,12). Both the activity and the selectivity were found to be higher than that of the parent
246
acids and salts of metallic cations 60~------------' (Fi g. 7). However more interestingly the product distribution with the 50 ammonium salt showed a considerably higher content in saturated hydro40 carbons, as contrasted with the largely olefinic products from the 9 30 parent acid and metallic salts. !!! > It should be noted that the 20 existence of the Keggin structure is apparently a necessary condition for 10 the catalysis of the methanol conversion. When the dime ric form of the anion was employed relatively little conversion was evident (9,10). Although evidence so far avail20 able suggests that the central atom in o the Keggin structure has relatively i&l 10 little influence on the catalytic pro> perties of the heteropoly oxometalates, the nature of the peripheral metal atoms is of considerable importance. For example. 12-tungstosilicic acid has a comparable activity Figure 7: A comparison of (NH~)s PW120~O and HSPW120~O in and selectivity to that of HPW in the the conversion of methanol conversion of methanol. while with to hydrocarbons. Catalysts pretreated in helium at (HPMo) 12-molybdophosphoric acid 400°C for 2 hr; W/F = 246 carbonmonoxide is the primary product mg-cat min/mt He. (9,10). Recent work with the ammonium and alkyl ammonium salts of 12-tungstophosphoric and 12-molybdophosphoric acids in the conversion of methanol and ethanol shows while the alkyl ammoni um salts show high activity in the conversion of methanol the selectivity to Cg-C s hydrocarbons is considerably reduced over that found with either the parent acid or its ammonium salt (14). However in the latter case deactivation is observed over a relatively short period of time. With N~PMo and Me~NPMo ethanol was converted to ethylene, ethane. aceta1dehyde, and trace amounts of di ethyl ether. The catalyst pretreatment conditions for this process were found to have a marked effect on this reaction. In general, with NH~PMo, activation in hydrogen at 623K or higher was
t
247
found to be the preferred method, yielding catalysts which showed little or no change in activity with ethanol reaction time and a preponderance of ethylene in the reaction product. In contrast, with catalysts pretreated in air, acetaldehyde was the primary product. The product distribution from ethanol with NH 4PMo also varies with the activation temperature (in hydrogen), the selectivity to ethylene increasing while that to ethane and acetaldehyde decreases as the activation temperature is increased from 613 to 733 K. The effects of activation temperature on the methyl ammonium salt are similar to those observed with the ammonium salt, although the activity of the former was approximately 25-50% of that of the 1atter. Increase of the reaction temperature was observed to alter the selectivity in the conversion of 7 0 t-- - - - - ----""- - - - - - -+ ethanol. The quantity of acetaldehyde decreases with increase of react ion te""erature, while that of ethylene increases. In contrast the amount of ethane passes through a maximum (Fig.8). The selectivity for 5+' -3 -------.-------4 568 ethane appears to be the result of 623 REACTION TEMP. IKI .... the hydrogenation of ethylene by the hydrogen obtained from the dehydrogenation of the alcohol. The observation of increased ethane and Figure 8: Effect of reaction temperature on the selectivity in decreased ethylene production when ethanol conversion on hydrogen is introduced to the feed NH 4PMo 1 20 4 0 . Continuous flow reactor, W/F 0.49 g.-cat stream provides support for this min m~_l ETOH. Catalyst hypothesis. pretreated in hydrogen at 400°C. (.) Ethylene, Recent work in thi s 1aboratory ( .. ) ethane, and (0) ace(15) with PAS FTIR has shown that the taldehyde. first step in the methanol conversion process on HPW involves the protonat i on of the a1coho1 fo11 owed by the c-o bond scission and the methylation of outer oxygen atoms of the heteropoly anion.
°
PORE STRUCTURES The effect of the cations on the pore structures of heteropoly oxometalates has been shown in recent studies from this laboratory (4,17-19).
248
Ce rta in of the heteropo ly oxometa 1ate salts have been shown to possess micro200,--------------, o PW pores which appear to be intrinsic to the " PMo Ii> SiW structure rather than a ammonium salt of 12-tungs tophosphori c aci d pes sesses an 150 activity and selectivity superior to that _ o of the parent acid in the conversion of ~'" • E methano 1 to hydrocarbons pr-ovi ded the<[ first implication of the existence of a ~ '00 <[ o porous structure probably consisting I-.J largely of micropores. ~ IX Studies of a series of monovalent ii: 50 salts of HPW, HPMo, and HSiW have shown that the surface areas of the parent solid acids and those salts, where the cation such as sodium and silver is small, are o 234 5 also small (Fig.9). As the size of the CATION DIAMETER IAl cation increases the surface area increases so that for the salts of cations such as potassium, ammonium. rubidium and Figure 9: BET surface areas of monovalent salts of cesium, the values for area are as high as H3PW1Z040. H3PMolZ040 2/g. 200 m However, the surface areas of and H4SiW1Z040. the methyl ammonium and tetramethylammonium salts are again small. The adsorption isotherms of the high area salts display a substantial adsorption at low equilibrium pressures, characteri st i c of the presence of micropores. High values ()300) of the C parameter in the BET equation and reduced ranges of linearity. both expected where micropores are present. were also found. In some cases adsorption-desorption hysters t s , characteristic of the existence of mesopores. was also observed. Analyses of pore size distributions was facilitated by applying the Roberts (20) and MP (21) methods for mesopores and micropores, respectively, to the measured nitrogen adsorption-desorption isotherms of the high area salts. Standard isotherms required in the latter case were obtained from a CSET matching process (22) verified by conoar-ison with those obtained for the parent, low area, solid acids. The t-plots for the high area heteropoly salts were found to extrapolate through the origin and to deviate negatively from linearity with increasing relative pressures. the latter as expected where porosity exists.
.
249
Distributions of micropores in the ammonium salts of HPW, HPMo and HSiW show a maximum in the range from 8-10 A and a second maximum centered at 13 A with NH 4PMo (Figs. 10-12). In the latter case the onset of mesoporosity is also evident. Areas calculated from the t-plot and BET methods are sUfficiently in agreement to provide evidence of the internal self-consistency of these methods, although of course no contention that either method is absolute is implied. It woul d be expected, if the surface areas of the high area salts are primarily the result of the existence of microporosity, that the number of adsorbed nitrogen layers woul d be considerably restri cted as compared with the infinite number of layers allowed, at least in principle, by the usual derivation of
80,---------------,
13
PORE RADIUS
6
(A,
Figure 10: Micropore distribution for ammonium 12-molybdophosphate.
60,------------,
60,--------------,
6
6
Figure 11: Micropore distribution for ammonium 12-tungstophosphate.
Figure 12: Micropore distribution for ammonium 12-tungstosilicate.
250
the BET equation. That this is indeed so with the high area monovalent heteropoly o PW 0 salts is evident from an application of o PMo 14 I::. SiW the n-layer BET equation (23) to the adsorption data. The optimumn values 13 were found to be small « 5) for all heteropoly compounds displaying microporosity and to vary linearly with o~ 12 the mean micropore radius, providing E further support for the val idity of the ... 11 micropore analysis (Fig.13). With the monovalent salts of the to heteropoly oxometalates there is a semiquantitative similarity between the 9 variation of micropore volumes and cation di ameter for the three anions of Keggi n 0 2 3 4 5 6 structure (Figs.14-16). However, it is n evident from both mi cropore volurnes and surface areas that the micropore structtures are not independent of the composi- Figure 13: Mean micropore radius and n value calculated tion of the anions. Data obtained from from the n-layer BET equation for the monoX-ray diffraction analysis are also shown valent microporous salts in Figs.14-16. The intensity of the ~
t5~-_N N
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~~ :::>....1
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....1-
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Figure 14: Micropore volumes and XRD [110]/[222] intensity ratios for monovalent salts of H3PW120~O.
CATION
Figure 15: Micropore volumes and XRD [110]/[222] intensity ratios for monovalent salts of H3PMo120~o •
251
[110] reflection relative to that of the [222] plane, the latter of which is usually the most intense reflection, displays an approximately inverse relationship to the micro-pore volume. The intensity is relatively high for salts of all three anions with negligible micro-porosity, but decreases s t gnifi cant ly when the cat ion diameter increases and
I
I
III -: 7.0
~5:!
::) ...J ...J -
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~
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I
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.=.~
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microporosity occurs, again i ncreasi ng as the mi croporosity decreases with sti 11 increase in cation size.
Ag NH: MeNH~ Na+ \K+ !Rb+ c.+
H+
0
0.1
0.2
0.3
0.4
O.?o
CATION DIAMETER (nm)
further
As discussed earlier, Brown and coworkers (2) have shown that
Figure 16:
Micropore volumes and XRD [110]/ [222] intensity ratios for monovalent salts of H4SiW12040o
the secondary structure of HPW contai ns two different i nterpenetrat i ng cubt c substructures (both of whi ch contain cations and anions) which are related to each other by inversion through the origin. In those heteropoly oxometalates where microporosity is absent, interstitial voids, separated from one another by the terminal oxygen atoms of the anions, are apparently present. These voids align in directions parallel with or normal to the [110] plane of the crystal. Consequently the X-ray intensity in the [110] plane appears to be related to the electron density associated with the terminal oxygen atoms. Since the cubic lattice parameter of the heteropoly salts increases with increasing cation size (although that for the parent acids is larger), it is expected that the interstitial voids would be widened and the terminal olO'gen atoms of the anions may be reoriented to generate a continuum of voids and consequently a decrease of the XRD intensity in the [110] plane. Of course pores of size such that nitrogen molecules will not enter will not be measured. A number of factors may be i nvol ved in determi ni ng the exi stence of a microporous structure. Both the size, shape, and nature of the cation may be important. While the size of the methyl ammoni urn salts may be the most important parameter in eliminating the pore structure, the nonspherical shape of these cations cannot be disregarded. Work on salts of divalent cations is currently in progress in thi s laboratory (24). Preliminary evidence suggests that the absence of measurable microporosity may
be
related to
the
deviation
from stoichiometry which
252
frequently occurs with such salts. ADSORPTION PROPERTIES One of the intriguing questions which naturally arises in considering heteropoly oxometalates relates to the ability of sorbate molecules to penetrate into the bulk or secondary structure of these sol ids. As has been noted earlier heteropoly oxometalates are ionic solids and as a consequence sorption into the solid must presumably involve the surmounting of rather substantial potential energy barriers. In this laboratory photoacoustic FTIR spectroscopy has been employed to answer this and related questions (4,25-27). Spectra obtained for 12-tungstophosphoric acid after exposure to successive aliquots of ammonia at 150'C illustrate the observations (Fi g.l7) (4,25,26). Above 3 molecules/KU no further uptake of ammonia was evident, implying penetration into the bulk and formation of a stoichiometric salt. The spectra show development of bands at ~3200 and 1420 cm-1 attributed to the tri ply degenerate asymmetri c stretching (V3) and bending (VII) fundamenta1s of the NH 4 +ion, respectively, and ultimately a spectrum resembling that of the bulk salt. Thus the PAS spectra confirm (a) H:3~2Qu)(473") the observation from adsorption 2400 1600 studies that ammonia is capable of rapid sorption into the bulk salt. Supporting evidence from XRD shows Figure 17: Sorption of ammonia by stepwise dosing (molecules that the cubic structure is retained sorbed/KU) at 150°C on H3PW12040, pre-evacuated while the lattice parameter shrinks at 200°C. to that of the ammoni um salt. Pyridine has also been shown by PAS to be capable of penetrating into the bulk of heteropoly oxometa1ates (27). However, important differences in behaviour were noted as compared to that of ammonia. After exposure of HPW to excess gaseous pyridine at 25'C, a rapid initial uptake, followed by a slow
253
continuous sorption occurs, reaching a limiting value of approximately 6 pyridine molecules per Keggin anion in one hour. The PAS spectrum after evacuation at 25'C displays new bands associ ated with sorbed pyri di ne (17001100 cm-l) and the band envelope (1100600 cm- l) characteri st i c of the Keggi n unit (Fig.18). However, it is evident that the formation of the pyridinium ion is i nhibited. In contrast if the heteropoly acid is dosed in controlled aliquots, the PAS spectrum exhibits strong bands at 1640, 1610, 1537, and 1485 cm-l, characteristic of protonated pyridine. After dosing the stoichiometri c quantity of pyri di ne, bands associated with the pyridinium ion are evi dent but the 1540 cnr' band is suppressed and a band characteristic of H-bonded pyridine is also present at Figure 18: Sorption of pyridine on 1443 cm-l. These and other observations H3PW12040 pre-evacuated were attri buted to the presence of the at 250°C. (PY2H)+ ion the existence of which apparently prevents access of all pyridine molecules to the available protons. However, when the HPW with stoichiometric quantity of sorbed pyridine was heated under static vacuum to 100'C a spectrum similar to that of the pyridinium salt resulted (Fig.18). Thus, although differences do exist between the sorption behaviour of ammonia and pyridine both are capable of penetrating into the bulk of the heteropoly oxometallates and interacting with all protons present. Studies of the diffusion and sorption of various organi c and inorgani c molecules in heteropoly oxometalates are presently underway in this laboratory with the aid of microbalance techniques (28). As expected in view of earlier results (4,17-19) the sorption capacities of heteropoly acids for aromatic hydrocarbons are considerably lower than those of the ammonium salts, but heats of adsorption are similar for both acids and salts. Diffusivities in both the solid acids and salts depend, not surprisingly, on both the size and molecule weight of the sorbate molecules, but also on the nature of the cation and the elemental composition of the anion.
254
NONSTOICHIOMETRY OF SALTS The abil ity of ammoni a and pyri di ne to penetrate into the bulk of heteropoly oxometalates has been described. Consequently it is possible to employ ammonia and PAS FTIR spectroscopy to determine the number of residual protons and hence the nonstoichiometry of a heteropoly salt (4,25-27). For this purpose the peak area ratio A1420/A1080 was measured for various amounts of sorbed NH 3/KU. Measurements of the former for the salt of interest then permitted the calculation of the latter and hence the number of residual protons present. It is of interest to note that of the three salts examined, that of aluminum contained the largest number of residual protons per anion, followed by the ammonium salt and lastly the sodium salt. The presence of residual protons can also be detected by the use of temperature-programmed desorption techniques. It has been found (4-7, ~ 29-30) from TPO on the heteropoly acids that two peaks emerge, both due ~ to water (Fig.19). However, while the low temperature peak is evidently the result of water hydrogen-bonded within the structure, the high temperature peak results from water formed by the extraction of anionic (presumably terminal) oxygen atoms by the protons. Recent work from this laboratory (31) has shown that nonstoichiometric divalent salts also exhibit such high temperature evolution of water the quantities of Figure 19: Temperature-programmed desorption profiles for which can be correlated with the HSPWIZ040, H4SiWIZ040, HsPMoIZ040 after pretreatment residual proton content. at 25°C for 16 hrs.
255
REFERENCES 1 2
3 4 5 6
7 8 9
10 11
12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31
M.T. Pope, Heteropoly and Isopoly Oxometalates, Springer, Berlin, 1973. G.M. Brown, M.R. Noe-Spirlet, W.R. Busing and H.A. Levy, Acta Cryst., B33 (1977), 1038. J.B. Moffat, J. Mol. Catal., 26 (1984) 385. J.G. Highfield, B.K. Hodnett, J.B. McMonagle and J.B. Moffat, in: Proceedings of the 8th International Congress on Catalysis, Berlin, 1984, Dechema, Frankfurt, 1984, P. 611. M.F. Portela (Ed.), Proceedings of the 9th J.B. Moffat, in: Iberoamerican Symposium on Catalysis, Lisbon, 1984, Jorge Fernades, Lisbon, 1984, P. 349. J.B. Moffat, in: S. Kaliaguine and A. Mabay (Eds.), Catalysis on the Energy Scene, Studies in Surface Science and Catalysis, Vol. 19, Elsevier, Amsterdam, 1984, P. 165. J.B. Moffat, in: B. Imelik, C. Naccache, G. Condurier, Y. Ben Tarrit, J.C. Vendrine (Eds.), Catalysis by Acids and Bases, Studies in Surface Science and Catalysis, Vol. , Elsevier, Amsterdam, 1985, P. 17. See, for example, To Hughbanks and R. Hoffmann, J. Am. Chem. Soc., 105 (1983) 1150. ---H. Hayashi and J.B. Moffat, J. Catal., 77 (1982) 473. H. Hayashi and J.B. Moffat, in: R.G. Herman (Ed.}, Catalytic Conversion of Synthesis Gas and AlcoholS to Chemicals, Plenum, New York, 1984, P. 395. C.D. Chang, Catal. Rev. -Sc i , Eng., 25 (1983) 1. H. Hayashi and J.B. Moffat, J. Catal., 81 (1983) 61. H. Hayashi and J.B. Moffat, J. Catal., ~ (1983) 192. 91 (1985) 132. J.B. McMonagle and J.B. Moffat, J. Cata~, J.G. Highfield and J.B. Moffat, J. Catal., 95 (1985) 108. J.G. Highfield and J.B. Moffat, J. Catal., Tn press. J.B. McMonagle and J.B. Moffat, J. Colloid Interface Sci., lQ.!. (1984) 479. D.B. Taylor, J.B. McMonagle and J.B. Moffat, J. Colloid Interface Sci., 108 (1985) 278. ~. Moffat, Polyhedron, in press. B.F. Roberts, J. Colloid Interface Sci., 23 (1967) 266. J. Colloid Interface Sci., 26 R. Sh. Mikhail, S. Brunauer and E.E. Bodo~ (1968) 45. -A. Lecloux and J.P. Pirard, J. Colloid Interface Sci., 70 (1979) 265. S. Brunauer, LvS, Deming, W.E. Deming and E. Teller, ~ Am. Chern. Soc., 62 (1940) 1723. ~. McGarvey and J.B. Moffat, to be published. J.B. Moffat and J.G. Highfield, in: S. Kaliaguine and A. Mabay (Eds.), Catalysis on the Energy Scene, Studies in Surface Science and Catalysis. Vol. 19, Elsevier. Amsterdam. 1984, P. 77. J.G. Highfield and J.B. Moffat, J. Catal •• 88 (1984) 177. J.G. Highfield and J.B. Moffat, J. Catal., ~ (1984) 185. V.S. Nayak and J.B. Moffat. to be published:B.K. Hodnett and J.B. Moffat. J. Catal •• 88 (1984) 253. B.K. Hodnett and J.B. Moffat. J. Catal., 9I (1985) 93. L. Zhang and J.B. Moffat, to be published:-
256
ACKNOWLEDGEMENTS The financial support of The Natural Sciences and Engineering Research Council of Canada is gratefully acknowledged. The kind permission to reprint various figures was granted by the International Union of Crystallography, Academic Press, Inc., Elsevier Sequoia S.A., and Pergamon Press.
257
DISCUSSION B. DELMON : Indeed, the picture you gave was coherent, ... and interesting. My question concerns the extent of the reorganizations that a heteropoly oxometallate might undergo during high temperature (300-550°C) calcination or pretreatment or catalytic work. In one particular case, namely the Bi salt of phosphomolybdic acid (PMo I2), we observed a complete decomposition to BiP04+Mo03, and it turned out that the the cooperation between BiP04 and Mo03 was central in explaining catalytic synergy in formamide dehydration (not the existence of any heteropolyoxometallates). On the other hand, you observed textural reorganizations (micro/macropores), and it is known in solid state chemistry that such reorganizations are often coupled with solid state phenomena like segregation. What information have you, from your work and literature, on the stability of the structure? J.B. MOFFAT: As with many heterogeneous catalysts, structural changes with increase in temperature are important also in the case of heteropoly oxometallates, but alterations in the Keggin structure of the anion are of particular interest. Work in this laboratory, em~oying photoacoustic (PAS) FTIR (1) has shown that in the case of 12-tungstophosphoric acid, for example, the Keggin structure is retained up to at least 450°C. Results from temperature programmed desorption (TPD) experiments (2) show that protons in the same solid begin to be lost as water at approximately 400°C, the oxygen having been extracted from the anion. Deuterium exchange in temperature-programmed reduction (TPR) with this particular catalyst begins at approximately 300°C and reaches a maximum at 400°c (3). The results from TPD, PAS FTIR, and X-ray diffraction measurements suggest that 12-tungstophosphoric acid, for example, is capable of losing a portion of its lattice oxygen without collapse of its primary Keggin structure, but with some slight rearrangement of its secondary structure. Decomposition of 12-tungstophosphoric acid to its constituent oxides, W03 and P205' apparently occurs after heating to 500°C (2). However, the thermal stability of these solids is quite evidently a function not only of the elemental composition of the anion but also of the nature of the cation (1-4). Differential thermal analysis shows that, for example, the ammonium salt of 12-tungstophosphoric acid is apparently stable to 550°C, while that of 12-molybdophosphoric acid shows no evidence of decomposition below 600°C. That the salts of the heteropoly acids are generally more thermally stable than the parent acids is also evident from PAS FTIR, TPD, and TPR results (1-4). The existence of microporous structures in certain of the heteropoly oxometallates appears to be predominantly dependent on the nature of the cation (5). Thus, for example, the sodium salt of 12-tungstophosphoric acid shows no evidence of microporosity while the ammonium salt of the same acid contains pores in the 9-13 A range. These micropores are not dependent on either the presence or absence of a thermal treatment and the Keggin structure is clearly evident in the microporous heteropoly oxometallates. Work is currently underway in our laboratory to correlate the thermal stability of a variety of heteropoly oxometallates with the composition of the anion and the nature of the cation. G.M. PAJONK : In order to give more insight in the microporosity of your solids why didn't you try the Dubinin's micropore filling treatment? J.B. MOFFAT: This is certainly a most constructive suggestion which we intend to pursue. Since the MP method (7) has been well-tested by many workers, this was employed for all of our micropore calculations in the present work. Comparison of the values calculated independently for BET surface areas and for the number of adsorbed layers from the finite-layer BET equation were found to be in agreement with results from the MP method, providing some confidence in the validity of the method.
258
G.R. LESTER: The large range of the C parameter of the BET equation which you report for the salts of different cations of the heteropoly oxometallic acids illustrates the importance of using multiple point N2 adsorption isotherms for the measurement of the surface area and pore structure in catalysis research, and reminds us of the fallacy of the assumption of a single C parameter which is inherent in so-called "single-point surface areas". J.B. MOFFAT: This comment emphasizes the importance of measuring full pressure range adsorption-desorption isotherms, with which we wholeheartedly agree. It is interesting to note, in this regard, that the observation of a wide range of values for the C parameter, with the microporous compounds showing the largest values, can at least qualitatively, be rationalized from the definition of this parameter in terms of the difference between the energy of adsorption in the first layer and that for the bulk condensation of the adsorbate. It may be anticipated that the global attractive forces operating on the sorbing molecule will be dependent on the field of interaction impinging on the molecule. After diffusing into a pore, the molecule will evidently find itself in a field whose magnitude will be reciprocally related to the radius of curvature of the pore. Consequently C values would be expected to be higher with solids possessing microporous structure. 1. J.G. Highfield and J.B. Moffat, J. Catal., 88 (1984) 185. 2. B.K. Hodnett and J.B. Moffat, J. Catal., 88-r1984) 253. 3. B.K. Hodnett and J.B. Moffat, J. Catal., ~ (1985) 93. (1985) 132. 4. J.B. McMonagle and J.B. Moffat, J. Catal.~91 5. J.B. McMonagle and J.B. Moffat, J. Colloid Tnterface Sci., 101 (1984) 479. 6. J.G. Highfield and J.B. Moffat, J. Catal., 95 (1985) 108; ibTO., 98 (1986) 245. - 7. R.Sh. Mikhail, S. Brunauer, and E.E. Bodor, J. Colloid Interface Sci., 26 (1968) 45. J. B.-NAGY: You have explained how the surface of heteropoly oxometallates was methylated. It is quite clear and it is similar to what is known on other supports. Did you carry out experiments on the formation of the first C-C bond and what is the most probable mechanism? J.B. MOFFAT: Photoacoustic FTIR measurements (6) of sorbed species obtained from a variety of alcohols on 12-tungstophosphoric acid have shown that the alcohols are first protonated, then suffer a c-o bond scission, followed by an alkylation of the terminal oxygen atoms of the heteropoly anions. On elevation of the temperature, deprotonation from the alkyl group appears to occur, leading to the formation of what is most probably a carbene intermediate.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
259
PREPARATION OF HETEROPQLYVANADOPHOSPHATE CATALYSTS SUPPORTED BY SILICA AND AN ACTIVE CARBON FIBRE FELT AND THEIR CATALYTIC PROPERTIES Y. KERA, Y. ISHIHAMA, T. KAWASHIMA, T. KAMADA, T. INOUE, and Y. MATSUKAZE . Applied Chemistry Department, Science and Engineering Faculty, Kinki University, Higashiosaka, Osaka 577, Japan
SUMMARY Alkali metal and ammonium heteropolyvanadophosphate(HPA) catalysts supported on silica and an active carbon fibre felt (KF-fel t ) were prepared, and isopropanol decomposi tion over the catalysts was examined as a test reaction. AI~ the alkali metal: HPA-silica catalysts worked stably at 300 -400 C, but tee ammonium:HPA catalysts were deactivated above 300 - 320 C. The activity of the KF-felt-supported catalysts was several times higher than that of the silica-supported catalysts. Both the activity and stability of ammonium:HPA catalysts were greatly improved by the use of a silica-coated KF-felt as the carrier. INTRODUCTION Heteropoly-molybdenum and -tungsten phosphates have been widely used as catalysts for various reaction systems, especially for acid-base and oxidation-reduction reactions(ref. 1). Roozeboom et al.(ref. 2) prepared an effective vanadium oxide catalyst, a so-called monolayer catalyst, by impregnating the oxyvanadium ion in various carriers at pH 4, where it exists stably as isopolyanions. We preliminarily found that the activity of the catalyst deposited with a heteropolyvanadophosphate(HPA) was much higher at low temperatures than that deposited with the isopolyanions. In this work, in order to prepare more effective catalysts, the thermal properties of HPA and the preparation conditions of HPA catalysts were investigated in detail. The HPA catalysts are shown to work more favourably in an amorphous state. Also, it is shown that alkali metals:HPA and ammonium:HPA are more effectively dispersed on KF-felt than on silica, and the activity and stability are greatly improved by previously coating the KF-felt surface wi th sili ca.
260
EXPERIMENTAL Prepara tion of alkali me.~al __ he~El_I.'()poJ'yvanadopho spha tes ~!'!~ltPA, M= Li - Cs) Na - Cs:HPA were prepared according to the procedures of Preuss and 8chug(ref. 3). Na:HPA was initially prepared by the reaction of NaV0 with H 1). K, Rb and Cs:HPA were obtained from 3P0 4(Eqn. 3 the Na:HPA by replacing the alkali metal ion(Eqns. 2 and 3). NaV0 3 + H3P0 4 + HN0 3 ----"'" 3Na20'P205,13V205,nH20 (1) Na:HPA + CH 3COOK ~ 6K;:P·P205'13V205.nH20 (2) Na:HPA + MCI(M: Rb and Cs) --~ 6M20'P205,13V205,nH20 (3) Li:HPA was prepared by adding H to V20 and Li 2C0 solution 3P0 4 5 3 according to the procedure of Hagenbruch and Hahn(ref. 4). Guaranteed-grade reagents from Kanto Chern. Co. Ltd. were used. Preparation of alkali metal and ammonium heteropolyvanadophos£~ate. catalysts supported on 8i0 2 and an active carbon fibre felt Amounts of 0.1 - 1.0 g of Li - Cs:HPA were dissolved in 50 ml of water, and 1.5 g of 8i0 2(silicic anhydride, precipitated; Kanto Chern. Co.), which was previously sieved between 100 and 170 mesh, were added to the solution and stirred for a few hours to impregnate the M:HPA. The concentration of the V(V) ion in these M:HPA solutions and the amount deposited were determined by the usual Mohr titration method. The 8i0 2 powder impregnated with the M:HPA was dried at 110 0C and then calcined at 280 0C and 400 0C for 2 h under streaming air to preparB the alkali metal:HPA catalyst supported on 8i0 2• A 1.5 g amount of 8i0 2Powder previously sieved between 80 and 100 mesh was also added to 50 ml of 2 - 20 gil NH solution, the 4V03 pH of which was maintained at 3 - 4 with H to impregnate 3P0 4, NH In order to impregnate NH in active carbon fibre 4:HPA 4:HPA. felt(KF-felt; Toyobo Co.), four sheets of the felt cut to a diameter of 12 mm were immersed in NH solution. As a trial for 4V03-H 3P0 4 coating the KF-felt surface with silica, KF-felt was immersed for 15 h in a solution suspended with 8umiceram(8umitomo Chern. Ind.; 8-18A), which has usually been used as an inorganic coating reagent, composed mainly of fine 8i0 2 particles, and was then dried at 110 0C. As another trial for the coating, hexamethyl~isiloxane(HMD8; Wako Pure Chern. Ind., guaranteed grade) was adsorbed on KF-felt from the vapour phase and then decomposed at 460 0C for 3 h. Ni metal particles were previously deposited on the KF-felt(ref. 5) in order to decompose HMD8 effectively.
261
Alkali metal:HPA and ammonium:HPA catalysts supported on this KFfelt coated with silica were prepared similarly by the procedure mentioned above. Measurements The thermal properties of the alkali metal:HPA were examined by the use of a laboratory-constructed DTA apparatus. During the temperature elevation, the sample was removed from the furnace for IR, ESR and X-ray diffraction measurements. The IR spectra of the M:HPA were recorded in the frequency range of 400 4000 cm- 1 by the normal KBr disk method using a JASCO Model A-102 spectrometer. The X-ray powder diffraction patterns were measured in the range 26 = 5 - 60 0 by the use of a GF-Rad-oA diffractometer (Cu·K~ radiation and Ni filter) (Rigaku Denki Co.). The ESR spectra were recorded at 77 K in a magnetic field of 2500 - 4200 G using Nippon Denshi Model JES-PE sp e c t.r-om e t e rf Xc ban dj )I = 9.3 GHz and 100 kHz modulation). The surface areas of the catalysts were determined by N2 gas adsorption at 77 K according to a convenient method(ref. 6). The reaction of isopropanol decomposition as a test reaction over these catalysts were examined with a closed cyclic system(0.5 1). The catalytic activity was determined from the total pressure change between 10 and 30 min in Li:HPA the initial stage using an MD-3300 digital manometer(Cosmo Keiki Co.). The vessel containing isopropanol was immersed in an icewater bath during the reaction in order to keep the vapour pressure constant. With the Si0 2-supported ~ catalysts 20 mg were evenly .~ distributed over the bottom of ~ tho reaction vessel and with the KF-felt supported catalysts one sheet of the felt was placed perpendicularly in the reaction tube. The reaction rate is expressed in units of Torr/h. 4000 1500 1000 500 The main reaction was confirmed Frequency/ cm- 1 as dehydration by gas chromato_graphy and mass spectrometry.
Fig. 1. IR spectra of alkali metal:HPA
262
RESULTS AND DISCUSSION IR, X-ray diffraction and ESR spectra of the M:HPA and their thermal properties In the IR and Raman spectra of heteropoly-molybdenum and-tungsten phosphates, the 1050 - 1070 cm-1 bands have been ascribed to the P=O stretching vibration, the 950 - 1000 cm- 1 bands to the Mo=O stretching and the 700 - 900 cm- 1 bands to the Mo-O-Mo or Mo-O-P stretchings(ref. 7). IR spectra Li:HPA of Li - Cs:HPA are shown in Fig. 1. The bands characteristic of the heteropolyanion appear in all of the samples. In fact, the spectrum Hb: for Li:HPA corresponds well with that obtained by Hagenbruch and Cs: Hahn(ref. 4), except the sharp band at 1390 cm- 1• With other samples, however, no comparable spectrum has o 200 400 t/OC been reported. No X-ray diffraction Fig. 2. DTA curves of alkali metal:HPA data for the alkali metal:HPA have been reported. DTA data for the compounds are shown in Fig. 2. Large endothermic peaks accompanied by dehydration appear around 150°C in the DTA curves. However, with Cs:HPA the dehydration peaks shift towards higher temperatures, e.g •• 180°C and 220°C. An exo-thermic peak appears at 330°C 360°C with Li:. Rb: and Cs:HPA. The change in the X-ray diffract-ion pattern of Rb:HPA with temperature is illustrated in Fig. 3. The patterns at 150°C and 280°C correspond to an amorphous state but that at 400°C 10 20 30 40 50 60 to a crystalline state. Hence, 29/deg. the sharp exothermic peak is Fig. 3. Change in X-ray as?ribed to the crystallization diffraction of Rb:HPA with process. A similar change in the temperature
263
(b) Na:HPA
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,
,
2500
.
,
,
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,
t
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,
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4000
Fig. 4. Changes in ESR spectra of M:HPA with temperature. (a) Rb:HPA; (b) Na:HPA X-ray diffraction was seen with other compounds, even if the exothermic peak did not appear. With Rb:HPA a very weak ESR signal appeared only at 280 oC, whereas with Na:HPA a very weak signal appeared first at 280 0C and grew considerably at 400 oC, as illustrated in Fig. 4 a and b, respectively. The g value is estimated to be 1.96, which can be ascribed to the V(IV) ion formed by some reduction process(ref. 8). Surface areas of the deposited alkali metal:HPA and ammonium:HPA catalysts, amounts deposited and surface coverages The surface areas. amounts deposited and the surface coverages of the alkali metal:HPA catalysts supported on Si0 2 are summarized in Table 1. The surface area of the Si0 2 itself sieved between TABLE 1 Surface areas. amounts deposited and surface coverages( B ) of alkali metal:HPA catalysts supported on Si0 2 Catalyst M:HPA Surface Amounts ) Surface b) concentration area deposi ted a coverage ( e ) (mg/ g carrier) ( gil ( m2 / g 0.20 Na:HPA-Si 10 58 22.5 0.10 2.6 102 20.6 K:HPA-Si 0.16 2.6 102 Rb:HPA-Si 27.4 0.16 2.6 90 Cs:HPA-Si 26.9 0.16 Li:HPA-Si 3.6 104 31. 3 a) M:HPA was assumed to be deposited in the V20~ state. b) It was further assumed that the V 0 deposlt~d extends twodimentionally and forms a rletwork over the surface in which the v-o distance is 2.0 ~.
V-o-?
264
TABLE 2 100 and 170 mesh is 240 m2/g and that sieved between 80 and 100 mesh Surface areas of the Na and NH 4:HPA catalysts supported on is 180 m2/g. The surface areas of Si0 2 and KF-felt carriers. the Si0 2-supported catalysts decrease with increasing Catalyst Surface area concentration on impregnation, as 2 ( m /g ) shown in Table 2, in which Na- and Na:HPA-Si0 2 NH-Si-1, -2, -3, -4 and -5 denote catalysts prepared by impregnation· Na-Si-1 191 Na-Si-2 192 with 20, 16, 12, 6 and 2 gil of Na-Si-3 192 Na:HPA and NH solutions at pH 3, Na-Si-4 205 4V0 3 Na-Si-5 213 respectively. NH-Si-1' denotes NH :HPA-Si0 2 a catalyst impregnated with 20 gil 4 NH-Si-1 152 NH solution at pH 4. NH-KF-Si1 NH-Si-1' 153 4V0 3 and -Si2 denote NH catalysts NH-Si-3 157 4:HPA NH-Si-4 158 deposited on KF-felt that had NH-Si-5 167 previously been treated with NH 4:HPA-KF-felt Sum iceram diluted 10- and 100-fold NH-KF-Si1-3 139 with water, respectively. The NH-KF-Si2-3 133 NH-KF-3 surface area of the KF-felt-suppcrted ~~~~~-----__164 ~~ ___ catalyst also decreases with surface pretrea tmen t , Catalytic properties of alkali 100 metal:HPA and ammonium:HPA catalystssuppcrted on Si0 2 I ..t:1 and KF-fel t In order to compare the catalytic activity of normal vanadium oxide with that of heteropolyvanadophosphate, preliminarily we prepared a vanadium oxide (VO) catalyst and an NH catalyst supported 4:HPA on Si0 2, which were impregnated 2.0 1.5 kK/T with 6 gil NH solution, the 4V03 pH of which was maintained at Fig. 5. Arrhenius plots of reaction rates over Si0 2-suppcrted catalysts: 3.0 with HN0 and at 3.7 with 3 (1) NH + HN0 3; (2) NH 4V03 + H Arrhenius plots for 4V0 3 3P0 4• H (3) NH + H2S0 , the isopropanol decomposition 3Po 4; 4 4v0 3 rates over catalysts are shown
265
in Fig. 5. Curves 1 and 2 indicate the results on the VO and NH catalysts, respectively, and the open and closed circles 4:HPA denote the data for the NH catalysts calcined at 280°C and 4:HPA 400°C, respectively. The open and closed squares denote similar meanings for the VO catalyst. A low activation energy is obtained with the NH catalyst pretreated at 280°C, whereas the VO 4:HPA catalyst does not show a variable activation energy. Curve 3 shows the results on the Si0 2-supported catalyst impregnated with NH 4V0 3 solution, the pH of which was maintained at 0.7 with H2S0 The 4, existence of SO~ion has an adverse effect on the catalysis. Arrhenius plots for the 100 isopropanol decomposition rates Rb:HPA over the M:HPA catalysts, listed I Na: ~ in Table 1, are shown in Fig. 6. Rb:, Na: and Li:HPA show a high activity relative to K: and Cs:HPA. Cs' The activity of the Rb:HPA changes clearly with the amount deposited. The activation energy decreased considerably on calcination at 280°C. The results are illustrated 1 I . . - - - L . _ . L -......._...L-_I-...J in Fig. 7, in which the open circles 1.5 2.0 denote the data measured after kK/T pre-treatment at 280 0C and the Fig. 6. Arrhenius plots of closed circles those at 400°C. reaction rates over M:HPA-Si0 2 As an amorphous state formed on catalysts. heating at 280°C, as mentioned above, it is now emphasized 10 that the Rb:HPA catalyst works ~ more favourably in the amorphous ~ H state kept atably at temperature 8 3 below 280°C; the very weak ESR ~ 10 signal at 280°C in Fig. 4a may ~ be connected with such good ~ activity. The initial rates of 2.0 1.5 2.5 isopropanol decomposition over kK/T the catalysts Na-Si-1 - 5 were Fig. 7. Arrhenius plots of reaction rates over Rb:HPA-Si0 2 catalyst: measured in the temperature Changes in activity with amount range 300 - 400°C. Arrhenius deposited and with pre-treatment plots of the reaction rates temp. 1: 3.6, 2: 2.6, 3: 3.6 gil.
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1.4
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1.5 kK/T
Fig. 9. Arrhenius plots of reaction rates over NH - 5 4:HPA-Si0 2-1 catalysts.
are shown in Fig. 8. The slopes of the lines are almost the same and the activity increases with increasing concentration of the Na:HPA on the impregnation, especially between 16 and 20 gil, as shown in curve 1 in Fig. 10. Arrhenius plots of the reaction rates similarly measured over the catalysts NH-Si-1 - 5 are shown in Fig. 9. These data were measured above and below 300 0C after pre-treatments at 400°C and 280 oC, respectively. In the NH-Si-5 catalyst, which was impregnated at the lowest concentration, the plot bends at about 300 oC, and the data at temperatures below 280 0C correspond to a low activation process similarly to the Rb:HPA catalyst. On the other hand, in other catalysts the activity increases sharply at 280 - 300 oC, but deactivation begins above 320°C. The rates at 280°C plotted against concentration are shown by curve 2 in Fig. 10. Hence, in NH-Si catalysts the maximum activity is obtained on the impregnat-ion at a concentration of 12 gil. A pH of the solution of 4.0 seems to be more favourable than 3.0. Catalytic properties of NH catalyst supported on KF-felt and 4:HPA on KF-felt pre-treated with Sumiceram and HMDS KF-felt has usually been applied as an absorber for industrial drains, and has been confirmed to be a excellent material for the electrodes of fuel cells, because of its effective pore structure for the dispersion of noble metal particles(ref. 5). Arrhenius
267
plots of the isopropanol 100 decomposition rates over the catalysts NH-KF-3, NH-KF-Si1-] 1 and NH-KF-Si2-3, e.r e shown as .£: curves 2, 3 and 4 in Fig. 11, ~ H c 50 respectively. The result for f-' <, Q) the NH-Si-3 cataly~t is also +" co given as curve 1 for comparison. ~ Below 280°C the apparent activity 0 of the NH catalyst supported o 10 20 4:HPA on KF-felt is about double that Concentration/g.1- 1 supported on Si0 2• Further, Fig. 10. Relationship between the the KF-felt supported catalyst activity of Na:HPA and NH 4:HPA is not deactivated up to 340°C, catalysts and the concentration whereas the Si0 2-supported on the impregnation catalyst is deactivated at 320°C. When the KF-felt was pre-treated with Sumiceram diluted 10-fold with water, the activity further doubled relative to the 100 untreated catalyst. HowevAr, it was not affected by treatment with Sumiceram diluted 100-fold with water. The activity of the NH catalyst 4:HPA supported on KF-felt also seems to increase sharply qt 280 0 300 e. The apparent activity of the KF-felt-supported catalyst 1 impregnated with 1.0 gil of 2.0 1.5 Rb:HPA was about four times kK/T higher than that of the Si0 2Fig. 11. Arrhenius plots of -supported catalyst prepared reaction rates over various NH 4:HPA at the same concentration, catalysts: comparison of the as shown in Fig. 12. When activity among catalysts supported we used KF-felt that had on Si0 2, KF-felt and Si0 2-coated previously been treated KF-felt. 1: NH-Si-3, 2: NH-KF-3. with HMDS as the carrier, 3: NH-KF-Si2-3, 4: NH-KF-Si1-3. the activity decreased, as shown by curve 3. Curve 4
268
shows the activity of the Ni metal deposi ted on KF-fel t. Hence, as 100 a preliminary conclusion, 1 silica-coated KF-felt prepared ..0; by treatment with HMDS does not h h 0 give any advantageous effect as 10 E-< <, a carrier. QJ +' oj In conclusion, KF-felt treated co with Sumiceram was found to be an excellent carrier for NH 2.0 1.5 4:HPA catalyst. The catalytic ability kK/T of NH appeared to increase Fig. 12. Arrhenius plots of 4:HPA sharply at 280 - 300 0C. reaction rates over various Rb:HPA This change may be related to catalysts: comparison of the the appearence of an amorphous activity among Rb:HPA supported on state similar to that with silica(curve 1), KF-felt(curve 2) Rb:HPA. However, we could and silica-coated KF-felt(curve 3). not confirm this point because curve 4 denotes the activity of the NH has not been Ni-KF-fel t. 4:HPA removed from the solution in a pure crystalline state. Therefore, it is now required to prepare NH and to examine its physico-chemical properties in order to 4:HPA develop its catalytic ability further.
REFERENCES 1 Y. Izumi and M. Otake, Kagaku Sosetsu (Chern. Soc. Japan,Ed.), 34 (1982) 116 - 1412 F. Roozeboom, T. Fransen, P. Mars, and P. J. Gellings, Z. Anorg. Allg. Chem , , 449 (1979) 25 - 40. 3 F. Preuss and H. Schug, Z. Naturforsch., 31b (1976) 1585 - 15914 Von R. Hagenbruch and H. Hahn, Z. Anorg. Allg. Chern., 438 (1978) 273 - 287; ibid., 467 (1980) 126 - 130. 5 1. Kera and M. Yokota, Denki Kagaku (Japan), 52(8) (1985) 601 - 605. 6 V. M. Samoilov and A. N. Ryabov, Kinet. Katal., 19(1) (1978) 250 - 252. 7 C. Rocchiccioli-Deltcheff, R. Thouvenot, and R. Franck, Spectrochim. Acta, 32A (1976) 587 - 597. 8 1. Kera and K. Kuwata, Bull. Chern. Soc. Japan, 52(5) (1979) 1268 - 1274.
269
DISCUSSION B. DELMON : 1/ I suppose your catalyst is very strongly hydrophilic. Have you results indicating that this imparts unusual intermolecular or intramolecular selectivity to your catalyst? 2/ Pyrene is an hydrogen transfer agent, and "oligomers" of pyrene would behave the same way. If we suppose that the strong acidity of your catalyst promotes condensation of pyrene to some coke precursor, the higher the acidity, the larger the quantity of "pre-coke", and, hence the larger the hydrogen-transfer activity. That hydrogen transfer activity would control the flux of hydrogen from Pd to the adsorbed reactant (pyrene "monomer") could explain the correlation activitiy-acidity. What do you think of this possibility? H.P. STEPHENS: 1/ We have not performed any experiments to investigate this possible catalytic property. 2/ Pyrene does behave as a hydrogen transfer agent, but only at much higher temperatures (> 350°C) where the rate of free-radical transfer is high. The experiments described were performed at a much lower temperature (lOO°C) where "pre-coke" is not expected to be formed. In addition, if hydrogen transfer did occur by "pre-coke" it would be expected that the solvent used for the pyrene (n-hexadecane) would be degraded and also products of pyrene condensation would be evident. We observed that the solvent is inert and that pyrene condensation products are not formed. H.A. van't HOF : 1/ Does reduction to metallic Pd require a separate reduction step, or does reduction take place in situ? 2/ What is the temperature stability of the catalysts thus obtained? H.P. STEPHENS: 1/ Because reduction was found to occur rapidly under the reaction conditions used, for these experiments the Pd was reduced in situ. 2/ TGA and DTA measurements indicate that hydrous titanium oxide is stable to 550°C. We have performed hydrotreating reactions with these catalysts to temperatures of 425°C. J. KIWI: Have you tried to substitute the Ti-isopropoxide with another Ti-alko-
xide as a starting hydrolysis reagent, since the size of the alkoxide will determine the Pd-exchange capacity of your catalyst? At which temperature do you dry your catalyst. Do you eliminate the isopropoxide residue totally at this drying temperature?
H.P. STEPHENS: 1/ For the syntheses of the catalysts, we have only used titanium tetra-isopropoxide as the starting agent. However, as described in the text, we performed a series of experiments in which the Ti-isopropoxide was converted to the methoxide by transesterification in methanol, before reaction with the hydroxide. The ion-exchange capacity for these materials is fixed by the ratio or the amount of hydroxide to the amount of titanium as can be seen from Fig. 1-3 in the text. 2 & 3/ The Pd catalysts prepared for this study were dried under vacuum at room temperature (20°C) for 24-48 hours. A typical batch of material contains about 25% volatiles, most of which is water. There are only small amounts of alcohols (methanol and isopropanol). More than 90% of the volatiles may be removed by drying at 100°C and 100% removed by drying at 350°C. R. CAHEN : At high temperatures, say above 200°C, Bronsted acidity will be lost to some extent. Do you have any evidence that your Bronsted acidity versus activity plot remains valid? H.P. STEPHENS: We have performed a few kinetic measurements to temperatures of 300°C which indicate that the higher activities of the acidified catalysts are maintained at elevated temperatures. In addition, we have heated an acidified Pd hydrous titanium oxide catalyst to 400°C, then measured its acidity by the technique described in the text, and found that it maintained its acidity.
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B. Delman, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
271
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
CATALYST PREPARATION VIA HYDROUS METAL OXIDE ION-EXCHANGERS* H. P. STEPHENS and R. G. DOSCH Sandia National Laboratories, Albuquerque, New Mexico, 87185 USA ABSTRACT Preparation of catalysts via a synthesis route which involves the use of inorganic ion-exchangers based on a1koxide derived Ti, Zr, Nb, and Ta hydrous oxides is described. These materials exhibit several unique properties, including Bronsted acidities that may be varied over many orders of magnitude by a facile ion-exchange process and the capacity to uniformally disperse any metal or mixtures of metals. The activities of Pd-exchanged hydrous titanium oxide catalysts, evaluated for the hydrogenation of pyrene, were found to dramatically increase with Bronsted acidity. The most active of the catalysts was an order of magnitude more active than gamma alumina supported Pd catalysts prepared by incipient wetness techniques. A mechanism is proposed to account for the high activity of the acidified catalysts. INTRODUCTI ON The search for new catalysts, comprised of several metallic components and promoters, has created a need for novel methods of dispersing multiple catalytically active components on a support. We have identified a group of hydrous metal oxide ion-exchange compounds of Ti, Zr, Nb and Ta which offer significant advantages for the preparation of catalysts which may be used for a wide variety of reactions. These ion-exchange compounds exhibit a number of properties which make them desirable as substrates for active metals: 1) ions of any active metal or mixture of metals can be atomically dispersed over a wide range of concentrations by an easily controlled process; 2) the ion-exchange capacity of the materials is large, permitting high loadings of active metals; 3) solution chemistry can be used to provide control of the oxidation state of the active metal; 4) catalyst acidity can be modified by ion-exchange; 5) the materials have high surface areas; 6) the ion-exchanger substrates are stable in oxidizing and reducing atmospheres, and over a wide pH range in aqueous solution; and 7) the ion-exchangers can be prepared as thin films on a wide variety of supports. The latter property offers the potential for tailor-made catalysts of chosen chemical, physical, and mechanical characteristics.
* This work
supported by the U.S. Dept. of Energy at Sandia National Laboratories under contract DE-AC04-76DP00789.
272
These hydrous metal oxide ion-exchangers were originally developed at Sandia National Laboratories for the preparation of ceramic materials (ref.1) and have also been applied to decontamination and solidification of aqueous nuclear waste (ref.2,3). Their potential for one catalytic process, direct coal liquefaction, has recently been reported (ref.4,5). The promising results obtained in that preliminary investigation of their use as catalysts, prompted us to begin a systematic study of the chemistry and catalytic properties of this class of inorganic ion-exchangers. Initial efforts have focused on the preparation of Pd-exchanged hydrous titanium oxide (HTO) and its application to the hydrogenation of polynuclear aromatic hydrocarbons at low temperatures and pressures. In this communication we describe the relation of the HTO acidity to its synthesis chemistry, and report the significant impact that acidity has on the hydrogenation activity of the Pd catalysts.
ALKOXIDE-DERIVED HYDROUS METAL OXIDE ION-EXCHANGERS The hydrous metal oxides described here are amorphous inorganic ion-exchange compounds synthesized in the form of salts of weak acids represented by the empirical formula C(M,OJ/;Iz)n, where C is an exchangeable cation and M is Ti, Zr, Nb, or Ta. These materials are made by hydrolyzing the products of reactions between metal alkoxides and alcohol soluble hydroxides. In contrast, group IVB and VB hydrous metal oxide ion-exchangers investigated by other researchers (ref.6,7) are typically prepared by hydrolysis of halides, oxyhalides, esters or salts, and often contain co-precipitated anionic impurities which are difficult to remove due to the nature of the precipitated material. Preparation Alkoxide-derived hydrous metal oxides are prepared by a two-step method, reported in detail elsewhere (ref.1-5). The procedure is briefly illustrated below for the synthesis of sodium HTO, the least expensive ion-exchanger, and the one we have most extensively characterized. Procedures for the Zr, Nb, and Ta compounds are similar. Sodium HTO. In the first step of the preparation, titanium tetraisopropoxide (TTIP) is slowly added to a 10 to 20 weight percent solution o f sodium hydroxide in methanol. (Other alcohol soluble hydroxides can be used, potassium, barium or tetramethylammonium hydroxide, for example, to provide alternative exchangeable cations.) This initiates a complex set of
213
t r a n s e s t e r i f i c a t i o n and h y d r o x y l a t i o n r e a c t i o n s , r e s u l t i n g i n a s o l u b l e intermediate.
I n t h e second step, t h e s o l u b l e i n t e r m e d i a t e i s h y d r o l y z e d b y
a d d i t i o n t o a s o l u t i o n o f t e n volume p e r c e n t w a t e r i n acetone.
The h y d r o l y s i s
p r o d u c t i s a w h i t e p r e c i p i t a t e which, a f t e r d r y i n g under vacuum a t room temperature, c o n t a i n s v o l a t i l e components, p r e d o m i n a t e l y water, a l o n g w i t 4 ininor amounts o f a l c o h o l s .
A t y p i c a l b a t c h d r i e d f o r 24 t o 48 h o u r s c o n t a i n s
25 wt.% v o l a t i l e s and 9 wt.% exchangeable sodium i o n s , e q u i v a l e n t t o a c a t i o n 9
exchange c a p a c i t y o f 4 meq/g,
and has a s u r f a c e a r e a o f a p p r o x i m a t e l y 150 m'/g.
Exchange o f sodium i o n s . HTOs c o n t a i n i n g c a t i o n s o f o t h e r m e t a l s can be e a s i l y prepared b y c o n t a c t i n g t h e Na ion-exchanger w i t h aqueous s o l u t i o n s o f inetallic cations.
I f d e s i r e d , a l l o f t h e Na can b e exchanged.
However, l o w e r
m e t a l l o a d i n g s can be e f f e c t e d b y v a r y i n g t h e amount o r c o n c e n t r a t i o n o f t h e s o l u t i o n , o r t h e s t o i c h i o m e t r y o f t h e Na ion-exchanger.
Contact o f t h e
ion-exchangers w i t h a c i d i c s o l u t i o n s r e s u l t s i n exchange o f hydrogen i o n s , which a l t e r s t h e a c i d i t y o f t h e m a t e r i a l .
The m a t e r i a l s can a l s o f u n c t i o n as
a n i o n exchangers i n a c i d i c s o l u t i o n s , a l l o w i n g a d s o r p t i o n o f m e t a l s which e x i s t as aqueous oxygenated anions. These p r o p e r t i e s o f t h e ion-exchangers
w i l l be d e s c r i b e d i n d e t a i l i n t h e f o l l o w i n g s e c t i o n s . Synthesis chemistry The s t o i c h i o m e t r y o f t h e r e a c t i o n o f TTIP w i t h NaOH i n methanol was i n v e s t i g a t e d b y p e r f o r m i n g a s e r i e s o f t h i r t y r e a c t i o n s r e s u l t i n g i n Na:Ti mole r a t i o s r a n g i n g f r o m 0.0 t o 1.5.
The t r a n s e s t e r i f i c a t i o n r e a c t i o n
i n i t i a t e d b y a d d i t i o n o f TTIP t o methanol r e s u l t s i n p r e c i p i t a t i o n o f more t h a n 90% o f t h e t i t a n i u m as t i t a n i u m t e t r a m e t h o x i d e .
This precipitate
d i s s o l v e s upon r e a c t i o n w i t h h y d r o x i d e . The c o u r s e o f t h e r e a c t i o n was f o l l o w e d b y a n a l y s i s o f t h e supernates o f these mixtures.
The c o n c e n t r a t i o n s o f T i and Na' were determined b y a t o m i c
a b s o r p t i o n s p e c t r o s c o p y (AAS) , and OH- b y g l a s s e l e c t r o d e measurements s t a n d a r d i z e d w i t h a l c o h o l i c NaOH s o l u t i o n s o f known c o n c e n t r a t i o n . A l t h o u g h a l l o f t h e Na'
i s found i n solution, t h e hydroxide concentration i s n i l
u n t i l a Na0H:TTIP r a t i o o f 0.33 i s reached.
Over t h i s same i n t e r v a l , t h e
amount o f t i t a n i u m d i s s o l v e d i n c r e a s e s l i n e a r l y , u n t i l a l l o f i t i s f o u n d i n solution.
Upon a d d i t i o n o f NaOH t o g i v e Na:Ti mole r a t i o s g r e a t e r t h a n 0.33,
t h e h y d r o x i d e i o n b e g i n s t o appear i n s o l u t i o n .
However, t h e amount f o u n d i n
s o l u t i o n i s l e s s t h a n t h e amount added i n excess o f t h e mole r a t i o o f 0.33, i n d i c a t i n g consumption o f a p o r t i o n o f t h e h y d r o x i d e . R e a c t i o n o f OH- w i t h TTIP i s d e p i c t e d i n F i g . 1, a p l o t o f t h e moles o f OH- consumed i n t h e r e a c t i o n p e r mole o f T i as a f u n c t i o n o f t h e amount
added.
The i n i t i a l p o r t i o n o f t h e c u r v e i n F i g . 1 i s l i n e a r , w i t h a s l o p e o f
274 0.5
.., •E .,
0.4
::J
l: 0
0.3
o 1= 0.2 <, I
:x: 0
0.1
0.1
0.2 OH
Fi~ure
0.3
-In
0.4
0.5
0.6
0.7
Added
1. Stoichiometry of the reaction of TTIP with NaOH in methanol.
unity, typical of a quantitative reaction. At a OH-:Ti mole ratio of 0.33 the slope changes, decreasing to zero at a value of 0.5, which indicates the onset of a secondary reversible reaction. Although Fig. 1 shows a maximum of 0.42 moles of hydroxide consumed per mole of titanium at a OH-:TTIP added ratio of 0.5, 100% recovery of the NaOH and Ti is achieved after hydrolysis and collection of the precipitate. Thus further reaction of the hydroxide in solution must occur during hydrolysis. For reactions in which OH-:Ti exceeds 0.5, unreacted NaOH is found in the acetone-water filtrate following the hydrolysis reaction. Composition and structure Based on elemental analysis, valence considerations and the results of the previously described experiments, the simplest empirical formula which adequately represents the Na HTO ion-exchanger for the Na:Ti mole ratio of 0.5 is NaTi 2040H. However, by varying the ratio of NaOH:TTIP for the hydroxide addition reaction, Na HTOs with a continuous range of Na:Ti mole ratios of 0.25 to 0.5 can be prepared. Hydrous metal oxides of Ti, Zr, Nb, and Ta prepared from quarternary ammonium hydroxides, such as tetramethylammonium hydroxide, are very soluble in water. Solutions of these materials have high conductivities, indicative of their existence as ionic moieties in polar solvents. Vapor pressure osmometry and light scattering measurements gave average molecular weights of 5000 to 15000 for these species, which suggests that they are polymeric clusters with empirical formulas such as [Me4NTi205H]n where n ranges from about 20 to 60.
275
Although structures for these complex amorphous hydrous metal oxide compounds have not been elucidated, the solution chemistry, stoichiometry and molecular weight measurements are consistent with hydrated polymeric clusters of metal atoms held together by metal-oxygen and metal-hydroxide bridges. For compounds with Na:Ti mole ratios of O.S, each Ti2040H "monomer" unit has a net negative charge which can accommodate the positive charge of a mobile cation. Our experience has shown that this inorganic polymer appears to have the characteristics of a Donnan membrane (ref.S), a binary polyelectrolyte which exhibits dissociation of the ionic components to yield diffusible cations and fixed collodial anions. In the following text. the ratio of the cation to the titanium-network collodial anionic species will be represented by formulas such as CxT, where x is a fractional number. For example, the ion-exchanger of stiochiometry NaTi 2040H is represented as Naa.ST. Acidity Titration with acid. Because Na HTO behaves as the salt of a weak acid, the acidity of the ion-exchanger can be investigated by titration with acid. Samples of NaO.:>~T were equilibrated with 100 ml of aqueous HCl solutions ranging in concentration from a to 0.1 M. Following measurement of the pH, the supernates were analyzed for Na+ by AAS and for chloride by titration with silver nitrate solution using a chloride specific electrode for endpoint detection. The results of these experiments are given in Figs. 2 and 3. Fig. 2 shows the titration curve. pH as a function of rneq of HCl added per meq of Na+(H+/Na+}. The titration curve exhibits two inflection points, the first at a H+/Na+ value of 0.33, and the second at 1.0. This is consistent with the behavior of the hydroxide addition reaction shown in Fig. land demonstrates the existance of two weakly acidic cation exchange sites. The equivalents per mole of Ti as Nao.sT required to titrate the two exchange sites is 0.17 and 0.33. These values are identical to the ratios of OH-:Ti required by the reversible and quantitative reactions to form the soluble intermediate to NaO.ST. Thus the stoichiometry of the acid form of NaO.ST can be represented as HO.ST, composed of the two weak acids HO• 17T l and HO• 33T2• The stronger conjugate base Ti, is associated with the hydroxide added to titanium tetramethoxide during the secondary reaction to form the soluble intermediate precursor to NaO.ST. The weaker conjugate base T; is associated with the initial quantitative reaction. Fig. 3 shows the behavior of NaO.ST for the exchange of H+ for Na+ (left-hand curve) and for adsorption of HCl (right-hand curve). The cation exchange (left-hand) curve exhibits two discontinuities, consistent with the
276
12 ~---------,
1.0
100
10
..
't:I
tll
8
75
0.75 ]
50
0.5
C
0
s: u
~
pH 6
+ 0 z
4
N
2
13 IT
o
0.5 1.0 1.5 2.0 meq H + /meq Na+
0.25
25
o
oL.--......l_---L_.....L.._...L..----l
E
0
o
2.5
Figure 2. Titration of Na O. 5T with HC1.
.,...0 "-c I
0.5 1.0 1.5 2.0 meq H + Imeq Na+
2.5
Figure 3. Exchange of H+ for Nat (left-hand curve) and adsorption of HCl (right-hand curve).
two equivalence points of the titration curve. At the second discontinuity, hydrogen ion has replaced all of the Na+ on the ion-exchanger, which has been converted to its acid form. At that point, the HO.ST begins to adsorb HC1, as an ion pair. The curvature of the right-hand curve indicates that adsorption of HCl is reversible. For contact with the solution of highest concentration, 0.77 meq/g was adsorbed. In other experiments with phosphoric acid, up to 3.0 mmol/g have been adsorbed. Hydrolysis and acid ionization constants. Upon contact with water, Na HTOs yield alkaline solutions resulting from Donnan hydrolysis, that is, the ion-exchangers behave as soluble salts of weak acids. The hydrolysis behavior of Na O• 5T was studied at ambient temperature (22°C) by equilibrating weighed amounts ranging from 0.5 to 5.0 g with 100 ml of water, followed by analysis of the amount of NaOH in solution by alkalimetry and AAS. Because the large separation of the equivalence points in the titration curve indicates a large difference in the ionization constants of the two weak acids, the hydrolysis constant for the Na form of the weaker acid (stronger conjugate base Til can be determined from these experiments. For the hydrolysis reaction,
1 + H20 ~
T
HT l + OH-
the hydrolysis constant
(1) ~l
can be expressed as: (2 )
277
By noting from the titration curve that the fraction of stronger base sites is 0.33, the ratio HT 1/T1can be calculated from the total amount of exchange equivalents for the experiment, yielding Khl equal to 0.0079 ~ 0.0002. The dissociation constant Kal for HT l was calculated to be 1.3 X 10-12 from values for Khl and the ionization constant for water. The pKa2 for the weak acid form of the second base site T2 is equal to the pH at the point of the titration curve at which the second site is half neutralized, 7.6. Thus Ka2 is 2.5 X 10-8 • The hydrolysis constant for T calculated from the value of Ka2 and the ionization constant for water ,. s 4 .0 X 10 -7 • As noted previoUSly, the acidic form of the ion-exchanger can adsorb acid as an ion pair. For the experiments represented in Figs. 2 and 3, an expression for the equilibrium constant for the reversible adsorption of HC1,
2,
(3 )
can be written as (4)
The denominator of equation (4) contains the chloride, instead of the hydrochloric acid concentration, because of the excess chloride present and the constraints of Donnan membrane equilibrium. From the data of the titration experiments, a value of 2.3 ~ 0.4 was calculated for Kad• The reciprocal of Kad gives an estimate of the dissociation constant Ka3 for HC1·H O• 5T in water. The value for Ka3, 0.36, classifies HCl HO• 5T as a strong acid. Thus on the basis of the three acid ionization constants, Kal = 1.3 X 10-12, Ka2 = 2.5 X 10-8, and Ka3 = 0.36, the acidity of these materials can be varied over many orders of magnitude. Slurry pH as a measure of acidity. A convenient and quantitative measure of the Bronsted acidity of these materials is provided by contacting them with water and measuring the pH of the supernate. This is equivalent to determining the J o acidity function (ref.9), often used for investigation of the mechanisms of acid catalyzed homogeneous reactions. The acidity of ion-exhangers, and catalysts derived from them, ~as determined by measuring the pH of slurries containing one percent by weight of the solid. Table 1 gives the examples of measurements on materials representing a range of compositions and acidities. As can be seen from the pH values, the acidity of these materials can be varied by more than eight orders of magnitude.
278
TABLE 1 ~cidities
of several hydrous titanium oxides. HTO Acidity HTO Formulation (pH) Formulation Nao.ST BaO.2ST Ca O• 2ST* Ni o•2ST* HO• ST·HC1*
11.7
7.9 8.2 9.1
Na O•3T Na O• 2ST Cr O• 17T* HO.ST*
Acidity
--iP!ll 10.4 10.4 8.S 4.6
3.0
* Derived from NaO.ST.
Several observations regarding the adjustment of acidity in these materials can be made from Table 1: 1) As predicted by the value of Kal, the NaO.ST material was found to be the least acidic. 2) Materials of stoichiometry Naa.2ST and NaO• 3T, which contain only one basic site Ti, of larger Ka, were found to have higher acidities (lower pH values). 3) Replacement of Na+ with multivalent cations, either by ion-exchange (eg. with Ca+ 2, Ni+ 2, or cr+3) or by use of another hydroxide in the synthesis (eg. Ba(OH)2)' resulted in material of higher acidity. This is undoubtedly due to stronger adsorption of the multivalent cations on the ion exchange sites, which significantly reduces the ability to accept protons and thereby changes the hydrolysis constant. 4) The material containing adsorbed acid, HC1·HO.ST, exhibited the highest acidity as expected from the value of Ka3•
CATALYSIS BY HYDROUS METAL OXIDE ION-EXCHANGERS Catalyst formulations The reagents and methods used for preparing the Pd catalysts have been reported previously (ref.4,5). Catalysts exhibiting a wide range of acidities were formulated by use of the techniques described in the previous sections of the text, including: 1) modification of the stoichiometry of the Na HTO ion-exhanger (ie •• the Na:Ti ratio), 2) varying the Na:Pd ratio for catalysts prepared from Na o 2ST, 3) use of a Ba o 2ST ion-exchanger synthesized from • +' +2 +2 +2 TTIP and Ba(OH)2' 4) exchange of Na for Ca , Co or Ni • prior 2 2, to exchange of Pd+ S) exchange of Pd+ from acidic solutions, and 6) contact of the Pd catalysts with aqueous hydrochloric and phosphoric acids,
279
after the Pd ion-exchange step. Following preparation of the catalysts, their Bronsted acidity was measured by the previously described procedure. In addition to the catalysts described above, several conventional catalysts, comprised of Pd on various high surface area (100 to 200 m2/g) gamma aluminas, were used to provide a controlled comparison for the activity measurements. These included catalysts from commercial sources and those prepared by incipient wetness techniques in our laboratories. Measurement of hydrogenation activity pyrene hydrogenation kinetics. The catalytic activities of the preparations were evaluated for the liquid phase hydrogenation of pyrene (Py), a polynuclear aromatic hydrocarbon (PAH). Pyrene, a significant component of coal-derived liquids, was chosen for study because of our interest in coal liquefaction process chemistry (ref.4). The catalytic hydrogenation of pyrene (ref.10) proceeds by an irreversible reaction path at temperatures below approximately 125°C. At hydrogen pressures below one MPa, two major products, 4,5-dihydropyrene (H2Py) and 4,5,9, 10-tetrahydropyrene (H4Py) are formed by a series reaction path: (5 )
The kinetics can be accurately modeled by pseudo first-order expressions. For the first step of the hydrogenation: (6 )
where [PYlo and [pYlt are the initial concentration and concentration at time t. The Pd weight basis rate constant kl is a quantitative measure of the catalyst activity. Apparatus and procedure. The techniques used to perform the hydrogenation reactions are only briefly desribed here. A more detailed description is given elsewhere (ref.10,11). Batch reactions with 10% by weight pyrene in n-hexadecane, an inert solvent, were performed in stainless steel microreactors equipped with thermocouples and pressure transducers. Catalysts were added to the reactors without prior reduction, as the Pd was found to be rapidly reduced in-situ under the reaction conditions used. After the reactors were charged with pyrene, hexadecane, and powdered (-200 mesh) catalyst, they were pressurized to 690 kPa and heated to 100°C in a fluidized sand bath (heat-up time < 1.0 min) while being agitated with a wrist-action motion at 160 cycles/min. Temperatures and pressures were recorded with a digital data acquisition system for the duration of the experiments, which
280
ranged from 10 to 30 min. Following the heating period, the reactor vessels were quenched (time of quench < 15 sec) and the reaction mixture was removed for analysis by capillary column gas-liquid chromatography to determine [PY]t' Results. The results of these experiments are shown in Table 2, which describes the catalyst formulations and lists the Bronsted acidity and hydrogenation activities as quantified by pH and kl, respectively. Examination of the data in Table 2 shows the striking range in activity for TABLE 2 Results of activity measurements Catalyst Formulation
~Jt
Pd/Naa.5T
15.8 9.9 4.9 7.7 3.2 7.6 7.3 7.6 7.5 7.1 4.8 3.7 4.3 3.4 8.7 3.0 1.2
Pd/Naa.5T Pd/Naa.5T Pd/Naa.25T Pd/Naa.25T
Pd/Naa .25T'HCl Pd/Naa.25T· HCl
Pd/Naa .25T'H3P04 Pd/Naa .25T· H3P04 Pd/Naa.25 T'H 3P04 Pd/Ni/Naa.25 T*
Pd/Ni/Naa.25 T• HC1* Pd/Co/Naa .25T* Pd/Co/Naa.25 T• HC1* Pd/Ca/Naa .25T* Pd/Baa. 25T Pd/Baa.25T· HCl Pd/alumina-l Pd/alumina-2 Pd/alumina-3 Pd/alumina-4 Pd/alumina-5
Pd
%
0.5 0.5 0.5 3.3 2.9
* Contains less than 0.02 wt. % Na.
wt. % metal
Measured pH
Actiyitr sec- g-
2.3 5.1 7.3 1.4 3.6 1.1 0.2 1.5 1.3 0.7 4.4 2.9 4.8 3.3 6.4 20.7 5.2
7.2 7.6 10.0 7.1 7.7 6.5 4.8 5.9 5.3 4.6 5.8 5.2 5.7 5.2 6.4 6. 1 4.5
0.27 0.10 0.05 0.41 0.24 1.14 2.09 1.07 1.87 2.45 2.03 2.74 1.87 2.21 0.64 0.69 1.61
Co-cation ----
0.09 0.14 0.08 0.10 0.22
281
the various catalyst preparations. Values of kl ranged from 0.05 to 2.74 sec-lg- l, well over an order of magnitude. For comparison, values of kl for the Pd/alumina catalysts ranged from 0.09 to 0.22. Thus the most active of the Pd/alumina catalysts was found to be an order of magnitude lower in activity than the most active Pd HTO catalyst. Many attempts were made to correlate the activities of the Pd HTO catalysts with various experimental parameters, including, for example, Pd loading, type of co-cation, acidity, and type of acid used for post-exchange treatment. Only one parameter, Bronsted acidity, provided a good correlation. Impact of acidity on catalytic activity There are myriad examples of reactions homogeneously (ref.12) and heterogeneously (ref.13,14) catalyzed by Bronsted acids. The authors are not aware of any reports of metal supported hydrogenation catalysts containing mobile protons which contribute to the catalytic mechanism. However a novel method of hydrogenating benzene and polynuclear aromatic hydrocarbons with the aid of liquid-phase super acid catalysts such as HF/TaF 5 has recently been reported (ref.15,16). The mechanism for these super acid catalysts involves protonation of the aromatic ring, followed by hydride transfer from isopentane. Gas-phase hydrogen reduces the isopentyl cations formed, thereby regenerating the isopentane and a proton. This report suggested to us that the increased catalytic activity of the acidified Pd HTO catalysts might be due to a similar mechanism involving transfer of mobile protons from ion-exchanger materials to pyrene. If this hypothesis were true, catalyst activity would correlate well with the measured Bronsted acidity of the catalysts. To test this hypothesis, a Bronsted plot (ref.17) of 10910 (k l) as a function of 10910[H+] (equal to -pH), shown in Fig. 4, was constructed. A good fit of the data to a linear curve (R = 0.89) was achieved. The slope of 0.34, and intercept of 2.02 correspond to the parameters a, and C in the following form of the Bronsted law (ref.17) for general acid catalysis: (7 )
where kl is the reaction rate constant, C is a constant characteristic of the reaction, and [H+] is the measured acidity. Although there are several phenomena which might cause the correlation between hydrogenation activity and acidity, the above observations support the inference that pyrene hydrogenation catalyzed by Pd HTO proceeds via a mechanism which includes a
282
0.5
~
-
0.0
]' -0.5 -1.0
-1.5
L..-_L..-_.l.....-_.l.....-_......_
......_
.....- - '
-4
-6
-10
Figure 4. Bronsted acidity plot for the activity of Pd HTO catalysts. protonation step as illustrated in Fig. 5. The catalytic hydrogenation of pyrene, and other PAHs. on Pd HTO catalysts could be initiated by protonation of an aromatic ring by a proton donated by the ion-exchanger support, followed by transfer of a hydride ion associated with the dispersed Pd metal. The catalytic cycle could be completed by adsorption of gas-phase hydrogen, which is dissociated into protons and hydride ions upon dissolution in the Pd (ref.18).
<$ ~ H2 (
I
Pd
~ I'
~
\
~
~
Q
H HH H
H- H+
) H + HTO -
(
I
I
Pd
0'
( Pd ) H + HTO-
HTO -
I
I
I
Figure 5. Proposed mechanism for the hydrogenation of pyrene on Pd HTO. Thus according to this mechansim, Pd HTO would behave as a bifunctional catalyst, with sites for protonation associated with the hydrous oxide ion-exchanger, and sites for hydride transfer associated with Pd metal.
283
Preliminary experiments with catalysts prepared from Zr and Nb hydrous metal oxides, as well as catalysts prepared by coating films of Pd hydrous titanium oxides on a variety of supports, have demonstrated that these catalysts exhibit a similar increase in hydrogenation activity with Bronsted acidity. Efforts to more precisely characterize the effect of acidity and the nature of the catalytic mechanism, including structural and surface chemistry studies, have recently been initiated.
REFERENCES 1 R. G. Dosch and W. M. O'Neill, U.S. Patent 3,699,044, Oct. 17, 1972. 2 R. G. Dosch, Sandia National Laboratories Report SAND 80-1212, NM, Jan. 1981. 3 R. G. Dosch, T. J. Headley, and P. F. Hlava, J. Amer. Chem. Soc., 67 (1984) 354. 4 H. P. Stephens, R. G. Dosch and F. V. Stohl, Ind. Engr. Chem. Prod. Res. Dev., 24 (1985) 15-19. 5 R. G. Dosch, H. P. Stephens and F. V. Stohl, U.S. Patent 4,511,455, Apr. 16, 1985. 6 A. Ruvarac in A. Clearfield (Ed.), Inorganic Ion Exchange Materials, CRC Press, Inc., Boca Raton, FL, 1982, Ch. 6. 7 M. Abe in A. Clearfield (Ed.), Inorganic Ion Exchange Materials, CRC Press, Inc., Boca Raton FL, 1982, Ch. 5. 8 F. G. Donnan, Z. Physik Chem. (Leipzig), A162 (1932) 346. 9 E. J. King, Acid-Base Equilibria, MacMillan, New York, 1965, pp. 320-323. 10 H. P. Stephens and R. J. Kottenstette, Prep. of Papers, Amer. Chem. Soc. Fuel Div., 30, No.2 (1985) 345-353. 11 R. J. Kottenstette, Sandia National Laboratories Report SAND 82-2495, Albuquerque, NM, Mar. 1983. 12 J. W. Moore and R. G. Pearson, Kinetics and Mechanisms, John Wiley &Sons, New York, 1981, pp. 334-389. 13 K. Tenabe, Solid Acids and Bases, Academic Press, New York, 1970, pp. 103-158. 14 H. Pines, The Chemistry of Catalytic Hydrocarbon Conversions, Academic Press, New York, ]981, pp. 1-155. 15 J. Wristers, J. Am. Chem. Soc., 97 (1977) 4312-4316. 16 J. Wristers, J. Am. Chem. Soc., 99 (1979) 5051-5055. 17 J. W. Moore and R. G. Pearson, Kinetics and Mechanisms, John Wiley &Sons, New York, 1981, pp. 353-356. 18 S. E. Livingstone in J. C. Bailar, Jr., H. J. Emeleus, R. Nyholn, A. F. Trotman-Dickenson (Eds.), Comprehensive Inorganic Chemistry, 3, Ch. 43, pp. 1274-1275.
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B. Delmon, P. Grange. P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
285
PHOTO-ASSISTED DEPOSITION OF NOBLE METALS : INVESTIGATION OF A NEW ROUTE FOR METALLIC AND BIMETALLIC CATALYST PREPARATION J.-t,l, HERRMANN 1, J. DISDIER 1, P. PICHAT l, and C. LECLERCQ2 lEcole Centrale de Lyon, Equipe CNRS "Photocatalyse", BP 163, 69131 Ecully cedex, France 2Institut de Recherches sur la Catalyse, CNRS, (conventionne avec 1'Universite Claude Bernard, Lyon I) 69626 Villeurbanne cedex, France SUMI·1ARY The photocatalytic deposition of noble metals on photosensitive supports is proposed to synthesize well dispersed metal catalysts in mild conditions of preparation. The effects of various parameters governing the kinetics of photodeposition (light flux, solution, concentration, temperature, nature of ions) have been studied. Noble metals, such as Pt, Pd, Ag, Rh, Au, Ir have been deposited on titania. Other photosensitive supports such as oXides (ZnO, Nb 205, ZrOZ, Th02, etc ... ) or sulfides (CdS) have also been employed. Bimetalllc catalysts (Pt-Pd, Pd-Ag, Pt-Ir) on Ti02 have been prepared with a probable alloy formation according to STEt,1 analysis. Although common insulating supports are theoretically unsuitable for photodeposition, a limited deposition of Pt was observed on alumina. A photodeposited 0.5 wt% Pt/TiOZ was used in two bifunctional photocatalytic reactions. It exhibited a higher activity than a well dispersed analogue obtained by impregnation and H? reduction, possibly because the support did not undergo a severe reduction treatment. INTRODUCTION The photodeposition of noble metals on semiconductors has been extensively studied in connection with potential photographic processes (see, for instance, ref(1) and refs therein). It has been previously used for preparing either bifunctional photocatalysts (refs 2-5) or supported metal catalysts (refs 6-8). In this last case, photodeposition was always carried out ~Iith sacrificial additives like acetate (ref.6), methanol and formic acid (ref.B), or various alcohols (ref.?) • In the present study, attempts have been made to prepare noble metal catalysts on transition metal oxide supports by photodeposition without such "sacrificial" agents in order to avoid (i) surface contamination by carbon-containing compounds which can poison either the metal or the support and (ii) undesirable and uncontrolled side reactions. The aim was to obtain as pure as possible metal catalysts prepared in mild conditions without thermal treatment or reduction, which can sometimes be detrimental for the surface properties of the support. Most experiments were carried out on titania because (i) it exhibits the highest photo-efficiency, (ii) it produces the so-called strong metal-support interactions and (iii) it is actually a promising support for interesting
286
catalysts in various reactions such as Fisher-Tropsch synthesis. EXPERIMENTAL Photoreactor 3 The photodeposition was carried out in a Pyrex flask of 100 cm with an optical window of the same material transparent to wavelengths > 300 nm. UVlight was provided by a Philips HPK 125 watts mercury lamp. Materials Most experiments were carried out with Degussa P-25 non-porous titania powder (S = 50 m2g-1 ; mainly anatase). The other supports used are mentioned at relevant places in the text. The metal compounds were pure grade re~qents. Procedure 3) The slurry (support amount: 50 to 500 mg ; volume: 10 to 40 cm was first evacuated for 2 min, then stirred under static vacuum for 15 min and evacuated again for 2 min in the dark. The reactor was thus chromatographically oxygen free. Analyses The kinetics of photodeposition was followed by analyzing (i) the concentration of the remaining cations by atomic absorption spectroscopy (AAS) ; (ii) the amount of deposited metal also by AAS after dissolution of the solids in aqua regia-HF mixture; (iii) the number of protons released in the solution and (iV) the quantity of oxygen evolved in the gas phase by on-line Qas chromatogrnphy. RESULTS AND DISCUSSION 1- Principle of the method The overall equation of metal photodeposition on a semiconductor SC + n n hv ( t4 n+) + -2H20 eSC) ~ MO + n H + T 02(g) can be decomposed in the following steps: - generation of electron-hole pairs by light absorption (SC) + hv ~ e- + p+ (hv ~ E G) where EG is the band-gap energy of SC ; - adsorption of metal cations or complexes on SC (M n+)
--
(M n+)
aq ads followed by their reduction by the semtconductor free electrons e
'":
( Mads + ne-
Wads - simultaneous oxidation of water by the photoproduced holes (H + OH-) + p+---.H+ + OW_H+ + 1/2 H + 1/4 O (g) 2 20 20;tH+ -----?-
(1)
(2)
(3) (4) (5)
287
Zerovalent metal atoms ~,o agglomerate then spontaneo~sly in small crystallites with a mean diameter depending on the nature of the metal and on experimental conditions. This agglomeration may occur by two possible ways, either with metal atoms mW(ads) - + Mm (supported) or via cathodic-like reduction at primary metallic nuclei: n W + Mn+_ (M )n+ ~ W ~ (M +ne-; M0 +M : ... 2 2 3)n+ 3
(6) -+
M (7) m
2- Photocatalytic deposition of various noble metals on titania Under identical conditions (50 mg Ti02 in 10 cm 3 la- 3M metal salt solution), the following reactivity pattern was found: Ag > Pd > Au > Pt» Rh » Ir » Cu = Ni = the corresponding metal salt or complexes being AgN0 3, Pd C1 2, AUC1 3, H2PtC1 6 (or Na 2PtC1 6) , RhC1 3, H2IrC1 6 ' CU(N03)2 and Ni (N03)2' Initially, the metal is deposited as homodispersed crystallites regularly distributed on all the particles of the support. Fig.1 and 2 present micrographs of Pt/Ti0 2 and Ag/Ti0 2 catalysts obtained after illumination for 10 min and which correspond to ~ 1 and 2 wt %of zerovalent metal respectively. Fig.3 shows a micrograph of Ir/Ti02(~ 1 wt %)obtained after 11 h illumination. In this 'latter case, the metal deposition stopped after 7 h of exposure but was not complete. The mean diameter of metal crystallites depends upon the nature of the metal (pt ~lnm ; Ag -v 3-5 nm ; Ir tv 1 nm).
°.
3- Effects of various parameters on the kinetics 3-1 Influence of wavelength and light flux Wavelengths shorter than that of the absorption threshold of the ,supporting oxide are required. For well divided supports, the tail of the absorption spectrum towards long wavelengths enables some photo-activity with photons of energy significantly smaller than the band gap determined for single crystals. On the other hand, the metal ion or complex must not absorb and be decomposed in the same spectral region as the semiconductor, which is generally the case in the near UV region. Measurements performed for Pt and Ag showed that the photocatalytic deposition rate r was found proportional to the light flux ~, at least under initial conditions, for Pt and Ag. This means that the photocatalytic reaction is efficient and dominates the electron-hole recombination (eqn 8), + e + p - - Heat (8) otherwise, r would be proportional to ~1/2. 3-2 Influence of the initial concentration Co The initial deposition rate r o varies with Co according to a Langmuir-
288
Hinshelwood mechanism ( 9) + KCo r 0 = k e = k. 1KCo which implies a true heterogeneous catalytic process, occurring in the adsorbed phase. The metal adsorption sites are probably constituted by the OH groups of the surface, since a decrease in the pH of the metal solution is induced by the addition of Ti02. 3-3 Influence of temperature The variations of r vs Tbetween 273 and 353 K gave only small changes for o Ag and Pt deposition, with an apparent activation energy of 1 - 1.5 kcal mol-I (~ 5 kJ mol-I). The deposition is nearly unactivated as expected for a true photocatalytic process. However, at high temperature, the photodeposition rate can decrease because of a smaller surface coverage by the metal compound, the adsorption (eqn 3) becoming the rate limiting step of the whole process. This was observed above 333 K for silver. 3-4 Influence of air Most experiments were carried out in an evacuated liquid phase reactor to follow the oxygen formation (eqn 1). Some experiments carried out in air showed that r was almost unaffected. This means that the metal deposit is stable in the presence of oxygen above the suspension and that electron capture by 02 chemisorption does not perturb the metal ions reduction. This is obviously of great interest for a practical application. 3-5 Influence of the nature of the metal complex Noble metals are often stable in solutions only as complexes. Some platinum compounds have been compared and exhibited the following reactivity order: IV IV IV II H2Pt C1 6 NaPt C1 6 = Na 2Pt (OH)6 > Pt (N0 2)2 (NH 3)2 One would have expected a higher activity for the last one because of the lower Pt valency according to eqns (2) and (4). The opposite result could be accounted for by a lower surface coverage, possibly linked to the absence of ion city of this complex. 4- Examples of various metal catalysts photodeposited on different supports The feasibility of photocatalytic deposition of metals on a semiconductor SC depends on the relative position of the flat band potential Ef b of SC with respect to the redox potential EO of the Mn+/M o couple. Taking into account this requirement, several catalysts, besides those indicated in the preceding paragraph, were prepared and are listed in Table 1.
289
Fig.2. Ag/Ti0
Fi g.1. Pt/Ti 02
Fig.3. Ir/Ti0
2
2
Fig.5. (Pt-Ir)/Ti0
2
Fig.6. (Pt-Pd)/Ti0
2
290
TABLE 1 Examples of photodeposited catalysts CdS Zr02 Th0 2 Support fuO Nb 20S Pt Rh Rh Metal Pd Ir 0.3 % 0.5 % 1 % Selected wt % 1 % 1 % Case of common metals on Ti02 The incompatibility between E and EO (Ni 2+/Ni o) positions prevents f b(Ti02) the preparation of Ni/Ti0 2 catalysts. For copper, thouqh the positions are favourable to obtain Cu·, the reduction is limited to Cu+ ions which are reoxidised on opening the reactor. The presence of a reducing agent during illumination, either in the solution (CH 3COO-) or in the gas phase (H 2), was inefficient for obtaining metallic copper. This feature illustrates one of the limitations of the method. Case of insulating supports Metal catalysts are commonly supported on wide band-gap oxides (insulators), such as Si0 2 (E = 11 eV), A1 (E > 7 eV) or combinations of both oxides, 203 G G .which are theoritically unsuitable for metal photodeposition with UV light (eqn 2). However, since alumina was found somewhat photo-active for oxygen ionosorption (ref.g) as well as for alcohol dehydrogenation in the presence of deposited Pt (refs 10 ; 11), an attempt was made to deposit platinum with the hope that the weak photosensitivity woUld be sufficient, at least for a small metal loading « 1 wt %). Indeed, UV-illumination of a suspension of Al?03 2 -1 ~ (Degussa, 100 m g ) induced a pH decrease as expected from eqns 1 and S but the release of protons ceased after ~ 1 h. The solid remained yellow which qualitatively indicated that part of PtCl~ions were not reduced in agreement with pH variations. AAS analysis of the supernatant solution indicated that 80 % of the Pt IV ions have disappeared. A calculation based on pH variations showed that 32 % of platinum ions were reduced as metal atoms which corresponded to a loading of ~ 1.2 wt %. This was confirmed by TEM examination which revealed that platinum was present as sm~ll crystallites, regularly distributed on the particles of alumina, with a mean diameter close to 1 - 2.5 nm (see Fig.4). Consequently, the limited photosensitivity of alumina, probably connected with surface sites or defects, enabled a partial photodeposition of platinum (ref.5), This deposition is not photocatalytic but is a photo-assisted stoichiometric surface reaction, since it ceases when the photo-activable species initially present on alumina are exhausted. This limitates the quantity of the metal deposit.
291
Case of a sulfide as support Since sulfides are known to be photosensitive, they can constitute suitable supports for photodeposition (ref.5). This was confirmed by depositing 1 wt % platinum on a CdS sample (Fluka). The interesting point with sulfides is that many of them are sensitive in the visible (CdS: E = Z.4 eV ; A ~ 515 nm). G 5- Photodeposition of bimetallic catalysts Three different metal couples (Pt-Pd ; Pd-Ag ; Pt-Ir) were prepared under similar conditions: ZOO mg Ti02 was dispersed in the required mixture of 10- 3M metal solutions to obtain a loading of 0.5 wt % of each metal, assuming complete depositions. Analysis of the supernatant solution showed that the removal of Ag, Pd and Pt was total in the case of (Pt-Pd) and (Pd-Ag) bimetallic catalysts, whereas, in the case of (Pt-Ir), only 80 %of the Pt and 20 % of the Ir ions were removed. The incomplete deposition of Ir is not unexpected, since this metal has a low photodeposition rate. Moreover, it induces a decrease of the Pt deposition rate, probably because of competitive adsorption. STEM analysis indicated for both (Pd-Ag) and (Pt-Pd) bimetallic catalysts a weight and molar percentage close to the expected values. (Pt-Ir) was not examined by STEM because of the lack of sensitivity resulting from the very high dispersion of the bimetallic phase. A thorough examination of the metal deposits at different locations on the support gave systematically the same metal composition. This is in favour of the possible formation of alloys, whose existence can be easily understood from eqns 6 and 7. In the case of (Pt-Pd) which was prepared from chlorine-containing compounds (H 2PtC1 6 and PdC1 Z)' STEM examination revealed no traces of Cl in the final bimetallic solid, which demonstrates that clean catalysts can be obtained by this method. Transmission electron microscopy showed for (Pt-Ir) very small particles (d ~ 1 nm) well distributed on all the surface of titania (Fig.5) and for (Pt-Pd) a majority of small crystallites (d ~ 1 nm) and some Digger ones of 2-3 nm (Fig.6). On the contrary (Pd-Ag) was poorly dispersed with metal free and metal-rich zones. A shorter illumination time (5 min) with a conversion of about 26.3 % did not allow a better distribution of the bimetallic phase on the support: the metals were present as not well defined crystallites of 2.5-3 nm gathered in spherical agglomerates from 7.5 to 15 nm wide. The impossibility of getting a well dispersed (Pd-Ag) bimetallic catalyst seems linked to the nature of this couple and is still under study.
292
6- Comparative photocatalytic tests To compare the catalytic activity of a photodeposited catalyst and that of a thermoreduced homologue, a typical 0.5 wt % Pt/TiOZ catalyst was chosen, since around this metal content, an optimum of photo-activity was observed for the hydrogen involving reactions we selected: alcohol dehydrogenation (ref. 10) or deuterium-cyc1opentane isotopic mono-exchange (ref.ll). For both types of reactions, the photo-prepared platinum catalyst was more active than its homologue reduced under HZ' For these reactions, the active part of the bifunctional photocatalyst is the semiconductor support, on which occurs the photo-activation process, the metal acting as a co-catalyst necessary to (i) evolve Hz (or HO) coming from the support by reverse spillover or (ii) dissociate Oz into atoms to regenerate by direct spillover the 00 species of the support required for the isotopic exchange. Consequently, two complementary explanations can be proposed to account for the higher activity of the photo-prepared catalyst: (i) a larger Pt-TiOz interface or perimeter due to a greater number of Pt crystallites (higher dispersion) for the same metal loading and (ii) a higher reactivity of the support due to the absence of hydrogen reduction at high temperature during the preparation. This second reason seems confirmed by the smaller difference in the mean activity of both catalysts for propanol dehydrogenation than for isotopic exchange, since in the first reaction the OH population is less affected, because of its possible partial reconstitution by the alcohol. CONCLUSION A photo-assisted method of preparing metal catalysts has been presented. Some limitations or drawbacks have been pointed out: impossibility of depositing common metals on certain semiconductor supports (for instance Ni or Cu on TiO z) ; necessity of using photosensitive semiconductors, thus rejecting insulators although a limited deposition of Pt on A1 z03 was obtained. By contrast, several advantages can be underlined: 1°) Use of mild conditions of preparation (room temperature, aqueous medium) which preserves the support. 2°) Existence of a large number of candidates for metal photodeposition (Pt, Ag, Au, Pd, Rh, Os, Ir, RU ... ) and of possibilities for supports (oxides: TiO z' ZnO, Nb 20S' Ta 20S' ZrO z' Th0 2• snO z' Ceo z' Sbz04 etc ..• or chalcogenides : CdS. ZnS, MoS Z' MoSe Z' MoTe z' etc ... ). 3°) High purity of the resulting catalyst without special cleaning treatment (for instance: absence of C1 on titania for a Pt-Pd/TiO z catalyst issued from chloride-containing compounds ; absence of gaseous poisonous reduction products like HzS in the case of a sulfide support). 4°) Possibility of depositing a metal on a photosensitive allotropic phase only
293
stable at low temperature or easily reducible in more severe conditions. 5°) Possibility of preparing bimetallic (or polymetallic) catalysts with the chance of getting alloys. 6°) Possible localization of the metal deposition. For instance, on a biphasic support, the noble metal will exclusively be reduced on the photosensitive phase. One can thus envisage a selective metal deposition on a supported support. Moreover, on a porous material the metal deposit could be localized in the vicinity of the pore mouths since the deep part of the narrow pores are unaccessible to light. Complex polyphasic catalysts could be designed by this method with the fixation of the metal at selected places, provided they are photosensitive. From the energetic point of view, the use of photons as the activation mode of the process does not seem redhibitory, since substantial advances in photochemical engineering and in UV-(or visible) lamp technology (with light power output> 20 %), have been accomplished, and the energy cost can be compensated by the absence of a subsequent reduction treatment at high temperature.
REFERENCES 1 F. MCJllers, H.J. Tolle and R. Memming, J. Electrochem. Soc., 121 (1974) 1160. 2 J.M. Lehn, J.P. Sauvage and R. Ziessel, Nouv. J. Chim., 4 (1980) 623. 3 D. Duonghong, E. Borgarello and M. Gr~tzel, J. Am. Chern. Soc., 103 (1981) 4685. 4 A. Mills and G. Porter, J. Chern. Soc. Faraday Trans. 1, 78 (1982) 3659. 5 N. BUhler, K. Meier and J.F. Reber, J. Phys. Chern. 88 (1984) 3261 and 5903. 6 W.W. Dunn and A.J. Bard, Nouv. J. Chim., 5 (1981) 651. 7 K.H. Stadler and H.P. Boehm, Proc.8th Intern. Congr. Catal., Berlin, Verlag Chemie, Weinheim, Vol. IV, 1984, p.803. 8. S. Sato, J. Catal., 92 (1985) 11. 9 unpUblished results. 10 P. Pichat, J.-M. Herrmann, J. Disdier, H. Courbon and M.-N. Mozzanega, Nouv. J. Chim., 5 (1981) 627. 11 C.P. Lafrance, S.Kaliaguine. P.C. Roberge and P. Pichat, Studies Surf. Sci. Catal. Vol.19, Elsevier, Amsterdam, 1980, pp. 309-318. 12 H. Courbon, J.-M. Herrmann and P. Pichat, J. Catal., 72 (1981) 129 and 95 (1985) 539.
294
DISCUSSION J.T. SCHWARTZ: My question concerns your analysis by STEM of the bimetallic particles. I assume you are using a high resolution STEM for your analysis; after doing a particle by particle analysis did you then do the same analysis on areas of the support in which no metal particles were found? The reason I am asking this is, very often, there are small metal particles which are not visible on the support, yet give rise to a signal in the X-ray emission spectra. Thus by scattering effects one might incorrectly assign a particle as bimetallic when in fact it was not. J.M. HERRMANN: Yes, of course, the bimetallic particle analysis was carried out with a high resolution STEM (V. G. HB 501 Model), whose probe size is of the order of 0.5 nm. Analysis carried out on each particle showed the simultaneous presence of both metals in proportions close to those expected. The same analysis, carried out under identical conditions, on the support in proximity of the metal deposit, detected no metals but only titania (titanium signal). G.R. TAUSZIK : 1/ Is the thermal stability of these catalysts comparable to that of conventional catalysts ? ' 2/ Since you propose that reduction of the metal takes place on the primary metallic nuclei, have you taken into consideration the possibility of depositing the metals in sequence, in order to achieve bimetallic particles with a controlled surface composition? J.M. HERRMANN : 1/ Some of our Pt/Ti0 2 catalysts were heated at 300°C in hydrogen and TEM examination gave the same textural aspect. 2/ Indeed, if a subsequent reduction of the metal takes place at primary metallic nuclei as we proposed for a possible explanation of the development of metal particles, it would be possible to sequentially deposit a metal on a previously deposited particle of another metal. But only experience with a well-adapted analytic tool can definitely answer this question. D. REINALDA : Many noble metal compounds are light sensitive. Have you studied the spectral dependence of the effects observed to ascertain that it is only due to activation of the support and not of the metal compounds ? J.M. HERRMANN: The noble metal comp~exes, of course, absorb in the visible. For instance, we observed that a 10- MH2PtC16 solution exhibited a high absorption peak at 420 nm and no transmittance below a wavelength of 380 nm which just corresponds to titania's absorption threshold. However, when the concentration was decreased from 10-2 Mto 10-3 M, a substantial transmittance for A > 310 nm was observed. With an optical pathway of 1 em solution, this transmittance is equal to 40% at 385 nm. Practically, since the photodeposition reaction occurs on the Ti02 grains when they are close to the bottom optical window of the photoreactor, the light absorption by the reactant is much smaller than that of the powder. Chromatographic analysis of the gas phase indicated a homogeneous photochemical reaction less than 1% after illuminating for several hours. Consequently, we are actually dealing with a true heterogeneous process, whose parameters have been studied in details for platinum (J.M. Herrmann, J. Disdier and P. Pichat, J. Phys. Chem. 90 (1986), to appear in Sept. issue). R. BICKLEY: 1/ I presume ,that you are referring to the more simple salts of copper and nickel e.g. N031·s, etc .•• when you indicate 'that these metal ions cannot be reduced on the Ti02 surface. Have you tried the copper or nickel salts of organic acids where it may be more feasable ? 2/ Do you think that the sensitising of A1203 and Si02 by adsorbed dyes (for example) would facilitate the deposition of noble metals on these solids?
295
J.M. HERRMANN: II We found no photodeposition of Ni nor Cu on titania. It is not surprising for the former metal since Ni has an unfavourable redox poten-tial with respect to the flat band potential of titania. For copper, which has been found to photodeposit on Ti02' at least partially (see H. Reiche, W.W. Dunn and A.J. Bard, J. Phys. Chem. 83 (1979) 2248), we have never been able to deposit it, even by adding an organic acid, like acetic acid, which could have helped to reduce andlor stabilize metallic copper nuclei or crystallites. 21 Sensitizing A1203 or Si02 by an adsorbed dye is a good suggestion but I am afraid that the support (and the metal) would be subsequently polluted by carbon-containing residues, which is detrimental for a clean preparation. Note that presently the solids are rather clean after full deposition and washing, since, for instance, no Cl was detected on (pt - Pd)/Ti02 from chlorinated complexes. J.G. van OMMEN: You suggest to put Ti02 on A1203 or another-insulator in order to photodeposit the metal on the top of it. 1/ Do you have experience in doing this? 2/ Don't you think that the properties of the metal on the top of Ti02 on for instance A1203 will be different from the same metal on A1 203 ? J.M. HERRMANN: II My suggestion of depositing Ti02 or other photosensitive supports on an insulator support in order to photodeposit the metal on top of it, was given to illustrate and to take advantage of the selectivity of the photocatalytic deposition of noble metals. We have not yet experience in doing it, since we neither found nor prepared well characterized TiOZ/support solids. 2/ Concerning the second part of your question, I am sure that the properties of the metal on top of finely divided Ti02 particles deposited on a support will be strongly modified as compared to those of the same metal deposited directly on alumina, because titania provides always special metal-support interactions, either weak or strong. Moreover, since titania is present as small islands, its limited quantity will give rise to different intensities in those interactions. B. DELMON : To what depth in a catalyst pellet could the attachment of catalyst particles by photo-assisted deposition take place? J.M. HERRMANN: The photocatalytic deposition of noble metals depends simultaneously upon 3 factors: the presence of adsorbed metal salt, the absorption of light and the mobility of the (photo-)electrons necessary to the metal reduction. The first point depends on the texture, whereas the two last ones are inherent to the nature of the photosensitive solid. In a semi-infinite single crystal, light absorption depends on the wavelength and follows BeerLambert's law. On a powder sample, there are multireflections on various particles which allow light to penetrate several layers of individual particles. For instance, a calculation has shown, from the saturation of light absorption in an uncompressed fixed catalytic bed of powdered pure titania (Degussa P-25, a - 25-30 nm) working in various gas phase oxidation reactions, under4identical conditions of illumination, that light penetration depth was - 4 10- cm, corresponding to about 150 layers of Ti02 particles. In a pellet, this depth of light penetration will vary conversely to the compression used. We have no experience on such a relationship. But it seems reasonable to think that most of the reduced metal will be located in the external part of the pellet.
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297
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
ThE PREPARATION AND ThE CH1\P.ACTBRISll.'IICN OF GXIDE PIiOTOCATALYSTS. 1
R.1. BICKLEY , '1'.
2
GONZALEZ-C1IRRU;C
SOlJlE
TERNARY
TITlilllm~
3
l;,lJC L. PALmSANO •
1
School of Chemistry, University of Bradford, BD? lDP, U.K. Instituto de Fisica-Quirr~ca Mineral, C.S.I.C., Madrid 6, Spain. 3 Istituto di Ingegneria Chimica, Universita di Palermo, Palermo, Italy.
2
SDrolll.ARY Ti0 2 catalysts, of moderate surface areas and minimal porosities, containing various nOffiinal quantitiA~ of iron have been prepared. E.s.r spectra indicate that Fe 3 is "'ell dispersed in coprecipitated samples, and in impregnated Tioxide samples, when the iron concentration is <2% and the firing temperature does not exceed 923K. Reflectance spectra shew increased absorption of visible light for all specimens relative to pure titania, a dependence upon increasing iron content and upon the firing temperature. Larger iron concentrations exhibit characteristics which suggest the presence of additional phases at the highest temperatures. INTRODUCTION Titanium dioxide, in both the anatase and the rutile crystallographic forrr~, has long been known to possess photocatalytic activity, a characteristic which, in its most conunonly used roles as pigments in paints, in paper and in plastics, is undesirable (ref 1). The ability of Ti0 2 to capt.ur c.. uv v, photons and to transmit the excitation energy to chemical reactions has been studied extensively (refs. 2 & 3). Intense interest in this topic was stimulated by the work of Fujishima and Honda (ref 4), when they demonstrated the photo-splitting of water. Associated develop~ents with other photocatalytic reactions have identified important commercial areas in which photocatalysts will De of great value, e.g. in the synthesis of fine cheF.icals for pharmaceutical uses, in agriculture, and in controlling environmental pollution in the atmosphere and in waste water effluents. One of the major limitations in photocatalysis is the relatively low values of the overall quantum efficiencies, combined with the necessity of using near ultra-violet radiation. Some success in enhancing the efficiencies of photocatalysts has been
298
achieved through the use of finely divided metals supported upon e.g. Pt, Pd, Rh, II' and Ni. Further benefit would be Ti0 2· deri ved if the need for u.v. radiation could be avoided whereby solar energy in the form of visivle light could provide an inexpensive alternative.
It is with this latter objective in view that
the present work has been undertaken with a view to producing a photocatalyst which will respond effectively to visible light and which can be prepared in a high degree of dispersion in order to optimise its light collecting properties. The photocatalyst sy s t.e m, which has been examinea a n the present study, involves t:-le dispersion of iron (as Fe 3+) in Ti0 matrices.
2 The choice of syster. has been dictated by our interest
in the photochemical fixation of dinitrogen (ref. 5). Moreover the solid state chemistry of the system, Fe 3+/Ti0 is interestin~, 2, per se, because of the irreversible crystallographic change from the anatase phase to the rut.ile phase encountered at the elevated preparation temperatures, a reported difference of solubility of iron in these phases, and the f ozmat.Lon of the compound, pseudobrookite, Fe 2Ti05 (ref.
6).
IJlE'l'HODS Preparation of the Catalysts Three series of photocatalysts have been prepared, each containing nominal concentrations of iron, (expressed as percentages of the total cationic contents of the ing to 0.5%, 1.0%, 2%, 5.0% and 10.0%. been prepared by coprecipitation (CP)
matrices.) correspond-
~'iC2
One series of solids has of the appropriate quantities
of ferric hydroxide and ti tani urn hyc roxt de ,
while tvlO further
series of solids have been prepared by impregnation of commercially availaDle powdered forms of Ti0 2 ('Iioxide CL/DD/1628/2 and Degusse. P25) using ferric nitrate solutions. ~hroushout the subsequent discussion, these solids will be identified by the following types of codes:-
e.g. TF/CF/0.5/773, TF/TCx/1/923, TF/Dg/5/1273 etc.
to indicate their method of preparation, the nominal content of iron and the temperature at which the catalyst has been fired in air for 24h. (i)
Catalysts prepared by Coprecipitation (TF/CP)
Iron and titanium hydroxides were coprecipitated in the appropriate amounts by reacting an aqueous solution of TiC1 3 (15 wt% Carlo Erba) containing the required quantity of Fe 3 + ions (ex
norr~nal
299
Fe(N03)3.9H20, Merck) with an aqueous solution of ammonia (25 wt% Merck) , which was added drop-wise at room temperature with vigourous stirring owing to the exothermicity of the reaction.
After
standing for 24h. at room temperature, the solid was filtered and washed repeatedly for several days to remove residual Cl- ions (Tested as AgCl(s».
The resulting solids were dried at 393K for
24h. before aliquots were fired in air (24h) at 773K, 923K, 1073K and 1273K respectively. (ii)
Catalysts prepared by Impregnation
(TF/Dg and TF/Tox
series) The commercial Ti0 2 preparations were added at room temperature with stirring to aqueous solutions of ferric nitrate containing the required amounts of iron in the minimum volume of bidistilled water.
After standing at room temperature for 48h., the liquid
phase was evaporated at 393K for 24h., after which aliquots of the dried-solids were fired in air for 24h at 773K, 923K, 1073K and 1273K respectively. Characterisation of the Catalysts Sixty separate catalysts emerge from the three distinct methods of preparation, the five nominal c0Lcentrations of iron, and the four firing temperatures.
Each solid has been subjected
to characterisation by some (or by each) of the following methods:a)
Surface Area measurement (BET method) using a dynamic method
of Argon adsorption at 78K. b)
(ref. 7)
Temperature Programmed Desorption T.P.D.
(5K.min- 1)
in a
greaseless high vacuum system, using a temperature programmable furnace (Stanton Redcroft Ltd.) and a small magnetic deflection mass-spectrometer (Vacuum Generators Micromass 2A) to identify the desorbing species. c)
Diffuse Reflectance Spectroscopy at 300K in the wavelength
range 380-700 nm.
(Macbeth MS 2000) using BaS0 4(s) as the refer-
ence. d)
E.S.R. using a Varian 4500 Spectrometer operating at X-Band
frequency (9.4 GHz.) RESULTS Surface Characterisation by Adsorption/Desorption Methods a) Pure Ti0 2 Specimens The surface areas of a specimen of pure Ti0 2 prepared from the ammoniacal hydrolysis of TiC1 and 3,
300
of twc ccrme r cf a.l Ly prepared. s ar-pLe s we re measured to assess the influence exerteo by the added iron during the firing of the (continued overpage)
.' ·0".
..
40
/.,'~"
'
.. 0" .
0 · · ' / .....- - -
_
(
.
I
o '
/
'
/ __ ........
I
-;r:
.
·c",""
')' ,
'.
',/'
<, '"
,,_/
~'
I ,I 1--- ,'. I ' ,
,
2
6
' •
--.. -I
I
20
373 /
/
473
573
673
773
T/K Surface Areas (0, Total; • "External"; Left Ordinate) and Water 'I'.P. D. Profile (---; Right Ordinate) of Ti0 2 (Degussa P25) as a function of the Outgassing Te~perature. .1.1.<;. 1.
'I'ABLE I
~ T/K
773 923 1073 1273
0,50/ 0 27,2 (27,0) 29,9 (22P) 16,6(14,5) 17,3 (14,9)
1,0%
2,0%
5,0%
10,0%
57,0(37,3) 24,2 (19,9) 33,6(27,6) 54,4(41,8) 30,0(26,0) 33,0(247) 29,5(22,2) 38,9 (29,2) 10,6 13,6 12,0 9,4 8,7(7,2) 7,9 (6,7) 7,9 11,4 1,6 2,1
3,1 0;;;1
2,9 E;;:1
3,2 E;;:1
E;;:1
~1
~1
~1
.E;;:1
~1
.E;;:1 E;;:1
~1
2,4 7,6
E;;:1
Surface Area measurements of Fe/Ti0 photocatalysts prepared by 2 the Coprecipitation ~lethod (cr , upper figures) and the Impregnation of Degussa P25 (Dg, lower figures) as a function of the nominal iron contents and the firing temperatures. Figures in brackets refer to "External" Surface (see text for procedure) • All specimens were outgassed initially at 423. K.
301
Fe/Ti0 2 preparations.
The TiCl 3 hydrolysate, after firing at 773K, the Tioxide specimen (CL/DD/1628/2) and the Degussa P25 specimen were found to possess specific areas of 38.5 m2g- 1 1.9 m2g- 1 and 44 m2g- 1 respectively.
The commercially prepared
specimens were used subsequently as starting materials for the preparation of the impregnated catalysts. porosities was obtained by
An estimate of their
a second argon adsorption
~easuring
isotherm at 78K, after outgassins the solids for 30 mi.nuce s at 78K following the completion of the initial adsorption isotherm. Desorption of argon from
pores at 78K is slow and the sur-
narr~l
face area which is determined from the second isotherm may be significantly smaller than the value obtained from the initial measurements.
The Tioxide specimen had no porosity, whereas the
results shown in Fig. 1, for Degussa P25, reveal unusual variations between the two surface areas as a function of the outgassing temperatures.
Also shown in Fig. 1 is the T.P.D. profile of water
desorbing from Degussa P25.
'Ihe development of porosity with
increased outgassing temperatures appears to be related to the loss of adsorbed molecular water in the temperature range, 373-573K. Loss of water at higher temperatures from the interactions of surface hydroxyl groups (ref. 8) at
600K and 700K respectively
corresponds to the commencement of sintering.
For the Ti0 2 specimen prepared by hydrolysis, the second adsorption isotherm gave a specific area of 30.7 r;,2 g-1, wh i ch again indicates the existence of internal surface. b) Iron-containing Ti0 2 specilclens Each of the specimens was characterised in a manner analogous to that used for the pure specimens of Ti0 2• Data for the coprecipitated specimens (TF/CP) and for the impregnated Degussa specimens (TF/Dg) are given in Table I.
Marked variations in the total surface areas of the
coprecipitated specimens, fired at 773K, as a function of the nominal iron content (upper figures) are not shown by the Degussabased specimens (lower figures).
In general there is a
loss of surface area (and porosity) as the firing temperature increases for all compositions.
The surface areas of each of the
7ioxide-based specimens was 10vI, but the relative undulatory variation in area, as a function of the increasing iron content for samples fired at 773K, was similar to that for the Degussa series, there being a maximum area of 2.7
In
2 -1
g
at 1% Fe
3+
.
302
Diffuse Reflectance Spectra.
Spectra of the solids containing
0.5%, 1.0% and 5% of iron, as a function of the firing temperatures, are shown in Fig. 2.
Spectra of 2% and 10% of iron in
Ti0 2 have been omitted from the present discussion because they reveal only changes which may be expected by considering them in relation to the 1.0% and 5.0% systems respectively.
With every
specimen containing iron, the absorbances are increased relative , to the pure Ti0 2 specimen at the same wavelength. Absorbances also increase with increasing iron content at a fixed firing temperature, although this effect is less evident at 1273K. The variations in the spectra of solids, with fixed nominal iron contents, as a function of the firing temperatures are complex; some apparent increases in absorbances arise through the increasing particle sizes
(~
10% R'oo maximum), but larger changes in
absorbances must be attributable to distinct surface chemical differences between the specimens, which, in the impregnated speci3 mens, is probably related to the rates of diffusion of re + from the initially iron-rich layer into the bulk structure of the Ti0 particles.
2
At 1273K, the large absorbances are almost indepen-
dent of the nominal iron contents with respect to the wavelength, suggesting that the surface layer of the particles has a composition quite different from that attained at lower firing temperatures. Electron Spin Resonance Spectra In general the e.s.r. spectra of the iron-containing specimens are complex.
The spectra of selected specimens are shown in
Fig. 3, from which varicus features can be identified.
The
resonance absorption on the high field part (right) of the Figure corresponds to the transitions for which
nm s
ponding g-value bring approximately equal to 2.
1, with the corresThe prominent
features in the low field region are due to the half-field transitions for which nm = 2. It is noteworthy th2t the spectra of s TF/CP/0.5/773, TF/TOx/0.5/773 and TF/Dg/0.5/923 have several features in common and they
reserr~le
spectra of TF/CP/l and 2/773
and TF/TOx/10/773 which are not shown.
TF/CP/5.0/773 contains
less spectral features than with the lower concentrations of iron in this series, and the transition, ent.
nms 2, is much less prominA marked change in the spectrum of this specimen is created
by firing it at 1073K. The original features, present at lower temperatures, are now dispersed under a very strong broad signal
303
A/nm
600
500 .~
o
\
o
\
\ \
A
•
A.
D
\ \ \
,,
\
60
o
...
..
C) ~
..
0
'"
• '"
•
o
•
A •
o
'-__
•
TF/CP <»
II
.
<:>
.11 ~ ()
..
<:>
.-
o
0
C)
" A
---"'___
+ II
•
" A
--_
II
0
•
•
"
• A
0
()
•..
o
o
_ _
11+
••
0
A
0
--------_E__ 600
500 TFlDg
-. •
, ", 40\,
%R~
0 0
, \
0
A
o
\
\
.~
, "
\ \
,
\
-,
4aO
...
~+<»
•
0•
• '" o
[I
A
•
Q
A
\
0
\
\
-40
II
\
0
• •
+ <»
A
• ", Q
A
0
-.- -
II
\
•o A" 600
[I
• ••
TFITOx
+ A
•
0
Q
.. • •
•-",
A
0
\
•
II
0
\
0
\
\
\
, ...
II
<»
0
.'" OA
... -
'"
--
A
0
\
\
60
'" ..
500
O. 4 0'"A Q \
00
.
0
--- --- ----- ---- - -
20\ \ \ O/oR'
'"
OAD
\
60
-. II
•
A
o
'" ..
II
--- - -- - ------ - -.'--
• A
[I
• • 0
rf 0
- - -- - -
Fig. 2. Diffuse Reflectance Spectra of Fe!Ti0 2 photocatalysts containing 0.5%, 1.0% and 5.0% Fe 3+ as a functl0n of the firing temperatures (for key to spectra see overpage).
304
Key to Fig. 2. %Fe 0.5% 1. 0% 5.0% with
~h
773K
923K
1073K
1273K
~ ()
II
IJ. 6 &
0
[J
(peak to peak) %2kG.
F~
•
--- Pure Ti0 2
~
anomaly to the characteristics
aescribed previously is presented by the specimens, TF/Dg/0.5 5.0/773, which exhibit much stronger resonances at g % 2 than for any of the other specimens, wh L le retaining similar intensities for the half-field transitions to the others; line shapes are quite different.
in acdition the
It is to be noted however that
by increasing the firing temperature to 923K brings the spectra of
these specimens into closer register with those of the two other preparations.
TFICPIO,51773
TFITDxIO';J/773
TFA: PIS,01773 TFIDg/5.. 0L773 -r : .... - ...-
1kG
,, ,,
Fig. 3. ESR Spectra of Fe/Ti0 photocatalysts showing variations 2 arising from different p rep ar at.Lve methods and firing conditions. DISCUSSION AND CONCLUSIONS At 773K the formation of TF/CP/0.5 and TF/Dg/0.5 results in modest decreases of surface area relative to the respective pure titania. For larger iron concentrations the surface areas of both series of solids exhibit maxima (TF/CP/l, 57rn 2g- 1; TP/Dg/2, 2g- 1) 33m and since the preparative methods differ a single reason cannot account for both rnaxf.r-a ,
'I'he precipitate structure of
'I'F'/CP/l may collapse less easily than the undoped precipitate, whereas the increase of surface area of TF/Dg/2 relative to
305
may be due to a porous surface layer forming on the
~F/Dg/O.5
'1'i0 2 particles.
For the largest iron concentrations the increas-
es of surface areas, and porosities, suggest that these solids are inhomogeneous. The diffuse reflectance spectra of TF/Dg/0.5, fired at 773K, 923K and l073K, are very similar despite there being a f our-ef o Ld decrease of surface area. increases in the
II
The growth of particles should cause
apparent" absorbances, but this effect is appar-
ently compensated by other factors wh i.ch decrease them. e.g. changes of surface concentration of Fe 3 + , and the elimination of surface irregularities (see Table I). of
TF/CP/O.~073
The reflectance spectra
and TF/Dg/O.5/1073 are similar also, while for
the former specimens the absorbances are less and almost superimposable for TF/Cr/0.5/773 and TF/CP/O.5/923.
TF/TOx/O.5/773-
1073 have similar reflectance spectra, as have TF/TOx/l and TF/ TOx/5.
In these solids, spectral variations cannot be ascribed
to changes of particle size, but to possible changes of surface concentrations of Fe 3 + , and to p:::et'('c·-brookite formation at the highest temperatures (ref. 6).
The similarities between the re-
flectance spectra of TF/CF/O.5/1073 and TF/Dg/O.5/1073 are ascribable to each solid being mainly rutile. has been shown to be
~80%
anatase and
~20%
Pure Ti0 2 (Degussa P25) rutile by x-ray diffra-
ction, and its impregnation and firing at 773K produces anomalies in the e.s.r. spectra relative to the other preparations which may be associated with the location of the rutile phase relative to anatase in the particles, e.g. the possibility of a surface layer. No ambiguity exists with Tioxide preparations which are ~99% rutile 3 and dissolution of Fe + is facile even at 773K. However the asymmetry in the high field resonance of TF/TOx/O.5/773 is evidence of stronger magnetic interactions "',dch may j ndicate a limit of solu<0.5% Fe 3 + in rutile. This phenomenon is not apparent
bility of
in TF/Dg/O.5/923 in accord with the suggestion that
iron is more
soluble in anatase, and the conversion of anatase to rutile being slow at 923K.
For larger iron concentrations, fired at 1073K
and 1273K, the strong magnetic interactions are considered as being due to pseudo-brookite, or to the precipitation of iron oxide as the solubility limit of the anatase phase is exceeded during its conversion into rutile. Specimens prepared by coprecipitation, or by impregnation of pure Ti0 2 of high specific area, with iron at concentrations <1%
306
car. prcduce solids which have a uniform convosition, are non-porous, and have moderately large surface areas when fired at 923K; they show also greatly enhanced light absorption relative to pure Ti0 2. With larger concentrations (>2%) the solids, fired at 923K, are probably saturated solid solution supporting surface layers of iron oxide.
The photocatalytic behaviours of these solids are
expected to be quite distinct since the latter are effectively two se@i-conductors in contact.
Specimens fired at the highest temp-
eratures should be inactive because of their low specific areas and from the existence of pseudo-brookite (Fe
2TiOS) ted to be inactive photocatalytically (ref. 5(b».
which is repor-
ACKNOWLEDGEMENTS The authors thank Dr. L. Shields for his help with a part of this work. Thanks are also offered to the following bodies for financial support. The Ramsey l~morial Fellowship Trust, and the Consejo Superior de Investigaciones Cientificas, Madrid (T.G.), the Consiglio Nazionale della Ricerche, Rome, and the British Council (L.P/R.I.B) REFERENCES 1
R.I. Bickley, Photo-induced Reactivity at Oxide Surfaces. Specialist Periodical Report of the Royal Society of Chemistry. Chemical Physics of Solids and their.' El1rfe.ces 7 (1978) 118-1S6 2 H. Formenti and S.J. Teichner, Heterogeneous Photocatalysis. Specialist Periodical Report of the Royal Society of Chemistry, Catalysis 2 (1978) 87-106 3 R.I. Bickley ibid. 5 (1982) 308-331 4 A. Fujishima and K. Honda, Nature, 238 (1972) 37 Sa)R.I. Bickley and V. Vishwanathan ibid. 280 (1979) 306-308 b)M. Schiavello, L. Rizzuti, R.I. Bickley, J.A. Navio and P.L. Yue, Photo-assisted dinitrogen fixation over titania catalysts in a flow reactor. 8th International Congress of Catalysis. Berlin 1984 III (383-394) 6 D. Cordischi, N. Burriesci, F. D'Alba, M. Petrera, G. Polizzotti and M. Schiavello. J.Solid State Chemistry 56 (1985) 182-190 7 H. Bosch and A. Peppelenbos, J.Physics E. 10 (1977) 605 8 G. Munuera and F.S. Stone, Discussions of the Faraday Society 52 (1971) 205-214
307
DISCUSSION J.M. HERRMANN: I would like to make two comments on the very interesting paper of R. Bickley and co-authors concerning the preparation and characterization of Fe-Ti-O photocata1ysts. First, the maximum solubility of Fe 3+ in Ti02 of ar~und 5% seems very probable, since we have found a complete dissolution of Fe + into Sn02 which is known to have the rutile structure. Secondly, if Fe3+ is dissolved into Ti02 structure, it can be considered as a doping agent (acceptor centres). We have observed for Cr3+ that this heterocation acts as recombination centres (Phys. Chern. Lett., 1984). Consequently, I am afraid that Fe 3+ will decrease the photo-efficiency of titania. Since you do not have presently the catalytic results, I formulate the wish that the iron will compensate the loss in photo-efficiency due to recombination by a strong catalytic aptitude to activate nitrogen. R.I. BICKLEY: Your comment concerning the solubility of Fe3+ in Sn02 (rutile structure) is inde~d interesting. Further investigations of the relative solubilities of Fe + in Ti02 (Anatase) and Ti02 (Rutile) are in progress currently, and the solids based upon Degussa P25 (Ti02) present us with an important opportunity to clarifY3this question. In aadition, I am inclined to agree that the presence of Fe +in the Ti02 lattice will create electron acceptor centres but this may also enhance the ability of photo-holes to reach the surface and facilitate the oxidative fixation process. J. KIWI: 1/ Does the Fe203-Ti02 coloured, powdered, semiconductor have photocatalytic activity in the light induced N2 process under study? 3+ 2/ What compensation effects have you observed when introducing 5% Fe into Ti02 ? R.I. BICKLEY: 1/ The iron oxide/titanium oxide powdered materials are currently undergoing tests for their photocatalytic activity in the laboratory of Professor M. Schiavello (Istituto di Ingegneria Chimica, Universita di Palermo). The activity of these materials, initially in a reduced state, is being determined in a flow reactor, using a stream of gaseous N2 containing a small partial pressure of water vapour to produce ammonia under irradiation with near UV/visible light. It is too early to indicate the findings of the kinetic investigations. 2/ In the solid state chemistry of the materials which contain nominally 5% of cations as Fe 3+, ions enter the lattice of Ti02 (anatase and rutile) substitutionally, with the creation of vacant anionic positions in the oxygen lattice (Va = 2 Fe3+). Some evidence suggests that the limit of iron solubility in anatase is T1much greater than in rutile, so that as anatase is converted into rutile when the solids are heated in air at temperatures of ~ 923 K, the solubility limit of Fe3+ in the rutile phase is exceeded, and an iron oxide forms as a separate phase. At the highest temperature, 1273 K, we presume that some of the saturated solution of iron oxide in rutile reacts with the excess iron oxide to form w-brookite, Fe2Ti05' J.G. van OMMEN": This is more of a comment than a question relating to the surface structure of P25 Ti02 of Degussa. From Me-OH oxidation experiments performed on V205 monolayers of P25 (J.G. van Ommen, H. Bosch, P.J. Gellings et al., Proc. 8th International Congress of Catalysis, Berlin. 1984, 4, 297) we have evidence that the surface of P25 is more of the rutile type tnan the anatase type. So we agree with you that P25 is possibly a rutile covered anatase Ti02. R.I. BICKLEY: We are encouraged by your observation about the surface structure of Degussa P25 Ti02' XRD of our specimen P25' by one of us (T.G.). has shown that the material which we are using is - 80% anatase and -20% rutile. Additionally we believe that our diffuse reflectance spectra. after taking account of particle size differences, show that the surface of each particle
308
of P25 (Ti02) is rutile. E.V.W. GRITZ: 1/ What is the photon yield during the photoreaction? 2/ How is the reactivity of the catalyst linked to the different Fe surface species after calcination? R.I. BICKLEY: 1/ The photon yields in these processes are usually very small indeed. Most of the energy absorbed is dissipated thermally through electron-hole recombination processes. Only those photons which produce free charge carriers within the surface space charge layer have any real chance of participating in the chemical processes at the surface. 2/ In the case of these solids, several possibilities exist for the species to enhance the activity. 3In single phase material, (e.g. Fe 3iron + in anatase or Fe 3+ in rutile) the Fe + centres can act as electron traps so freeing photo-holes in the valence band of the solid, and they can provide also alternative absorption centres for reagent molecules. In biphasic materials, the differences in work function (Fermi level) between the phases will enhance the space charge layer of the contact region and encourage electronhole separation, which is an essential pre-requisite for heterogeneous photocatalysis.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
309
DEVELOPMENT OF METHODS FOR REGULATING THE CHARGED SURFACE GROUPS OF V-A1 203 IN AQUEOUS SOLUTIONS. L. VORDONIS, A. AKRATOPULU, P.G. KOUTSOUKOS AND A. LYCOURGHIOTIS* Department of Chemistry-Research Institute of Chemical Engineering and High Temperature Processes, University of Patras, Patras, Greece SUMMARY Two methods for regulating the lero Point of Charge (l.P.C.) and the concentrati on of the charged surface groups (ATOH2+, ATO-) of v-A120.3 at each pH are described. The first method involves doping of v-A1203 with varlOUS amounts of Li+, Na+ and F- ions. The second method involves the variation of temperature of the impregnating solution in the range 10-50 0 C. Doping of V-A1203 with very small amounts of Li+ and Na+ ions increases its l.P.C. from 5.30 to 9.80 and consequently the pH range in which this support can adsorb negatively charged species. At pH lower than 9.00 increase in the dopant concentration increases the concentration of the ATOH2+ groups and the extent of adsorption of negative species as well. On the contrary, doping of v-A1203 with minute amounts of F- ions decreases the Z.P.C. from 5.30 to 3.40. This ion therefore is suitable in order to extent the pH range, where positive species could be deposited on v-Al203 By adsorption, at lower pH values. Increase in the concentration of the F- ions induces a considerable increase in the concentration of ATo- groups at pH> 3.40 and thus in the extent of adsorption of positive species. An extremely precise regulation of the l.P.C. and the concentration of the ATOH2+ and ATO- groups can be achieved by variation of temperature of the impregnating solution in the range 10-500 C. Increase in temperature brings about increase in the l.P.C. and ATOH2+ groups concentration whereas it causes decrease in the concentration of ATo- groups. This method can therefore be used in order for a fine regulation of the sorptive ability of v-A1203 for negative or positive species to be achieved. INTRODUCTION It is well known that v-Al 203 is the most important carrier used in contact catalysis. The deposition of an active ion on its surface, which is the most critical step for the preparation of supported catalysts based on this carrier, ts- usually performed using two methods: (i) Pore volume or wet impregnation followed by drying. (ii) Impregnation using quite dilute solutions followed by filtration. By the first method we can deposit on v-A1 203 surface quite large amounts of active ions but the dispersity of the supported phase is quite low because the deposition takes place mainly by non-controlled precipitation during drying. (1). However, a satisfactory value of active surface [(active surface)=(dispersity)x(amount of supported active ion)] can be obtained
310
in this case by increasing the amount of the supported phase. This method, is therefore suitable when a cheap active ion should be deposited. The second method results in high dispersity of the active phase because the deposition takes place by adsorption of a suitable species, containing the active ion, on the surface of the v-A1 203• (1). Therefore, this technique is used for the deposition of noble expensive metals, where quite low concentration of the active phase is usually required. Unfortunately, by this method we can't prepare supported catalysts with high active surface because the low concentration of adsorption sites limits the amount of the active ion that can be deposited by adsorption. Recently, it has been recognized that the adsorption sites for negative (positive) species are ~OH2+ (~o-) surface groups (2-7) resulting from the protonation-deprotonation equilibria of the surface hydroxyls of the v-A1203 in aqueous suspensions (8): int int K ~OH ~~O+ H+ ~OH + ~ A10H + H + 2 s' S
HS+ and H+ denote the hydrogen ions on the surface of the v-A1 203 and in the aqueous solution, respectively. The abo:ve equil ibria show that the concentration of the adsorption sites depends on the pH of the impregnating solution as well as on the values of the surface acidity .constants K1~nt and K2i nt• It has been demonstrated that the halfsum of the PK11nt and p~lnt is equal to a pH value, defined as lero Point of Charge (l.P.C.) (3-8), at which the concentration of ~OH2+ is equal to that of ~O- and negligible as compared to the concentration of ~OH. At pH values higher and lower than l.P.C., the adsorption sites for positive and negative species respectively is predominant. From the above, it is evident that increase in the surface concentration of is necessary when an enhancement in the adsorption of a negathe ~OH2+ (~O-) tive (positive) species is sought. The method followed in practice involves a proper regulation of the pH of the impregnating solution (9-12). This method, seemingly very simple,may become problematic in some cases.More specifically, at the pH value where adsorption is enhanced,the species to be deposited may be unstable and the support may be partially dissolved.Moreover,deposition by spontaneous precipitation may occur in considerable extent,resulting to low dispersity of the supported phase.We need therefore,alternative methods for regula-' ting the l.P.C. value as well as the concentration of the charged surface groups of v-A1203,throughout the pH range. In the present communication we describe two methods recently developed in our laboratory. The first method involves the doping of v-A1 203 with various
311
amounts of a basic [Na+,Li+J or acidic /}"- J modifier (3,4,6). The second method consists in varying the temperature of the impregnating solution in the range 10-50 0 C (7). METHODS Preparation of the samples The Li+, Na+ and F- doped specimens were prepared by pore volume impregnation of y-A1 [Ho 415, Houdry, 100-150 mesh, SSA = 123 m2.g- 1 ] with aqueous 203 solutions of LiN0 3 /}"errak, reinst J, NaN0 3 [Merck p.a. ] and NHi [Carlo Erba R.P. J, respectively. Following impregnation, the specimens were dried at 110 0 C for 2.5 h and then air-calcined at 600 0 C for 12 h. The specimens thus prepared are denoted by Li-X-y'A1 203, Na-X-Y·A1 203 and F-X-y.A1 20 3 where X represents the nominal composition expressed as mmol of the dopant ion per g of y-A1 203• The specific surface areas of the specimens studied were determined by a multiple-point B.E.T. method. Full details are reported elsewhere (3,4,6). Determination of Z.P.C. and the concentration of the ArOH; and AlOThe determination of the title parameters was achieved using the method of the potentiometric titrations. The setup and procedure used and the method followed have been reported elsewhere (3,4,7). RESULTS AND DISCUSSION Regulation of the Z.P.C. The variation of the Z.P.C. values with Li+, Na+ and F- content as well as with the temperature of the impregnating solution is illustrated in figures 1 and 2 respectively. Inspection of figure~,shows that both Li+ and Na+ ions cause an increase in the Z.P.C. of the y-A1 203. However, the effect of Na+ ions is stronger as compared to that of the Li+ ions. Thus, an amount of Na+ as low as 0.392 mmol Na+ per g of y-A1 203 is sufficient to cause an increase in the Z.P.C. from 5.30 to 9.70 whereas 0.621 mmol Li+ per g of y-A1 203 are required in order for the same result to be obtained. The above, clearly demonstrate that by doping y-A1 203 with relatively small amounts of Li+ or Na+ ions a significant enlargement of the pH range. from 1.00-5.30 to 1.00-9.80. where negative species could be adsorbed on its surface in considerable extent, can be achieved. Moreover, it can be observed that modification of y-A1 203 with 0.125 mmol Fper g of y-A1 203 is sufficient to decrease the Z.P.C. value from 5.30 to 3.40. Therefore. it can be used to exploit this pH range for the adsorption of positive species on y-A1 203 surface. Figure 2 shows, that a precise regulation of the value of Z.P.C. from 4.45 to 8.95 can be achieved. by simply changing the temperature of the impregnating
312
\I
10 9
8 7 c..> 0N
6 5 4 3 0.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
4.0
X(mmollg Y-A2~)-
FIGURE 1: Variation of the experimental values of Z.P.C. with content of the dopant ion. [Li+(-), Na+(O) and F-(l»J. 1\
10
/
/0
9
t o 0-
8
0
/
7
N
6 5
- 0 -0
/0/0
4 0
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30
40
50
60
8/ c _ o
FIGURE 2: Variation of the experimental values of Z.P.C. with temperature of the v-A1203 suspension.
313
solution in the range 10-500 C. The practical consequences from the view point of the preparation of supported v-A1 203 catalysts are obvious: A simple increase of temperature from 25 to 50 0 C is sufficient in order for an enlargement of the pH range, in which v-A1 203 can adsorb negative species, from 1.00-5.30 to 1.00-8.95 to be achieved. A decrease, on the other hand of the temperature from 25 to 10 0 C allow, for an extension of the pH range, in which v-A1 203 can adsorb positive species, by almost one pH unit. Moreover, the regulation achieved, facilitates the selective adsorption of eventually desired negative species stable in a narrow pH range. Regulatiun of the concentration of the charged surface groups. Figures 3 and 4 illustrate the variation of the concentration of the charged surface groups of v-A1203 with pH for the Li+ and Na+ doped specimens, respectively. From figure 3 one can observe that doping of v-A1203 with 0.392 mmol Li+ per g of v-A1203 has no influence on the concentration of the charged surface groups over the entire pH range studied, whereas modification with larger amounts of Li+ causes a marked increase in the consentration of ATOH 2+ groups at pH lower than 9.00. In the same region the Li-doping results in an almost complete disappearance of the negatively charged surface groups, Similar, but more strong, is the action of the Na+ ions. (Fig. 4).
3
t
2
+N I
0
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~I
,
~
~
~
I
]0
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4
5
6
7
8
9
10
II
12
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FIGURE 3: Dependence of charged sites (ATOH 2+ and ATo-) on pH for the Li+ doped samples. Numbers correspond to the nominal composition of the modifier expressed as mmol per g of v-A1 203,
314
4
t
2
+C\J
:I:
0
l
<,
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IJ')
I
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l
t
.S:! ~
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7
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FIGURE 4: Dependence of charged sites (~OH2 + and ~O - ) on pH for the Na + doped samples. Numbers correspond to the nominal composition of the modifier expressed as mmol per g of v-A1 03• Na+ originates from the v-A1 203 used. 2
*
In fact, comparison of figures 3 and 4 shows that quite smaller amount of sodium than lithium is required, in order for the same extent of the effects mentioned before to be achieved. These results show that the Na+ and Li+ doping can be used in order for a promotion of the adsorption of the negative species on the v-A1 203 surface at pH lower than 9.00,to be obtained. The extent of this promotion, being more considerable at relatively low pH values, increases almost linearly with the dopant concentration at a given pH (Fig. 5). Moreover, the complete disappearance of the sites for the adsorption of positive species at pH lower than 9.00, after the Na+ or Li+ doping, could be proved extremely usefull in the cases, where a desired macrodistribution of an active ion on an alumina pellet has to be achieved.(1). Figure 6, illustrates the variation of the concentration of the charged surface groups of v-A1 203 with pH for the F- doped specimens. One can observe that the modification by F- ions increases the concentration of the ~O- groups and, thus, promotes the adsorption of positive species on the v-A1 203 surface at pH higher than 3.40. The dependence of the concentration of Al0~groups on the concentration of the F- ions at pH = 4.50, taken as a typical example, is illustrated in figure 7.
315
3
t
2
'Ec <,
u: (I)
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+ C\J
::r:
o
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ol.--"":::::::::::.......L..!:::L--.L..-----L----'----~
0.0
1.0
0.5
2.5
2.0
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FIGURE 5:Con~rationofpositively charged groups (~OH2+) at pH = 8.50 as a function of the nominal composition of Li+ (0) and Na+ (A) . 4
2
t
+N
::r:
0
0 I~~
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7
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FIGURE 6: Dependence of charged sites (~OH2+ and ~O-) on pH for the F- doped samples. Numbers correspond to the nominal composition of the modifier expressed as mmol per g of y-A1 203•
316
7
-'"
E
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-b (/) Q)
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.l.-
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FIGURE 7:Concen'~ation ofnegatively charged groups (ATo-) at pH function of the nominal composition of F- ions.
, + C\J
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3
4
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6
7
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FIGURE 8: Dependence of charged sites (ATOH 2+ and ATo-) on pH with the temperature of the v-A1203 suspension.
317
Figure 8, illustrates the variation of the concentration of the charged surface groups of V-A1 203 with pH obtained for various temperatures in the range 10-50 0 C. One can observe that the variation in the temperature of the impregnating solution is an attractive method not only for regulating Z.P.C. but also for obtaining a desired concentration of charged surface groups at a given pH. In general, it can be seen that increase (decrease) in the temperature of the impregnating solution increases the concentration of the sites for the adsorption of negative (positive) species. Detailed interpretations concerning the effect of the Li+, Na+ and F- ions as well as the charrqe of temperature of the impregnating solution on the Z.P.C. and the concentration of the charged surface groups have been reported elsewhere (4,6,7). CONCLUSIONS The main results of this work can be summarized as follows: (a). The method of doping of v-A1203 with Li+ and Na+ ions can be employed in order to enlarge the pH range in which this oxide can adsorb negative species as well as in order to promote the adsorption of these species at pH lower than 9.00. It is important to notice that the enlargement of the pH range is achieved after the doping with relatively small amounts of Na+ and Li+ because the presence of these ions at higher concentrations usually leads to catalytic deactivation (13-16)• (b). Doping of v-A1 203 with F- ions allows the extension of the pH range, in which this oxide can adsorb positive species and increases the extent of adsorption of these species at pH higher than 3.40. Since F- ion is a promotor of many supported catalysts based on v-A1203 (17-20), its use for regulating the sorptive capasity of this carrier is not expected to create serious problems concerning the catalytic activity. (c). The change of the temperature of the impregnating solution it was proved to be a very simple method in order a fine regulation of the sorptive ability of v-A1203 for negative or positive species to be achieved. Evidently, this "clean" method is not expected to cause deactivation problems. REFERENCES 1 A.V. Neimark, L.I. Kheifer and V.B. Fenelonov, Ind. Eng. Chern. 20(1982) 439. 2 J.P. Brunelle, Pure and Appl. Chern., 50(1978) 1211. 3 L. Vordonis, P.G. Koutsoukos and A. Lycourghiotis, J. Chern. Soc. Chern. Commun. (1984) 1309. 4 L. Vordonis, P.G. Koutsoukos and A. Lycourghiotis, J. Catal, in press 5 L. Vordonis, P.G. Koutsoukos and A. Lycourghiotis, Langmuir, in press 6 L. Vordonis, P.G. Koutsoukos and A. Lycourghiotis. submitted for puplication. -
318
7 8 9 10 11 12 13 14 15 16 17 18 19 20
A. Akratopulu, L. Vordonis and A. Lycourghiotis, submitted for publication. C.P. Huang and W. Stumm, J. Coll. Interf. Sci. 43(1973) 409. W.K. Hall and L. Wang, J. Catal. 77(1982) 232. M.J. Jr.D'Aniello, J. Catal. 69(1981)9. M. Houalla, C.L. Kibby, L. Petrakis and D.M. Hercules, J. Catal. 83(1983) 50. M.S. Heise and J.A. Schwarz, J. Coll. Interf. Sci. 107(1985) 237. A. Lycourghiotis, D. Vattis, G. Karaiskakis and N. Katsanos, J. Less-Common Metals 86(1982) 137. Ch. Kordulis, A. Lycourghiotis and S. Voliotis, Applied Catalysis 15(1985) 301. A. Lycourghiotis, C. Defosse and B. Delmon, Bull. Soc. Chim. Belg. 91(1982) 613. A. Lycourghiotis and D. Vattis, React. Kinet. Catal. Lett. 21(1982) 23. P.M. Boorman, J.F. Kriz, J.R. Brown and M. Ternan, Proc. Climax 4th Int. Conf. Chemistry and Uses of Molybdenum, Ed. H.F. Barry and p.e.H. Mitchell [Climax Molybdenum Co. Ann. Arbor. Michigan. 1982, p. 192J. M. Ternan, Can. J. Chern. Eng. 61(1983) 133. C. Muralidhar, F.E. Massoth and J. Shabtai, J. Catal. 85(1984) 44. P.M. Boorman, J.F. Kriz, J.R. Brown and M. Ternan, 8th Int. Congr. on Catalysis, Berlin 1984, paper 11-281.
319
DISCUSSION B. DELMON : This is a beautiful work, especially the fine "tuning" with temperature. 1/ It is certainly more convenient (practically) in many cases, to use surface additives, than to change solution temperature. This may also be necessary for other purposes (regulating activity/selectivity of the support). In your case, how far do the surface dopants dissolve back during impregnation? Did you try to measure this solubility? Are these additives very strongly attached to the support? 2/ Data which you did not present, but certainly have. How many Lit, Na+, .•. are attached for 1 mm 2? A. LYCOURGHIOTIS : Thank you for these questions which provide me the occasion to clarify two important points: (a) We have measured the solubility of the dopant ions and we have found that it increases with the amount of the additive deposited (1, 2). However, even in the case of the specimens with the maximum nominal dopant concentration, the amount of the additive dissolved is extremely small. Thus, in the ionic strength where the concentration of the charged groups is reported, the percentage of the dopant ion remaining on the support surface is 93.0% for the Na-2.470-yA1 203, 98.6% for the Li-2.470-yA1203 and 97.4% for the Fe-3.818-yA1203. Therefore, one can conclude that due to the air-calcination at 600°C, the additives' are very strongly attached to the support surface. (b) Taking into account the initial specific area of y-A1203 used (~130m2.g-1) and the concentration of the surface hydroxy1s reported in the literature (8-10 sites.nm- 2) the dopant concentrations were selected to cover, approximately, the range 0-100% of the monolayer concerning the alkali ions and 0-150% of the monolayer concerning the fluoride ions. However, due to the method of dopant deposition (pore volume impregnation), one may expect the formation of dopant aggregates resulting in real coverage much lower than that predicted. I think that it will be useful to report the concentrations of Li+, Nat and Fions corresponding to the abrupt change observed for Z.P.C. : (Na-O.392-yA1203: 1 ion .nm-~, Li-O.621-yA1203 : 3.1 ions.nm-2, F-O.125-y-A1203 : 0.4 ion.nm-Z). Taking into account the concentration of the surface hydroxyTs mentioned before and the formation of aggregates on the support surface, one can conclude that in these critical concentrations only a small part of the surface is covered by the dopant ions. (1) L. Vordonis, P.G. Koutsoukos and A. Lycourghiotis, J. Cata1. 98 (1986) 296 (2) L. Vordonis, P.G. Koutsoukos and A. Lycourghiotis, J. Cata1. lUl (1986) 186. J.W. GEUS : My question is concerned with the range of interaction of the modifying ions, such as Nat or F-, on the adsorption of the active precursors to be applied onto the support. The changes present on colloidal particles and determining their electrokinetic behaviour are exhibiting a very low surface density. Therefore the ratio of the charge-modifying ions and the precursor ions is very important. Is that ratio of the order of unity or of the order of 10-6 or 10-3? Take for instance application of Ni(II) at a pH below 4 onto A1203, what is the ratio between F- and Ni(II) that can be applied at a pH below 4 onto the aluminium? A. LYCOURGHIOTIS : I agree with you that the ratio of the concentration of the dopant ions to the concentration of the charged sites, which adsorb the precursor ions, is very important. But for a given dopant loading it depends on pH. In order to obtain an idea for this ratio, I shall give you two examples corresponding to critical dopant concentrations at which an abrupt change in Z.P.C. was observed: (a) Na-O.309-y.A1 203, pH = 3, [Na+]/[positive sites]= 0.3; (b) F-O.125-y.A1203, pH = 4, [F-]/[negative sites] = 0.1.
320
B. NOTARI : You propose the preparation of A1 203 based catalysts with a pretreatment with alkali and/or fluoride. It is well documented in the literature that in many important applications where A1Z03 catalysts are used, A1203 does not simply act as a support, but particlpates in the chemical transformations by catalyzing acid reactions : moreover the best performances of the catalysts are obtained with a close control of the acidity of alumina. It is well documented also that very small amounts of alkali reduce drastically the acidity of alumina and small amounts of fluorine increase the acidity. I believe that this is a serious drawback of the method you propose since it would put the acidity of alumina out of control. A. LYCOURGHIOTIS : I generally agree with you when you say that doping may change the physicochemical properties of the carriers including acidity. But in some cases, the changes in the physicochemical properties obtained are desirable from the catalytic point of view. For instance, the doping with Fions has been proved to promote catalytic acidity for many reactions. Typical examples are the hydrotreating catalysts based on Co - Mo or Ni, sulfide supported on 1-A1203 (1-3). Analogous examples can be found in the literature for Na+ or Li doping. In these cases, one can use the method of doping to achieve both increase of the active surface of the supported catalyst and catalytic promotion. In contrast to that, there are many cases where doping results in catalytic deactivation (for example see ref. 4 and 5). Obviously in these cases, doping with 1arfe amounts of the modifiers is not suggested and one should use the metho 0 changing the temperature of the impregnating solution to regulate the sorptive capacity of y-A1203 at a given pH. (1) M. Ternan, Canad. J. Chern. Eng. 61 (1983) 133 (2) C. Mura1idhar, F.E. Massoth and ~ Shabtai, J. Cata1. 85 (1984) 44 8th Int. (3) P.M. Boorman, J.F. Kriz, J.R. Brown and M. Ternan in "~oceedings Cong. on Catalysis" Berlin, 1984, paper 11-281 (4) A. Lycourghiotis, C. Defosse and B. De1mon, Bull. Soc. Chim. Be1g. ~ (1982) 613 (5) A. Lycourghiotis, D. Vattis, N.A. Hatsamos and G. Haraiskatis, J. Less Common Metals 86 (1982) 137. P. VILLA: You propose a mechanism of adsorption which is essentially due to electrostatic interaction and therefore non-localized. In the past Iannibel10 and co-workers, considering the pH increase observed during adsorption, and the Raman spectra of Mo adsorbed on alumina, proposed that adsorption occurs through a mechanism of exchange of surface OH- and Mo0 leading at least, at low Mo loadings, to isolated Mo on alumina. Could you comment?
4
A. LYCOURGHIOTIS : The main goal of this work is to show that one can change the concentration of the charged surface groups of y-A1203 throughout the pH range by doping this oxide or changing the temperature of the impregnating solution. Obviously, this regulation could be used to control the sorptive capacity of y-A1203 at a given pH provided that the predominant mechanism of deposition of a given species is its adsorption on the above mentioned charged sites. Concerning the adsorption of the Mo(VI) species, one cannot, for the moment, ex£lude that a portion of the Mo(VI) is deposited via exchange of surface OH with MOOa-. However, there are strong evidences-in the literature which show that the predominant mechanism of deposition of the Mo~Og-+species on the y-A1203 surface is the adsorption of these species on the A1~H2 groups. For instance, Boue11e et a1. (1), who recently reported on adsorption studies of various species, including Mo(VI) species, have calculated, on the basis of adsorption isotherms, the concentration of the sorptive sites and found it equal to that predicted by our model. However, I think that further experimentation is needed to determine the relative contribution of each mechanism in the whole deposition process. Experimental work is now in progress to investigate this point. (1) G. Meunier, B. Mocaer, S. Kaszte1an, L.R.L. Coustumer, J. Grimb10t and J.P. Boue11e, Applied Cata1. ~ (1986) 329.
321
D.M. HERCULES: What is the physicochemical origin of the large temperature effect on zero-point charge? A. LYCOURGHIOTIS : This is an important question. An inspection of the protonation-deprotonation equilibria illustrated in the text shows that a considerable change of the Z.P.C. (Z.P.C. PK1 + PK2/2) should be expected provided that the surface protonations-depronotations are accompanied by considerable enthalpy and entropy changes. This is the case for y-A1203' In fact, we calculated the thermodynamic parameters of these equilibria, from experimental 1nk vs. liT curves, assuming that the surface of the y-A1203 particles can be Simulated with a two-dimensional ideal solution (1). The
~~6u~s5g~~!T~~1-\6~~d=6s~8~
~9lgK~·~g~~~'m~~Zl]
j~~~iTy2~~~·~~~~~tio~S!f=Z.P.C.
with-temperature observea in the present work. (1) A. Akratopulu, L. Vordonis and A. Lycourghiotis, J.C.S. Faraday Trans. 1, in press. C. WUNDE : Do you have also made experiments by treating y-A1203 with ammonia or amines as doping agents. A. LYCOURGHIOTIS : No, we did not.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV 1987 Elsevier Science Publishers B.V .. Amsterdam - Printed in The Netherlands
UNDERSTANDING THE MORPHOLOGICAL TRANSFORMATIONS THAT OCCUR IN THE PREPARATION OF ALUMINA SUPPORTS W. C. Conner and E. L. Weist1 and L. A. Pedersen 2 IDepartment of Chemical Engineering, University of Massachusetts, Amherst, Massachusetts 01003 (USA) 2ALCOA Laboratories, Alcoa Center, Pennsylvania 15069 (USA)
SUMMARY Using mercury intrusion and extrusion pressures from 0 to 60,000 psi, we have studied the effect of processing steps on the textural morphology of alumina powders. The void morphology is bimodal and different portions of the structure may be subject to separate analysis. The surface area increases after rehydration of an amorphous alumina and our analysis indicates that the macrostructure of the solid seems to undergo a transformation from agglomerated microparticles to a platelet-like structure. SEM was used to confirm this interpretation. Within the platelets themselves, additional voids, possibly cracks, develop with thermal treatment and then begin to increase in size. The rehydration creates a narrow distribution of pores around 2 nm in radii. INTRODUCTION Alumina is one of the primary supports used in catalysis. Several, sequential processes are involved in ALCOA's production of high surface-area aluminas: precipitation. calcination. rehydration. and activation. Though the pore structure of the final alumina often controls the subsequent catalysis, the changes in the void morphology (pore structure) that occur during these processes are not understood. The main problem in understanding the changes in void morphology has been the inability to characterize pore structure taking into account the threedimensional nature of the void network. Our recent studies have focused on the analysis of pore structure utilizing mercury porosimetry (refs. 1-5) and nitrogen sorption techniques (ref. 6). Indeed. if the pores are more realistically viewed as a three-dimensional network. these techniques give considerable insight into the internal void structure. These studies indicate that the porosimetric data (combined intrusion and extrusion processes) differ for agglomerated particles with, for example, different geometries (spheres, platelets, needles. etc.) (refs. 7,8). It may be possible to infer the internal void shapes within the network from porosimetry and/or nitrogen adsorption-desorption. Porosimetry can. therefore, be used to monitor any
323
324
changes in morphology that occur during the steps in the preparation of high surface-area aluminas. Our recent studies of a series of agglomerated microparticles with differing particle shapes (refs. 7.8) suggest that it is possible to infer the shapes from detailed analysis of the intrusion and extrusion curves. The ratio of the radii of extrusion to intrusion seems to be characteristic of the shape of the agglomerated particle and by extension the shape of the void morphology created (ref. 7). More specifically. the radii ratios are not necessarily constant as a function of the fraction of the void that is filled; Le •• the radii ratios at 10, 20, ••• 80, 90% filling may differ. This change in the ratios is a further indication of the nature of the void shape. is found (a For spherical agglomerates a constant ratio of 2.3 ~0.3 straight line on the plot of Re/R i versus Phi. percent filling). For needle-like solid agglomerates ratios of over 5 (and often >10) are found and the plot has a pronounced negative slope. For platelet-like agglomerates ratios between 1.4 and 2 are found and the plots have little slope in the intermediate values of Phi (0.3-0.8). For a sample consisting of cylindrical non-intersecting pores a constant value of 1 is found. More general than an absolute determination of the void shape. the dependence of the ratios may be used to characterize any changes that are occurring as a high surface area solid is subject to different processing conditions. So, for example, a shift to higher values of the radii ratio might be indicative of agglomeration of microparticles in one dimension as a "needle-like" string and shifts to values below 2 might be due to two-dimensional agglomeration as in plates or sheets (flat or not). With this perspective we have studied a series of alumina samples in order to deduce the morphological changes that may be occurring. We will briefly describe the porosim~tric studies that were conducted and our preliminary analysis of this data. EXPERIMENTAL Porosimetry was conducted using Quantachrome's high and low pressure Autoscane porosimeters. The samples were first evacuated to <50 p pressure before immersion into mercury. The low pressure porosimeter scanned between vacuum and 500 psia while the high pressure porosimeter was used to intrude and extrude mercury up to 60.000 psig. Both instruments were directly interfaced with an Analog Devices Macsyme 350 computer which was used to collect. analyze. and plot the data.
N2 desorption was measured with an Omicron Omnisorp· 360. SEM micrographs were obtained with a lSI Model DS-130. The starting materials identified as CP-X (X is the median size of the particles in microns) are commercial and developmental activated alumina
325
powders produced by the Aluminum Company of America. They are X-ray amorphous with BET surface areas over 250 m2/g. These powders are rehydratable and form a strong hydroxyl bond upon contact with water (ref. 9). The experimental samples designated as rehydrated CP-X have been heated in water at 90°C for 1 hour in a manner similar to Vorobiev et al. (ref. 10). Surface areas of the oven dried materials are typically 300 m2/g. Heat treatment of rehydrated powders to 425°C in air produces materials that are labeled "activated." Over a dozen samples of CP aluminas were studied. For the sake of comparison, these can be subdivided into three groups. CP alumina powders before rehydration: CP-1, CP-2, CP-5 CP alumina powders after rehydration: CP-1, CP-2, CP-5 CP-2 alumina powders: before and after rehydration, and after 425°C activation RESULTS Because we are concerned with the size and shape of the void morphology, the discussion below will focus on the relative pressures of the intrusion and extrusion. Since the total volume of a powder measured by Hg porosimetry will include both inter- and intraparticle voids, the values derived from the intrusion plots are not absolute. Further, a discussion of total pore volume is not germane to this analysis. The differences in the porosimetric data are depicted in Figures 1-3. As two regions of pressures were studied for each sample (0-500 psia and
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191171
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200 300 Pressure (PSI A)
400
Figure 1: Comparison of CP Alumina Powders before Rehydration
500
326
0-60,000 psig), over a dozen curves of intrusion and extrusion could have been included. Within a series those curves that are essentially identical are not represented. The figures will be described in sequence below. In Figure 1, the general form of the intrusion and extrusion curves is similar for samples of CP-5, CP-2, and CP-1. As expected there is a shift to higher pressures (smaller dimensions) for both intrusion and extrusion as the particle size decreases (5+2+1). For each of the samples, considerable "intrusion" occurs at the onset of the measurement. This seems to be an artifact due to the compaction and filling of the void structure exterior to the primary particles.
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F'l9ure 2: Comparison of Rehydrated Alumina Powders Both the low (0-500 psia, left) and high (0-60,000 psig, right) pressure measurements are similar in Figure 2. As above, there is a general shift to increasing intrusion at higher pressure as the series moves to smaller particles (CP-5+CP-2+CP-1). 2.0
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Figure 3: Effect of Rehydration on CP·2 Alumina The dramatic influence of rehydration and activation of CP-2 alumina, for
50
327
example, are depicted in both low and high pressure regions of porosimetric analyses in Figure 3. With rehydration, the intrusion and extrusion at low pressures shift to lower pressure, and the difference between intrusion and extrusion decreases dramatically. Activation at higher temperature (425°C) tends to increase the extent of intrusion at lower pressure in the low pressure region and to higher pressures in the high pressure region. Also note that the total volume of the pore network is increasing during the sequence.
ANALYSIS AND DISCUSSION Figure 1 shows that the macrostructure of the series of CP alumina powders are all similar. There is a general trend to both intrusion and extrusion at higher pressure as the nominal particle sizes decrease; i.e., CP-5 to 1. With rehydration, as shown in Figure 2, the pressures of intrusion and extrusion both decrease for each sample. This indicates that in this region of larger void dimension (>0.1 ~ = 100 nm) the particles are agglomerating to form larger sUbparticles. Further, note that the shift between intrusion and extrusion decreases dramatically. This represents a change in the general shape of the particle morphology, as discussed below. At high pressures the porosimetry is essentially the same for each sample after rehydration. Focusing on a single sample, CP-2, the effect of rehydration and subsequent "activation" at 425°C is seen in Figure 3. As discussed above, there is a substantial shift in these low pressure porosimetric curves due to rehydration. The activation is seen to represent a slight shift to increased intrusion across the range of pressures. The high pressure region shows that rehydration and the subsequent activation also develops a noticeable mesoporosity in CP-2 alumina powders. The development of porosity in the range of 10 to 50 A radius was confirmed by N2 desorption results. Figure 4 is an overlay of the first derivative of the desorption curves for the three CP-2 samples. The change in the height and shape of the peak centered around 20 A indicates that the pore volume and pore size distribution is modified as the alumina goes through the processing steps. The shift of the skewing from left to right confirms that the broadening seen in the Hg data takes place at the low end of the mesopore range especially in the activated CP-2 sample. As N2 desorption corresponds to Hg intrusion (ref. 7), desorption gives added detail to the picture obtained with the Hg technique. The doublet observed for the 20 A radius peak of each sample would not be easily measured during Hg porosimetry. The differences in the relative intensities of the doublets in the distributions may be an artifact of the measurement technique.
328
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70 60 50 40 30 20 10 0
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20
30
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Pore radius, RP (A)
Figure 4: Comparison of N2 Desorption Curves for CP-2 Aluminas Morphological Analysis It is possible to represent the intrusion and extrusion curves in a manner that is better able to compare agglomerated particles of different shapes (refs. 7,8). A portion of the void volume over which intrusion and extrusion occur is compared by plotting the measured ratio of the radii of extrusion and intrusion (Re/Ri=Pi/Pe) against the fraction of voids filled within this portion of the void volume (defn Phi). It should be noted that each plot tends to have negative slopes at high and low fractions of filling (O.8
ii:
5.0 4.5
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Phi Figure 5: Low Pressure Poroslmetry ofAlcoa CP Alumlnas before Rehydration Similar representations for this data are depicted in Figures 5 and 6. From Figure 5 we can see that in the low pressure measurements the alumina
329
powders behave as spherical agglomerates. The specific reason for the slightly higher values for CP-5 is not understood, but the trend is evident. Rehydration results in a shift to a plot similar to the platelet-like
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Initial
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0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 Phi
Figure 6: Low Pressure of Alcoa CP·2 Aluminas agglomerates as shown in Figure 6 for CP-2. As a result of rehydration, the ratios of the radii of intrusion to extrusion decrease to less than 2.0, and the ratio becomes flat above 50% void filling; i.e., Phi> 0.5. The effect of rehydration seems to be the formation of sheet-like structure from the initial microparticle agglomerates. This may be the result of the subparticles congealing and rearranging to form two-dimensional structures. This may involve a partial dissolution of the alumina surface during the rehydration. Due to the increase in surface area, this may also involve a roughening of the surface of the primary particles. In consideration of the shifts found for the intrusion and extrusion data, the resultant surface may involve a two-dimensional, rib-like roughness (as opposed to stalactite-like formations). Within the subpartic1es, cracks or fissures seem to develop that are evident in the sorption data. If these hypotheses are correct, the evidence for these changes in the larger void structure may be visible by SEM. The micrographs contrasting CP-2 alumina before and after rehydration are shown in Figure 7. The top photograph (a) is a 5K magnification of the original CP-2 powder. The material has the appearance of agglomerated hexagonal structures. Upon rehydration (Figure 7b), a two dimensional-like surface texture has been created. This change is observed for all of the rehydrated samples used in thi s study.
330
Figure 7:
SEM Micrographs of CP-2 Aluminas, X 5K - (a) Initial (b) Rehydrated
CONCLUSIONS The rehydration of alumina seems to reflect a rib-(or plate-)like roughening of the surface of primary particles. These subparticles are also cemented together during the migration or dissolution and redeposition of the alumina solid. During the process, the roughening and redistribution of the alumina results in an increased surface area and the concomitant creation of pores less than 10 nm. This may be due to cracks and fissures interior to the subparticles. This analysis is offered as a hypothesis; whereas, the structural
331
assignments may not be precise, the changes do reflect morphological differences. Further analyses are needed to ascertain the modifications that are taking place and eventually to control the processes involved in the formation of these and other aluminas and the effect of subsequent treatments. ACKNOWLEDGMENTS One of us, W. C. Conner, wishes to thank NSF for the donation of the porosimeters and Analog Devices for the donation of the computer and Gulf Oil for their support of morphological studies. All authors acknowledge the generous support of ALCOA for this research. REFERENCES 1
W. C. Conner, A. M. Lane, K. M. Ng, and M. Goldblatt, J. Catal., 83, (1983) 336. 2 W. C. Conner, A. M. Lane, and A. Hoffman, J. Call. & Interface Sci., 100, (1984) 185. 3 W. C. Conner and A. M. Lane, J. Catal., 98, (1984) 217. 4 K. M. Ng, A. M. Lane, and W. C. Conner, Chem. Eng. Comm., 38, (1985) 33. 5 A. M. Lane, W. C. Conner, and N. Shah, J. Coll. & Interface Sci., 109, (1986) 235. 6 W. C. Conner, J. F. Cevallos-Candau, E. L. Weist, J. Pajares, S. Mendioroz, and A. Cortes, Langmuir, 2, (1986) 151. 7 J. Neil, W. C. Conner, C. Blanco, and J. Pajares, J. Phys. Chem., submitted for publication 1986. 8 W. C. Conner, J. Neil, C. Blanco, and J. Pajares, in preparation 1986. 9 Product Data Sheet CHE 945, Alcoa Chemicals Division. 10 J. K. Vorobiev, et al. (USSR), U.S. Patent 4,166,100 (1979).
332
DISCUSSION G. LESTER: As more studies of the basically non-equilibrium curves generated as mercury enters and leaves the porous structure of catalysts and supports, it is more important to keep in mind the physical processes involved. In particular, mercury extrusion is not a proper form to describe the removal of mercury from the structure by reducing the imposed pressure on the system. I suggest to use, for a more deflnltlVe form, the word "extraction" to describe this process. L. PEDERSEN: We agree wholeheartedly with your comment. In searching for a better term to describe the process, we have found that J.R. Anderson and K.C. Pratt (in "Introduction to the Characterization and Testing of Catalysts", Academic Press, 1985, p 136) use the term "retraction" which we are now tending to favor. By the way our studies have shown that Hg prosimetry is a static process. S. TAMHANKAR : Could you comment on dead-end pores, bottle necks and lastly possible distortion of pores at high mercury pressures, particularly with respect to using mercury porosimetry as a general tool? L.A. PEDERSEN: As the paper implies, bottle neck shaped pores are not needed to explain hysteresis. If bottle neck and dead-end pores exist then the surface tension of the Hg is normally able to force HG out of those pores during retraction. If the process, retraction, is performed too rapidly, "snap -off" may occur. Snap-off would depend on the rate at which mercury is intruded and retracted. By varying the rate of intrusion and retraction, we have found that there is no noticeable difference in the volume of Hg retained under our conditions for different rates. For some few samples, not normally catalyst supports, Hg porosimetry may affect the pore structure. But in general the technique seems to be universally applicable.
B. Delmon. P. Grange. P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
333
© 1987 Elsevier Science Publishers B.V.• Amsterdam - Printed in The Netherlands
MICROPOROUS AMORPHOUS ALUMINA OF A ZEOLITIC TYPE FOR CATALYTIC REACTIONS WITH METHANOL G. TOURNIER, M. LACROIX-REPELLIN, G.M. PAJONK and S.J. TEICHNER Laboratoire de Thermodynamique et Cinetique Chimiques, UA 231 au CNRS, Universite Claude Bernard Lyon I, 43 bd du 11 novembre 1918 - 69622 Villeurbanne Cedex. SUMMARY A new type of amorphous alumina of a high surface area and a monomodal pore size distribution in the microporous range, exhibiting a zeolitic texture has been prepared and some of its properties stUdied. The preparation methods are described concerning pure aluminas as well as copper or cobalt supported on these aluminas. Methanol conversion and Fischer-Tropsch synthesis were selected as catalytic tests in a preliminary study of their catalytic properties. INTRODUCTI ON A microporous amorphous alumina exhibiting a high surface area was already prepared in this laboratory and its main textural properties published elsewhere (ref. 1-3). A brief account of its preparation is given here for the purpose of comparison with a new method of obtaining a similar alumina but in a shorter time. As both kinds of alumina exhibit a zeolitic texture (but in an amorphous state) it is interesting to attempt at the methanol conversion and also CO/H 2 reaction since when Co is deposited on porous supports like alumina a specific selectivity in hydrocarbons is recorded depending on the size of the pores as it was shown by Vanhove et al. (ref. 4). The amorphous, microporous aluminas were synthetized through the sol-gel process performed in a non aqueous medium (methanol). EXPERIMENTAL Reactants and techniques used The first type of microporous alumina, labelled AC, was prepared with A1C1 3, 6 H20 while the second one, noted AN, was elaborated with A1N0 3, 9 H20. Both aluminium salts are reactant grade and are soluble in methanol (water-free) at DoC which is the temperature of the corresponding alcogels (see below). Pure methanol, gaseous ammonia (in a cylinder) and nitrogen were used respectively as a reaction medium, a precipitating base and an adsorbate for textural determination through physical adsorption experiments. XRD patterns were registered with the Cu Ka radiation.
334
Adsorption-desorption isotherms as well as surface areas and porosity were measured in a dynamic chromatography system (ref. 3) using N2 at 77 K. Purity tests were performed with AgN0 3 in the case of AC ~bsence of Cl-) and with Fe S04 (brown ring) in the case of AN (absence of N0 3-). In the latter case the detection limit of anions is within the ten ppm of N0 3- ions range. The Cu on alumina catalyst (named Cu AN) was prepared (see below) with copper acetate in solution in methanol, whereas the Co on alumina catalyst was obtained using diocty1carbony1 cobalt Co 2(CO)8 dissolved in n-heptane (it is labelled Co AC). Gas chromatography analyses of products were performed during the catalytic tests performed in flow conditions, under a total pressure of 1 atm. in a microreactor. A thermal conductivity detector TCD chromatograph equipped with a Porapak Q column was used for the methanol conversion tests while for the Fischer-Tropsch reaction a TCD unit with a SE 3n column and a temperature linear programmation were used. Preparation and properties of pure microporous aluminas The preparation of sample AC is only briefly recalled here since it was already described in a previous series of papers (ref. 1-3). Equation (1) represents the formation of this alumina sample:
NH 4C1 was eliminated from the alumina alcogel by extraction with boiling methanol in a soxh1et unit (ref. 1-3). Due to the relatively low solubility of NH 4C1 in hot methanol this washing step was rather long and tedious to perform. For instance it took several weeks to purify from C1- ions a sample of 20 g of A1 203 in a 500 cm 3 soxhlet. Finally, the purified alumina (absence of Cl- ions) was dried at the ambient temperature under reduced pressure to remove CH 30H and gaseous ammonia. The second alumina sample AN was prepared with Al(N0 3)3' 9H 20 following the same method as before according to equation (2) :
Gaseous dried NH 3 was flown slowly in the aluminium nitrate methano1ic solution at DoC until the final pH reached a value of 8-9. The precipitated alcogel was transferred into the soxhlet unit (500 cm 3) and NH 4N03 was extracted by hot methanol. As the solubility at 25°C of NH 4N03 (171 g/l) in methanol is much higher than that of NH 4C1 (~ 6 g/l) the purification period was considerably reduced. To prepare a 20 g sample from aluminium nitrate required roughly only
335
half of the time used to purify the alumina AC, all other conditions beeing the same. Finally, the alumina alcogel was dried under reduced pressure at the ambient temperature. The AN sample showed an amorphous XRD pattern which did not change when the sample was heated step wise in air up to 750°C where some XRD lines belonging to the n phase began to be recorded. Under vacuum its amorphous structure remained unchanged up to 800°C where some of the XRD lines due to the n phase appeared : amorphous AN in750°C air, 15h ) amorphous AN 1n800°C vacuum, The behavior of the sample AN is very similar to that of the sample AC under the same heating conditions (ref. 2). Since the highest temperatures used for the catalytic tests never exceeded 400°C the aluminas remained always amorphous during these experiments. Concerning now the textural properties of the sample AN, table 1 summarizes the results, whereas figure 1shows the physical N2 adsorption-desorption isotherms for the initial alumina (curve A) and for samples, heated in air at 400°C (curve B) and in vacuum at 400°C (curve C). It can be seen that the initial sample and that heated in vacuum at 400°C exhibit a Langmuir isotherm (type I) revealing their microporous texture while the sample heated in air at 400°C is characterized by a type IV isotherm showing a transition to a mesoporous texture. This behavior is again in full agreement with that of sample AC in the same conditions (ref. 2, 3). Figure 2 shows the corresponding t-plots for the initial sample AN and those heated at 400°C in the same conditions as above (ref. 2, 3). The corresponding Dubinin linear transforms are shown on figure 3 and the pore radius distributions according to the MP method (ref. 3, 5), are represented on figure 4 which clearly shows a rather monomodal pore radii distribution, at least for curves A (the initial sample) and C (sample heated at 400°C in vacuum). A perfect agreement is registered between the SBET and St values and the values of r p and the maximum of the MP distribution curves(fig. 4). Although the surface area decreases rather severely for the sample heated in vacuum at 400°C with respect to the initial one, its pore distribution is still monomodal with a radius corresponding to the maximum of the MP distribution unchanged (~ 8,8 A). This radius value of the maximum is raised for the sample heated in air at 400°C (~ 14,4 AL
336
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Fig. 1. Nitrogen adsorption-desorption physical isotherms at 77 K A - Initial sample B Sample treated in air at 400°C C - Sample treated in vacuum at 400°C.
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B
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Fig. 2. Corresponding t plots (see fig. 1).
15
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337
> _0.9 en o
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Fig. 4. MP pore volume distribution curves corresponding to Fig. 1.
338
TABLE 1 Textural properties of sample AN Vbad Treatment Type of SBET Sta 2/g) isotherms (m2/g) (m (cm~/g)
c Vdads Vads r- = 2V/S r (MP) max. 3/g) 3/g) (cm (cm (A) (A)
as 548 176,6 176,6 0.273 10.1 540 obta ined heated in 286 148.7 294 95.6 0,230 15~7 air at IV 400°C heated in 231 239 85.3 vacuum 77 .3 0.132 11 400°C (a) specific area determined from the t-plots (fig. 2) (b) total adsorption NTP (c) micropore filling volumes NTP (d) total pore volume.
8.8 14.4 8,,8
It is interesting to observe that the variations of the pore size distribution after the heat treatments in vacuum or in air are exactly the same as for sample AC heated in the same conditions, prevjously described (the pore radii values are also very similar (ref. 2, 3)). The only significant difference between samples AC and AN is a higher decrease of the surface area of the sample AN after the heat treatments (in air or in vacuum)(ref. 1-3). In conclusion the properties exhibited by the microporous amorphous alumina AN are very similar to those of the microporous amorphous alumina AN. Preparation of Cu and Co supported alumina xerogels In both cases either AN or AC aluminas were used as supports for the metal. The Cu AN sample was obtained by impregnation of alumina AN sample dispersed in methanol with a solution of copper acetate dissolved in methanol. The copper percentage was of 6 %. The stoechiometric amount of water required to hydrolyze the copper acetate was added to the solution of copper acetate and methanol. The alcohol was evacuated under hypercritical conditions (with respect to methanol) in the autoclave in order to preserve the porous texture of alumina (ref. 6) and to precipitate copper hydroxyde or oxide. Copper oxide on alumina (aerogel) was reduced at 383°C in flowing hydrogen for 3 hours prior each catalytic test. The Co AC sample was prepared by impregnation of alumina AC with 002(CO)8 in solution in n-heptane and a subsequent reduction in flowing H2 at temperatures increasing gradually from the ambient to 200°C in order to obtain finely divided Co crystallites in the micropores (ref. 4). The weight percentage of Co was 6.8 %.
339
Catalytic test: methanol conversion A microreactor was used under dynamic conditions with partial pressures of CH 30H varying from 45 to 200 torrs, complemented to 1 atm. by He. The temperatures of the reaction were set between 350 and 400°C and the mass of catalyst was fixed to 200 mg. In all tests the main products were dimethyl ether, methane, CO, CO 2, H20, C2H4 and C2H6. Over the alumina sample AN, after 1 hour of reaction near 400 uC, the conversion of CH 30H was in the range 80-90 % giving essentially dimethyl ether. In the presence of copper catalyst Cu AN, 70 mg of catalyst, at 383°C, and in comparable conditions the CH 30H conversion was also close to 80 % and gave essentially dimethyl ether. Further studies are now in progress with the aim for a modification of the selectivity by an increase or a decrease of the acidity of the catalysts (Bronsted and even Lewis type acidities). Catalytic test : CO/H 2 reaction The cobalt containing Co AC was tested in a microreactor under flow conditions, at 200°C for a HZ/CO ratio equal to 2. It was observed that the product distribution (from C1 up to C7' saturated hydrocarbons only) did not follow the Shulz-Flory law and presented a cut-off for C1-C2 hydrocarbons in agreement with the results concerning the influence of the pore texture of the support containing cobalt (7). According to Vanhove et al (ref. 4) such a cut-off is typical of Co particles finely dispersed in a microporous matrix. The slope a of the Schulz-Flory distribution was rather small (a = 0.5) if compared with that found for cobalt supported on conventional carriers (a = 0.8). This behaviour shows that the polymerization of hydrocarbon initial radicals is favoured by the presence of microporosity in alumina. CONCLUSION The preparation of amorphous aluminas by a sol-gel process in a non aqueous medium is the only one which leads to a solid with the microporous, "zeolitic" type, texture. Indeed, no microporous aluminas prepared by conventional methods are known. The monomodal distribution of the microporous volume in the range ° is not sensitive to the heat treatment. The amorphous of a pore radius of 10 A state is preserved up to 800°C. REFERENCES 1 S.J. Teichner, C.R. Acad. Sci. 246 (1958) 1429. 2 G. Pajonk, M. Repellin and S.J. Teichner, Bull. Soc. Chim. Fr., 1333 (1976). 3 G" Pajonk, B. Pommier, M. Repellin and S.J. Teichner in "Characterization of Porous Solids" S.J. Gregg, K.S.W. Sing and H.F. Stoeckli Eds, SCI London 135 (1979).
340
4 D. Vanhove, Z. Zhuyong, L. Makambo and M. Blanchard, App. Catalysis, 9 (1984) 327. 5 R. Sh. Mikhail, S. Brunauer and E.E. Bodor, J. Coll. Interf. Sci. 26 (1968) 45. 6 M. Astier, A. Bertrand, D. Bianchi, A. Chenard, G.E.E. Gardes, G. Pajonk, M.B. Taghavi, S.J. Teichner and B.L. Villemin in "Preparation of Catalysts" B. Delmon, P.A. Jacobs and G. Poncelet Eds., Elsevier Amsterdam, 315 (1976). 7 D. Vanhove, unpublished results.
341
DISCUSSION B. CORMACK: Did you have a mechanism for the production of microporous aluminas from methanolic solutions of aluminium nitrate: in view of our earlier results with aluminium nitrate dissolved in various alcohols, which showed that by altering the alcohol and concentration of aluminium nitrate in solution we could prepare macroporous (mostly tertiary alcohols), mesoporous (some secondary alcohols) and microporous aluminas (primary alcohols) by this method. G.M. PAJONK : It has to be emphasized that in our method alcohol is the only solvent of the Al hydrated salt. Al hydrated salts are mainly soluble in CH30H and C~5CH also we haven't tried any other alcohol as a solvent. The role of the alcohol is to produce an amorphous precipitate of alumina when gaseous NH3 is added. The size of molecule of alcohols does not appear clearly in this process neither its surface tension properties. E.K. POELS : As you claim that methanol conversion is almost limited to production of dimethyl-ether because of weak Bronsted acidity of the alumina surface due to the anhydrous conditions during its preparation : did you observe an induction period or change in product distribution with time due to modification of the alumina surface by HZO formed during dimethyl-ether formation from methanol ? G.M. PAJONK : The conditions of the preparation of amorphous alumina are not entirely anhydrous because Al(N03)3 crystallizes with 9 molecules of HZO and A1Cl~ with 6 molecules. Now it is possible that the strong acidity of A1Z03 remalns neutralized by NH3 even after the decomposition of NH4N03 or chloride at 400°C. However, there is a change in product distribution with time, the first product is the DME then appear CZH4 and CZH6. F. DI RENZO : In your preparation, alumina needs to avoid any contact with water. Once the catalyst is activated, is its microporous structure stable in the presence of the amount of water formed by the catalytic reaction? G.M. PAJONK : It has been shown previously that when amorphous Al alkoxide is hydrolized by HZO vapour at ZOO°C and above, the alumina which is formed remains amorphous (Imelik et al., J. Chim. Phys., 51, 1954, 51). Therefore, in the case of formation of water vapour at reaction temperature, in the present work the amorphous structure of A1Z03 is probably preserved. This point is being studied now. V. TWIGG: The importance of removing potentially explosive ammonium nitrate from precipitates before they are calcined should not be overlooked in practice. Similarly, care should be taken with solutions of nitrates in organic solvents. G.M. PAJONK : The NH4N03 is not liable to decompose in the presence of liquid methanol. The amount of NH 4N03 which remains on the dry extracted alumina is negligible. XU Xiaoding : I suppose you reduced the catalyst before reaction with MeOH. Do you see CO and HZ (products of MeOH decomposition) besides DME ? G.M. PAJONK : In the case of Cu deposited on amorphous A1 Z03 products other than DME are CO, CZH4 and CZH 6 in a small amount. The reduction of Cu was carried out at 383°C in flowing HZ' CO and HZ were not detected during activation in NZ or He.
342
A. HOLT: Have you examined the preparation of alumina by first precipitating in aqueous media, washing out nitrate or chloride ions and finally boiling in methanol ? G.M. PAJONK : It is not possible to convert a crystallized A1 203 into amorphous A1 203 by boiling it in CH30H.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
THE RELEVANCE OF KNEADING AND EXTRUSION PARAMETERS IN MANUFACTURE OF ACTIVE POROUS ALUMINAS FROM PSEUOOBOEHMITES
343
THE
A. DANNER AND K.K. UNGER Institut fur Anorganische Chemie und Analytische Chemie Johannes Gutenberg-Universitot, 6500 Mainz, FRG
ABSTRACT In order to study the effect - on crushing strength, attrition resistance, specific pore volume, and pore volume distribution of extrudates - of kneading and extrusion parameters in the manufacture of active porous alumlnas, a commercial pseudoboehmlte (Pural SB from Condea Chemie, Brunsbuttel, FRG) was employed. A minimum kneading torque was required to prepare an extrudable paste, predetermined by extrusion conditions. The torque had an upper limit with respect to certain extrudate properties. The amount of water and the concentration of nitric acid as binder solutions were found to have a decisive bearing on the mechanical properties and the distribution of the pore volume. Crushing strength and attrition resistance were primari Iy controlled by the amount and distribution of macropores. On increasing the speed of revolution of the kneader, the necessary kneading time was reduced. Extrudates with smooth surfaces were obtained at a low speed of the extruder screw and employing dies with a bore diameter of one third of their length. INTRODUCTION The shaping of catalysts, carriers and adsorbents has today become a powerful method far adapting active solids to the specific operational condit ions required in catalyt ie, isolation and purification processes. Moreover, the confection processes allow a tal loring of the pore structure and the chemical composition, and thus have a direct impact on the desired catalytic activity and selectivity. Unlike other methods of powder agglomeration, extrusion gathers together the knowledge of several disclpl ines, namely powder rheology, colloid and sol id state chemistry, mechanical engineering, etc. As a result, the final properties of the extrudates are dependent on many variables and are known to be control led at the stages of powder formation, paste processing, extrusion, and activation of extrudates to active solids.
344
In a previous paper we reported on the role of the parent alumina in the manufacture of active porous aluminas by extrusion (ref. 1). The present paper focusses on the bearing that kneading and extrusion parameters have on the crushing strength, attrition resistance, specific pore volume, and pore size distribution, employing a commercial pseudoboehmite as parent material. EXPERIMENTAL M~1~Ll~1~~_~gYlQm~nl_~ng_QL2~~gYL~~
The starting material was Pural S8, kindly suppl led by Condea Chemie GmbH, 8runsbuttel, FRG, with the fol lowing properties: specific surface area (as) 250 m2g-1, specific pore volume (vp) 0.58 ml g-1 total water content 23 ~ (wjw) and mean particle size 34 ~m (refs. 2,3). Water and nitric acid solutions were employed as binder solutions. The kneader was composed of a 3 I container (Erweka. Heusenstamm, FRG) with two twisted kneading arms and counter-rotating motion, driven by an Infinitely variable mechanism. Torques of up to 60 Nm could be appl ied at constant speed of revolution in the range of 0 to 100 rpm. The torque. proport ional to the paste viscosity, was recorded as a function of time by an instrument from Kipp und Zonen, Kronberg. FRG. The extruder was a single screw extruder (Col I in, Ebersberg, FRG) with the fol lowing screw measurements: length 500 mm and diameter 40 mm. The speed of the extruder was Infinitely variable in the range of 0 to 125 rpm. The die plate employed had 6 circular openings, each with a diameter of 8.45 mm. The discharged green extrudates were cut into pieces of 10 mm length by means of a granulator constructed In our laboratories. The paste was prepared accord Ing to the fo II owIng st andard procedure: 750 g of 0.5 ~ (wjw) nitric acid solution were kneaded with 1000 g Pural S8 (water content adjusted to 23 ~ (wjw» at 40 rpm until a torque of 20 Nm was reached. Then, portions of 50 g of 0.5 ~ (w/w) nitric acid solution were added (up to 1100 g of peptizing liquid) while maintaining a constant torque of 20 Nm. The paste was extruded at 10 rpm. and the green extrudates dried in an oven for 24 hours at 383 K. They were then subjected to a heat treatment in a muffle furnace (from Naber. Lilienthal, FRG) at 773 K for 3 hours. ~b~Lg~1~Ll.~112n
The crushing strength (cs) of the activated extrudates was measured with equipment (T8 24) from Erweka, Heusenstamm. FRG.
345
The cs-value was averaged from 40 measurements. The attrition resistance (as) was determined as the amount of fines (in grams) formed when 20 g of activated extrudates were subjected to rotat ion at 100 rpm far 24 hours in a 250 ml round flask. The specific pore volume and the pore volume distribution were obtained by means of mercury poroslmetry in a home mode device which allowed pressures up to 450 MPa (ref. 4). In order to calculate the pore volume distribution according to the Washburn equation, a contact angle of mercury of a - 140' at 293 K and a surface tension of mercury of 480 mN m- 1, also at 293 K, were chosen (ref. 5). RESULTS AND DISCUSSION
Pural S8 was chosen because of Its excellent kneading and extrusion properties. It consists of porous particles of 7 nm overage pore size, which represents aggregates of primary crystallites of 10 nm size. During kneading, shear forces partially disrupt the aggregates; this yields submicron fines which envelop the remaining particles and function as solid interpartlculate bridges. With nitric acid solution, or even water, a thin liquid film is formed, which adheres to fines and aggregates. The rapid access of the liquid is accelerated by the high surface area of fines and primary particles. When nitric acid is employed. a peptization of the particles takes place at the surface. On removal of the peptizlng liquid during the calcination step a recrystal I izat ion occurs, thereby strengthening the extrudate by forming interparticle sol id contacts. It should be emphasized. however. that for each peptizing acid - and a given alumina - a critical concentration exists, beyond which a noticeable dissolution of the particles storts, and extrudable pastes cannot be obtained (refs. 1.6). For Pural S8, the critical concentration of nitric acid is about 3" (wjw) (ref. 1). Eii~~1 __2i_12Lgy~_gl_~2nIlgnl_IQ~~g_2i_L~~21Yl12n~ In order to achieve reproducible properties at extrusion, the prepared paste should possess a minimum viscosity, corresponding to a certain minimum torque at the end of the kneading process. The minimum torque under the conditions appl ied in these extrusion experiments was around 15 Nm. At a lower torque the particles did not cohere sufficiently, but adhered to the wall of the extruder screw. The green extrudates then showed the typical (spl it) Christmas tree-pattern at the outer surface. Furthermore. It was
346
necessary to adjust the Pural S8 to a constant water content of 23X (wjw) before use, because during storage the water content increased sl ightly as a result of water adsorption from concurrently
the
pseudoboehmite according to
kneading conditions
powder
of
23"
(w/w)
had
to
water
be
content
the standard procedure with 70% of
the air;
adjusted.
the
The
was
mixed
total
binder
solution and kneaded at constant 40 rpm. When a torque of 20 Nm was reached, the next portion of binder solution was added. The change
in
torque
as
a
function
of
the
successive amounts of binder solution, two h 0 u r s
0
Differences
f
k n e ad i n g. in
the
not
time,
adding
con s ton t pas t e vis cos i t Y was a chi eve d .
0
properties
of
the
extrudates -
portionwise or by continuous addition of were
kneading
is shown in Fig. 1. After whether
by
the binder solution-
observed.
30,----------
--,
~
1500 1400
r
0 L OJ
c "0
1300
'"
"U "U
0
E
z
c
1200
.,
1100
..
1000
0
20
0
:J
c II
:J
tr
L
o
I-
900
0
z'"
.. .
::c
<,
-!
800.
I f)
700
ci
... 0
10
20
30
40
120
50
130
600
...
500
<
C :J 0 E
KnQodlngtlmQ In mlnutQQ
Fig. 1: Variation of the torque (in Nm) of the paste at kneading as a function of the kneading time in minutes at constant speed of revolution (40 rpm). AI I other conditions correspond to those described under standard procedure in the experimental section. Raising
the
torque
from
revolution did not affect further
increase
to
65
20
to
30
Nm at
constant
speed
of
the extrudate properties. However,
Nm gave
rise
to
drastic
changes.
a
The
347
crushing strength increased by roughly a factor of 4, whi Ie the attrition resistance concurrently increased by a factor of 10. Whereas the total specif Ie pore volume did not chonge, the relative proportion of the three pore sizes (7, 15 and 1300 nm) did. At a torque of 65 Nm, the macropores (mean pore diameter 1300 nm) disappeared completely, in controst to a torque of 20 to 30 Nm. The frequency of pores with a maximum at 7 nm decreased sl ightly, whi Ie the peak of those pores with a maximum at 15 nm nearly doubled. The changes in the pore volumes of the individual pore sizes are probably associated with the destruction of secondary aggregates (peak at 1300 nm) and primary particles (peak at 7 nm) and the formation of primary aggregates (peak at 15 nm). The disappearance af the macropore volume at a torque of 65 Nm undoubtedly causes the Increase in crushing strength and attrition resistance. ~11.~1_21_!h._gm2Yn!_2i_~lng.L_~21Yl12n~ In one series of experiments the amount of water as binder solution was Increased from 950 9 to 1300 g, employing the standard procedure. The torque was kept at 30 Nm at 40 rpm. The results are collected In Table 1. Table 1: The crnount of water and the concentration of nitric acid solution as binder solutions, and their effect on the properties of the extrudates (conditions see text). extrudate properties
Variable
pore maxima of the relative distribution
crnount of water in g (standard procedure)
cs/N mm-2 amount of f lnes/g
vP/ml g-1
950 1100 1300
143
0.06 0.46 1.95
0.52 0.53 0.67
7.15 7.53 7.15
13.94 15.38 16.03
1355 1082 2300
7.40 0.89 0.22 0
0.73 0.61 0.57 0.52
7.94 7.54 8.03 7.54
17.05 14.19 13.58 11.44
1559 1358 1546
97
31
pd1 /
1T11
pd2/1T11 pdy'1TII
concentration of nitric acid solution in 911000 9 Pural 58 0 2.75 5.5 11
39 59 77
100
-
348
An increase In water reduced the crushing strength from 140 to 40 N m~2. Simultaneously, the attrition was raised from 0.06 g to 2.0 g. The total specific pore volume increased slight Iy. A detailed analysis of mercury Intrusion data showed that the specific pore volume from the 7 nm pores decreased, while the specific pore volume due to 15 nm pores nearly doubled. Furthermore, the pore size distribution of the macropores (of about 1300 nm) become extremely wide. These results again demonstrate that the specific pore volume of the macropores controls the mechanical stabi I ity of the extrudates, i.e., the crushing strength and the attrition resistance. The decline of the 7 nm pore size pores in specific pore volume was probably caused by the lengthening of the kneading period from 32 minutes (at 950 g of water) to 160 minutes (at 1300 g of water) to achieve a torque of 30 Nm, causing a pronounced attrition of the aggregates. The results obtained with water were simi lar to those obtained with 0.5~ (w/w) of nitric acid solution. t11~£1_21_1bl_~20£~ni£~i12n_21_nllLl~_g~lg_121Yi12n. Under otherwise constant conditions (torque 20 Nm, speed of revolution 40 rpm) the amount of peptizlng acid was varied from zero to 11 g HN0 3 per 1000 g of Pural 58. With neat water the crushing strength was relatively low (about 40 N mm2) and rose to 100 N mm 2 at maximum amount of acid (see Table 1). A high crushing strength was accompanied byhighattrition resistance, and vice versa. The amount of acid also affected the total pore volume, which decl ined from 0.73 to 0.52 ml g-1. Related to that there was a change in the pore size distribution: with an increasing quantity of acid the proportion of macropores became smaller and smaller and finally disappeared completely at 11 g HN0 3/1000 g of Pural 58. Concurrently the pore size maximum of the primary aggregates shifted significantly to lower values. With neat water as binder solution, the shear forces during kneading form fines which cause only loosely bound aggregates with low mechanical stabl I ity. The acid, however, chemically attacks the surface of the particles through peptizatlon. The dissolution of the surface accelerates with higher acid concentrations. As a consequence, the primary particles are attacked more strongly; they then fill the interst ices between the part icles of the primary and secondary aggregates. The recrystal I ization of the peptlzed particles during calcination thus leads to a stronger interparticle bond than in the ease of a low acid concent rat ion.
349 ~11~£i_21_1h~_~e~~g_21_L~~21YilQn_21_1h~_~n~Qg!L~ In two sets of experiments with water and 0.5% (wjw) nitric acid solution as binder, corresponding to a torque of 20 and 30 Nm, respectively, the speed of revolution was adjusted to 30, 40 and 50 rpm. The speed of revolution was not observed to exert a significant influence on the mechanical properties of the extrudates. There were sl ight changes in the specific pore volume, which might be attributed to the difference in the binding mechanisms of water and nitric acid. The only effect worth noting was that the greater the speed of revolution, the shorter become the period for obtaining a paste with the desired viscosity, i.e. torque. Doubling the speed of revolution reduced the necessary kneading time by roughly half.
~11!£1_21_1h~_~e!!g_21_L~~21Yl12n_2i_lh!_~~lLyg~L_~~L!~~ Extrusion of the paste at various speeds of revolution did not couse remarkable differences In the properties of the extrudates. With increasing speed, the attrition resistance and also the specific pore volume were reduced. Simi lor to the situation met at varying binder acid concentrations, the proportion of macropores was reduced at high speed, but to a much smaller extent. The sl ight decrease of the macropore volume was probably responsible for the change in the attrition resistance. ~!1!£1_21_1h~_gl~~ Various constructions of dies - in terms of length to diameter ratio - and tests with these dies under otherwise constant conditions showed that they borely affected the properties measured. The only notable inf luence observed was on the surface pattern of the extrudates. The smoothest extrudates were obtained with bore diameter to length ratios of between 1 3 and 1: 6. Extrudates with smooth surfaces were usually obtained at low speed of the extruder and employing dies with bores of a diameter of one third of the length.
REFERENCES 1. W. Stoepler and K.K. Unger, in: G. Ponce let, P. Grange and P.A. Jacobs (Eds.), Preparation of Catalysts III, Elsevier, Amsterdam, 1983, p. 643. 2. Company broschure, Condea, Br unabu t t e l , FRG. 3. W. Stoepler, Ph.D. thesis, Johannes Gutenberg-Universitat, Mainz, FRG. 1983.
350
4. E. Schadaw and K. Unger, High Temperatures - High Pressures 9 (1977) 591.
5. S.J. Gregg and K.S.W. Sing, Adsorption, Surface Area and Porosity, Academic Press, London, 1982, pp. 173-194. 6. K. Jiratova, L. Janacek and P. Schneider, In: G. Poncelet, P. Grange and P.A. Jacobs (Eds.), Preparation of Catalysts Ill, Elsevier, Amsterdam, 1983, p. 653. ACKNOWLEDGEMENT This project was supported by the Arbeitsgemeinschaft industrleller Forschung, Koln, FRG, grant No. 5588.
351
DISCUSSION ZHAO JIUSHENG : It will be more practical if you determine the attrition resistance in fluidized-bed way instead of rotation method. K. UNGER: Pellets applied in fluidized bed reactors usually have diameters between 0.2 and 0.5 mm while the extrudates produced in this study were much more larger and are preferably employed in fixed bed reactors. Attrition of extrudates under these conditions occurs during transportation, filling the reactor, etc .• Hence the rotation test used simulates much better this kind of attrition than one based upon the fluidized bed mode.
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B. Delmon, P.Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
353
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in TheNetherlands
AMORPHOUS ALLOYS AS CATALYSTS OR CATALYST PRECURSORS M. Shibata1 and T. Masumot0 2 1Central Researoh Laboratries, Idemitsu Kosan Co., Ltd., 1280 Kamiizumi, ~odegaura, Kimitsu-gun, Chiba 299-02 (Japan) Researoh Institute for Iron, Steel and Other Metals, Tohoku University, Sendai 980 (Japan)
SUMMARY The applioation of amorphous alloys in oatalysis has been investigated. Reoent information is reviewed in terms of the advantages and disadvantages of amorphous alloys, espeoially with pre-treatment. INTRODUCTION Amorphous alloys have attraoted attention owing to their superior eleotrical, magnetic, mechanical and chemical properties. However, reports on their chemical properties are relatively rare, and most studies are related to corrosion, electrocatalysis and catalysis. In corrosion studies, Masumoto et al. (ref. 1) compared the corrosion rates of amorphous Fe-Cr-13P-7C alloys with those of crystalline Fe-Cr alloys in 1 mole dm-3 HCl. Although the amorphous Fe-P-C alloy had a lower resistanoe to corrosion than pure Fe, when Cr metal was added to the amorphous Fe-P-C alloy its resistanoe to corrosion was oonsiderably improved, and on addition only 8 at.% of Cr the amorphous alloys were completely passivated. On the other hand, orystalline Fe-Cr alloys were not passivated even by the addition of up to about 13.5 at.% of Cr. From the results of X-ray photoelectron spectroscopic (IPS) analysis of these alloys, Hashimoto et al. (ref. 2) conoluded that a passive film which oonsists of mainly a hydrated chromium oxyhydroxide is rapidly formed on the surfaoe and provides a high resistanoe to oorrosion beoause of a lack of lattice defects and segregations. In electrocatalysis studies, an amorphous Pd-Ni-P alloy was reported to have a relatively high electrocatalytic actvity in the oxidation of methanol after activation treatment of the surface (ref. 3). The surface activation was effected as follows: a zinc-plated amorphous alloy was heated below the crystallization temperature in order for zinc to permeate into the alloy, then the zinc was extracted with a concentrated aqueous alkaline solution.
354
This process provided the electrode with a higher surfaoe area. More detailed results on corrosion and electrocatalytic studies have been reviewed recently (refs. 4-6). Therefore, in this paper, we shall discuss the principal motivations for applying amorphous alloys in heterogeneous oatalysis. In the first section, we describe the preparation teohniques for amorphous alloys, in the second section the advantages and disadvantages when amorphous alloys are used as catalysts and in the third seotion the application of amorphous alloys in catalysis. PREPARATION The techniques for the preparation of amorphous alloys are classified into (L) vapour and sputter deposition, (ii) eleotroplating and chemioal plating, (iii) melt-quenching, and so on. Eaoh method has both advantages and disadvantages. Vapour and sputter deposition This technique is usually applied for the production of thin films. Vacuum evaporation is the conventional method for produoing films with a production rate of 0.5-1.0 nm/s at 10-3-10- 4 Pa. The advantages of this method are the simple apparatus required and easy operation. This is also suitable for the formation of an amorphous single-component metal because of easy cooling of the substrate. The disadvantage is that it is unsuitable for mass production because of its low producti vity. Also, the precise control of the composition of the deposited alloy is diffioult, especially when the difference in vapour pressure between the oomponent metals of the target is relatively large. The kinetic energies of atoms deposited on the substrate by sputtering are higher (over 10 eV) than those in vacuum evaporation (ca. 0.1 eV). Therefore, in the case of sputtering, the degree of contaot between the deposited film and the substrate is high and the composition of the film is almost the same as that of the alloy target, even if the differenoe in vapour pressure between the component metals of target is relatively large. When using plasma or magnetron sputterring the production rate is relatively high (1 ~m/min). However, the temperature of a substrate rises during sputtering and a small amount of gas enter the film because of the higher pressure (1-0.1 Pa) in the sputtering system. The range of amorphous phase formation by sputtering depositions is wider than those given by plating and melt-quenching methods, whereas the productivity of amorphous alloys is lower. The coating of the inner surface
355
of reactors and porous substrates with amorphous alloys is attractive, if it is possible. Some amorphous alloy films prepared by sputtering have been examined as catalysts. Amorphous Ni-B films prepared by RF glow discharge sputtering were applied to a catalytic reaction (ref. 7). This system was discharged between the inside nickel and outside electrodes of a reactor cell. Highpurity Ar and B2H6 diluted with H2 were used as a sputtering atmosphere. Ni-B films with various compositions were readily obtained and the Ni-B films containing boron above 15 at.% changed into an amorphous state. This range of amorphous phase formation in the Ni-B binary system is wider than that prepared by the melt-quenching method (ref. 8). However, these films contained 15-20 at.% of oxygen because the preparation system was evacuated to only 0.1 Pa before glow discharge. Recently, Ohnuma et al. reported that amorphous ultra-fine alloy particles (10-100 nm) can be produced by sputtering on the rough substrate obtained by Ar+ ion etching (ref. 9). Plating In the electroplating technique, an amorphous film is deposi ted on an anode by electroreduction of metallic cations in an electrolytic bath. Electrolysis for preparing an amorphous Ni-P or Co-P film is carried out in an aqueous solution consisting of metal chloride (e.g., NiC1 2' 6H 20 or CoC1 2' 6H20) , phosphorous acid (H3P03) and phosphoric acid (H3P04) at 320-360 K and 0.5-4.0 A/cm2 using a copper or carbon anode (ref. 10). In another plating method (chemical plating), an amorphous film is deposited on a substrate by the reduction of metallic cations with a reducing agent. For example, an amorphous Ni-P film is prepared by reducing NiC1 2'6H20 in an aqueous solution with the help of sodium hYpophosphite as a reducing agent (ref. 11). Some kinds of amorphous alloy films have been prepared by these methods, but there have been few reports related to catalytic studies. However, it is interesting that a film with a large surface area can be easily prepared by these methods if an appropriate substrate is chosen. Me1t-guenching The melt-quenching technique using a gun method was introduced by Duwez et al. (ref. 12) in 1960. The "splat" produced was non-uniform and only a small amount was obtained (a few hundred milligrams). In the 1970s mel tspinning, single-roller and twin-roller type quenching methods were developed (ref. 13). The first "tapes" produced were 1-2 mm wide, 0.02-0.03
356
mm thick and several metres long. Later, by improving these methods, a tape 20 cm 11 \ wide, 0.1 mm thick and several kilometres long was prepared using a singleroller type quenching method. Fig.1 shows the meltquenching apparatus for preparing amorphous alloys. In each method, a melting alloy is made to spout from the nozzle of a crucible by gas pressure and then the Fig. 1. Schematic diagram showing amorphous alloy is quenched the principles of systems for on the surface of the manufacturing sheet amorphous products by continuous meltrotating cooling substrate. quenching. In single-roller type (1) Centrifugal; (2) single roll; (3) twin roll; (4) and (5) others. quenching the apparatus is simple and a large amount of amorphous alloys can be produced. As the productivity of the melt-quenching technique is the highest, the results of practical catalytic tests have been mainly obtained using amorphous alloys prepared by this technique. Using another technique by melt-quenching, amorphous alloy powders can be produced (ref. 14). At present two types of production technique are known: one is an atomization method in which amorphous powder particles are produced by rapid solidification of discrete droplets of liquid metal, and the other is a communition method in which powders are produced by chopping or comminuting ribbons or filaments of amorphous alloys. With these methods, an average diameter of powders in the range between several micrometres and several hundred micrometres is obtained, but the surface areas of these powders are not greater than 1 m2/g.
~
ADVANTAGES AND DISADVANTAGES OF AMORPHOUS ALLOYS IN CATALYSIS It is known that in heterogeneous catalysis on a metal-based catalyst, clusters of metal atoms on the bulk solid play an important role. In industrial catalysts, the bulk solid is·usually a carrier that modifies the
357
electronic properties of the metal clusters, stabilizes the desired high dispersion of the metal clusters or orders the structure around the metal clusters. Further, promotors that control electronic properties or prevent metal clusters from sintering are sometimes added. Therefore, delicate preparation procedures are required in order to obtain reproducible industrial catalysts that have such complex characteristics. An amorphous alloy seems to consist ideally of similar metal clusters having uniform electronic properties and a uniform environment. There is no porous structure on the surface of an amorphous alloy, and each metal cluster on its surface may be dispersed homogeneously and uniformly. Alloying is the most effective method for controlling the electronic structure of active sites because of the wide selectivity in the composition of amorphous alloys. For example, ultraviolet photoelectron spectroscopic (UPS) measurement of CO adsorbed on an amorphous Ni-Zr alloy shows that alloying of zirconium to nickel increases the amount of dissociative adsorbed CO only by addition of 5 at.% of Zr (ref. 15). On the other hand, the surface of an as-prepared amorphous alloy, especially produced by the mel t-quenching method, is covered by an oxide layer. This oxide layer has to be removed by a pre-treatment procedure for exposing acti ve metal species. Further, the surface area of an amorphous alloy is small, i.e., 0.1-0.01 m2/g, which is too small to use as an industrial catalyst because of the low producti vity per unit weight of a catalyst. An amorphous phase is a non-equilibrium state, so that crystallization of an amorphous alloy occurs above the crystallization temperature (Tc). Further, Tc decreases under a reactive atmosphere such as hydrogen (ref. 16). Therefore, the determination of catalytic activities over an amorphous alloy has to be carried out at temperatures sufficiently below Tc• However, crystallization caused by oxidation of an alloy results in an active catalyst, as described later. APPLICATION IN CATALYSIS Table 1 summerizes reports on catalysis over amorphous alloys. Pretreatment for removing a surface oxide layer were usually oarried out before oatalytic reactions. In some instances (refs. 28,29,32), oatalytic aotivity was not observed when no pre-treatment was carried out. Therefore, in this seotion, examples of the use of amorphous alloys as catalysts are olassified with respect to the pre-treatment prooedures.
358
Table 1 Applications of amorphous alloys for catalysis Reaction
Alloy system
Reaction conditions Pressure Temperature
IMPa
Main product
ref.
IK
CO hydrogenation
Fe-Ni-P-B Pd35Zr65 AU25Zr75
0.1 0.1 6.0
473-623 533 500-550
Methanol synthesis
CU70Zr30
6.0
500
methanol
42
Ammonia synthesis
Fe91Zr9
600-690
ammonia
48
CO disproportionation
NiZr2
0.1
513-593
CO 2, carbon
24
HYdrogenation of olefins
Ni-P,Ni-B Ni-P
0.02 0.013
373-523 373-523
butenes, butane butenes, butane
28 7
HYdrogenation of propenal
Ni28Ti n
0.1
333-39S
propanal
32
Isomerization and deuteration of cis-cyclododecene
PdSOSi20
0.1
29S
transcyc10dodecene, cyclododecane
17
Isomerization and deuteration of (+)-apopinene
PdSOSi20
0.1
(-)-apopinene, apopinane
20
Isomerization and hydrogenation of 1-hexene
PdSOSi20
0.13
hexenes, hexane
19
323
C1-C4 hydrocarbons 26 methane 37 methane 38
Table 2 Phases produced in splat-cooled and conventionally produced palladium and their catalytic behaviour Alloy PdSOSi20 PdSOSi20 Pd77Ge23 Pure Pd Pure Pd
Phases present Amorphous Amorphous and incipient crystallization Amorphous and incipient crystallization f.c.c. f.c.c. (large crystalline)
Activity (specific ratea)
Selgctivitl %d 1 %d 2
15.8
s.o
52.6 44.6
31.7 25.7
8.2
54.7
35•.3
12.1
40.6 .3.3.S
24.1 27•.3
a Product molecules per surface Pd atom. b Per cent monodeutero-trans-cyclododecene. c Per cent dideuterocyclododecane.
359
No special pre-treatment before reaction Brower et al. applied amorphous Pd80Si 20 and Pd77Ge23 alloys to isomerization and deuteration of cis-cyclododecene (ref. 17). Quenched amorphous alloys produced by a shock-tube type splat cooling device were loaded into a reactor. These amorphous alloys showed different selectivities with respect to trans-cyclododecene and dideuterated cyclododecane from their crystalline Pd system, as shown in Table 2. After cleaning by brief Art ion bombardment, examination of the surface electronic structure by both Auger electron spectroscopy (AES) and UPS showed no difference in the spectra between the amorphous alloys and the alloys fully crystallized in situ (ref. 18). This indicates that the difference in the selectivity of this reaction between amorphous alloys and the crystalline alloys is due to the difference in the surface topography. However, no significant differences between amorphous and crystalline phases of Pd80Si 20 prepared by the melt-spinning method were found with regard to catalytic selectivity in the cis-trans isomerization and double bond migration of n-hexenes and stereochemical hydrogenation of «-pinene under the reaction conditions as shown in Table 1 (ref. 19). Over Pd catalysts, the isomerization and the deuteration of (+)apopinene as a probe molecule for distinguishing the relative percentages of terraces, ledges and kinks available on the surface were reported by Smith et a1. (ref. 20). Fig. 2 shows the plots of (+)-apopinene isomerization vs deuteration over Pd-containing catalysts. The crystallized alloy showed a higher ratio of isomerization to deuteration than the parent amorphous alloy. The isomerization of (+)-apopinene reflects the total number of ledge, kink and terrace sites, whereas hydrogenation reflects only the number of kink sites (refs. 21-23). From the above results, it was concluded that the surface structure of an amorphous alloy is not twodimensionally random (flat), but is three-dimensionally random (hilly or rolling). Smith et al. (ref. 24) also reported that amorphous Fe80B20 and Zr2Ni alloys appear to be more active in the dissociation of CO than the supported nickel catalyst as shown in Table 3. Further, M6ssbauer spectroscopy, electron microscopy and differential scanning calorimetry (DSC) have shown that the quenched alloy structure is sensitive to the melt-quenching technique used to produce the alloy. The activation energies of amorphous Fe80B20 alloys for crystallization determined by the Kissinger method based on peak temperature shift are 2.83 eV for a shock-tube flake and 2.44 eV for
360
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a
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Fig. 2. Plot of (+)-apopinene isomerizaton vs addition over catalysts with different surface characteristics. <)jamorphous PdSi alloy, • jcrysta1lized PdSi alloy, V jO.03% PdfA1 203 ' e ;0.05% PdfA1 203 ' o ;0.05% PdfA1 203 sintered, • ;1.0% Pdf A120]' [) ;1.0% PdfA1203 sintered.
a
a
50
Boron content / % Fig. 3. Initial product distribution in the hydrogenation of 1,3-butadiene over amorphous Ni-B films as a function of the boron concentration. Oj1-butene,Ajtrans-2-butene, O;cis-2-butene,ejbutane.
Table 3 Dissociation kinetics of CO as measured by thermogravimetric analysis for amorphous alloys and a conventional catalyst Catalyst Ni/SiO rAl 203 (60% N~) Zr2Ni (amorphous) FeSOB20 (amorphous)
Reaction temperature fK
100
Rate fg m-2 min
593
4.2 x 10-6
595 595
5.2 x 10-4 7.0 x 10- 5
361
a tape produced by the single-roller method. The same authors modelled the bulk amorphous structure on the basis of the procedure of Gaskell (ref. 25), who used a trigonal prism to model a short range order in an amorphous PdaoSi 20 alloy. These prisms are connected at edges and corners, but never faces. Energy is minimized by the method of conjugate gradients. For the bulk FeaOB20' such a model requires that the prism is shared so as to bring the Fe and B atoms into contact. There were many protuberances on the FeaOB20 glassy surface, which was simulated by a computer, as estimated in ref. 20. The hydrogenation of 1,3-butadiene over amorphous Ni-B films prepared by the RF glow discharge sputtering system was carried out by Imanaka et aL, (ref. 7). The films were transferred from the sputtering system to a conventional closed circulation system in a glove-box exchanging the atmosphere with nitrogen. The initial product distribution in this reaction depends on the boron concentration of Ni-B films as shown in Fig. 3. In Ni2p 3/2 IPS spectra a satellite peak approached the higher binding energy side of the main peak. With increasing concentration of boron, the intensity of the satellite peaks increased and the main peak became less asymmetric. This indicates that unoccupied d-holes of Ni are partly filled by the electrons transferred from B to Ni, increasing the electron density of nickel. The change in product distribution could be associated with the electron density of nickel, that is, the hydrogenation ability of nickel increased with increasing electron density of nickel. These results indicate that the catalytic activity of a quenched amorphous alloy is sensitive to the preparation method, because its surface structure and the electronic structre of metals on the surface are sensitive to the preparation method. Reduction with hydrogen or carbon monoxide Yokoyama et a1. first demonstrated catalysis over amorphous alloys (ref. 26). They studied fifteen amorphous Fe-Ni base alloys prepared by the single-roller type quenching method to determine the activities of amorphous alloys for CO hydrogenation. As shown in Fig. 4, pre-treatment of an amorphous Fe-Ni-P-B alloy with hydrogen at 593 K increases the activity during the initial period of operation, but the activity rapidly decreases and approaches a constant value. Pre-treatment with helium containing 1.26% of CO decreases the initial activity, but the same constant activity is attained. The catalyst without any pre-treatment exhibits a change in activity intermediate between the H2 and the CO pre-treatments, attaining
362
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Reaction time / h Fig. 4. Activity of amorphous Fe40Ni~OB~P16 catalyst at 593 K as a fUnc'\;ion of time following various pre-treatment procedures. t:::. jin H2 at 593 K for 0 h, jin H2 at 593 K for 1 h, [Jjin H2 at 593 K for 2 h, ~ jin CO at 593 K for 2 h.
I I
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o
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the same activity. These results indicate that the pre-treatment procedures affect the initial surface conditions, presumably related to the reduction state but, regardless of the pre-treatment, similar surface conditions are attained under the reacting gas stream. With the exception of only one composition, the stable activities of the amorphous alloys are from several to several hundred times higher than those of crystallized alloys of the same composition, as shown in Fig. 5. Both the amorphous and crystallized catalysts show an activation energy and a rate law. These results suggest that the active sites are of a similar nature but different in number as for the different structures and compositions of the catalysts. Carturan et ale (ref. 27) applied amorphous and crystalline Fe-Ni-Cr-P-B alloys prepared by using the single-roller type quenching method to catalysts for the hydrogenation of ethyne. A relatively mild H2 pretreatment (3 h at 563 K) was chosen to avoid structure relaxation and
363
nucleation. Catalytic activities over the amorphous alloys were lower than those over the crystallized alloys except for Fe-Ni-P-B, but the selectivities towards ethene over the amorphous alloys were higher than those over the crystallized alloys. The results show that the surface is not completely covered by an inacvtive boron or phosphorus oxide layer and/or a part of them is easily reduced only by hydrogen treatment for several hours in the case of the amorphous Fe-Ni-P-B system. Further, it is thought that the surface of FeNi-P-B was partially oxidized during CO hydrogenation by the CO 2 or H20 produced. Treatment with acid followed by reduction Yoshida et al. reported the application of amorphous Ni81P19 and Ni62B38 alloys prepared by the single-roller type melt-quenching method in the hydrogenation of olefins (refs. 28,29). However, the quenched alloys showed no activity for hydrogenation of olefins even after hydrogen reduction of the alloys at 573 K for 6 h. Succesive treatments of alloys with dilute HN0 3 (1.5-6 mole dm-3), oxygen at 373-523 K and hydrogen at 570 K led to high catalytic activity. The effect of the pre-treatments on amorphous alloys was studied by observing the IPS spectra, which revealed that (i) treatment with dilute nitric acid was effecient in removing the stable surface oxide layer, (ii) treatment with oxygen resulted in partial oxidation of Ni, P and B and (iii) treatment wi th hydrogen reduced Ni ions to the metallic state but did not reduce the oxidized P and B. The electron density of Ni atoms after the above treatments was lower than that of pure Ni metal. These results indicate that the electron transfer from electron rich Ni atoms to more electronegative P or B atoms by alloying may be important in enhancing the catalytic activity of Ni atoms, which may be promoted by partial oxidation of P or B atoms. The high activity of amorphous alloys may be ascribed to the specific formation of partially oxidized P or B atoms on the catalyst surface. In other work by the same group (refs. 30,31), the effect of oxygen pretreatment on the activation of amorphous Ni-P alloys was studied by determining the number of surface nickel atoms by temperature-programmed reduction (TPR) of the oxidized species and recording the IPS spectra of the surface layers. Oxygen treatment followed by reduction with hydrogen created surface metals that interacted with phosphorus, oxidized nickel and phosphorus species. The number of surface nickel atoms and the strength of interaction increased with increasing temperature of oxygen treatment and
364
was correlated with catalytic activity. The decrease in catalytic activity caused by over-oxidation and crystallization was concluded to be due to the disappearance of homogeneity of the amorphous alloys and reconstruction of atomic structure involving aggregation of nickel species. Funakoshi et ale reported the hydrogenation of propenal over amorphous and crystalline Ni28Ti72 alloys, ultra-fine nickel particles and aluminasupported nickel (ref. 32). The Ni-Ti alloys were covered with a titanium oxide layer and were inactive even with hydrogen treatment at 573 K. They were washed either with hydrofluoric acid solution or sodium hydroxide solution, and then were reduced in a hydrogen flow at 473 K before use. The Ni-Ti alloys showed a high selectivity for the hydrogenation of propenal to propanal, but over other nickel catalysts 1-propanol was produced together with propanal. They attributed this unique selectvity of Ni-Ti alloy by alloying to electronic effects. After the treatment of amorphous alloys with an acid, the surface oxide layer is almost completely removed, but the surface metal atoms are partially oxidized. Subsequent oxidation provides a catalyst in which the Ni is highly dispersed in the metalloid oxides. other pre-treatments The amorphous alloy ribbon prepared by using the melt-quenching method has a small surface area and an inactive oxide layer on the surface. Therefore, hydrogen or acid treatment followed b,y oxidation and reduction is neccesary for the generation of catalytic activity. Yamashita et ale solved these problems by mechanical pulverization of amorphous alloys using a vibratory rod mill (ref. 33). As shown in Fig. 6, with powders of amorphous Ni62B38 alloy under 300 mesh, the number of surface Ni atoms, the BET surface area and the initial rate of hydrogenation of ethene were naturally greater than those of powders over 300 mesh. Further, the number of surface Ni atoms per unit BET surface area and relative turnover frequency were greater than those of powders under 300 mesh. These results and the IPS spectra of pulverized powders indicate that the electrondeficient nickel atoms exposed to the surface from the matrix of the amorphous alloy are more active than the nickel atoms on the quenched surface. Amorphous alloys as catalyst precursors Experimental observations of the in situ activation and formation of oxides such as Si0 2, Ce02' Zr02 and Th02 during the hydrogenation of CO over intermetallic compounds as catalyst precursors have been reported (refs. 34-
365
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366
Table 4 T.F. (turnover frequency) values for methanation catalysts
ECH b
T.F' CH a
Catalyst
ref.
4 4 /sec- 1 x 10- 3 /kJ mole- 1
100 635 Pd35Zr65 112±7.5 0.32 Pd7Si02 98.6±2.9 27 Pd/Ti0 2 130 Raney Ni 45 117 91 Ni/ Zr 02 528 Ni/Ti02 181 181 Ru/Al203 a at 548 K, PCO=25 kPa, PH =76 kPa.
37 49 49 50 51 52 52
2
b Activation energy for methanation.
Table 5 Results ~f CO hydrogenation at 6.0 MPa (CO/H2/Ar=31/64/5) at a flow-rate of 2.1 1 h" at the stationary state and the surface area after the reaction Oatalyst Temperature/K Rh25Zr75 Pd25Zr75 OS25Zr75 Ir25Zr75 Pt 25Zr75 AU25Zr75
670 473 653 653 663 523
00 conver- 00 base selectivity/% OH4 MeOH CO2 Others a sion/% 0.0 11.7 85.8 2.5 8.7 21.7 15.0 79.0 2.1 3.9 66.7 13.8 0.0 19.5 2.5 0.0 34.5 3.8 51.9 13.6 61.7 29.1 0.0 2.6 9.2 3.0 47.1 33.0 49.5 0.4
Sur~ac~1area
/m g <1 36 <1 <1 <1 51
a O2+ hydrocarbons, O2+ alcohols and dimethyl ether.
Table 6 Results of methanol synthesis reaction over Ou-Group IV metals as catalyst precursors at 6.0 MPa (CO/H2/Ar=31/64/5) Catalyst
Temperature/K
C~ozr~o
500
CO conversion /mole% CH30H CH3OCH3 CH 4 CO2 Total 8.4 0.2 0.0 0.3 8.9
cU70zr3~
505 553 541
4.4 6.1 7.9
0.1 0.8 1.2
0.1 0.7 0.7
0.0 0.6 0.7
4.6 8.2 10.5
CU60Ti~0
600
0.0
0.0
0.1
0.1
0.2
(amorp ous)
(crysta line) CU63.5Zr36 5 (amorphous' (amorp ous)
367
36). The same in situ activation was observed in amorphous alloys containing zirconium. A highly active methanation catalyst was prepared in situ from an amorphous Pd35Zr65 alloy prepared by using the single-roller type mel tquenching method in a reaction at 533 K and under atmospheric pressure (ref. 37). The catalyst prepared in situ was an unknown oxide composed of Pd. Zr and o. As shown in Table 4. the turnover frequency was greater than or comparable to those of the most active methanation catalysts known. We applied amorphous MZ5Zr75 alloys (M=Rh. Pd. Os. Ir.,Pt. Au) to CO hydrogenation at 6.0 MPa (ref. 38). Under these conditions, amorphous Pd25Zr75 and AUZ5Zr75 alloys were converted into ZrOZ and metallic Pd or Au. having high surface areas due to oxidation. and then attained the same properties as the supported catalysts. The oxidized gold catalyst had a high activity for methanation. as shown in Table 5. although there are a few reports that gold has an activity for hydrogenation (refs. 39-41). The electronic properties of gold in this gold catalyst are perhaps different from those of conventional supported gold catalysts. Some other hydrogenation reactions were investigated over the oxidized gold catalyst. We observed a similar in situ oxidation of amorphous Cu-Zr alloys in CO hydrogenation (ref. 4Z). Over Cu-Zr alloys. methane was produced at a low CO conversion in the initial stage and. after stabilization. methanol was produced. The selectivity with respect to methanol over the amorphous alloy was higher than that over the alloy crystallized by heat treatment at 800 K under vacuum, as shown in Table 6. These in situ oxidations are attributed to an increase in activity for CO hydrogenation in the initial stage. for example, as shown in Fig. 7. We measured the phase changes in the oxidation process using X-ray diffraction (XRD) and high-pressure differential thermal analysis (HP-DTA) in various atmospheres at 1.0 MPa. An amorphous Cu-Zr alloy was simply crystallized in Ar or CO. whereas in HZ it was converted into Cu and ZrHZ and in CO/Hz into Cu and ZrOZ• The same crystalline phases as appeared after the CO hydrogenation were observed only on heating in CO and HZ' Fig. 8 shows an -increase in surface area with elapsed time after treatment of an amorphous Cu-Zr alloy with air. NZ' CO/HZ and HZO/NZ in a continuous flow type reactor (ref. 43). In air and NZ' no increase in the surface area of the alloy was observed. whereas in CO/HZ a gradual increase in the surface area was observed and in HzO/N z an increase in the surface area was observed for 1 h at 500 K. It is known that Zr-based alloys begin
368
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tQ
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to oxidize at an accelerated rate after the initial protective oxidation. This breakaway phenomenon also occurs in water vapour. It has been suggested that the surface monoclinic ZrOZ layer provides poor protection against oxidation of Zr-based alloys (ref. 44). Further, Sato and Shimada reported that water vapour considerably accelerated crystalline growth for monoclinic and tetragonal ZrOZ and facilitated the tetragonal-to-monoclinic phase transformation. Microcrack formation on the surface accompanied this phase transformation (ref. 45). In the initial stage of the reaction, the amorphous Cu-Zr phase swells owing to absorption of hydrogen, and then a number of cracks are formed in the surface ZrOZ layer. CO and HZ penetrate into these cracks, and methanation takes place on the amorphous Cu-Zr phase. Zirconium is then oxidized by water, which is a by-product of methanation. Owing to the oxidation of zirconium, the zirconium content in the amorphous phase decreases, which makes the amorphous phase unstable and the alloy changes
369
into crystalline Cu and Zr02' At the same time, new cracks may form in a newly formed Zr02 layer beneath the original Zr02 layer. This process provides the catalyst with a large surface area that contains fine Cu particles dispersed in Zr02' Herman et a L, pointed out that Cu+ disolved in ZnO phase is an active site in the methanol synthesis reaction (ref. 46). The observation of CU2P3/2 and Auger peaks in the XPS spectra of the catalysts derived from an amorphous Cu-Zr alloy after the reaction showed that Cu+ and Cu2+ existed on the surface. An active site of the catalyst derived from an amorphous CuZr alloy might be Cu+, similar to the Cu/ZnO catalyst. A porous structure of ample surface area as a catalyst was successfully developed in an amorphous Ni-Zr alloy by CO hydrogenation (ref. 47). An amorphous Ni67Zr33 alloy of surface area 0.05 m2/g was polished with sandpaper, reduced by hydrogen at 523 K for 2 h to remove surface oxide layer and then treated in CO/H 2 at 538 K for 3 h. After these pre-treatments, the alloy was almost glassy from the XRD results and had a surface area of 5.3 m2/g. The specific reaction rate observed on this catalyst was higher than that on a fully oxidized catalyst. These treatments are effective for investigating true catalysis of amorphous alloys, as a "well defined" surface amorphous alloy is obtained. The in situ oxidation of an amorphous alloy provides a highly acti ve catalyst. The same oxidation is observed with a crystalline alloy, but the catalyst derived from an amorphous alloy is more active or selective in CO hydrogenation than that from a crystalline alloy. It is not clear why the catalytic activities are different between catalysts derived from an amorphous alloy and a crystalline alloy. However, this in situ oxidation may be a new method for preparing supported catalysts. CONCLUSION As mentioned above, the pre-treatment of an amorphous alloy makes the surface of a catalyst acti ve for catalytic reactions, owing to effective control of electronic properties and dispersion of metal clusters in the surface oxide layer. If the crystallization of the alloy is followed by oxidation, the in situ treatment such as CO hydrogenation is a new method for preparing a supported catalyst. An amorphous alloy as catalyst precursor is more active and selective than crystalline alloys for catalytic reactions such as hydrogenation of olefins and CO and isomerization. However, it is not clear that the surface of an amorphous alloy is really more active than a crystalline alloy for catalytic reactions.
370
Alloying in the amorphous state makes a precursor that has controlled electronic properties and controlled dispersion, and then a reproducibly controlled catalyst is derived from this precursor. Therefore, some surfaces of amorphous alloys may be an ideal model surface for studying elementary catalytic reactions. At present there are many problems in the application of amorphous alloys as catalysts; such as their low surface areas, high prices, the presence of inactive surface oxide layers and limitation on the reaction temperature to prevent crystallization. Efforts to increase the surface area by making ultra-fine particles or depositing an amorphous alloy on a suitable carrier and to find more active alloy systems at lower temperatures must be continued. Pre-treatments for removing surface oxide layers by mechanical or chemical methods must be improved so as to expose the active surface. Further, it will be necessary to carry out more studies on the stability of amorphous alloys under reactive atmospheres by holding amorphous alloys for long periods at a suitable temperature. REFERENCES 1 T. Masumoto, K. Hashimoto and M. Naka, Proc. 3rd Int. Conf. on Rapidly Quenched Metals, Brighton, in: B. Cantor (Ed.), 1978, The Metals Soc., pp, 435-438. 2 K. Hashimoto, M. Naka, J. Noguchi, K. Asami and T. Masumoto, Passi vi ty of Metals, in: R.P. Frankenthal and J. Kruger(Eds.), Princeton, The Electrochemical Soc., 1978, pp. 156-. 3 A. Kawashima and K. Hashimoto, Proc. 4th Int. Conf. on Rapidly Quenched Metals, Sendai, in: T. Masumoto and K. Suzuki (Eds.), 1982, The Japan Inst. of Metals, vol.II, pp, 1427-1430. 4 R.B. Diegle, J. Non-Cryst. Solids, 61&62(1984)601-613. 5 H. Jones, J. Mater. Soc., 19(1984)1043-1049. 6 K. Hashimoto, Amorphous Metallic Alloys,in: F.E. Luborsky(Ed.), Butterworth, 1983, pp. 471-486. 7 T. Imanaka, J. Tamaki and S. Teranishi, Nippon Kagakukaishi, (1985)10641069. 8 A. Inoue, A. Kitamura and T. Masumoto, Trans. Japan Inst. Metal., 20 (1979)404-406. 9 S. Ohnuma, Y. Nakanouchi and T. Masumoto, Proc. 5th Int. Conf. Rapidly Quenching Metals, in: S. Steeb and H. Warumont (Eds.), Wuerzburg, 1985, pp. 1117-1124. 10 A. Brenner, Electrodeposition of Alloys, Academic Press., N.Y., 1963 pp. 457-483. 11 A.W. Simpson and D.R. Brambley, Phys. Stat. Sol., 43(1971)291-300. 12 P. Duwez, R.H. Willens and W. Klement, J. Appl. Phys., 31 (1960)362-371. 13 Metallic Glasses, in: J.J. Gilman and H.J. Leamy(Eds.), AST, 1977. 14 S.A. Miller, Amorphous Metallic Alloys,in: F.E. Luborsky(Ed.), Butterworth, 1983, pp. 506-521. 15 R. Hauert, P. Oelhafen, R. Schloegl and H.J. Guentherodt, Solid State Commun., 55(1985)583-586.
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16 K. Aoki and T. Masumoto, J. Mater. Soc., 21(1986)793-798. 17 W.E. Brower, M.S. Matyjaszczyk, and T.L. Pettit and G.V. Smith, Nature, 301(1983)497-499. 18 S.D. Bader, L. Richter, M.B. Brodsky, W.E. Brower and G.V. Smith, Solid State Commun., 37(1981)729-732. 19 B.C. Giessen, S.S. Mahmoud, D.A. Forsyth and M. Hediger, Mater. Sci. Soc. Symp. Proc., 8, (Rapidly Solidified Amorphous and Crystalline Alloys), in:B.H. Kear, B.C. Giessen and M. Cohen (Eds.), 1982, pp, 255-258. 20 G.V. Smith, W.E. Brower, O. Zahraa, A. Molnar, M.M. Khan and B. Rihter, J. Catal., 83(1983)238-241. 21 G.V. Smith and D.S. Desai, Ann. N. Y. Acad. Sci., 214(1973)20-35. 22 G.V. Smith and J.R. Swoap, J. Org. Chem., 31(1966)3904-3906. 23 M.J. Ledoux and F.G. Gault, J. Catal., 60(1979)15-20. 24 G.V. Smith, W.E. Brower, M.S. Matyjaszczyk, and T.L. Pettit, Proe. 8th Inter. Congo Catal., (1984), pp. IV-236-272. 25 P.H. Gaskell, Nature, 276(1978)484-485. 26 A. Yokoyama, H. Komiyama, H. Inoue, T. Masumoto and H.M. Kimura, J. Catal., 68(1981)355-361. 27 G. Carturan, G. Cocco, E. Baratter, G. Navazio and C. Antonione, J. Catal., 90(1984)178-181. 28 S. Yoshida, H. Yamashita, T. Funabiki and T. Yonezawa, J. Chem. Soc. Chem. Commun., (1982)964-965. 29 S. Yoshida, H. Yamashita, T. Funabiki and T. Yonezawa, J. Chem. Soc. Farady I, 80(1984)1435-1446. 30 H. Yamashita, T. Funabiki and S. Yoshida, J. Chem. Soc. Chem. Commun., (1984)868-869. 31 H. Yamashita, M. Yoshikawa, T. Funabiki and S. Yoshida, J. Chem. Soc. Farady I, 81(1985)2485-2493. 32 M. Funakoshi, H. Komiyama and H. Inoue, Chem. Lett., (1985)245-248. 33 H. Yamashita, T. Kaminade, T. Funabiki and S. Yoshida, J. Mater. Sci. Lett., 4(1985)1241-1243. 34 V.T. Coon, T. Takeshita, W.E. Wallace and R.S. Craig, J. Phys. Chem., 80 (1976)1878-1879. 35 H. Imamura and W.E. Wallace, J. Catal., 65(1980)127-132. 36 F.P. Daly, J. Catal., 89(1984)131-137. 37 A. Yokoyama, H. Komiyama, H. Inoue, T. Masumoto and H.M. Kimura, Chem. Lett., (1983)195-198. 38 M. Shibata, N. Kawata, T. Masumoto and H.M. Kimura, Chem. Lett., (1985) 1605-1608. 39 G.C. Bond, P.A. Sermon, G. Webb, D.A. Buchanan and P.B. Wells, J. Chem. Soc., Chem. Commun., (1973)444-445. 40 I.W. Bassi, F.W. Lytle and G. Parravano, J. Catal., 42(1976)139-147. 41 J. Schwank, Gold Bull., 16 (1983) 103-110. 42 M. Shibata, Y. Ohbayashi, N. Kawata, T. Masumoto and K. Aoki, J. Catal., 96(1985)296-298. 43 M. Shibata, N. Kawata, T. Masumoto and H.M. Kimura, in preparation. 44P. Kofstad, High-Temperature Oxidation of Metals, John Wiley & Sons, Inc., 1966, pp. 179-188. 45 T. Sato and M. Shimada, J. Am. Ceram. Soc., 68(1985)356-359. 46 R.G. Herman, K. Klier, G.W. Simmons, B.P. Finn, J.B. Bulko, and T.P. Kobylinski, J. Catal., 56(1979)407-429. 47 Y. Shimogaki, H. Komiyama, H. Inoue, T. Masumoto and H.M. Kimura, Chem. Lett., (1985)661-664. 48 E. Armbruster, R. Schoegl, R. Hauert, P. Oelhafen, H.U. Kuenzi and H.J. Guentherodt, Mater. Tech. (Duebendorf. Switz.), 13(1985)92-97. 49 8-Y. Wang, S.H. Moon and M.A. Vannice, J. Catal., 71(1981)167-174. 50 R.A. Dalla Betta, A.G. Picken and M. Schelef, J. Catal., 40(1975)173183.
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51 R.A. Dalla Betta, A.G. Picken and M. Schelef, J. Catal., 35(1974)54-60. 52 M.A. Vannice, J. Catal., 37(1975)449-461.
373
DISCUSSION B. DELMON : Amorphization promoters are necessary in many cases (P, B, Silo My question concerns the respective influence of (i) the amorphous state and (ii) the promoters on catalytic properties. Do you know many examples where those two factors have been isolated from each other in a clear way? More specifically, do you know of quantum mechanical calculations investigating the electronic influence of P, B, etc ... in alloys? On the experimental side, do you know of double experiments with an alloy of the same composition, one sample being amorphous, the other one crystallized? -rfiese questions might concern, for example, PdSi, FeB, NiB (substances mentioned in your lecture). M. SHIBATA: I know no examples where those two factors have been isolated from each other in a clear way. However, there are some reports to show that the stable activities of the amorphous alloys are from a few to several hundred times higher than crystallized alloys of the same composition (ref. 26 etc.). L. GUClI : I have a comment on the importance of surface restructuring on the catalytic activity. In your lecture, you showed the precipitation of a-Fe during the CO+H2 reaction (G. Kisfaludi et al., Appl. Surf. Sci. 24 (1985) 225) on the surface of Fe80B20 amorphous alloy. However, if one changes the preparation condition, e.g. changing the rotation speed during the RQM, different surface crystalline structure is present on the surface of the same alloy which shows a much higher catalytic activity than that prepared at high disk speed. That is, conflicting experimental data is likely due to the fact that so far not too much attention was paid to the surface microstructure of an amorphous alloy. M. SHIBATA: Thank you for your comment. B. WALl: In the first part of your talk, you claim that the surface of an active amorphous catalyst is 3-dimensional random. Since the structure of the surface is very important for the catalytic activity, have you any evidence by surface sensitive method (no XRD) that the active surface is amorphous? M. SHIBATA : We have no evidence by surface sensitive method that the active surface is amorphous. We cannot decide whether the surfaces of amorphous alloys under reaction conditions are amorphous or partially crystalline.
J. BARRAULT : During the CO-H2 reaction, you said that alloys are transformed into metal-support systems, i.e. Ni-Si02, Ni-lr02' etc ... Our results showed that some metal-carbon species were also formed. So, have you some measure~ ments on carbon content of your catalysts and if so, what are their values? What is your feeling about the possibility of metal-carbon formation? M. SHIBATA : We measured the carbon content of our catalysts after the CO hydrogenation but it was lower than the limit of detection. However, in some cases there was a small shoulder peak on the main peak (from contaminated carbon) of Cl s in XPS measurement of the catalysts after the reaction. It should be considered that this peak showed carbide formation.
J. B.NAGY : Do you know any example in the literature using microemulsion or polymer solution to prepare amorphous alloy particles? In this case one could prepare monodisperse colloidal particles. M. SHIBATA: I do not know the preparation method of amorphous alloys from aqueous or organic solvent solution except for the plating method. It was reported that refluxing of a methanol solution of rhodium(III) chloride and poly-vinylpyrrolidone with sodium hydroxide gave a stable colloidal dispersion of fine rhodium particles (H. Hirai et al.,Chem. Lett. (1978) 545). But the authors did not mention that the particles were amorphous.
374
J.B. MICHEL: I would like to comment on Dr Nagy's question. We report in our paper the preparation of alloy Pd-Au particles prepared in colloidal suspension (J.B. Michel et al., Controlled preparation ... ). These are relatively monodisperse and can be prepared with a full range of composition PdxAul_~ (O<x
B. Delmon. P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
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NICKEL CATALYSTS DERIVED FROM EUTECTIC AND PRO-EUTECTIC NICKEL ALUMINUM ALLOY C. S. BROOKS Recycle Metals, 41 Baldwin Lane, Glastonbury, CT 06033, U.S.A. SUMMARY A new type of catalyst precursor, eutectic and pro-eutectic alloys of nickel and aluminum, leads to promising configurations of nickel catalyst. The novel precursor alloys arise from recent developments in metallurgical practice for preparation of nickel aluminum structural alloys for aerospace applications by directional solidification for synthesis of in situ composites and by rapid solidification of powders for synthesis of structural composites by hydrostatic press. INTRODUCTI ON The nickel aluminum eutectic and pro-eutectic alloys as alloy ingots and alloy powders provide precursors for high activity nickel catalysts in the form of unsupported skeletal fibers and powders and supported configurations such as tube wall or with alumina. Evaluations of surface properties and catalytic performance demonstrate that caustic activation of these alloys produces high dispersion nickel with promising activity for hydrogenation in the liquid phase and heterogeneous reaction of hydrogen and carbon monoxide to methane. ALLOY SYNTHESIS Directional solidification (OS) was originally developed to produce high strength single crystal whiskers for fabrication of in situ composites for structural applications (ref. 1). Unidirectional solidification produces unique metal structures via solidification of a metallic melt by uniaxial removal of heat. Parallel fibers or lamellae of a disperse intermetallic compound, in the for the nickel aluminum eutectic system (Fig. 1), alee present instance Al 3Ni oriented in the direction of heat removal and their dimensions depend on tile o rate of solidification or heat removal. Fibers down to about 800A can be produced using solidification rates of the order of 1000 cm/hr (ref. 2). The synthesis procedure is also applicable to a wide range of eutectic alloys For the nickel aluminum eutectic composition AI including AI 4Pt, AI 3Pd, gC02• with 6.2 wt% (2.8 at%) nickel in aluminum OS, the intermetallic Al 3Ni (25 at% Ni) forms highly oriented parallel faceted single crystal fibers. These fibers solidify within a matrix of almost pure aluminum (-0.2 at% Ni) in solid solution. Nickel aluminum melts with nickel contents greater than the eutectic composition (up to the peritectic composition of 28.5 wt% Ni) produce faceted dendrites of Al3Ni not fully aligned parallel to the axis of heat removal.
376
GROWTH DIRECTION
.
) ./. • e
--e • !~. I....
•
9 .-
• •
-; • •
--.-\.-
..c • •
•
•
L----J lpm
Fig. 1. Rod-like microstructure of AI-AI 3Ni eutectic. The second metallurgical process used for nickel aluminum alloy powder synthesis involves melt solidification into highly dispersed particle sizes in controlled atmospheres under conditions where cooling rates of the order of 105oK/ sec (ref. 3) are achieved. Nickel aluminum alloy powders have been prepared by a commercial rapid solidification procedure involving hydrogen solubilization and solidification rates of the order of 10 2 oK/sec (ref. 4). Rapidly solidified nickel aluminum powders (RSR) were also prepared in an inert atmosphere by a process still undergoing development (ref. 3) where solidification rates of the order of 10 5oK/sec were attained. PRECURSOR ALLOY COMPOSITION The phase composition of the precursor alloy plays an important role in determining the activity of the nickel catalyst derived from these alloys. The common commercial nickel aluminum Raney® type alloys usually have nickel contents in the range of 45-50 wt%. The equilibrium phase compositions of these alloys can be expected to consist of several components, Ni 2A1 3, Al 3Ni and the eutectic consisting of Al 3Ni with aluminum with a small amount of nickel in solid solutions. No significant NiAl phase is present for nickel contents <50 wt%. The eutectic and pro-eutectic nickel aluminum alloys produced by DS and the rapidly solidified powders under consideration lead to the important result that all the nickel is present in the Al 3Ni phase and that has been demonstrated to be advantageous for catalyst preparation (refs. 5-7).
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EXPERIMENTAL PROCEDURES Catalyst preparation Powder Preparations. Alloy powders (RSR) were synthesized by an experimental process (ref. 3) in the following compositions: Alloy wt% Nickel wt% Aluminum RSR 587 28.5 71.5 RSR 588 50.2 49.8 For comparison, Homogenous Metals, Inc. (HMI) powder (ref. 4) containing 42 wt% nickel and W. R. Grace 2813 with 49 wt% nickel were also tested. The activation procedure is arbitrarily standardized for conditions established to optimize surface structure for commercial alloys (ref. 8) by extracting 1 to 2 grams with 200 cm 3 of 20 wt% aqueous NaOH in a water-cooled Teflon beaker at 50 to 70°C. After 4 hours, the caustic solution is decanted and the activated nickel powder washed with distilled water to neutrality. The wet, activated, water-washed catalyst powders are washed in isopropanol to displace the water solvent before the liquid phase hydrogenations. Supported Nickel. A typical preparation procedure for the supported nickel is the reaction of the aluminum-nickel alloy (AI + Al 3Ni) pro-eutectiC powder system of 28.5 wt% nickel synthesized by explosive atomization with solubilized hydrogen (ref. 4) for a period of 1 hr at 70 to 90°C with 20 wt% NaOH to dissolve the aluminum matrix components. Catalytically active unsupported "skeletal" nickel Raney® type materials are usually produced by simply washing away the water-soluble sodium aluminate reaction products. For generation of a supported catalyst, however, a second synthesis stage consists of precipitating the aluminum as an insoluble aluminum hydroxide phase by the reaction of the sodium aluminate with a dilute inorganic acid and neutralization to pH 7. The final steps in the preparation of these supported catalysts are (1) the washing away of all the water-soluble residual reactants from the caustic activation and acid precipitation with distilled water, (2) drying the insoluble residue at 100 to 150°C in air, (3) calcining at 500 to 600°C in air and hydrogen reduction at 400°C to reduce the catalytically active nickel component and (4) sieVing to obtain a particle size range appropriate for catalytic test evaluation (refs. 9, 10). Nickel Carpet. The OS nickel aluminum alloys can be formulated into an attached carpet structure for catalytic tests (Refs. 11-13). In preparing these DS alloys for the heterogeneous catalytic reactions the activated nickel fibers are formed as a carpet structure with basal attachment of the fibers to unactivated nickel aluminum alloy in two configurations (Fig. 2). In one the nickel carpet is produced on the top and bottom of cylindrical disks formed from transverse sections (1/8" thick) cut from the OS ingots (1/2" to 1" in diameter). Upper and lower surfaces of these disks are activated under controlled
378
conditions to obtain a fiber carpet deptn of the order of 15 to 25 um. shows calibration plots of activation versus exposure time.
"Oil) CD .. .. CD
20 • 2%NaOH
(,)~
~.!!
",c
=~ -.-
Fig. 2
e
4%NaOH
o
8%NaOH
10
oz II)C')
:g<
20
40 80 80 Time in hours
100
Fig. 2. Al fibers (carpet); fiber length as a function of caustic leaching 3Ni time for the production of a catalytic nickel carpet on an Al 3Ni eutectic substrate: 6., 2% NaOH; 0, 4% NaOH; 0, 8% NaOH. Another configuration (Fig. 3) is prepared by cutting a OS ingot transversely and drilling holes longitudinally to the axis of the ingot. A typical slug cut from a 1" OS ingot would be 7/8" in diameter and 1" in length with 7 holes 1/32" to 1/16" in diameter. The inner surface of this alloy ingot is caustic activated under controlled conditions to obtain a fibrous nickel carpet of desired depth. The nickel fibers of this carpet on the interior of the longitudinal holes then approximate a radial distribution in a tube wall configuration. Several ingots mounted in sequence are fitted snugly into a stainless steel tubular reactor for bench scale test evaluations of the heterogeneous reactions. Characterization These experimental catalytic materials are characterized by elemental chemical analyses, gas adsorption, :transmission elect.ron microscopy' (tUl) and X-ray diffraction analyses (XRO). Volumetric gas adsorption procedures are used to determine the BET surface areas by low temperature nitrogen adsorption and the
379
2 in. Ingot Eutectic fiber oriented vertically
(1)
(2)
(3)
Fig. 3. Catalytic nickel carpet in cylindrical tube configuration used for heterogeneous catalytic performance tests. available catalytic metal surfaces by hydrogen chemisorption at room temperature (22 + 2°C). X-ray diffraction analyses are used to establish the gross crystallographic character of the alumina support and possible metal Al 203 reaction products. The electron microscopy has been conducted with a Philips EM-300 ° and capable of a direct magnifica~ microscope with a resolution limit of 2.3 A tion of 500,000X with the standard stage. The TEM methods have included both the replica method and ultramicrotome sectioning. Catalyst Performance The liquid phase hydrogenation catalytic performance tests are conducted with selected organic compounds (acetone, nitrobenzene, butyronitrile and dextrose) in a shaker-type low pressure hydrogenation apparatus with concentrations from 0.8 to 5 molar in appropriate solvent. The heterogeneous catalytic performance of the Ni-AI and nickel carpet catalysts are demonstrated in a single-pass 203 flow test in a differential bench scale tubular reactor at a pressure of 1 to 4 atmospheres for the methanation of synthesis gas at conversions of 2 to 20% for 3 vol% H2 plus 1 vol% CO in helium. RESULTS AND DISCUSSION Powder Catalysts Characterization. RSR 587 typifies a pro-eutectic alloy with 28.5 wt% nickel and provides a fine dendritic distribution of the Al 3Ni precursor phase. RSR 588 has a composition equivalent to that of the commercial Raney® aluminum nickel alloy with 50 wt% nickel. The metallographic microstructures for RSR alloys 587 and 588 compared with the microstructure of the unactivated Raney®
380
powder obtained by crushing bulk casting show the latter to have intermetallic phase separations much coarser and more segregated due to the slower cooling rates characteristic of the synthesis of these materials. The efficiency of the caustic extraction activation based on chemical analyses of the initial and activated alloys is shown in Table 1. All the RSR alloys show a conspicuously more efficient aluminum removal compared with the commercial Raney® alloy with 49 wt% nickel. The BET areas of the activated RSR alloys are consistently greater than those of the W. R. Grace alloy. This enhancement of BET area is indicative of the more complete removal of aluminum from the activated alloy. TABLE 1 Properties of caustic activated nickel
Alloy RSR 587 RSR 588 W.R.Grace 49 wt% Ni
BET area (m2/g)
Ni area (m 2/g)
Porosity to 0 104A (cm 3Rad. /g)
Average PoreoRadius
112 101 56
22 18 19
0.111 0.079 0.098
19.8 23.9 35.0
(A)
Nickel Composition Initial Act. wt% wt% 28.5 50.2 49.0
94.9 95.6 89.5
The surface property most clearly related to catalyst performance is the available nickel surface area as measured by hydrogen chemisorption. The two RSR alloys yield nickel surface areas within the relatively narrow range of 18 to 22 m2/g. The RSR 587 alloy provides a nickel surface area marginally superior (~2 std. dev.) to the 49 wt% nickel W. R. Grace alloy. Table 1 also summarizes data on the pore structure of these activated alloy powders based on low temperature nitrogen adsorption isotherms with the microo pore volumes in pore sizes up to 104A radius (uncorrected Kelvin radius). The pore volumes fall within a relatively narrow range of 0.079 to 0.111 cm3/g with the RSR alloy providing the largest value. The average pore sizes of the RSR alloys calculated from the BET areas consistently yield tne smaller pore o 0 sizes, ranging from 19.8 to 23.9A, compared with the pore size of 35.0A for the conventional bulk cast commercial alloy. The cumulative pore volume versus pore demonstrate that size and the incremental pore size distribution plots of ~V/~r the major portion of the pore volume for the RSR pro-eutectic alloy and the o W. R. Grace alloy is less than 100A with a conspicuous maximum for pore sizes o below 20A in the case of the RSR alloy. Catalyst Performance. Liquid phase hydrogenation tests were conducted with 4 organic compounds of diverse functionality to establish whether or not these caustic activated RSR nickel aluminum alloys show any significant advantages in specific hydrogenation catalytic activity compared with commercial alloys (Table 2). The selected hydrogenation reactions consisted of (1) acetone to
381
isopropyl alcohol, (2) nitrobenzene to aniline, (3) butyronitrile to butyl amine and (4) dextrose (d-glucose) to sorbitol. The tests were conducted under very mild conditions to accentuate differences in catalytic reactivity. The first 3 reactions were conducted at 22°C with the fourth reaction at 80°C (Table 2). TABLE 2 Hydrogenation reaction rates at 22°C (micromoles H2 absorb/g cat. min. ) Alloy
Acetone*
Nitrobenzene*
Butyronitrile*
RSR 58; RSR 588 W.R.Grace 49 wt% Ni
226 .± 20 197 180
239 .± 20 46 66
242 40 .±. 5 11
Dextrose** 1246 271
* Isopropanol solvent The RSR 587 pro-eutectic alloy shows consistent superiority over the W. R. Grace alloy for all four reactions (30% for acetone to a factor of 24 for butyronitrile). The RSR 588 with 50.2 wt% nickel, which is the counterpart of W. R. Grace 2813, demonstrates outstanding superiority only for butyronitrile. Supported Nickel Catalyst Characterization. Surface properties of the supported nickel on alumina catalyst prepared from the pro-eutectic HMI alloy powder with 28.5 wt% nickel are shown in Table 3. TABLE 3 Surface properties of 6.1 wt% Ni-AI 20 catalyst 3 BET N2
Catalyst pretreatment history Fresh After 2 hr methanation test at 324-428°C + 4 hr exposure to steam at 318-660°C
surface area
Metal surface area
H2 chemisorption, 22°C
(m 2g- 1 ) 288 173
94.9 77.5
6.9 5.7
Dispersion ms1mt * 0.18 0.15
* ms' number of surface atoms; mt, total number of atoms Catalyst Performance. The methanation reaction rates (Table 4) at 275°C and 1 atm. pressure for the experimental catalyst expressed as micromoles reacted per gram of catalyst or as a turnover number (~) (molecules reacted per metal site per second) are an order of magnitude greater than the reaction rates
382
published for nickel supported on alumina at comparable metal loadings (refs. 14, 15) with reaction conditions in close fit to the temperature, pressure and space velocity of the published rate data. In addition, the experimental nickel catalyst demonstrates a superior maintenance of the state of dispersion compared with the published results of Vannice (refs. 14,15). TABLE 4 Methanation of 3H2-CO over Ni-AI 203 catalysts
Catalyst 6.1 wt% Ni-AI 203 5.0 wt% Ni-AI 203
Cat. activity Metal Reaction condo dispersion ((~mol reacted) (g cat. s)-l N** Source Temp.oC, 1 atm ms 1mt * 275 0.18 (fresh) 93.0 0.98 Ref. 9 0.15 (used 0.27 (fresh) 0.037 Refs. 14,15 275 4.3 0.11 (used)
* ms' number of surface atoms; mt, total number of atoms ** ~' turnover number Nickel Carpet Catalyst Characterization. The morphology of the nickel fibers derived from OS eutectic alloys was established by controlled extraction using dilute mineral acids to selectively remove the matrix aluminum from the OS alloy ingots, leaving intact the Al 3Ni fibers. BET areas of the fibers so prepared give a value of 8.6 m2 g-l which corresponds to the cylindrical fiber diameter of 0.6 ~m characteristic of the eutectic alloy prepared by OS. Caustic extraction reduces the residual aluminum to 2 wt% and creates hollow porous nickel skeletons with a diameter approximating the 0.6 um typical of the precursor Al 3Ni phase with a BET surface area of 111 m2 g-l. The H2 chemisorption nickel area of the activated fiber with 98 wt% Ni was determined to be 40 m2 / g. The scanning electron microscopy (SEM) examination of these caustic-activated nickel fibers provides further confirmation of the morphology (Fig. 4). Electron diffraction measurements of the fiber extracted in mineral acid show a well defined orthorhombic crystallography typical of Al 3Ni (Fig. 4), whereas the caustic-extracted fiber provides an electron diffraction pattern typical of the cubic structure of nickel (Fig.5). The resistance of the fibrous nickel catalyst to sintering, as measured by changes in total surface area and metal surface area, during the methanation reaction (3 H2-CO) is shown in Table 5. The total surface area decreases by approximately 16% after 4 hours of exposure to the reactant mixture at 309°C at 1 atm. and the nickel metal surface area decreases by approximately 24%, indicating a promising stability toward thermal sintering in the reducing atmosphere and at temperatures typical of the methanation reaction.
383
Fig. 4. Al 3Ni fiber and X-ray diffraction pattern of orthorhombic structure after hydrochloric acid extraction of aluminum matrix from Al-AI 3Ni eutectic. Fig. 5. Porous nickel fiber and X-ray diffraction pattern of cubic structure after caustic extraction of aluminum From Al-AI 3Ni eutectic. TABLE 5 Surface properties of carpet nickel methanation catalyst BET surface area Catalyst Ni-Al eutectic ingot
Pretreatment history
(m 2g-1)
Freshly act. 4 hr exposure in 3H2-CO, methanation test at 309°C
43 + 5 36 :; 5
Metal surface area by H2 chemisorption H2 cilernisorption Calculated area ((m 2 Ni) ((~mol H2 adsorbed) (g carpet)-1) (g carpet) -1 ) 28 20
49 37
Catalytic Performance. The results of the methanation tests (Table 6) for the activated DS nickel carpet catalyst demonstrate catalytic activities (expressed as micromoles of carbon monoxide reacted per square meter of active surface of nickel per second), comparable within the same range as calculated from published results for plasma-sprayed Raney~ nickel on a catalyst insert from the
384
Pittsburgh Energy Technical Center (PETC) hybrid Synthane reactor catalyst (refs. 16, 17). TABLE 6 Methanation rates of 3H 2CO at 300-350°C over nickel catalysts ReactlOn rate «~mol CO)(m2 activated Ni)-1s-1) Catalyst Activated Ni-AI eutectic alloy 5.7 - 16.3 PETC hybrid reactor catalyst* 9.5 - 23.0 * Refs. 16, 17 CONCLUSIONS It has been demonstrated that nickel catalyst prepared from OS and rapidly solidified eutectic and pro-eutectic nickel aluminum alloy precursors provide especially active nickel for hydrogenation reactions with wide flexibility in configurations ranging from powders for liquid phase reactions to nickel supported on alumina and to tube wall for heterogeneous reactions. The potential for preparing catalysts from eutectic and pro-eutectic alloys of binary and ternary alloys involVing aluminum and the transition metals remains largely unexploited. REFERENCES 1 R.W. Kraft, U.S. Patent 3,124,452, 1974. 2 F.O. Lemkey and G.S. Golden, U.S. Patent 4,069,171, 1978. 3 C.S. Brooks, F.D. Lemkey and G.S. Golden, Raney Type Nickel Catalysts from RSR Atomization of AI-Ni Powders in: B.H. Kear, B.C. Giessen and M. Cohen (eds), Rapidly Solidified Amorphous and Crystalline Alloys. ElseVier, Amsterdam, 1982, pp. 397-407. 4 J.M Wentzell, Advanced Fabrication Techniques in Powder Metallurgy and Their Economic Implications, 42nd Meet. of the Structures and Materials Panel (Adv. Group for Aeronautical Research and Development) Ottawa, 1976. 5 R. Sassoulas and Y. Trambouze, Bull. Soc. Chim. Fr. 5 (1964) 985. 6 Alvin B. Stiles, U.S. Patent 3,627,790, 1971. 7 A.B. Fasman, V.F. Timofeeva, V.N. Rechkin, Yu. F. Klyuchnikov and I.A. Sapukov, Kinetics and Catalysis (USSR), Eng. tr ans. 13(6) (1972) 1347-52. 8 J. Freel, W.J.M. Pieters and R.B. Anderson, J. Catal. 14 (1969) 247. 9 C.S. Brooks, G.S. Golden and F.O. Lemkey, Surf. Technol. 11 (1980) 333. 10 G.S. Golden, F.D. Lemkey and C.S •. Brooks, U.S. Patent 4,287,096, 1981. 11 C.S. Brooks, F.D. Lemkey and G.S. Golden, U.S. Patent 4,086,264, 1978. 12 C.S. Brooks, F.D. Lemkey and G.S. Golden, Fibrous Eutectic Alloy Catalysts in: Conf. on In Situ Composites--III, J.L. Walter, M.F.Gigliotti, B.F. Oliver and H. Bibring (eds.), Ginn Custom, LeXington, MA, 1979, pp. 221-231. 13 C.S. Brooks, F.D. Lemkey and G.S. Golden, Surf. Technol. 16 (1982) 67. 14 M.A. Vannice, J. Catal. 37 (1975) 449. 15 M.A. Vannice, op. cit., 462. 16 R.R. Schehl, H.W. Pennline, J.P. Strakey and W.P. Haynes, American Chemical Society 172nd National Meet., Fuel Div. Preprints (1976) 2-21. 17 H.W.Pennline, R.R. Schehl and W. P. Haynes, Ind. Eng. Chern. Process Des. Dev. 18(1) (1979) 156-162.
385
Figs. 1 - 5 reprinted by permission of the publisher of SURFACE TECHNOLOGY, Vol. 16 (1982) 67-85: Fig. 2, p. 69; Fig. 5 a, b and Fig. 6, p. 71; Fig. 8 and Fig. 9, p. 76, respectively. Copyright of Elsevier Science Publishers B.V.
386
DISCUSSION J.A. SCHWARZ: Can you comment on the possibility of different types of Ni that might be present on the "supported Ni catalyst"? For example, Ni metal, Ni oxide or Ni aluminate. C.S. BROOKS: Hydrogen chemisorption was the primary measure of available nickel surface after reduction and indicated a moderate degree of dispersion of the order of 20-30 percent for the "nickel supported on alumina". A much higher order of nickel availability was observed for the "nickel carpet", approaching 90 percent of the BET nitrogen surface area. In neither case were more sophisticated surface spectroscopic procedures used to identify compound formation between the nickel and the 2-5 percent residual aluminium. P.G. MENON: How will the cost of these special preparation techniques compare with that of typical commercial catalysts in use today? A high cost is, of course, still acceptable if very superior properties can be built into these novel catalysts described by you. C.S. BROOKS: Firm manufacturing costs cannot be cited but it is certain that either the directionally solidified eutectic alloys or the atomized proeutectic alloy powders will have higher fabrication costs than the currently available bulk cast Ni/Al alloy catalyst precursors marketed by W.R. Grace, Degussa, etc. The enhanced cost of these new preparation routes will have to be justified by a demonstrated superiority in specific catalytic activity or unique features such as an opportunity to formulate tube wall configurations. The atomized powders are the most likely to provide competitive fabrication costs. B. DELMON : Your "carpet" catalysts are very exciting and I cite them often when I mention innovative approaches to novel catalysts. My question refers to your other catalysts, namely the rapidly cooled powder catalysts. I notice their composition (28% Ni) is different from that of commercial catalysts. In your comparisons with commercial catalysts, what do you believe to be due to rapid cooling, corresponding phase composition, grain size, and to composition? Don't you believe that certain textural differences (pore volume) might be modified in a still more drastic way by just changing the leaching conditions? C.S. BROOKS: The principal advantage of using a Ni/Al precursor alloy with no more than 28.4 wt% Ni is that the Ni is present exclusively as the NiAl 3 phase which has been established as the most readily reactive with caustic to yield active nickel. The parameters which playa role in determining the efficiency of aluminum removal by caustic dissolution consist of powder particle size, NiA1 3 dendrite size and precursor phase composition. Rapid atomization provldes unique control over the critical parameters of dendrite size not available in conventional bulk casting procedures. The conventional bulk cast commercial Ni/Al alloys can certainly be more exhaustively reacted with caustic, if Al removal is the only objective. The experiments reported used only relatively mild caustic activation and were designed to demonstrate the higher order of reactivity obtained for powders combining small NiA13 dendrite size and exclusive NiA13 compositon producing enhanced micropore volumes and large available nickel areas. J.R.H. ROSS: Can the system of preparation be used for three component systems, for example for Ni-Al-La materials? Our experience is that La gives substantial promotion of Ni-A1 203 catalysts and it might also do so in your system if it could be used. C.S. BROOKS: The results reported covered only the binary Ni/Al system. The atomized powder alloys can be fabricated from binary and ternary alloys of almost unrestricted compositions, except that one metal (AI, Si, Zn, Mg) must be
387
susceptible to chemical dissolution by caustic or acid. The use of the directional solidification procedure restricts selection to eutectic alloys. This limitation still provides a wide latitude for alloy selection using m"l1y transition and noble metals. Examples that have been synthetized are eutectic alloys of Au, Pt, Rh, Ir, Pd and Co with Al. J.W. GEUS : My question is concerned with the extent of hydrogen adsorption exhibited by the Raney nickel catalysts. Before you can reliably measure the extent of hydrogen adsorption, you have to remove the adsorbed hydrogen generated by the reaction of aluminum and water during the leaching procedure. We always have had many difficulties with the desorption of hydrogen prior to measure the extent of hydrogen adsorption. We have to go to elevated temperatures where the catalysts are sintering rather badly. If you are removing the hydrogen at high temperatures, what is the effect of the thermal treatment on the BET surface area of the Raney nickel? C.S. BROOKS: You are quite correct in observing that the gas adsorption procedures such as the BET nitrogen areas and the H2 chemisorption nickel areas do not provide a measure of active available nickel area of an activated "wet" catalyst. We used these gas adsorption parameters only as relative measures of surface morphology, recognizing that they are undoubtedly transformations of the "wet" morphology. The liquid phase hydrogenation rates we cited were on a weight basis. We feel, however, that the use of the gas adsorption values for total surface area and available nickel area are quite appropriate for the degassed and reduced catalysts used for the heterogeneous reactions and are correct for reference of these reaction rates.
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B. Delmon. P. Grange. P.A. Jacobs andG. Poncelet (Editors). Preparation of Catalysts IV
389
© 1987 Elsevier Science Publishers B.V.. Amsterdam - Printed in The Netherlands
AMORPHOUS METAL ALLOYS AS PRECURSORS IN CATALYST PREPARATION: AMMONIA SYNTHESIS CATALYSTS FROM AMORPHOUS Ni-Zr SYSTEMS
1 2 E. ARMBRUSTER, A. BAlKER, H.J. GUENTHERODT 1 , R. SCHLOEGL 3 , and B. WALZ 1
lInstitute of Physics, University of Basel, CH-4056 Basel, Switzerland 2Department of Industrial and Engineering Chemistry, Swiss Federal Institute of Technology (ETH), ETH-Zentrum, CH-8092 ZUrich, Switzerland 3present Address: F.Hoffmann-La Roche AG, CH-4002 Basel, SWitzerland
SUMMARY Amorphous metal alloys with the compositions Ni were and Ni 4Zr applied as catalyst precursors in ammonia synthesis. 24Zr76 The amorphou~ pr~2ursors were activated by in-situ reaction with the synthesis feed gas at 673 K. The activity of the resulting catalysts exceeded the one of the amorphous precursors by about an order of magnitude. Catalysts from amorphous precursors led to stable steady-state activities, whereas crystalline precursors of the same composition were only initially active and deactivated rapidly. Structural and chemical changes of the amorphous precursors during their transition to the stable ammonia synthesis catalysts were investigated using BET adsorption, differential scanning calorimetry, X-ray diffraction, photoelectron spectroscopy and scanning electron microscopy. The resulting active catalysts consist of small strongly disordered nickel particles embedded in a zirconium oxide zirconium nitride matrix. INTRODUCTION In the past two decades amorphous metal alloys have been investigated extensively due to their interesting electrical, magnetic and chemical properties. However, it was only recently when researchers started to explore the catalytic properties of these materials. A variety of amorphous metal alloys were found to exhibit exceptional activities and/or selectivities in reactions such as isomerization [1], hydrogenation of carbon monoxide [2], hydrogenation of olefins [3], oxidation of methanol [4], and ammonia synthesis [5]. Several activation procedures such as oxidation with nitric acid [3], reduction with zinc [6] and in-situ reaction with the feed gas [5,7] were applied to improve the catalytic properties of the amorphous metal alloy precursors. Although in some investigations efforts have been untertaken to understand the deeper chemical and structural changes occuring during the activation, understanding of these processes is badly needed. With this in mind, we have
a
390
studied the structural and chemical changes occuring during the transformation of amorphous nickel-zirconium precursors into highly active ammonia synthesis catalysts. EXPERIMENTAL Amorphous and crystalline Ni 64Zr76 and Ni 64Zr36 were prepared from premelted ingots using the melt-spinning technique. The obtained ribbons of ca. 30 ~m thickness were milled in a zentrifugal mill under liquid nitrogen which prevented both crystallization and oxidation. The resulting flakes of about 0.5 - 1 mm size were used in the catalytic experiments. Their amorphous structure was retained during the milling process, as evidenced by powder X-ray diffraction (XRD). The material was analysed for its composition by wet chemical methods; the composition was found to be at its theoretical value within the experimental error of 1 %. Ammonia synthesis was performed in a continuous tubular microreactor of 4 mm inner diameter. The catalyst bed was about 20 mm long and packed between two inert packings of glass beads. The temperature of the catalyst bed was measured by a chromel-alumel thermocouple positioned in the center of the catalyst bed. The ammonia formed in the reactor was measured by gas chromatography. Hydrogen (99.99) and nitrogen (99.99) were taken from commercial cylinders and purified using zeolite traps and oxysorb filters. The surface areas (BET) of the precursor alloys and the resulting catalysts were determined using krypton adsorption at 76 K and assuming a cross-sectional area of 19.6 A2 for the krypton atom, Powder X-ray diffraction measurements (XRD) were carried out on a Siemens D 500 instrument with copper radiation. Textural and morphological changes of the precursor during the activation process were followed by scanning electron microscopy (SEM) using a JEOLex 200 TEMSCAN instrument equipped with an EDX Link 270 window less detector. Photoemission experiments (XPS) were carried out in a Leybold LHS 10 electron spectrometer operating at 10- 10 Torr base pressure. XPS data reported were adequately processed on a DS 5 computer system. Clean surfaces were obtained by Ar ion sputtering at 80 K. RESULTS AND DISCUSSION The characteristic activation behavior of amorphous Ni 24Zr76 and Ni 64Zr36 precursor alloys is presented in Fig. 1. Both precursors exhibited a relatively low activity at the start of the long term runs. Although the most significant increase in activity were observed for both precursors in the first two hours, the activities increased steadily up to about 1000 hours on stream. Note that the steady-state activities measured after this period exceeded the activities of the original precursors by almost an order of magnitude. Most
391
important is the fact that crystalline precursors of similar composition did not yield a stable active catalyst under the conditions used. This result demonstrates that the metastable state of the amorphous precursors is essential for the formation of highly active catalysts.
'I
,...,
'I
1
'I
5 r-
N
IE
I
CI'
C
E
s:::
......
1 ....
Q)
+> RS s,
0.5
s:::
c
+> u
•
,•
•
• •••
oo::fSP
•••
•• ••
0 0 00 0
RS
Q)
"-
1
«lOS>
-
0
0
,I
-,
cPOOO~
0
0.1
..'.•
I
,I
•
a- Ni64 Zr36
0
a- Ni24Zr76
,I
10
100
.1
1000
time [hJ Fig. 1. Change in ammonia synthesis activity of amorphous nickel-zirconium alloys during activation in stoichiometric nitrogen-hydrogen mixture at 673 K and 9 bar total pressure. Surface areas of the original precursor materials were 0.023 m2jg for Ni 24Zr76 and 0.016 m2jg for Ni 64Zr36. Surface area measurements carried out on samples taken from the reactor during activation revealed that the surface of both precursors increased by more than an order of magnitude during the activation period. Differential scanning calorimetry (DSC) measurements performed on Ni 64Zr36 samples taken from the reactor during the activation showed a slow but gradual increase of the crystallinity of the precursor with time on stream. In contrast. with Ni 24Zr76 crystallization was considerably faster and occured already during heating the precursor to the reaction temperature of 673 K. This is due to the considerably lower thermal stability of Ni24Zr76 as compared to Ni 64Zr36• In order to gain some insight into the chemical and structural changes which occured during the transition of the precursor into
392
the active catalyst, both samples were investigated with regard to their bulk and surface chemical compositions and structure. Visual inspection of the catalyst bed revealed that the color of the catalyst changed with axial distance from the bed entrance. After thousand hours on stream the catalyst at the bed entrance was black and partially deactivated whereas the color of the large down stream section of the catalyst bed was golden. Catalyst samples taken from both sections were investigated using XRD. SEM and XPS. Samples designated by (a) were taken from the golden colored active part of the catalyst bed and samples designated by (b) from the small black zone at the bed entrance. The black color of the catalyst at the bed entrance was due to the formation of a carbonaceous deposit, as evidenced by XPS. The carbonaceous deposit was probably formed by coking of hydrocarbons which were synthesized from traces of carbon monoxide in the feed.
Ni(111) I as prepared
I
(milled)
~---
after catalysis
I Ni(200) IZrN(311) I I
I
i
30
i
40
i
50
i
60
I
70
2 - Theta
Fig. 2 Bulk structural changes of amorphous Ni 24Zr76 precursor during transition to active catalyst measured by XRD. (a) active catalyst from down stream section of catalyst bed; (b) partially deactivated catalyst from black zone at catalyst bed entrance.
393
Figure 2 compares the XRD patterns measured for the original amorphous Ni 24Zr76 precursor and the resulting catalyst taken from the active (a) and deactivated zone (b) of the catalyst bed after thousand hours on stream. The XRD pattern of Ni 64Zr36 was found to be almost identical and did not exhibit significantly higher relative intensities for the nickel reflections. The XRD patterns of the original Ni 24Zr76 precursor (top trace in Fig. 2) indicate the existence of traces of crystalline nickel (Ni (111) reflection) in the original amorphous material. The bulk structural properties of the resulting active (a) and deactivated (b) catalysts are considerably different. The XRD patterns of catalyst (a) indicate the presence of metallic nickel particles coexisting with NiO, Zr02, ZrN and some metallic zirconium. Note that zirconium nitride was not found in the deactivated catalyst (b). A comparison of the intensity ratios of the Ni(111) and Ni(200) reflections of catalysts (a) and (b) reveals that the nickel particles in sample (a) were small and strongly disordered, whereas sample (b) contained comparably large well developed nickel particles. This is supported by the intensity ratio 1(111)/1(200) of sample (b) which agrees well with that of bulk nickel. XRD analysis of the precursor Ni 64Zr36 and the resulting catalyst led to results which were qualitatively the same as those presented for Ni 24Zr76 in Fig. 2. In order to complement the bulk structural information obtained by XRD and DSC, the amorphous precursor as well as the resulting catalysts were investigated with regard to the elemental composition and chemical structure of the surface using XPS. Surprisingly the amorphous Ni 24Zr76 and Ni 64Zr36 precursors exhibited both about similar surface composition with respect to the constituents. The surface contained about 90 ± 4 % zirconium existing prevalently as zirconium oxides, and 10 ± 4% nickel present as nickel oxide (ca. 70 %) and metallic nickel (ca. 30 %). More specific information about the chemical structure of the zirconium present at the surface is gained by the XPS spectra presented in Fig. 3. Two types of zirconium oxide can be distinguished, stoichiometric Zr0 and non-stoichiometric Zr02_x deficient in oxygen. The 2 non-stoichiometric oxide was prevalent on the surface of the original precursor and on the active catalyst (a). In contrast, the surface of the deactivated catalyst (b) contains mainly stoichiometric Zr0 Note that the surface 2• concentration of ZrN of the active catalyst (a) was small compared to the corresponding bulk concentration found by XRD (Fig.2). Gentle argon sputtering showed that in the subsurface region of the active catalyst (a) metallic zirconium was coexisting with oxygen deficient Zr02 . This result is in -x agreement with the bulk structural information obtained by XRD. The deactivated catalyst (b) exhibited a significant difference between the relative amounts of the different zirconium oxides on the surface and in the subsurface region. Zirconium dioxide was prevalent on the surface; however, with increa-
394
sing sputtering time the zirconium oxide became increasingly deficient in oxygen, leading to a prevalence of non-stoichiometric Zr0 • In contrast to 2-x the active catalyst, practically all of the metallic zirconium in the subsurface region was converted to zirconium oxides.
Zr 3d z
'N
-a ..... <1J E
N
o
'-
I 190
b.e.(eVl
185
Fig. 3 XPS core level spectra of Zr 3d measured for original precursor alloy, active catalyst (a) and deactivated catalyst (b). Vertical lines denote binding energies of the Zr 3d 5/2 line measured for reference materials. The el~ctron micrographs presented in Fi~ 5 illustrate the surface morphology of the active catalyst. The ragged surface structure depicted in the top plate originates from zirconium oxide and nitride in which small nickel particles were embedded. The presence of nickel in the zirconium oxide-zirconium nitride
395
matrix was confirmed by EDX. The bottom plate shows the stepped surface regions build up by the nickel particles, embedded in the electrically isolating zirconium oxide - zirconium nitride matrix. To gain some knowledge about the governing mechanism of ammonia synthesis performed on the nickel-zirconium alloys, we have investigated nitrogen adsorption on the amorphous Ni 64Zr36 precursor using XPS. Figure 6 presents the N ls spectrum after exposure of the clean precursor surface to a static pressure of nitrogen at 130 K.
Zr 3d
Zr 3d
Ni 24Zr76 used ob
0'
5'
45' 45' sputtered 190
185
b.e.leVI
180
190
185
b.e.leVI
180
Fig. 4 Comparison of XPS core level spectra of Zr 3d before and after argon sputtering of the active catalyst (a) and partially deactivated catalyst (b). The sputtering conditions were such that a homogeneous surface would have lost about 40 nm in 45 minutes. The central spectrum of catalyst (b) was obtained after 5 minutes of sputtering.
396
Fig. 5 Electron micrographs showing surface morphology of active catalyst. Top plate, morphology of surface build up by zirconium oxide and zirconium nitride; bottom plate, surface morpholoy of regions were nickel was prevalent (white lines arise from charging of the uncoated specimen surface),
397
N 1s 130 K IE -7 Torr N2 Ni64Zr36
Z I
Z
Z
I
I
I
j
410
405
i
400 leV]
i
395
precursor at Fig. 6 XPS spectrum of nitrogen adsorbed on amorphous Ni 64Zr36 130 K. Right hand peak which is due to dissociatively adsorbed nitrogen disappeared upon addition of a hydrogen pulse to the static nitrogen pressure. The spectrum indicates clearly the coexistence of molecularly adsorbed dinitrogen and dissociatively adsorbed nitrogen. The binding energies of the "end-on" bound dinitrogen and the atomic species are very similar to those determined by Grunze et. al. [8] for nitrogen chemisorption on Fe (111). However, the bonding interaction of the 1r-complex seems to be quite different, as judged from the different shake-up patterns (broadening of the 407 eV peak) observed on the two substrates. Note the absence of "edge-on" bound dinitrogen, which gave rise to a peak at 399 eV on Fe (111) [8]. Most interesting is the fact that the peak due to dissociatively adsorbed nitrogen disappeared completely after addition of a pulse of hydrogen into the static nitrogen atmosphere. This indicates that under the conditions used the most abundant reaction intermediate in ammonia synthesis over Ni 64Zr36 is dissociatively adsorbed nitrogen. Further evidence supporting this result has been obtained by analysis of kinetic data collected in the temperature range 570 700 K. The kinetics could be described well using rate expression (1) which is based on the assumptions that (i) dissociative adsorption of dinitrogen is the rate determining step; (ii) N* is the most abundant reaction intermediate;
398
and, (iii) the surface is uniform. k [N 2] a r = ---------
(1 + K [NH
(1)
3]/[H 2]1.5)2
ka represents the adsorption rate constant of nitrogen, K the constant of the equilibrium NH 3 + * ~ N* + 3/2 H2• and quantities in brackets denote partial pressures. CONCLUSIONS In-situ activation of amorphous Ni 24Zr76 and Ni 64Zr36 alloys generates highly active ammonia synthesis catalysts. The active catalysts contain strongly disordered nickel particles which are embedded in a zirconium oxide zirconium nitride matrix. The high activity of the catalysts results from a sequence of solid-state reactions which occur during exposure of the amorphous alloys to ammonia synthesis conditions. Crystalline alloys used as precursors did not yield a stable active catalyst under these conditions, which indicates that the metastable state of the amorphous alloy is essential for the formation of the highly active catalysts. XPS studies revealed that nitrogen is adsorbed in two forms on amorphous Ni 64Zr36, these are molecu1ary adsorbed dinitrogen and dissociated nitrogen species. Upon addition of hydrogen the latter species is rapidly converted to ammonia. ACKNOWLEDGEMENTS Financial support by LONZA AG, Switzerland and the Swiss National Science Foundation is gratefully acknowledged. REFERENCES
2
3 4 5 6 7 8
G.V. Smith. O. Zahraa, A. Molnar, M.M. Kahn, B. Richter and W.E. Brower. Jr, J. Cata1., 83 (1983) 238. G. Kisfaludi. K. Lazar, Z. Schay. L. Guczi, C. Fetzer, G. Konczos and A. Lovas, Appl. Surface Sci., 24' (1985) 225; A. Yokoyama. H. Komiyama. H. Inoue, T. Masumoto and H. Kimura, J. Catal., 68 (1981) 355; M. Peukert and A. Baiker, J. Chern. Soc., Faraday Trans. I, 81 (1985) 2797. S. Yoshida. H. Yamashita. T. Funabiki and T. Yonezawa. J. Chern. Soe•• Chern. Comm•• (1982) 964. K. Hashimoto and P. Kawashima, in: T. Matsumoto and K. Suzuki (Eds.), Proc. 4th Conf. Rapidly Quenched Metals, Vol. II, 1982, p. 1427. E. Armbruster. A. Baiker, H. Baris. H.J. Guntherodt, R. Schlagl and B. Walz, J. Chem. Soc., Chern. Comm.• (1986) 299. K. Hashimoto, A. Kawashima and K. Sakai, Jap. Patent No. 200'565 (1982). Y. Shimogaki, H. Komiyama, H. Inoue, 1. Matsumoto and H. Kimura, Chern. Let., (1985) 66l. M. Grunze, M. Golze, W. Hirschwald. H.J. Freund, H. Pulm. U. Seip. M.C. Tsai, G. Ertl and J. Kuppers, Phys. Rev. Lett., 53 (1984) 850.
399
DISCUSSION K. NOACK: Is there an equilibration of the N2 and the N adsorbed on the alloy? Is the ratio of the two N ls-signals connected to the activity of the catalyst? A. BAlKER: The dissociation of the molecularly adsorbed di-nitrogen N2 + 2* --+ 2N* is a slow process at l30K and consequently not equilibrated under the condition given. In contrast, the reaction of the dissociatively adsorbed nitrogen with hydrogen is comparatively fast and can be considered to be equilibrated even at l30K as indicated by our experiments. Concerning your second question, the ratio of the two N 1s-signa1s reflecting the amount of molecularly adsorbed dinitrogen and dissociative1y adsorbed nitrogen may be correlated to the catalyst activity, if the dissociation of dinitrogen is the rate determining step of the ammonia synthesis over our catalyst. Although this assumption is feasible, it has to be evidenced by isotope labelled experiments first. L. GUCZl : I have three questions. (1) Is specific rate referred to the total surface area of metal? (2) If yes, did you try - by measuring the surface composition (XPS) - to relate your specific rate to the nickel surface being active in nitrogen activation? (3) It seems to me that amorphous phase is not active but the active site formulation contains crystalline nickel particle. Was it not possible to correlate your activity to the nickel crystallites determined by XRS? A. BAlKER: The specific rates of ammonia synthesis quoted in the paper are referred to the total surface area of the active catalyst, as measured by krypton adsorption at 77K. These rates are certainly considerably lower than those which would be estimated if we would take into account the nickel surface area only. The elemental chemical composition of the active catalyst (after 1000 hours stream, see fig 1) as determined by XPS was: 4180 nickel, 19% zirconium, 97% oxygen, and 3% carbon. Thus a rough estimate shows that the ammonia synthesis rate referred to the nickel surface would be more than two times higher than the one referred to the total surface area. As concerns your third question, a correlation of the activity data with the nickel particle size as determined by XRD was not possible because we have to distinguish between at least two different nickel particles which could contribute to the observed activity; small disordered nickel particles and comparatively large crystalline nickel particles. Their contribution to the total nickel surface area could not be extracted from the XRD measurements. D.M. HERCULES: (1) How did you reference the binding energy scale in your ESCA spectrum of nitrogen? The value quoted for N2 is very high if the spectra are referenced to Au 4f7/2' (2) The chemical shift observed between N2 and N seems to be too large. Given that there is no charge difference between the two, I do not believe the difference in chemical shifts can be accounted for by final state effects which are usually ~ 2 eV for nitrogen. A. BAlKER : The binding energy of the EseA spectrum of nitrogen (fig 6) was referenced to Au 4f7/2' The chemical shift observed between N2 and N is indeed large. However, note that a similar large chemical shift for the two nitrogen species was reported by Grunze et al. (ref 8) for the adsorption of of nitrogen on Fe(lll). P.G. MENON: The equilibrium between adsorbed dinitrogen and atomic nitrogen may be slow because of the very low temperature of l30K used. At higher temperature (+ 300K), perhaps you may notice the fast transition of dinitrogen to atomic nitrogen, when the latter is removed from the surface by reactive scavenging using a pulse of H2.
400
A. BAlKER: I agree with your comment. The dissociation of dinitrogen is an activated process and should therefore be considerably faster at higher temperature. N. PERNlCONE : (1) You have concluded that you have obtained a highly active ammonia synthesis catalyst. Which is the term of comparison for this statement? (2) Which was the reason for the choice of a nickel alloy for your study, whereas nickel is not one of the most active metals for this reaction? A. BAlKER: We have recently reported on an iron catalyst for ammonia synthesis which was prepared from an amorphous FeglZr9 alloy by similar in-situ activation (ref 5). This catalyst which exhioltea a more than an order of magnitude higher activity for ammonia synthesis than polycrystalline iron was selected as a reference for the activity. The nickel catalyst presented in this paper exhibited an activity which was about 30% of the one of the FeglZrg catalyst, but considerably higher than the activity of po1ycrystal1ine iron. J.G. van OMMEN: Would you speculate a little on what actually the reason is why this Ni catalyst has such a high activity in NH3 synthesis. Is it for instance the loosening of N=N bond by the nickel by electron donation into the antibonding orbitals of the dinitrogen molecule? A. BAlKER: We attribute the high activity of the nickel in the Ni-Zr system to its structural properties only, and rule out electronic charge transfer effects. Evidence for this was obtained from microanalytical and photoelectron spectroscopy results. It appears that the mechanism of the dissociation of the dinitrogen on nickel is similar to the one suggested by Grunze et al. (ref 8) for the dissociation of dinitrogen on Fe(lll). Indication for this originates from the fact that the binding energies of the "end on" bound dinitrogen and the atomic species are very similar to those determined for nitrogen chemisorption on Fe(lll). However, the bonding interaction on the .-complex seems to be quite different, as judged from the different shake-up patterns observed on the two substrates.
401
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
The Production of Iron- and Nickel-Copper Alloy catalysts by means of Catalytic Decomposition of the Gaseous Metal Carbonyls A.F.H.
Wielers*, C.M.A.M. Mesters*, G.W. Koebrugge, C.J.G. van der Grift,
and J.W. Geus. Department of Inorganic Chemistry, State University of Utrecht, Croesestraat 77a, 3522 AD Utrecht, the Netherlands.
* present
address: Koninklijke/Shell laboratorium, Badhuisweg 3, 1031 CM
Amsterdam,The Netherlands. ABSTRACT Iron- and nickel-copper catalysts have been prepared by catalytic decomposition of
iron
order
or nickel carbonyl within a previously produced copper to
prevent
transport
limitations
during
iron
catalyst.
(nickel)
In
carbonyl
decomposition the reaction is carried out in a fluidized-bed reactor with small particles
(0.2 -0.4 mm in diameter) of the pure copper catalyst.
leads
exclusive
to
alloy formation between iron (or nickel)
This
and
assessed by means of infrared spectra of adsorbed carbon monoxide. (nickel)
is
completely taken up in the copper particles,
iron (nickel) ions with the support is prevented.
method
copper
as
As the iron
the interaction
of
As a result the reduction of
iron (nickel) is markedly facilitated if compared to the reduction of pure iron (ncikel)
catalysts.
Furthermore,
with
these
bimetallic
catalysts
deposition in the Fischer-Tropsch reaction is markedly suppressed, to
a better activity maintenance of the alloy catalysts in the
as compared to the monometallic iron (nickel) catalysts.
carbon
which leads
F.T.
reaction
402
INTRODUCTION Syngas reactions over monometallic nickel and iron catalysts often from
an
irreversible
Although
loss of activity due to
extensive
carbon
suffer
deposition.
the ability to form carbides is a prerequisite for these catalysts to
exhibit the desired activity,
extensive carbide formation leads to the
of
Owing to the mechanical strength of these carbon
filamentary carbon [1-4].
whiskers bed
the structure of the catalyst completely disintegrates and in
reactors reactor plugging occurs.
carbon
growth
proceeds
at
It has been proposed
the metal-support
metastable iron (nickel) carbide phase [2,3]. a
growth
carbide,
we anticipated
interface
that via
fixed-
filamentary
a
decomposing
As copper is not readily forming
that alloying iron or nickel with copper
may
be
effective in eliminating the growth of filamentary carbon. However, it has been established inactive
with
an
one (copper) leads to an appreciable loss of activity and changes
that
alloying an active metal (such as iron and nickel)
in
selectivity in e.g. copper
on
the Fischer-Tropsch reaction [5J. Therefore, the effect of
the activity and selectivity also needs to
be
investigated.
This
paper, however, mainly deals with the preparation of Fe-Cu and Ni-Cu catalysts. The
catalytic
especially
activity
of
these systems
will
only
briefly
in connection with the preparation procedure.
be
discussed
Full details on
the
catalytic properties of these alloy systems can be found elsewhere [6,7J. In the preparation of bimetallic catalysts an intimate contact between the atoms
of
alloy
formation
the constituent metals must be established in order to at a low temperature,
bring
about
thereby circumventing pretreatment
at
elevated temperatures that can induce sintering. Furthermore, alloying can only overcome
the unfavourable properties of the monometallic catalysts,
individual Finally,
alloy a
particles
uniform
have an
essentially
uniform
bulk
when
the
composition.
bulk composition is a prerequisite to obtain the
optimum
selectivity with alloy catalysts. As alloy formation with metals of group VIII and of group IB is difficult to achieve by conventional methods, preparation mentioned established and
properties
a new
method is described that meets as close as possible with the above requirements. that
both
With
copper
Cu-Fe as
single
compared
single crystal
surfaces
gaseous nickel and iron carbonyl
surfaces at elevated temperatures ( 423 Cu-Ni
generally
such as impregnation [8],
< T < 623
K) [9,10].
crystal alloy surfaces appeared to the constituting
metals.
For
it
decompose
had
been
on
these
The thus produced to
have
instance,
different on
the
403
Cu(110)-Ni markedly
and
Cu(lll)-Ni
surfaces the dissociation of
suppressed compared to that on pure nickel
molecular
adsorption
carbon
surfaces
monoxide
[11].
of CO proceeds on such an alloy surface in
is
However,
contrast
to
pure Cu surfaces [12]. Thermodynamics kPa)
show that at low nickel and iron carbonyl pressures ( <1
and at temperatures above room temperature both iron and nickel
can
be completely converted to CO and the respective metal
carbonyl
[13,14].
However,
due to kinetic hampering the decomposition rates are very low [15]. It has been found
that
in the presence of a metal surface the decomposition
reaction
is
markedly accelerated [16]. We have employed the catalytic decompostition of nickel and iron to
produce supported Fe-Cu and Ni-Cu alloy catalysts.
The reaction is carried
out within a previously produced supported Cu catalyst. homogeneous
distribution
particles, enables
the
us
of
nickel
and
iron
over
carbonyl
In order to achieve the
individual
copper
decomposition is carried out in a fluidised bed reactor.
to
use
small catalyst particles
and
thus
to
avoid
a
This
transport
limitations. In this paper we will discuss the experimental set-up utilized for the
preparation of the bimetallic catalysts and the most important
that
parameters
affect the decomposition of the gaseous metal carbonyl in the presence of
a supported copper catalyst.
Furthermore,
bimetallic
regard
catalysts
with
to
the advantages of the thus produced reduction
behaviour
and
catalytic
performance will be briefly dealt with.
EXPERIMENTAL We
will
catalysts
first
describe the experimental set-up
to
produce
by means of nickel (or iron) carbonyl decomposition on
bimetallic the
silica-
supported copper catalyst. The starting material is a copper-on-silica catalyst (30
wt.%
solution. suspension
copper)
prepared
by deposition-precipitation
from
a
homogeneous
Copper
ions are precipitated by means of urea decomposition in a of the silica carrier (Aerosil (Degussa) 200 m2/g) at 363 K. For
comparison purposes a pure iron catalyst (20 wt.%) was prepared in the same way [17,18]. to
obtain
relevant table 1
The loaded carriers were dried at 393 K for 16 h, a
fraction of particles with dimensions of 0.2
characteristics
grinded and sieved to
0.4
of the reduced parent copper catalyst are
mm. given
Some in
404 Table 1. Characteristics of the 30 wt.% CulSi02 catalyst
Weight percentage copper
29.1
(completely reduced catalyst) BET surface area (m2 / g)
450
specific copper surface area* per gram cat.(m 2) per gram eu (m2)
18.9 64.4
mean particle size (nm)
11
(average of the values obtained with electron microscopy and *extent of N20 chemisorption at 298 K)
A schematic production and
2.
nitrogen
representation
of the fluidized-bed reactor
used
for
of bimetallic copper catalysts with appendages is given in fluidising
The at
673 K,
catalyst was first calcined in a flow
of
the
figs. 1
oxygen
in
and subsequently reduced at 723 K for at least 16 h in
a
flow of about 5 vol.% hydrogen in nitrogen. The nitrogen flow necessary to keep the catalyst in a fluidised state varied between 30 and 160 mIls,
depending on
the amount of catalyst (5 - 20 g). Nickel or iron carbonyl were synthesized in a small (50 cm 3) stainless steel reactor containing either reduced nickel or iron
pellets
in
a carbon monoxide
prepared
mixture
of
admitted
to
gas-mixing
nitrogen.
the
The
atmosphere
at
0.5
± 0.1
MPa.
carbon monoxide and either nickel or iron chamber together with
a
flow
The
thus
carbonyl
of
was
hydrogen
concentration of the metal carbonyl in the gas flow fed to
in the
reactor was about 0.1 vo1.%. The hydrogen was used also to prevent oxidation of the
reduced
nitrogen
metal particles by a possible oxygen contamination in
stream
(see
results).
The temperature of the catalyst
the
large
during
the
decomposition of nickel and iron carbonyl was 473 - 523 K. To prevent the metal carbonyl
to
temperature
decompose below
markedly before reaching the catalyst due
the glass filter (which supports the fluidising
only the upper part of the reactor,
above the glass filter,
was
to
a
high
catalyst), heated.
The
405 r - - - - -
+- - - - - - I I L._
--
-
=::::::::::
s-
Figure 1: Schematic
representation
of
the experimental
set-up
for
the
preparation of bimetallic catalysts in the fluidized-bed reactor. 1. fluid-bed
reactor;
2. gas-mixing chamber; 3-S. devices
to
capture dust particles generated from the fluidising catalyst; 6. furnace to decompose the uncomp1ete1y reacted
nickel
(iron)
carbonyl; 7. nickel (iron) carbonyl synthesis reactor; 8. BASF R 3-11 catalyst for oxygen removal; 9. molecular sieve type SA; 10. Pti-tieoxo catalyst;
•••• ,.
Ni(CO)4-CO gas mixture.
LIl C")'
Mi
;
Figure 2: Representation of the fluid-bed reactor. Dimensions are in mm.
406 lower part was cooled by a flow of air. Using these precautions it was possible to
create
filter. the
a very sharp temperature gradient over the thickness of
This
reactor
bottom
the
glass
caused nickel or iron to be deposited only in the upper part above
the filter.
part of the reactor.
There was no visible decomposition
After regular time intervals samples
from the reactor for chemical analysis.
onto were
of the
taken
After fluidising for several hours the
metal carbonyl feed to the reactor was switched off. The catalysts were allowed to
cool
down to room temperature and subsequently statically
passivated
by
exposure to air. For by
comparison purposes a pure iron-on-silica catalyst has been
exposing pure silica to iron carbonyl.
In this
case,
prepared
however,
a
higher
temperature was needed to bring about iron carbonyl decomposition (i.e. 6Z0 K). This will be discussed below in more detail. The
metal loadings of the thus prepared catalysts have been determined by
means of AAS following standard procedures [19]. The
distribution
of
nickel
(or iron)
over
the
copper
catalyst
investigated by means of transmission infrared spectra of adsorbed CO.
was
A self-
supporting wafer was pressed and transferred into an in-situ infrared cell. The passivated
bimetallic
catalysts
were
reduced at 7Z3 K for
16 h
and
after
reduction evacuated at 673 K for one hour; the lower evacuation temperature was used to prevent oxidation by desorbing water molecules. allowed
to
cool
to
room temperature and
admitted.
Spectra
connected
to a 3500 CDS data-station.
were
13.3 kPa
Then the catalyst CO
(AGA,
99.997%)
was was
recorded using a Perkin-Elmer 580B spectrophotometer Gas absorption was compensated
for
by
means of an identical cell placed in the reference beam. Spectra were corrected for
background absorption of the silica support.
Details of the vacuum system
have been given elsewhere [ZO]. The syngas activity measurements were performed in an atmospheric bed reactor (i.d.
fixed-
10 mm). Details of the experimental set-up and the analyzing
system have been described by de Bokx et al. [Z]. The passivated Fe-Cu catalyst and the freshly prepared pure Fe/Si0 2 catalyst were reduced in a 10 vol.% HZ/N Z for 16 h at 773 K. Subsequently, the samples were flushed with pure nitrogen for
1 h at 723 K before cooling down to the reaction
temperature.
Next,
the
CO/HZ mixture was introduced in the reactor (p(CO) = p(HZ) = 0.1 bar and NZ as a balance, space velocity (S.V.) is 4000 h- 1). The amount of catalyst was adjusted
to keep the total CO conversion below 5% ensuring that rate data free
of mass and heat transfer effects were obtained.
The activity is expressed as
407 the
rate
of
Selectivity
CO conversion to hydrocarbons (rT) is
defined
as
per
unit
the fraction of CO converted
weight
catalyst.
the
particular
to
hydrocarbon. Carbon
deposition was investigated in the same reactor at 723 K and at
CO/H 2 ratio of
4
S.V. = 4000 h- l).
(p(CO) = 0.20 bar and
a
N2 as balance, such a run samples were subjected to a forced cool-
After
p(H2) =0.05
bar,
down in N2 to room temperature, typically within two minutes, and then analysed with regard to the deposited carbon. Temperature programmed hydrogenation (TPH) of these carbonaceous was
carried out in the same reactor. Its
methane. temperature
evolution
using
GC
was
discontinuously
analysis.
deposits
The only product observed during TPH was measured
as
Due to the time required
a for
function one
of
analysis
(7 min) linear heating rates were limited to 0.5 K/min.
RESULTS and DISCUSSION Fig.
3
catalyst passed
shows
a function of the amount of Ni(CO)4/CO gas (in moles)
as
through
deposition later on.
the reactor.
It can be seen that the initial rate
varying
of
had
nickel
the change in the rate of nickel deposition can be
either
of the gas mixture passing over the
catalyst or to a change in the rate of decomposition.
We
ascribe
deposition rate to the fact that before starting an experiment the
gas in the reactor,
containing the nickel pellets,
long period of time (3 - 20 days). Ni(CO)4
which
of
Since due to the way of preparation the amount of Ni(CO)4 in the gas
to a variation in the Ni(CO)4 content
copper
gram
is much higher than the more or less constant rate which is reached
mixture is unknown, due
an example of the amount of nickel deposited per
was produced,
the CO
contacted the nickel for a
During this period a considerable amount of
whereas the rate of formation of nickel carbonyl during
the deposition experiments is governed by the residence time in the synthesis reactor,
of about I minute
where the CO flows through the nickel pellets.
This
leads to about 0.1 vol.% Ni(CO)4 in the gas stream which is fed to the reactor. At
the experimental conditions employed complete conversion of Ni(CO)4 in
reactor
takes
place.
With a higher Ni(CO)4 content (above 0.5 vol.%
the
in
the
total gas stream) the decomposition of Ni(CO)4 on the fluidising particles
was
408
1S
1
10
Ni content
(mg/g cct.I
5
2
1
NilC0l4+ CO
(mole) -
Figure 3: Amount of nickel deposited onto the copper-on-silica catalyst via nickel
carbonyl
decomposition in a fluid-bed reactor
versus
the
amount of Ni(CO)4-CO gas mixture passed through the reactor. Bar indicates estimated error.
not
quantitative,
the
outlet
of
decomposition loading
as a nickel film developed onto the wall of the furnace
the of
reactor.
iron
Similar
carbonyl.
observations
were
made during
In order to increase the iron
in the thus prepared alloy catalysts,
(or
at the
nickel)
the decomposition reaction
was
carried out several times, every time with a new batch of freshly prepared iron or nickel carbonyl.
The thus prepared Fe-Cu and Ni-Cu catalyst investigated in
this paper contain about 2 at.% Fe and 4 at.% Ni. It was observed that the rate of iron carbonyl formation in the reactor 723 K)
strongly declined in time. restored
already
the original activity.
dissociates
at
synthesis
Re-reduction of the iron pellets (24 h at As it is known that
room temperature
on
a
pure
carbon
iron
monoxide
surface
[21],
deactivation by carbon deposition may account for the observed features. In fig. 4a the infrared spectra are shown that are recorded upon admission of
carbon
monoxide to a reduced eu-Fe catalyst.
exposure at room temperature is shown.
The effect of
Two bands are visible,
prolonged i.e.
CO
a band at
about 2130 cm-1 and a band around 2030 cm-1 which increases in intensity with increasing CO exposure time. The band at 2130 cm-1 is due to CO adsorbed on a copper surface [22] and the band at 2030 cm-1 arises from the CO molecules
409
linearly has
adsorbed on iron atoms [23].
shifted
adsorbed
on
alloyed
with
results
from the dilution effect [24].
increases
a
frequency
The
indicating that the
red shift of the
prolonged CO exposure,
vibration
iron
atoms
frequency
As the intensity of this
CO are
presumably latter
the iron atoms must indeed be
band
alloyed
CO-induced iron segregation can be understood by considering
difference
in
the
heat of CO adsorption
on
the
respective
the
metals
= -67
kJ/mol and ~H(CO-Fe) = -163 kJ/mol [25J). Room temperature for 0.5 h leads to a significant decrease of the Co-Cu (2126 em-I)
(~H(CO-Cu)
evacuation band,
pure iron surface [23], copper.
upon
with copper. large
The maximum of the vibration
to lower wavenumbers if compared with the frequency found for
whereas the intensity of the CO-Fe (2024 em-I) is not markedly
affected
(fig. 4b). Similar results were obtained with the Ni-Cu catalyst (not shown).
2032
1
0 05 .
1
after evoc uotion CO FeCu 2024
A
Fe!like'
15.1~
Cu 2133
1 A
1900 15min 2300
2000
----1l (cm-'l)
1900
Figure 4a: Infrared spectra recorded with the Fe-Cu catalyst illustrating the (left)
effect of prolonged CO exposure (13.3 kPa at 298 K).
Figure 4b: Infrared (right)
spectrum recorded with the Fe-Cu catalyst
temperature evacuation (0.5 h) of CO.
after
room
410 The
infrared spectra of adsorbed CO obtained with the pure iron-on-si1ica
catalyst silica
(fig. 5), support,
bimetallic for
are
Fe-Cu
prepared
prepared
by decomposition of iron carbonyl
on
bare
completely different from the spectra recorded with
catalyst.
Before admission of carbon monoxide
the
catalyst was re-reduced at 673 K for 16 h and subsequently
0.5 h
the
at 623 K.
CO adsorption on such a sample gives rise
the
freshly evacuated
to
a
strong
absorption band at 2170 cm- l• It can be seen that reduction of this catalyst at a more elevated temperature (773 K, the
infrared
adsorption in
more
detail
that
if
not lead to further changes in points
to
CO
In a previous paper [26] we have investigated
the reduction behaviour of
a pure Fe/Si0 2 catalyst. With recorded during the reduction of such a catalyst, it was
spectra,
support,
16 h) does
vibration frequency of 2170 cm- l
The
on Fe(II) ions [26].
MBssbauer shown
spectrum.
a highly dispersed iron precursor is applied onto
ferrous
ions
react
with the support
to
iron(Il)
the
silica
silicate.
Room
temperature CO adsorption on this compound gives rise to a vibration frequency at 2170 cm- l• As is also indicated by the present data the formation of this compound strongly reduced
metallic
retards iron
reduction
to metallic iron (CO
surface gives rise to bands
around
adsorption on a 2000 cm- l [23]).
However, the infrared spectra recorded with the bimetallic iron-copper catalyst show
that
iron
present in the copper matrix is
experimental conditions employed. into the copper particles, prevents
the
i.e.
completely
reduced
at
the
Thus, the exclusive deposition of iron atoms the formation of iron-copper alloy particles,
in-situ formation of iron(ll) silicate.
Therefore the
problems
encountered during the reduction of a pure Fe/Si0 2 catalyst can be effectively circumvented. Furthermore, a comparison of the infrared spectra recorded with the
pure
Fe/Si0 2 and with the Fe-Cu/Si0 2 catalyst shows that, within the detection limits, alloy formation indeed exclusively occurs when
experimental the
decomposition of the metal carbonyls is carried out in the presence
of
a
copper catalyst. With regard to the preferred decomposition on the copper particles we have established that the reaction temperature and the partial pressure of the metal carbonyl are essential parameters. pressures
It was observed that with carbonyl
partial
above about 0.5 kPa the decomposition reaction also proceeds on
the
bare silica support. The
effect
of the temperature on the decomposition of iron carbonyl
been investigated in more detail. does
not
has
Provided the iron carbonyl partial
pressure
exceed 0.1 kPa no measurable thermal decomposition of iron
carbonyl
411
C"l
A
r-,
N I
I I
Io01
I I
)
b
---.
....
N N
a
2300
2200
2100
2000 v-(cm-'l
Figure 5: Room temperature infrared spectra of adsorbed CO recorded with the reduced Fe/Si0 2 catalyst prepared by iron carbonyl decomposition on the bare support. Catalyst reduced (a) at 673 K for 16 h and (b) subsequently at 773 K for 16 h.
412
has
been
observed in an empty reactor up to temperatures
presence copper
of
670 K.
In
a fluidising pure silica carrier (identically pretreated
as
catalyst
carbonyl
of
the the
employed in the production of the bimetallic catalysts)
decomposition
occurred
at
temperatures
above
600 K.
Once
particles
were present (evidenced by the grey color of the then loaded
carrier),
the
decomposition reaction continued even at a
lower
iron iron silica
temperature.
Such an autocatalytic effect during the decomposition of iron carbonyl has been observed by other investigators [27]. In the presence of the reduced copper-onsilica catalyst, however, the decomposition markedly started above 350 K. As no deposition of iron was observed onto the bare support up to 573 K,
marked
deposition of iron during the production of bimetallic catalysts was carried
out
between
the
routinely
473 - 523 K in order to increase the efficiency
of
the
process. Furthermore, experiments on copper single crystal surfaces showed that at these temperatures complete decarbonylation of nickel carbonyl has occurred. With iron carbonyl, does
however,
complete removal of the carbon monoxide
not take place even up to temperatures of 623 K [10].
carbon
or
employed,
metal
carbides,
may induce
which can be easily formed
phase
segregation
in
As the presence of
at
iron-Cor
ligands
the
temperatures
nickel)-copper
alloy
particles [7], removal of carbon monoxide is essential. Therefore, the presence of
hydrogen,
rise to hydrogenation of
g~v~ng
carbidic
and/or
carbonaceous
deposits, is necessary to obtain truly alloyed iron (nickel) particles. Finally,
the catalytic performance of the Fe-Cu and the pure Fe catalysts
in the Fischer-Tropsch reaction has been investigated. The pure Fe/Si0 2 used in this experiment was prepared by homogeneous deposition-precipitation. In fig. 6 the
hydrocarbon production rates (rT) is plotted as a function of the time-on-
stream. of
As the activity of the bimetallic catalyst is markely lower than
the pure Fe catalyst,
that
the reaction temperature with the Fe-Cu catalyst
620 K and with the pure Fe catalyst 520 K.
Nevertheless,
is
the activity of the
pure iron catalyst is initially appreciably higher than that of the iron-copper alloy catalyst. The activity of the iron catalyst strongly decreases with timeon-stream,
whereas
the
Fe-Cu catalyst maintains its
activity
much better.
Selectivities measured after 5 h on stream are given in table 2. It can be seen that
the
decreased selectivity catalyst.
selectivity if
compared
of the Fe-Cu catalyst towards higher to
a
pure
Fe
catalyst.
hydrocarbons
Furthermore,
the
of the pure Fe catalyst is markedly higher than that of the
has
olefin Fe-Cu
413
'T
0.2 n mol
5.mglcat)
GFe/SiO z
0.15
xFeCu/SiO z
0.1
005~ 10
5
15 timelh)-
Figure 6: Activity of the Fe/SiO Z and the Fe-Cu/SiO Z catalyst as a function of the time-on-stream. It should be emphasized that the reaction temperature with Fe is 520 K and with Fe-Cu is 620 K.
r
n.mol _Fe/5i0 2 --Cu/SiO z
s.mglFeJ
3
.r'\
x 0.25
.' ;I
I
_._FeCu/SiOz
I
~
E c:
.~
"t>
2
..,
::J
I
0
I
C. 'I~
U
xO.005 ~L_
I I
\ \ \
\
"\ \
\
\
100
200
300
400
500
600
\
700 600 T("(;I-
Figure 7: TPH profiles of carbided catalysts previously exposed to severe carbiding
conditions (see text).
Multiplication with the
factors
414 Table 2 Selectivities of the Fe and Fe-Cu catalysts
Fe
Fe-Cu
0.54
0.95
0.33
0.05
0.90
o
0.07
1.0 0.06 Cn= is fraction of Cn product present as Cn=
* In
fig. 7 TPH profiles of catalysts which were previously
severe
conditions
with
shown.
The
velocity
deposition
space
regard to carbon deposition
(see
treatment were the same for all three catalysts
Fe/Si0 2 catalyst gives rise to a large peak centered
peak
areas
with
experimental)
and the time on stream (16 h) during
pure
measured
subjected
the pure copper
and
the
the
are
carbon
investigated. around
The
820 K.
bimetallic
to
The
iron-copper
catalysts are significantly smaller compared to the peak area observed with the pure iron catalyst. with
Thus,
the alloy catalyst.
carbided
carbon deposition appears to be markedly suppressed This is corroborated by electron micrographs of
samples which showed extensive filament growth with the pure
catalyst,
whereas
catalysts.
no whiskers
Similar
results
are
were
observed
obtained
with
the
Fe-Cu
with the Ni-Cu and
and the
the
Fe/Si0 2 Cu/Si0 2 pure
Ni
catalyst. Compared to the pure iron catalyst the alloy catalysts exhibit (i) a lower activity,
(ii)
significantly due
an
enhanced
selectivity
towards
better maintenance of the activity.
methane,
and
(iii)
a
The strong activity decline
to alloying of an active metal (iron) with an inactive one (copper)
is
a
well-known phenomenon and has been extensively discussed in the literature (see e.g. ref [5]). Our TPH results suggest that the increased stability is a result of a decreased carbon-deposition rate with the alloy catalysts. whisker Fe
The absence of
growth with the Ni-Cu and the Fe-Cu catalysts once more shows that the
or Ni deposited in the Cu-catalysts have completely reacted with Cu to
an
415
alloy.
Especially with regard to prevent the growth of these whiskers complete
take
up of the Fe or Ni into a copper alloy is a prerequisite,
growth
on
individual
structure number
unalloyed
iron
of the particular catalyst.
of
particles
completely
since
whisker
detoriates
It has been observed that a very
the small
pure Fe particles is sufficient to lead to a substantial growth
of
carbon and to disintegrate the catalyst structure. The and
enhanced production of methane suggests the application of the catalysts in the methanation reaction.
Ni-Cu
compared more
to
the pure Fe and Ni catalysts is low,
important.
Although
reaction can be advantageous.
For instance,
Fe-Cu
activity
if
the increased stability
is
Therefore application of these catalysts in
the the
methanation
the use of these catalysts allows
(i) the application of a higher reaction temperature and (ii) processing of COrich syngas mixtures.
The former is advantageous with regard to the production
of steam of a higher quality and the latter is preferred, since this eliminates the water gas shift reaction,
which is generally necessary to adjust the CO/H 2 ratio of the raw syngas mixture.
CONCLUSIONS The
production of Fe-Cu and Ni-Cu alloy catalysts by the decompostion
of
nickel or iron carbonyl in the presence of a reduced copper catalyst meets with the
most important requirements of alloy catalysts.
Provided the experimental
conditions as to the reaction temperature and the composition of the gas are
optimized,
phase
decomposition of metal carbonyl exclusively takes place on the
copper particles. That is, the deposited iron or nickel has completely taken up by
the copper.
fluidized-bed
Furthermore, reactor,
over the copper particles. particles
should
individual the
be
since the alloying reaction is carried out in
the iron (nickel) atoms are homogeneously Thus,
a
distributed
the bulk composition of the individual alloy
almost equal.
In principle
the
heterogeneity
in
the
bulk composition is determined by the particle size distribution of
parent
copper
catalyst.
Copper
catalysts
prepared
by
homogeneous
deposition-precipitation exhibit an almost mononodal particle size distribution (the
mean metal particle size being primarly determined by
the
total
metal
loading). Since reduction
copper
matrix,
of thus prepared alloy catalysts overcomes the difficulties
iron
(or
nickel) atoms exclusively reside in the
usually
416
encountered with pure iron catalysts. With the latter catalysts the interaction of iron ions with the support generally impedes reduction to zero-valent iron. Alloying
of
Fe or Ni with copper appreciably suppresses
the
growth
of
carbon in the catalysts during the reaction of CO and H2• That the activity is lower, is not a serious drawback. Working at a more elevated temperature and thus
producing
steam
of
a
better
quality
is
advantageous.
The higher
selectivity to methane is undesired when the production of higher hydrocarbons is aimed at. However, the selectivity can be much better using higher pressures of syngas.
REFERENCES 1.
M.E.
Dry, in "Catalysis, Science and Technology", (J.R. Anderson
and M. Boudart, Eds.), Vol.1, p.159. Springer Verlag (1981) Berlin. 2.
P.K. de Bokx, A.J.H.M. Kock, E. Boellaard, W. Klop, and J.W. Geus, J. Catal. 96 (1985) 454.
3.
A.J.H.M. Kock, P.K. de Bokx, E. Boellaard, W. Klop, and J.W. Geus, J. Catal. 96 (1985) 468.
4.
E. Boellaard, P.K. de Bokx, A.J.H.M. Kock, and J.W. Geus, J. Catal.
5.
V. Ponec and W.A.A. Van Barneveld, J. Catal. 89 (1984) 542.
6.
C.M.A.M. Mesters, Ph.D. Thesis University of Utrecht (1984).
7.
A.F.H.
8.
Cata!. J.R. Anderson,
96 (1985) 481.
Wielers,
G.W. in
Koebrugge, and J.W. Geus, submitted to J.
"Structure of metallic Catalysts".
Academic
Press, London (1975), p.176. 9.
C.M.A.M. Mesters, G.Wermer, O.L.J. Gijzeman, and J.W. Geus, Surf. Sci. 135 (1983) 396.
10. O.P. Van Pruissen, E. Boellaard, O.L.J. Gijzeman, and J.W. Geus, Appl. Surf. Sci., in press. 11. O.L.J.
Gijzeman, C.M.A.M. Mesters, F. Labohm, and J.W. Geus, J.
Mol. Catal. 25 (1984) 193. 12. C.M.A.M. Mesters, A.F.H. Wielers, O.L.J. Gijzeman, J.W. Geus, and G.A. Bootsma, Surf. Sci. 115 (1982) 237. 13. K. Kelley, Bureau of Mines Bulletin 584, Washington D.C. 14. L.E. Wichs and F.E. Block,
Bureau of Mines Bulletin 605,
417
Washington D.C. 15. C.F. Powel, I.E. Campbell, and B.W. Gonser, "Vapor Plating", Wiley, New York (1955). 16. H.E. Carlton and J.H. Oxley, A.I.Ch.E. 11 (1965) 79. 17. J.W.
Geus, in
"Preparation of Catalysts
III"
(G.
Ponce1et,
P. Grange and P.A. Jacobs, Eds.) p.1, Elsevier (1983). 18. A.J.
van Dillen, J.W. Geus, L.A.M. Hermans, J. van der Meijden,
in "Proc , 6th Int CongvCata l ;" VoL2, Society (1977). i
19. B.A.M.G. JUtte, A. Acta 110 (1979) 345.
p.677.
The Chemical
Heikamp, and J. Agterdenbosch, Anal. Chim.
20. A.F.H. Wielers, G.J.M. Aaftink, and J.W. Geus, Appl. Surf. Sci. 20 (1985) 564. 21. A.J.H.M. Kock and J.W. Geus, Prog. Surf. Sci 20 (1985) 165. 22. K.P. de Jong, (1980) 273.
J. Joziasse, and J.W. Geus, Appl. Surf. Sci. 6
23. N. Sheppard and T.T. Nguyen, in "Advances in Infrared and Raman Spectroscopy" VoL 5. Heyden and Son (1978). 24. F. Stoop, F.J.C.M. Toolenaar, and V. Ponec, J. Catal. 73 (1982) 50. 25. D. Tomanek,
S. Mukherjee, V.
Kumar, and K.H. Bennemann, Surf.
Sci. 114 (1982) 11. 26. A.F.H. Wielers, A.J.H.M. Kock, C.E.C.A. Hop, J.W. Geus, and A.M. van der Kraan, submitted to J.Catal. 27. T. Bein and P.A. Jacobs, J. Chern. Soc. Faraday Trans I 80 (1984) 1391.
418
DISCUSSION E.V.W. GRITZ: The short olefin selectivities in the product spectrum after FT-syntheses with your Fe-Cu/SiO? catalyst look not so favourable. Do you intend to promote your catalyst With K and which would be the effect on the copper-iron phase? J.W. GEUS : In this paper we have focussed on the preparation of Fe-Cu catalysts by means of catalytic decomposition of iron carbonyl. A full description of the catalytic properties of this system will be published shortly. The poor olefin selectivity of the Fe-Cu catalyst studied in this work can be improved by (i) raising the operating pressure, (ii) increasing the CO/H2 ratio in the syngas mixture, and (iii) increasing the Fe concentration. As yet we have not studied the effect of potassium promotion, however. L. GUCZI : By looking at your selectivity figures for iron- and iron-copper samples, it appears to me that in spite of the fact that iron occupies the outher layer, during the pretreatment or preparation Cu segregates to the surface and in fact, the bimetallic particles are covered by copper-like in the case of Ru-Cu or Os-Cu. Could you comment on it? J.W. GEUS : Our infrared data (see fig 4) show that the surfaces of the Fe-Cu alloy particles become enriched in iron in the presence of carbon monoxide already at room temperature. We have additional evidence (to be published) that under F.T. conditions, indeed such a segregation takes place: the surface consists of iron (-carbide) inlets in a copper matrix. With increasing iron content the copper phase becomes completely enveloped by the iron (-carbide) phase (Sachtler's cherry model). Thus, at low iron contents (as is the case with the catalyst studied in this work) a considerable amount of copper is present in the surface of the Fe-Cu alloys, thereby explaining the selectivity and activity changes compared to a monometallic iron catalyst. At high iron contents the Fe-Cu alloy behaves like a pure Fe catalyst, however. J. MARGITFALVI : (1) What is the driVing force for the exclusive formation of bimetallic surface species? (2) What was the ratio of H2/(metal carbonyl) in your preparation method? How would the change of this ratio alter the amount of metal introduced into the Cu/Si02 catalyst? Comment : I would suggest that the nydrogen chemisorbed at copper sites might be considered as driving force for Ni or Fe introduction. J.W. GEUS : Thermodynamics show that under the conditions employed iron- and nickel carbonyl can be completely converted to carbon monoxide and elemental iron and nickel, respectively. The homogeneous decomposition reaction is severely kinetically hampered, however. A reduced metal surface (such as e.g. copper) can act as a catalyst for this reaction leading to exclusive formation of alloys. As hydrogen is not involved in the decomposition reaction, the equilibrium is not affected by its presence. Hydrogen is used (i) to prevent oxidation of the reduced copper catalyst and (ii) to hydrogenate possible carbonaceous residues which might induce phase segregation in the Fe-Cu catalyst. M.A. VANNICE: (1) Have you employed chemisorption measurements (such as CO) to estimate Fe crystallite size in the Fe/Si0 2 and Fe-Cu/Si02 catalysts? (2) Where there any differences in the Fe carbide phases formed in the used Fe/Si0 2 and Fe-Cu/Si0 2 catalysts? J.W. GEUS : (1) The crystallite size of the pure Fe/SiOZ catalyst has been extensively studied by means of Transmission Electron Mlcroscopy. Reduction of the Fe/Si02 catalyst for a relatively short period (16h) at an elevated temperature (823-873K) gives iron particles of about 40-50 nm in diameter,
419
whereas prolonged reduction (150h) at a lower temperature (700-7Z3K) yields very small particles (dimater below 10 nm). With the Fe-Cu catalyst such studies are in progress. (Z) Mossbauer spectra show that in the fresh Fe-Cu catalyst iron and copper are intimately mixed forming a solid solution (homogeneous alloy). When this catalyst is exposed to a CO/H? stream (ratio = 1) at 553K alloyed iron readily reacts to iron carbide ( '-FeZ Z C and Fe3 C). These carbide phases are also observed with a pure Fe/SiO Z catalyst. These results once more show that in the presence of CO (or upon the formation of iron carbides) phase segregation occurs in the Fe-Cu alloy. The catalytic behaviour of such a system can then be entirely described by a cherry configuration (see answer to Prof. Guczi's question). R.A. SCHOONHEYDT inhomogeneous?
What is the distribution of Fe in Cu : homogeneous or
J.W. GEUS : The answer to this question is already discussed in the answer to the question of M.A. Vannice.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B,V., Amsterdam - Printed in The Netherlands
421
REACTIVITY AND STRUCTURE OF METAL CATALYST PARTICLES* C. LEEi', S. GAO i't and L. D. SCH~1IDTtt ttDept. of Chemical Engineering; Univ. of Minnesota; Mpls., MN 55455 tpresent address: Cray Research, Chipewa Falls, WI 5ur4~lARY
The correlation between catalytic activity and particle microstructure is examined for hydrocarbon hydroqenolysis and for CO hydrogenation on Rh and Ni supported on 5i02 using TEM, XPS, reactivity, and H2 chemisorption. It is found that C2H6 hydrogenolysis is more than 102 higfier on Rh after oxidation and low temperature reduction at 250°C than after high temperature reduction at 650°C. Catalyst activity can be restor~d repeatedly by oxidation and low temperature red~ction. TEM shows that 100A single crystal particles are converted into -lOA clusters of particles upon oxidation-reduction treatment. XP5 shows that oxide particles are comoletely reduced to metal by 200°C. Propane hydrogenolysis shows a comparable effect of oxidation-reduction on Rh. Ni on 5i02 exhibits a factor of two reversible decrease in activity for CO hydrogenation upon low temperature reduction. These results show that the properties of low dispersion catalysts can be altered reversibly by oxidationreduction treatments because of the variation in surface microstructure oroduced on the metastable surfaces created by low temperature reduction of oxide. INTRODUCTION The relationship between catalyst particle structure and reactivity dominates much of catalysis literature. This arises from the dramatic variability in performance of catalysts with nominally identical properties and from the need to create catalysts with desired properties. The definitive correlation of structure and reactivity has been elusive because the particle sizes and loadings are usually too low to permit quantitative determination of microstructure. Consequently, these correlations are frequently speculated to arise from electronic effects, defects, surface crystallography, edges and corners, support interactions, etc. We have developed techniques for transmission electron microscopy (TEM) and x-ray photoelectron spectroscopy (XPS) on metal particles between 10 and 200A diameter which permit examination of the time evolution of particle microstructures in various reactive gases. In particular, we examine the changes in individual particles following sequential treatment in different environments. We have used these techniques to monitor shape changes in Pt (ref. 1) surface enrichment in Pt-Rh alloys (ref. 2), oxidation and reduction of Rh (ref. 3), Pd (ref. 4), and Ir (ref. 5) and sulfidation of Pt and Rh (ref. 6). I
I
422
In this paper we summarize recent results on the effects of microstructure on ethane and propane hydrogenolysis C2H6 + H2 -+ 2CH 4 (1) and C3H8 + H2 -+ CH 4 , C2H6 on Ni or Rh on Si02 and CO hydrogenation
(2)
(3) CO + 3H 2 -+ CH 4 + H20 on Ni on Si02, Alkane hydrogenolysis has been studied extensively beginning with the finding of Carter, Cusumano and Sinfelt (ref. 8) in 1966 that the activity of C2H6 hydrogenolysis on Rh depended on particle size, with smaller particles being more active. Recent results by many investigators generally confirm these conclusions (ref. 7). CO hydrogenation has also received wide attention on Ni, both in supported form (refs. 9,10) and on single crystal planes (refs. 11,12).
EXPERIMENTAL As sketched in Fig. 1, samples for TEM analysis were prepared by depositing o 0 -lOA films of metal on 200A Si02 flakes and mounting these on Au microscope grids and heating in H2 to form particles (refs, 1-6). XPS samples were metal deposited on oxidized Si wafers. Reactivity and H2 chemisorption measurements were on 5% Rh on Cab-O-Sil and 15% Ni on Cab-O-Sil. All samples were prepared o to have approximately the same average particle sizes (150 to 250A) and all were treated identically in H2 and air environments. In some experiments the same sample was transferred between all systems as sketched in Fig. 1.
100i.
-I I-
metal
bC\C\-L sro, 200.4 T
-o
e
OVEN
f.
~
~ 51
&~ ¢::>
TEM
Fig. 1. Sketch of sample used for TEM and XPS and sketch of transfer capability between furnace, TEM, XPS, and reactor.
or
'\teo XPS
Au GRID
~
~
I
REACTOR
423
TEM and XPS procedures have been described in detail previously (refs. 2,3). Samples could be transferred between the TEM or XPS and the oven repeatedly to determine the evolution in shapes and the surface composition and chemical shifts. All treatments were for four hours at the temperatures specified in flowing H2 or air. Reactivity and H2 adsorption measurements were in a GC microreactor in flowing He (ref. 13). Metal surface areas were determined by H2 adsorption using the pulse desorption method of Butt et al. (ref. 14). We were able to obtain ±10% reproducibility on the same sample and -30% reproducibility for different samples. Reactivity was measured using 5% C2H6 or C3HS' 15% H2, and SO% He at 1 atm. RESULTS Ethane hydrogenolysis Catalysts were oxidized at 500°C in air, and then reduced in H2 at increasing temperatures. After each H2 treatment, the activity was measured at 230°C. The catalyst was then re-oxidized and the process was repeated. C2H6 H2 C2H6 H2 O2 H2 C2H6 H2 prep 500°C 300°C +H 2 400°C +H 2 500°C +H 2 600°C (4)
1
~
-_ _-----'
Figure 2 shows a plot of rate of CH 4 formation at 230°C versus total time of H2 treatment, each being 4 hours. The first point is after annealing in 10- 5
250 (5)
300
300
10- 6
C
'" '"
" 4>
.....en 10-7 en 4>
"0
,g 4>
C
600
a: r." 10- 8
u
(41 650
10- 9 0
8
16
24
32
Reduction Time (hr)
40
48
Fig. 2. Plot of rate of CH4 formation from C2H6 and H2 on 5% Rh on Si02 versus reduction time in H at temperatures indicated by num ters in figures. After oxidation (arrow with "02"), the activity increases by up to a factor of 1000. This is reversible in that the activity can be restored by oxidation and low temperature reduction.
424
H2 at 60Qoe while the second is after oxidation followed by reduction at 300 oe. The rate is observed to increase by a factor of 50. Heating in H2 at successively higher temperatures causes a monotonic decrease in rate, while oxidation followed by 300 0e reduction restores the rate to within a factor of 2 of its initial value. Reducing between 250 and 600 produces a change of almost a factor of 1000 as shown by the last points in Fig. 2. Activation energies were measured in each reaction measurement by measuring the rate versus temperature between 210 and 280 oe. Arrhenius plots yielded E = 37 kcal/mole on the active low-temperature reduced catalyst and E = 64 kcal/mole on the inactive high temperature reduced catalyst. 0e
0e
XP5 and TEM Figure 3 shows XP5 spectra of the Rh 3d doublet and the 0 Is peak following oxidation, reduction at low temperature, and reduction at high temperature. It is seen that the surface is totally oxidized by 400 and that surface reduction is complete by 200 oe. Therefore we conclude that all hydrogenolysis experiments were carried out on a totally reduced metal surface. 0e
Fig. 3. XP5 spectra of Rh 3d and a Is peaks following H2 treatment at 600 0e (upper), oxidation at 500 (middle), and H2 reduction at 300 oe. Rh?03 is totally reduced by 200 oe. Vertical dashed lines indicate peak positions of Rho, Rh+ 3, a from 5i02, and 0 from Rh 203 respectively. 0e
Figure 4 shows TEM micrographs of Rh on 5i0 2 following treatments indicated by the arrows. Micrograph (a) shows the reduced Rh particles after formation o in H2 at 600 0e with an average particle diameter of -120A. After oxidation (b) particles spread on the surface due to increased wetting of Rh 203 on 5i0 2. Low temperature reduction at 300 0e (c) produces clusters of 10-20A metal particles where single crystal particles existed in (a). Heating at 400 0e produces some coalescence (d), but only after heating to 600 0e do clusters disappear as particles return to roughly their original shapes.
H
2 ---.
..
•. ~
600·C •
41.
.·4~
17~17
Metal
.·4~
500·C • • •
'...
(a)
I
~4
- •• soo; - ••
300·C
~
7iii7iiiJ> Oxidized
",.
~
• • • H2
(b)
-
... .
• .~
•
425
..
(c)
~
/~7
Reduced
~i
(d)
--
~
1I~7
Coalesced
~
• •' (e)
-11.7 .
Sintered
Fig. 4. TEM micrographs and idealized sketches of particle cross sections following treatment in gases at tewperatures indicated. Micrograph c shows that particles are broken into -20A clusters of smaller particles by this treatment. Sketched below each micrograph is an idealized cross section of metal, oxide, clusters, and sintered configurations shown in the micrographs above. This breakup into clusters upon oxidation and reduction appears to be universal for metal particles on Si02, as we have observed it for Rh, Ni, and Ir and also for Pt, Rh, and Pd upon formation and decomposition of the sulfides (refs. 1-6). It is driven by the different interactions between metal and Si02 compared to metal oxide or sulfide and Si02 which produces a metastable cluster of smaller metal particles from a single larger metal particle. In results published elsewhere we have shown that the metal surface area changes by no more than a factor of two between low and high temperature reduction (ref. 14). These changes in reactivity must therefore be caused by alterations in surface structure rather than changes in surface area alone. Propane hydrogenolysis Propane hydrogenolysis on Rh exhibits a similar trend (ref. 15) upon oxidation and reduction to that observed for C2H6. Figure 5 shows a plot of rate versus oxidation or reduction treatment. The curve for C2H6 is from Fig. 2 and shows the extremes of the high rate immediately after oxidation and low temperature H2 reduction (marked "oxidation" on the graph) and after high temperature H2 reduction (marked "reduction" on the graph). For C3H8 the rate is higher than the C2H6 rate by a factor of ~10, and the variation between 250°C reduction and 650°C annealing is somewhat smaller, -20. In this reaction both CH 4 and C2H6 are produced in approximately equal amounts. Activation energies also vary with activity; Arrhenius plots yield E = 30 kcal/mole after reduction at 650°C. kcal/mole after 250°C reduction and 40
426
Fig. 5. Rates for all reactions followinq oxidation and low temperature reduction and following high temperature reduction. Rates increase and decrease dramatically as shown in text.
rR
moles ) ( g cat sec 7
16
16 9
REDUCED
OXIDIZED
REDUCED
OXIDIZED
REDUCED
OXIDIZED
OXIDIZED
C3HS hydrogenolysis exhibits some self-poisioning in that the rate decreases by about a factor of 3 over a period of several hour exposure to the C3H a + H2 mixture. These phenomena will be described in greater detail in a later publication (ref. 15). Ethane hydrogenolysis was also measured on Ni on 5i0 2. As shown in Fig. 5 the activity was much lower than on Rh and the variation was much smaller, less than a factor of 2. This is almost the variation predicted by the change measured in the metal surface area. CO hydrogenation on Ni The variation in activity of a 1/3 CO/H 2 mixture in CO hydrogenation to CH 4 on Ni was also measured in the same apparatus (ref. 14). TEM, XP5, and H2 chemisorption were also examined just as for Rh. As shown in Fig. 6, the rate'decreases upon oxidation and low temperature reduction, but the variation is much smaller than that for alkane hydrogenolysis on Rh. Metal surface area variations were also measured and found to increase by oxidation and low temperature reduction. The turnover frequency is therefore found to decrease by a factor of 2-3 upon oxidation and low temperature reduction for this reaction. Variations in effective activation energy are found to agree with variations in activity.
427
3.0 -
Fig. 6. Plot of methane formation rate from CO and H2 on Ni on Si O2 versus treatment temperatures listed on the figure. Also shown is the metal surface area measured by H2 chemisorption.
w 2.0 i-
f-
«
.
Ct: ¢
I U
1.0 -
10
20
30
40
TIME (HOURS)
DISCUSSION These results demonstrate that large changes in catalytic activity can be created by oxidation and low temperature reduction of supported catalysts. As summarized in Fig. 5, these changes can be as large as a factor of 1000 (for C2H6 hydrogenolysis on Rh) or as small as a factor of 2 (for C2H6 hydrogenolysis on Ni). Rates are higher after low temperature reduction for hydrocarbon hydrogenolysis but are lower for CO hydrogenation. These results are both reproducible and reversible. Activity variations ca~ be restored reproducibly by repeated oxidation and reduction treatment on a given catalyst, and we were able to reproduce the results shown on a different catalyst sample. The activity variations appear to be repeatable indefinitely, although we have never subjected a particular catalyst to more than six cycles. There is of course the expected irreversible sintering which occurs after extended heating at 600°C for both Rh and Ni. This amounts to approximately a factor of two after many hours of heating in H2. The correlation of activity with surface area, particle morphology, and particle oxidation state shows that morphology changes dominate the variation in activity. The area changes are very small and in fact go in the opposite direction to activity for CO hydrogenation. Surface reduction of NiO and Rh 203 occur by -200°C in H2 at one atmosphere, and this indicates that only metal catalysts are present in all situations. The dispersions are low (a few percent) in all of these systems so that we expect few isolated metal atoms which might have properties quite different from the bulk metals. Support interactions also appear incapable of influencing reactivities because most metal atoms are not in contact with the Si02, which is itself a very inert support.
428
We conclude therefore that the clusters of 10-20A° crystallites formed by low temperature reduction have much different catalytic activity than the lOO-200A polyhedra formed by high temperature reduction. We interpret this as resulting from a different activity on the low index (111), (100) and (110) crystal planes which one expects on an annealed surface compared to the high index (hkl) planes formed on the smaller crystals formed ~y low temperature reduction. One could of course associate this with higher fractions of edge, corner, and grain boundary atoms on the smaller crystals, but, since catalytic activity of metals probably depends on local ensembles of surface atoms, these arguments are probably equivalent. The equilibrium shape of a crystal is given by the expression I riAi = minimum (5) where the summation is over all metal surfaces, oxide surfaces, metal-oxide interfaces, and grain boundaries. In catalysis one is always dealing with metastable systems, and the micrographs of Fig. 4 suggest that the entire 200A° particle is equilibrated upon annealing at 650°C, but that equilibrium of Eq. (5) is only over a distance of -IDA upon oxidation and reduction at 250°C. One can regard the annealing process as producing equilibration over increasing distances in the metal particle, with long time annealing at high temperatures causing interparticle equilibrium, i.e. sintering. Superimposed on crystallographic variations caused by different crystal sizes are chemical effects caused by different adsorbed species (probably hydrocarbon fragments) on different planes. In fact we observe self-poisoning of Rh in C3H8 hydrogenolysis in that the activity after any H2 treatment decreases over a period of minutes. With C2H6, self-poisoning is probably too slow to be observed at these temperatures, while with C4H10 it is probably too rapid to be observed in a GC microreactor. We also note that the activation energy of the reaction rises markedly with H2 treatment. This is further evidence that the catalysts are chemically different rather than simply having different densities of a single type of active site. The selectivity also varies with treatment as we shall discuss in a later publication. These results are a clear demonstration of activity variations caused by microstructure variations. These appear to be quite predictable from the morphologies expected and observed from oxidation and low temperature reduction, and these morphologies can be stable over long periods of time at low temperatures. Such treatments have obvious relevance to the fundamental understanding and rational design of active and selective supported catalysts. *This research partially supported by NSF under Grant No. CPE8214048.
429
REFERENCES 1 T. Wang, A. Vazquez and L. D. Schmidt, Surface Science, 163 (1985) 181. 2 T. Wang and L. D. Schmidt, J. Catalysis, 71 (1981) 411. 3 T. Wang and L. D. Schmidt, J. Catalysis, 70 (1981) 187. 4 M. Chen and L. D. Schmidt, J. Catalysis, 56 (1979) 198. 5 T. Wang and L. D. Schmidt, J. Catalysis, 66 (1980) 301. 6 T. Wang, A. Vazquez, A. Kato and L. D. Schmidt, J. Catalysis, 78 (1982) 306. 7 E. H. Broekhoven and V. Ponec, Progress in Surface Science, 19 (1985) 351. 8 J. L. Carter, J. A. Cusumano and J. H. Sinfelt, J. Phys. Chern., 70 (1966) 2257. 9 V. Ponec, Catal. Rev. - Sci. Eng., 18 (1978) 151. 10 P. Biloen and W. M. H. Sachtler, Advances in Catalysis, 30 (1981) 165. 11 R. D. Kelley and D. W. Goodman in: "The Chemical Physics of Solid Surfaces and Heterogeneous Catalysis" (D. A. King and D. P. Woodruff, eds.) Vol. 4, p. 427 (1982). 12 H. P. Bonzel and H. J. Krebs, Surface Science, 117 (1982) 639. 13 C. Lee and L. D. Schmidt, J. Catalysis, to be rublished. 14 J. M. Amelse, L. H. Schwartz and J. B. Butt, J. Catalysis, 72 (1981) 95. 15 S. Gao and L. D. Schmidt, to be published.
430
DISCUSSION G. HORN: We reached similar results on supported nickel-methanation catalysts. During the high-temperature methanation, the Ni-catalyst was deactivated by sintering (not by coking). As we oxidized the deactivated catalyst, several hours at elevated temperature by air and reduced it after this treatment, the catalyst showed the same activity as the fresh catalyst. From hydrogen-chemisorption measurements we could conclude that the dispersion of nickel which was lowered by sintering could be restored by the oxidation/reduction procedure. L.D. SCHMIDT: This appears to be in agreement with our results, although in the experiments you describe coking and other complications associated with real processes are superimposed on the catalyst morphology changes. D.E. RESASCO : Both, activity suppression for hydrogenolysis reactions and enhancement for CO/HZ reactions have been many times ascribed to metal-support effects. In your case, you have chosen an inert support, such as Si02, which minimizes these effects. However, in recent works, some Si02 supports appear to be less inert as it could be expected. For instance, RobBins et al. (ACS Annual Meeting, N.Y. 1986) have found that the addition of SiOZ to unsupported Pt enhances the CO/HZ reaction by one order of magnitude. In our own laboratory, we have observed 1-Z orders of magnitude differences between Pt catalysts supported on different silicas, i.e. Davison and Cab-O-Sil. Would you expect that the SiO? support (or its impurities) may play any significant role in your experiments ? L.D. SCHMIDT: We chose SiO Z intentionally as an "inert" support, and the influence of support interactions probably decreases as particle size increases. On never knows whether trace contaminants may act as promoters or poisons for a particu1ar cata1yst system. P.G. MENON: You have given a very nice example of high-temperature reduction (HTR) producing a dramatic decrease in the hydrogeno1ysis activity of a supportedmetal catalyst, without invo1ving the so-called SMSI effect. Eight years ago, Guczi and coworkers had shown simi1ar effects for Ru/SiOz catalysts (Znd Symp. Prep. Heterogeneous Catalysts, Louvain-la-Neuve, 1978). For Pt/A1Z03 and Pt/SiOZ cata1ysts, similar effects of HTR on H-chemisorption and hydrogeno1ysis have been reported by us (P;G. Menon and G.F. Froment, J. Catal. (1979; Applied Catal. 1,31 (1981)). A1l these resu1ts serve to show that the effects of HTR are far more varied than what is impl ied by SMSI. L.D. SCHMIDT: The observation of activity variations with treatment is common in the literature, and I suspect that most laboratories have many unpublished examples of these effects. By our mechanism Pt should not exhibit this phenomenon because it is impossible to form Pt oxides by air treatment, at least for large partic1es. However, Pt particle shapes can change by treatment in different gases. We p1an to do experiments similar to those you reported and see if we can detect differences in structure by TEM. J.A. SCHWARZ: Subsequent to high temperature reduction, the catalyst particles appear in the TEM to divide into clusters of smaller particles. Can you rule out the possibil ity that some "support covering" could be occurring that makes it appear as if the particles get subdivided into sma11er c1usters ? L.D. SCHMIDT: Rh is notorious for forming "rafts", and I am not sure if we wou1d see a monolayer of Rh by TEM. However, we do not observe 1arge HZ chemisorption areas in situations where rafts might be present. A. FRENNET : In the study of ethane-hydrogeno1ysis, you mentioned very 1arge effects on the activity. Do you have some information concerning modifications of the kinetic parameters (orders, activation energies). Such modifications
431
should be expected as these large variations in activity are associated with a not very large variation in surface area derived from HZ chemisorption. L.D. SCHMIDT: You raise an important point. We measured only activity at a fixed pressure and composition, and some of the variations could be due to changing orders of reaction. L. GUCZI by TEM. However; which is
: Desintegration of large matal particle is convincingly demonstrated After mild hydrogenation, there is considerable increase in dispersion. an increase in dispersion should be indicated by XPS line broadening not seen at all. Could you comment on it ?
L.D. SCHMIDT: We saw no significant line broadening. This could be because most Rh is still in large particles. Experiments with much higher dispersion would have a better chance of revealing those effects. B. WALZ : Are the small metal (Ni, Rh) particles you observe 'finely ordered, i.e. show ordered crystal faces or are they strongly disordered exhibiting many defects and thus leading to higher activity? Did you check this by electron diffraction? L.D. SCHMIDT: We attribute the results to different crystal planes rather than deffects because vacancy defects in metals should attain equilibrium rapidly. Electron diffraction in TEM shows line broadening generally consistent with Debye SCherer line widths predicted from the average measured crystal size.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation oi Cetetvsts /V © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
433
INFLUENCE OF LANTHANUM OXIDE ON THE SURFACE STRUCTURE AND CO HYDROGENATION ACTIVITY OF SUPPORTED COBALT CATALYSTS
J. S. LEDFORD1, M. HOUALLA1, L. PETRAKIS2, and D. M. HERCULES1 10epartment of Chemistry, University of Pittsburgh, Pittsburgh, PA 15260 (USA) 2Chevron Research Company, P.O. Box 1627, Richmond, CA 94802-0627 (USA) SUMMARY Two series of La/alumina and CoLa/alumina catalysts have been characterized by several bulk and surface sensitive techniques. The chemical state and dispersion of the La and Co phases have been determined. The results were compared with the CO hydrogenation activity and selectivity of supported Col alumina catalysts. INTROOUCTl ON Rare earth and actinide metal oxides are often used as promoters in Ni, Co, and Fe based CO hydrogenation catalysts. Previous studies (refs. 1-9) have focused primarily on their effect on catalyst actiVity and selectivity. Little effort has been devoted to investigating systematically the effect of these promoters on the structure and reactivity of the active phase. This work is part of a broad study to investigate the effect of rare earth and actinide promoters on the structure and reactivity of Group VIII metalbased CO ~drogenation catalysts. X-ray photoelectron spectroscopy (XPS,ESCA), x-ray diffraction (XRO), Raman spectroscopy, H2 chemisorption, and gravimetric analysis are used to examine the influence of lanthanum on the state and dispersion of cobalt supported on alumina. The information derived from these techniques is compared with the CO hydrogenation actiVity and selectiVity of supported cobalt catalysts. EXPERIMENTAL The La modified alumina carriers (La/A1203) were prepared by pore volume impregnation of y-alumina (Harshaw 1401, nominal surface area: 200 m2/g) with solutions of lanthanum nitrate (Fisher). The samples were dried at 110°C for 12 hours and calcined at 400°C for 8 hours. Lanthanum content in the La/A1203 series was varied from a La/Al atomic ratio of to 0.08 (0 to 20 wt% La203). CoLa/Al203 catalysts were prepared by pore volume impregnation of the Lamodified alumina carriers using cobalt nitrate solutions (Fisher). Drying and calcination conditions were the same as those used for the La/A1203 series. The cobalt content of the samples was held constant at 10 wt% of the alumina support. The La/Co atomic ratios varied from 0 to 0.9. Increasing the
°
434
lanthanum content resulted in a maximum decrease of 25 percent in the BET surface area of the catalysts. Catalyst samples will be designated by "Col.ay" where y is the La/Co atomic ratio. Reduction was carried out at 400°C by flowing H2 (50 cc min- 1, 99.999%) over the catalysts. X-ray diffraction patterns were obtained using a Diano XRD-6 diffractometer employing Ni-filtered Cu Ka radiation (1.540 56 A). Raman spectra were recorded using a Spex Ramalog spectrometer equipped with holographic gratings. ESCA spectra of oxidic catalysts were obtained using a Leybold-Heraeus LHS-l0 electron spectrometer equipped with an aluminum anode (1486.6 eV) operated at 12 kV and 20 mAo ESCA measurements of reduced samples were performed with an AEI-ES200 spectrometer using an Al anode operated at 12 kV and 22 rnA. ESCA analysis of reduced catalysts was performed using a sealable probe which allowed the transfer of samples from a tube reactor to the spectrometer without exposure to air (ref. 10). The Al 2p line from the support was used as the binding energy reference (74.5 eV) for the catalysts. The binding energies for standard compounds were referenced to the C 1s line (284.6 eV). It has been shown by Defosse et al. (ref. 11) that one can calculate the theoretical ESCA intensity ratio (1m/Is) expected for a supported phase (m) atomically dispersed on a carrier (s). An extension of the Defosse model proposed by Kerkhof and Moulijn (ref. 12) has been used in the present investigation. The photoelectron cross-sections and the mean escape depth of the photoelectrons used in these calculations are taken from Scofield (ref. 13) and Penn (ref. 14), respectively. Gravimetric analysis was carried out using a Cahn 113 microbalance. Hydrogen chemisorption was performed using a volumetric chemisorption apparatus. Measurement of CO hydrogenation activity was performed in a flow microreactor with a 25 cc min- 1 flow of H2/CO/He (9%/3%/88%) at 185°C. RESULTS Lanthanum Modified Aluminas XRD patterns of the lanthanum modified aluminas showed only the lines which were characteristic of the y-A1203 carrier. The ESCA binding energies of the La 3d5/2 peaks measured for the La/A1203 series (836.2 eV) were significantly higher than the value measured for La203 (834.6 eV). The variation of the La 3d5/2/Al 2p intensity ratio as a function of La/Al atomic ratio is shown in Figure 1. The theoretical line calculated for monolayer dispersion of lanthanum is shown for comparison. The intensity ratios measured do not deviate significantly from the calculated monolayer line. It must be noted, however, that for the high La loadings (La/Al > 0.06) the calcu-
435
lated value for monolayer dispersion must be considered as a lower estimate because of the observed decrease in the surface area and indication of preferential deposition of the La species at the outer parts of the catalyst particle. 15.0"T"""--------------------.
»:
;"
~./
,....
7.5
./
./
;"
;" ;"
Monolayer
,,;"
./
".,""
5 La/AI Atomic Ratio (X10 2)
10
Fig. 1. Variation of ESCA ILa3ds/2/IA12p Intensity ratio as a function of La/Ai atomi c rati o, Cobalt-Lanthanum/Alumina Catalysts The XRD pattern of the unmodified Co/A1203 catalyst showed the diffraction lines characteristic of C0304 (see Fig. 2). La addition to the A1203 carrier up to La/Co = 0.3 had little effect on the intensity of the C0304 lines. However, for higher La contents the C0304 peaks were reduced sharply and practically absent in the XRD pattern of the La-rich catalyst (CoLaO.9). 693
CoLaO.7
~
~
39
Degrees (29)
34
Fig. 2. X-ray diffraction spectra of selected cobalt-lanthanum catalysts.
450
Raman Shift (cm-1)
750
Fig. 3. Raman spectra of selected cobalt-lanthanum catalysts.
436
Similarly, Raman spectra of CoLa/A1203 catalysts (see Fig. 3) showed that the intensity of the C0304 lines (peaks at 485, 524, and 693 em- 1) remained essentially unchanged for La/Co < 0.3 but decreased drastically for higher La contents (see Fig. 4). No Raman peaks other than those associated with C0304 were observed. Figure 5 shows the variation of the C02P3/2/Al 2p intensity ratio as a function of La/Co atomic ratio measured for the oxidic catalysts. Up to La/Co atomic ratios of 0.3, the measured intensity ratios are approximately 20 percent of the value calculated for monolayer dispersion (3.4). A drastic increase in the measured intensity ratio is observed for catalysts with higher lanthanum contents. 10
• ....
>.
'iii
c II)
.E
•
5.0
Co
' -I
•
•
A1
5
2.5
e
III
•
E
III
II:
O+------+-----~
o
....
• 0.5
1.0
La/Co Atomic Ratio
Fig. 4. Variation of the intensities of C0304 Raman peaks vs. La/Co atomic ratio.
O+------t-----~
o
0.5
1.0
La/Co Atomic Ratio
Fig. 5. Variation of ESCA IC02P3/2/IA12p intensity ratio vs. La/Co atomic ratio.
The reducibility of cobalt as a function of the La/Co atomic ratio was determined from both ESCA and gravimetric measurements. For catalysts with La/Co atomic ratios < 0.3 the ESCA and gravimetric data indicate that the reducibility of the cobalt phase (ca. 85%) is not affected by the addition of lanthanum. A decrease in the extent of reduction from 85% to ca. 50% is observed with further increase in lanthanum loadings. Table 1 shows the variation of the cobalt metal dispersion as a function of lanthanum content calculated from H2 chemisorption data. For La/Co atomic ratios < 0.15 the addition of La has little effect on the dispersion of the cobalt metal. An increase in dispersion is observed for higher La loadings. For catalysts with La/Co atomic ratios < 0.3 the cobalt metal dispersion calculated from ESCA data did not show any significant variation with La content.
437
TABLE 1 Cobalt Metal Dispersion Calculated from H2 Chemisorption La/Co Atomic Ratio Dispersion
o 4.0
0.06 3.5
0.04 4.0
0.15
3.5
0.3 6.3
0.7 7.6
0.9 8.1
For higher La loadings the ESCA data show a greater increase in Co dispersion. as compared to H2 chemisorption results. However. a part of the increase in dispersion indicated by ESCA must be ascribed to a small degree of inhomogeneity observed for La-rich catalysts. Catalytic Activity and Selectivity Figure 6 shows the variation of turnover frequency (TOF) for CO hydrogenation. based on H2 chemisorption data. as a function of lanthanum content. It can be seen that for La/Co ratios ( 0.3 the TOF is not significantly affected by lanthanum addition. For higher La loadings a drastic decrease in the TOF is observed.
•
•
O+---------+------------t
o
0.5 La/Co Atomic Ratio
1.0
Fig. 6. Variation of the turnover frequency (TOF) for CO hydrogenation with La/Co atomic ratio. Table 2 shows the the olefin/paraffin ratio and the weight fraction of methane for various La/Co atomic ratios. Over a twofold increase in the selectivity to olefinic products is observed with the addition of lanthanum. The weight percent methane decreases from 35% to 17% as the lanthanum content of the catalyst increases.
438
TABLE 2 Effect of Lanthanum on the CO Hydrogenation Selectivity of Co/A1203 Catalysts La/Co Atomic Ratio Methane wt% Olefin/Paraffin (mole fraction)
0 35 .28
.04 33 .32
.06 33 .31
.09 33 .34
.15 31 .37
.3 27 .42
.7 19 .63
.9 17 .72
DISCUSSION State and Dispersion of Lanthanum A comparison of the La/Al ESCA intensity ratios calculated for monolayer dispersion with the values measured for La/A1203 catalysts having La/Al atomic ratios < 0.06 suggests that the lanthanum is dispersed as a monolayer-like phase over the alumina carrier. The presence of a surface lanthanum phase is consistent with the XRD data and may be attributed to the formation of a La-A1203 interaction species. This is in agreement with Doesburg et al. (ref. 15) who reported that lanthanum aluminate (LaA103) is formed on La/A1203 samples having similar La loadings. State and Dispersion of Cobalt XRD and Raman data for CoLa/A1203 catalysts with low lanthanum contents (La/Co < 0.3) indicate that addition of La does not affect the amount of C0304 on the catalysts. The decrease in intensity of the C0304 diffraction peaks observed for higher La loadings may be attributed either to an increase in the fraction of highly dispersed cobalt oxide or to a decrease in the amount of C0304 on the catalysts. The latter hypothesis is confirmed by the parallel decrease in the Raman intensity of the C0304 lines for high La loadings (see fig. 4). The observed decrease in the amount of C0304 for catalysts with high La loadings can be ascribed to the formation of a surface Co-La interaction species. This is consistent with the observed decrease in the reducibility of the Co phase at high La loadings. The variation in Co/Al ESCA intensity ratios with La content observed for oxidic catalysts indicates that for low loadings of La (La/Co < 0.3), the cobalt phase is not well dispersed over the alumina carrier. These results are consistent with XRD patterns that showed crystalline C0304 on the catalysts. For higher La contents, the increase in the Co/Al intensity ratio suggests that the addition of La increases the dispersion of the oxidic cobalt phase. This may be attributed to a stronger interaction of the cobalt species with the La-rich alumina during the cobalt impregnation step. It is known (ref. 16) that the adsorption of C02+ on y-A1203 is enhanced in mildly basic solutions (6
439
This enhanced adsorption has been attributed to hydrolysis of the metal ion to give a species which is more strongly adsorbed on the alumina carrier. Thus. the increased dispersion of the cobalt phase observed on catalysts with high La contents may be ascribed to the hydrolysis of the Co2+ ion during the impregnation of basic. La-rich alumina supports. A similar mechanism has been proposed for the influence of Na on the dispersion of Co supported on y-A1203 (ref. 17). For reduced catalysts with high La loadings the apparent discrepancy between the cobalt dispersion calculated from H2 chemisorption and ESCA data suggests that the chemisorption of H2 may be suppressed. This is consistent with the work of SChaper et al. (ref. 8) which showed that H2 chemisorption on lanthanum modified Ni/A1203 catalysts was much slower than on the unpromoted nickel catalysts. In addition. Bartholomew et al. (ref. 18) have reported that H2 chemisorption is an activated process on dispersed Co/A1203 catalysts. This would result in a low estimate of the cobalt metal dispersion when H2 chemisorption experiments are conducted at room temperature. Catalytic Activity and Selectivity The observed decrease in the TOF for high La loadings may be tentatively ascribed to a blocking of the active sites by excess La moieties. Lower TOF values would also be expected from the observed increase in the dispersion of the Co phase for high La contents. This is due to the reported structure sensitivity of CO hydrogenation reaction over Co/A1203 catalysts (ref. 19). However. the observed changes in the product distribution were not consistent with the reported behavior (ref. 19). The decrease in methane fraction and increase in olefin/paraffin ratio observed on La addition are similar to the reported changes in selectivity when Ru is deposited on La203 instead of Si02 (ref. 20). Such variation in the product distribution usually reflects a lower H2 concentration on the catalyst surface. This can be ascribed to the proposed decrease in the ability of Co to chemisorb H2 with increasing La content. The presence of dispersed Co phase (ref. 18). the decrease in the extent of reduction of Co phase for high La contents (ref. 21). and the interaction of the Co phase with La species (ref. 8) will contribute. in principle. to the postulated suppression of H2 chemisorption. CONCLUSIONS ESCA data indicate that for low loadings of La (La/Al < 0.06) lanthanum is present as a monolayer-like phase over the alumina carrier. XRD, Raman, and ESCA data suggests that for low loadings of La (La/Co < 0.3) the amount and dispersion of C0304 is not affected by La addition. For catalysts with higher La loadings, C0304 cobalt phase is suppressed in favor of the formation of dispersed CoLa/A1203 species.
440
ESCA and chemisorption data indicate an enhanced dispersion of the metallic cobalt phase for high La loadings. Lanthanum addition up to a La/Co atomic ratio of 0.3 does not affect the TOF of the catalyst for CO hydrogenation. Further addition of lanthanum results in a substantial decrease in the TOF of the cobalt catalyst. The addition of lanthanum results in an increase in the olefin/paraffin ratio of the reaction products and an increase in the selectivity for higher hydrocarbons. ACKNOWLEDGEMENT We acknowledge the National Science Foundation (Grant No. CHE-8401202) for their financial support. REFERENCES (1) (2) (3)
(4) (5) (6)
(7) (8)
(9)
(10) (11)
(12) (13)
(14) (15)
(16) (17)
(18) (19) (20) (21)
G. N. Sauvion, J. F. Tempere, M. F. Guilleux, G. D. Mariadassou, and D. Delafosse, J. Chern. Soc., Faraday Trans. I, 81, (1985) 1357. M. R. Gelsthorpe, K. B. Mok, J. R. H. Ross, and R. M. Sambrook, J. Mol. Catal., 25, (1984),253. B. Denise and R. P. A. Sneeden, React. Kinet. Catal. Lett., 26, (1984) 265. V. U. S. Rao, R. J. Gormley, A. Shamsi, T. R. Petrick, J. M. Stencel, R. R. Schehl, and R. D. H. Chi, J. Mol. Catal., 29, (1985) 271. T. Inui, K. Ueno, M. Funabiki, M. Suehiro, T. Sezume, and Y. Takegami, J. Chern. Soc., Faraday Trans. I, 75, (1979) 787. T. Inui, K. Ueno, M. Funabiki, M. Suehiro, T. Sezume, and Y. Takegami, J. Chern. Soc., Faraday Trans. I, 75, (1979) 1495. R. E. Hayes, W. J. Thomas, and K. E. Hayes, J. Catal., 92, (1985) 312. H. Schaper, E. B. M. Doesbury, P. H. M. De Korte, and L. L. Van Reijen, Appl. Catal., 14, (1985) 371. Y. C. Xie, M. X. Qian, and Y. Q. Tang, China-Japan-U.S. Symposium of Heterogeneous Catalysis Related to Energy Problems, B.10C (1982) Dalian, China. T. A. Patterson, J. C. Carver, D. E. Leyden, and D. M. Hercules, J. Phys. Chern., 80 (1976) 1700. C. Defosse, D. Canesson, P. G. Rouxhet, and B. Delmon, J. Catal., 51, (1978) 269. F. P. J. M. Kerkhof and J. A. Moulijn, J. Phys. Chern., 83, (1979) 1612. J. H. Scofield, J. Electron Spectrosc. Relat. Phenom., 8, (1976) 129. D. R. Penn, J. Electron Spectrosc. Relat. Phenom., 9, (1976) 29. E. B. M. Doesburg, G. Hakvoort, H. Schaper, and L. L. van Reijen, Appl. Catal., 7, (1983) 85. P. H. Tewari and W. Lee, J. Colloid Interface Sci., 52, (1975) 77. C. Defosse, M. Houalla, A. Lycourghiotis, and F. Delannay, in "Studies in Surface Science and Catalysis," T. Seijama and K. Tanabe (Eds.), Elsevier, 1981, 7(A), 108-121. R. C. Reuel and C. H. Bartholomew, J. Catal., 85, (1984) 63. L. Fu and C. H. Bartholomew, J. Catal., 92, (1985) 376. Y. W. Chen and J. G. Goodwin, React. Kinet. Catal. Lett., 26(3-4), (1984), 453. s. H. Moon and K. E. Yoon, Appl. Catal., 16 (1985) 289.
441
DISCUSSION B:G. BAKER: Cobalt reacts with alumina supports to form solid solution spinds C0304-CoA1Z04' At light loadings, cobalt is lost and very low activity results. When lanthanum is present, there is the possibility of forming LaCo03' also a semiconductor. In your catalysts with 10% loading of cobalt, do you have evidence that metallic cobalt is present in the reduced catalyst? D. HERCULES : ESCA spectra of the Co Zp3/Z region measured for the reduced catalysts indicate that a substancial amount of cobalt metal is formed after reduction at 400°C. In addition, the weight loss observed in gravimetric experiments cannot be attributed exclusively to a partial reduction of C0304 to CoO. Thus we conclude that a significant fraction of the cobalt species has been reduced to cobalt metal. J.G. van OMMEN: In your talk, you use the concept of a theoretical monolayer. On what concept did you calculate this monolayer? Is it based on the structure of LaZ03 or on the structure of the support (A1 Z03) ? D. HERCULES: The basis for the monolayer calculation is given in references (11) and (IZ). J.R.H. ROSS: I find your paper extremely interesting, in particular in relation to the fact you see no proportional effect of lanthanum on the turnover numbers for your Co-La-Al system. As you will be aware, in our paper which you quote, we see a significant promotional effect for the Ni-La-Al system when the catalyst is prepared by coprecipitation. I should like to suggest that if you prepared a Ni-La-Al catalyst in the same way as you have prepared your Co catalysts, you might also not see a promotional effect. In later papers and also in unpublished work (R. Paalman, M.Sc. Thesis, 1986), we have shown that the Ni-La content of the coprecipitates is crucial to the promotional effect and we infer that the lanthanum must be closely associated with the nickel to get the improvement of turn-over numbers. Have you looked at the Ni-[a-Al system and do you still get similar results? Have you also looked at other methods of preparation, for example coprecipitation or impregnation by La after Ni ? D. HERCULES: We are currently investigating the properties of La and Th promoted Ni/A1Z03 catalysts. Preliminary results for NiTh/A1Z03 catalysts prepared by coimpregnation and sequential impregnation of nickel first and then thorium suggest that there is no significant promotion of CO hydrogenation activity on addition of thorium. E.B.M. DOESBURG : Do you see any formation of La-aluminate and could that influence your results ? D. HERCULES : We do not have any direct evidence for the formation of a La-aluminate species; however, the formation of a La-A1Z03 surface compound is consistent with the high dispersion of the La phase measured by ESCA. The interpretation of our results for low La loadings (La/Co ~ 0.3) is in fact based on the assumption that La forms a surface compound with the alumina support. A. VANNICE: I presume thei:turnover frequencies were based on your HZ chemisorption measurements. Recent studies by Bartholomew et al. have shown that HZ chemisorption on Co is activated, particularly on smaller crystallites, and higher temperatures near ZOO°C are required for equilibration. At what temperatures were your measurements made and were they made on the used catalyst samples? D. HERCULES: The TOF's reported in the paper are based on HZ chemisorption data. ESCA data for the reduced catalysts were also used to calculate the
442
dispersion of the cobalt metal. The trend in TOF calculated from the ESCA data agrees with that reported using chemisorption data. Our H2 chemisorption measurements were made on the reduced samples at room temperature. We realize that for highly dispersed cobalt catalysts the chemisorption of H2 is an activated process. In fact, as mentioned in our paper, we believe that this activated process may explain why for catalysts with high La contents, the H2 chemisorption results suggest a much lower dispersion than ESCA data. V. PERRICHON : It has been shown in the case of rhodium (for instance by Ichikawa) that lanthanum has a promoting effect on alcohols formation. Did you observed such a formation of alcohols in your conditions? Did you perform catalytic measurements under pressure? D. HERCULES: All the catalytic experiments were performed at atmospheric pressure. Under these conditions, we do not expect any significant formation of alcohols. It would indeed be interesting to perform experiments at elevated pressure to compare the effect of lanthanum to that reported by Ichikawa for Rh-based catalysts. R.A. SCHOONHEYDT : Why does La 3+ inhibit the reducibility of C0 2+/ 3+y ? D. HERCULES: We do not necessarily imply that La 3+ inhibits the reduction of the cobalt phase. As discussed in our paper, we believe that the La species present on La/A1203 samples with high La loadings increase the interaction of the cobalt phase with the La/A1203 support and thus decrease the reducibility of the cobalt. J. BARRAULT : 1/ In your paper, you present results obtained with catalysts calcined at 400°C. Have you studied the effect of calcination temperature and did you obtain the same results? In our laboratory, we observed that high calcination temperature decreased the reducibility of cobalt oxide in cobaltlanthanum oxide catalysts. At the opposite, the reducibility is favoured by low temperature treatments. 2/ Could you give some details about the carbon monoxide conversion and on catalyst duration? D. HERCULES: 1/ We are currently investigating the influence of several preparation variables on the structure and reactivity of CoLa/A1203 catalysts. One would indeed expect a decrease in the reducibility of the cobalt phase at high calcination temperature. This can be attributed to an increase in the formation of Co-La mixed owides or to an increase in the interaction of the active phase with the support. 2/ Activity measurements were performed at conversions less than 5%. After approximately 12 hours on stream no significant change was observed in the activity of the catalysts up to 30 hours. XU Xiaoding : 1/ What is the form of C0304? Is it in a spinel form or ~~~s?
~
2/ Is C0304 reduced to Co or Co metal? Is the decreasing reduction degree of C0304 ' by chance, caused by encapsulation of C0304 by La ? D. HERCULES: 1/ XRD data indicates that the C0304 present is in a spinel form. 2/ ESCA results for the reduced catalysts indicate that cobalt metal is formed after reduction. However, one cannot categorically exclude the possibility that a minor fraction of the C0304 phase is reduced to CoO. We do not believe that there is a significant amount of the C0304 phase encapsulated by La since this should not lead to the observed decrease in the amount of C0304 at high La loadings. We attribute this decrease to the formation of a surface Co-La interaction species. In principle this would decrease the reducibility of the cobalt phase.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
443
EFFECTS OF THE TECHNIQUE FOR THE PREPARATION OF SUPPORTED COBALT CATALYSTS ON SELECTIVITY IN THE FISCHER-'EROPSCH SYNTHESIS F. LIU Department of Chemical Engineering, Hebei Chemical Engineering Institute, Shijiazhuang, People's Republic of China SUMMARY A model for controlling the distribution and maximum length of hydrocarbons in the Fischer-Tropsch synthesis was developed. It follows from the model that all preparation parameters that affect the chain growth probability and maximum length of hydrocarbons must act through the medium of activity, geometry and energy factors. In addition, the maximum chain length was limited by the chain growth probability. The model can be used in both cases, obeying or deviating from the Schulz-Flory distribution. The calculated data based on the model and the experimental results obtained from different preparation techniques agreed well. INTRODUCTION The Fischer-Tropsch synthesis (FTS) has recently attracted much attention as a possible means for the production of fuels and chemical feedstocks. Classical FTS is far less selective (refs. 1-3). As the product of this process consists of a broad range of hydrocarbons between C1 and C30, or even as high as Cl05, studies have been focused mainly on the improvement of catalyst selectivity (refs. 1-4). A number of stUdies have shown that the preparation techniques, particularly for heterogeneous catalysts, ·are the most promising field for this objective, and in some instances remarkable selectivities have been obtained. The chain growth and CO dissociation in FTS require a mUltiple atom site, suggesting that FTS is structurally sensitive (ref. 5). As is known, the techniques of catalyst preparation may be much more critically important for structure-sensitive than for structure-insensitive reactions (ref. 6). Jacobs and co-workers prepared zeolite-supported catalysts with controlled structure that can be used to monitor the carbon number of hydrocarbon products in FTS, and a model based on the particle size effects has been postulated (refs. 7-10). Robert and Bartholomew (ref. 11) observed that the product selectivity varied with the dispersion. Another aspect of this field is the catalyst derived from the thermal decomposition of supported carbonyls,
444
which will cause an increase in the production of olefins of lower mass (ref. 12). Liu and Bartholomew (ref. 13) reported that the average molecular weight of hydrocarbons produced during FTS on Co/alumina was a function of metal loading, reduction temperature and method of preparation. As indicated above, the selectivity of the catalysts in FTS proved to be highly dependent on the methods of preparation. The emphasis in this investigation was placed on the scientific basis for the preparation of supported metal catalysts, and a study of how the preparation techniques affect the selectivity in FTS is reported. CHAIN-CONTROLLING MODEL Reaction mechanism The mOdel advanced in this paper is based on a reaction mecha nism that is generally accepted: the intermediates ci formed from dissociation of CO and sUbsequent hydrogenation of the carbon species add to the growing hydrocarbon chains (refs. 14-17):
cA
C~ + Cl rp' k p C~+l ~ rt, kt
+.
Cn
where C~ and C; represent CnHm and CHx adsorbed on the catalyst surface, Cn is the hydrocarbon product, rp and rt are the rates of propagation and termination of the chain and k p and kt are the rate constants of propagation and termination of the chain, respectively. The product distribution in FTS is generally described by the Schulz-Flory equations (refs. 18-21): (1)
_
rp (2) rp + rt where a is the chain growth probability and Wn is the weight fraction of hydrocarbon containing n carbon atoms. However, it has recently been noted (refs. 22,23) that product distributions in FTS often deviate from Schulz-Flory kinetics. Therefore, the aim of this model is to find a rule for controlling growth probability and maximum length of hydrocarbons in which both cases, obeying or deviating from the Schulz-Flory equation, can be used.
(1
-
Equations for controlling chain growth probability Two reasonable assumptions were made: (1) steady-state condi
445
tions can be reached and in this instance CO will only be converted into hydrocarbon products and C02; (2) the reaction mechanism stated above. Hence the rate equations of chain propagation and termination may be written as follows: rp =
kpCC~CCt
rt = ktCc;'l where CC~ and Cct are the concentrations of C~ and ct. Combining Eqns. (2), (3) and (4) and using the Arrhenius equation, we obtain
where Ap and At are the frequency factors for the chain propaga tion and termination and Ep and Et are the activation energy for the chain propagation and termination, respectively. Suppose Ap and At are the frequency factors including concentration factors (CO conversion and H2/CO were maintained constant), then r p = Ap e-Ep/RT
( 6)
rt = Ate-Et/RT
(7)
~e-(Ep-Et)/RT C1
= +
( 8)
A ite-(Ep-Et)/RT
can be calculated from Eqn. (8), Ap, At, Ep and Et can be obtained from a least-squares linear regression of the logarithmic forms of Eqns. (6) and (7), and rp and rt can be derived from equations described elsewhere (ref. 24).
C1
Equations for controlling the maximum length of a hydrocarbon chain Consider an extreme situation, that is, if a = 1, from Eqn, (2), rt = O. Based on the law of mass conservation, we have Total number of ct produced on the surface of unit weight of catalyst = N(CHx)N(site)t = NCOSHN(site)t
(9)
Total number of hydrocarbon chain produced on the surface of unit weight of catalyst = N(:ite) (10) g
446
where N(CHx) is the turnover frequency factor for the production of CH x' N(site) is the number of active sites on the surface of unit weight of catalyst, t is the propagation time of the hydrocarbon chain on catalyst surface, NCO is the turnover frequency for CO conversion, SH is the selectivity for the hydro carbon product and Ng is the number of active sites required to produce a single hydrocarbon chain, and is probably the number of active sites existing on a single metal particle or a single ensemble. According to the definition of a molecule chain and combined with the above reaction mechanism, the maximum length of the hydrocarbon chain, nm, will be given by nm
= NCOSHNgt = N(CHx)Ngt (a = 1) If a = 0, from Eqn , (2), r p = 0,
(11)
then nm = 1 (that is, methana tion). If O
(12)
It is Eqn. (12) that we are searching for in order to control nm, in which the parameters imply that N(CH x) is an activity factor that provides building blocks CH x for chain growth, Ng is a geometric factor that supplies locations for chain growth, t is an energy factor that represents the strength of binding between the intermediates and catalyst surface and a is a selectivity factor for propagation rate to termination rate that will determine the probability of chain growth. If one of the four factors in Eqn. (12) is much smaller than the others, this factor will determine nm and may be called a controlling factor. For example, if Ng is the smallest factor with respect to others, Ng may be expressed by the number of active sites existing on an ensemble, DN, which will be a controlling factor for Dm, and hence ( 13) where a is a constant. Eqn. (13) means that a linear relationship exists between particle size and the maximum length of the hydro carbon chain when a approaches unity. This conclution is identical with the statement of Jacobs and co-workers (refs. 7-10). EXPERIMENTAL The catalysts used in this study were prepared by one of four techniques: (1) impregnation with an aqueous solution of cobalt nitrate for 3%, 10% and 15% Co/A1203 (refs. 25,26); (2) pH-con
447
trolled precipitation for 3% CO/A1203 (refs. 27,28); (3) evapora tive deposition for 10% Co/C (refs. 25,29); and (4) decomposition of C04(CO)12 on dehydroxylated alumina during cyclohexane reflux for 3% Co/A1203 (refs. 13,30). A modified catalyst (15% Co/A1203) was treated by sintering at 873 K for 24 h in flow of H2. Hydrogen adsorption uptakes were measured statically at 298 K after equilibration for 45-60 min. Catalyst samples of about 1 g were placed in a Pyrex flow-through cell and reduced in a flow of H2 at a space velocity of 2000 h- 1• Measurements of CO/H2 synthesis activity and selectivity were performed in a Pyrex differential fixed-bed reactor at 1 atm, H2/CO = 2 and 443-523 K. The space velocity was varied with the reaction temperature in order to maintain CO conversions within the range 5-10%. All data were obtained after reaction for 30 h, i.e., after steady-state conditions had been reached. The apparatus, procedure and calculation methods used have been described in detail elsewhere (refs. 11,25,26).
RESULTS AND DISCUSSIONS Hydrogen chemisorption uptakes, dispersions, particle sizes and TABLE 1 Characterization of supported cobalt catalysts DC Catalyst dd H2 uptake a N(site)b (l&mol/g) (1/gCoxl0- 20) (%) (nm) 3% Co/A1203 prepared by 11.0 0.21 15.0 6.4 impregnation 3% Co/A1203 prepared by 31.0 30.0 3.2 precipitation 3% CO/A1203 prepared from 0.29 20.0 4.8 C04(CO)12 0.22 10% Co/A1203 prepared by 9.0 10.6 impregnation 10% cole prepared by 121.0 1.0 36.0 2.7 evaporative deposition 15% CO/A1203 prepared by 38.0 0.45 7.8 12.0 impregnation 15% CO/A1203 modified by 0.18 2.2 43.6 sintering a. Based on total H2 uptake at 298 K (refs. 25,26). b. Number of active catalytic sites per gram of catalyst and based on the assumption of an adsorption stoichiometry of 1 hydrogen atom per active catalytic site. c. Cobalt metal dispersion (percentage exposed) (ref. 26). d. Surface mean crystallite diameter (ref. 26). number of active catalytic sites for all catalysts studied are listed in Table 1, and their catalytic activities and selectivi
448
TABLE 2 Catalytic activities and selectivities of supported cabalt catalysts in Catalyst
NCO
b
(5- 1 x103)
3% Co/A1203 prepared by impregnation 99.0 0.73 14 3% Co/ A120 3 prepared by precipitation 2.00 0.84 98.6 0.71 13 2.88 3% CO/A12 0 3 prepared from C04(CO)12 2.00 99.6 0.77 15 8.73 10% Co/A1203 prepared oy impregnation 10% Co/C 0.19 92.8 0.47 10 0.17 0039 prepared by evaporative deposition 15% CO/Al203 14.10 10.92 14.00 99.5 0.78 18 :prepared by J.mpregnation 7.00 99.4 0.78 18 15% CO/A1203 modified by sintering a. Catallsts reduced at 648 K for 16 h in a flow of H2 and reacted at 473 K, H2/CO = 2, 1 atm. b. Turnover frequency for CO conversion, i.e., number of CO molecules converted per catalytic site per second. c. Turnover frequency'for production of CH x ' i.e., number of CH x produced per catalytic site per second. d. rp and rt are the rates of chain propagation and termination, i.e., number of CH x used in the chain propagation and termination per catalytic site per second. e. Selectivity with respect to hydrocarbon, i.e., the ratio of CO converted into hydrocarbon to CO converted into hydrocarbon + C02. f. The chain growth probabilities were obtained from a least-squares linear regression of the logarithmic form of Eqn. (1). g. Maximum length of hydrocarbon chain. ties in FTS are listed in Table 2. There are two cases for the relationship between the particle size and product selectivity: (1) the collision frequency, binding energies between the carbon chain and the catalyst surface and the turnover frequency are directly proportional to the particle size (refs. 5,6); (2) the maximum hydrocarbon chain length is strictly limited by the particle size (ref. 7). Catalysts having the same support and method of preparation but different metal loadings belong to the former case, the chain growth probability increasing linearly with increasing particle size (see Tables 1 and 2). Bartholomew and
449
TABLE 3 Influence of reaction temperature on chain growth probabilities a Catalyst
443 K
463 473 K K - 0.70
488 498 K K - 0.62
508 K
518
523 K
K
3% Co/A1203 prepared by impregnation 3% Co/A1203 prepared - 0.74 - 0.70 by precipitation 3% Co/A120, prepared - 0.73 - 0.66 from C04(CO)12 10% Co/A1203 prepared - 0.78 0.75 0.72 0.65 0.57 by impregnation 10% Co/C prepared by - 0.48 - 0.49 evaporative deposition 0.82 0.79 0.78 0.74 15% Co/A1203 prepared by impregnation 15% Co/A1203 modified 0.79 0.80 0.76 0.75 by sintering a. Reaction conditions: HZ/CO = 2, 1 atm and the reaction temperature shown above.
0.50 0.60 0.49
0.44
co-workers (refs. 24-26) observed that CO adsorption on 3% Co/A1203 prepared by pH-controlled precipitation was greater than that on 3% Co/A1203 prepared by impregnation. In contrast to catalysts having different metal loadings, the particle sizes have opposite effect on the selectivity behaviour of catalysts prepared by different methods, which can likewise be explained by differences in the relative distributions of different adsorption states (refs. 24,25). 1.0 0.8
a
70
......
'*
60.;
a
iI
1.0
60
30
0.4 C2- C4
20
0.2 C12+
.. .. ..
10
0 0.0 468 473 488 498 508 518 528 T (K)
=
0 ..0
~ o
......
'* ..... iI
~
S-tl
IS
70
0.8
50 .... u ::s 40 'g 0.6
0.6
b
50 .... o ::s '0 40 0 ~
P,
30
0.4
= 0
.0
20
e
0.2
-=
0.0 468 473
~ o
0
1O,a
'0
»
498
» 0 -= 523 528
T (K)
Fig. 1. Variation of selectivity with reaction temperature. (a) 10% Co/A203 prepared by impregnation; (b) 10% Co/C prepared by evaporative deposition.
450
TABLE 4 Kinetic parameters of supported cobalt catalysts in FTSa Catalyst ECO ACO Ep Ap Et At (KJ/mol) (KJ/mol) (KJ/mol) 86 7.8 X10 6 2.4 x105 107 74 3% Co/Al 2 0 3 prepared by impregnation 82 6.1x107 72 3% Co/Al 2 0 3 97 prepared by precipitation 118 80 3% co/Al203 95 prepared from C04(CO)12 122 90 10% Co/Al203 99 prepared by impregnation 10% Co/C 130 135 133 prepared by evaporative deposition 15% Co/Al203 115 112 1.9 xl0 10 127 prepared by impregnation 6.2xl0 10 116 1.o-io 10 117 115 15% Co/Al203 mOdified by sintering a. ECO, Ep, Et, ACO, Ap and At are the activation energies and frequency factors for CO conversion, chain propagation and termination. Data based on a least-squares linear regression of the logarithmic form of Eqns. (6) and (7). The variation of the product selectivity with reaction temperature is shown in Table 3 and Fig. 1. The influence of reaction temperature on the product selectivity is consistent for most catalysts, i.e., as the operating temperature increases the product selectivity shifts to lower molecular mass compounds (refs. 31,32). However for la% Co/C prepared by evaporative deposition the chain growth probability increased with increasing reaction temperature in the range 473-498 K. This unusual behaviour was confirmed by the same phenomenon which occurred on the 15% CO/AlZ03 sintered catalyst. Kinetic parameters for all catalysts are listed in Table 4. It is evident that 10% Colc prepared by evaporative deposition has an extremely small value of -(Ep-Et) (5.0 kJ/mol), and the 15% CO/AlZ03 sintered catalyst even has a negative value (-1.0 kJ/mol). This is the reason why the chain growth probabili ties of the two catalysts indicated above increased with increasing reaction temperature in certain ranges. Based on Eqn. (8), if Ep = Et, e-(Ep-Et)/RT equal to unity, and in this event
451
TABLE 5 Comparison of chain growth probability obtained from Schulz-Flory kinetics and the equation presented here a Catalyst Chain growth probability obtained from Eqn. (8) Eqn , ( 1 ) 3% Co!AIZ03 prepared by impregnation 0.70 0.69 3% CO!AIZ03 prepared by precipitation 0.74 0.73 3% CO!AIZ 0 3 prepared from C04(CO)lZ 0.71 0.73 10% CO!AIZ03 prepared by impregnation 0.78 0.77 10% Co!C prepared by evaporative 0.48 0.47 deposition 15% Co!AIZ03 prepared by impregnation 0.78 0.78 15% CO!AIZ03 modified by sintering 0.76 0.78 a. Reaction conditions: 473 K, Hz/CO 2, 1 atm.
=
1.0
0
0.8
'" *'" '"
0.6
A
e
0.4
*
A& A
* 1
-2 -4 ,.... ><
&
:::r::
-6
o
z 1=1 ,...;
0.2
-8
In nm Fig. 2. (*) Relationship between In nm and« and (1) relationship between In nm and In N(CH x). the reaction temperature has no effect on the chain growth probability. The chain growth probabilities obtained from Eqns. (8) and (1) agreed closely, as can be seen in Table 5. The rela t10nships between ln nm and a and between ln nm and ln N(CHx), derived from the logarithmic form of Eqn. (lZ), are plotted in Fig. 2. The product of N(CHx)Ngt gives the possibility of forming a maximum hydrocarbon chain length and a provides the probability of forming a realistic hydrocarbon chain length. Eqn. (12) can give some important clues for improving the selectivity in FTS. The chain growth probability of the 15% Co!Al203 catalyst mOdified by sintering surpassed that of the fresh catalyst under
452
certain conditions. This novel characteristic is very interesting, although the crystal growth was outside the range where particle size effects are important. REFERENCES
1 M.E. Dry, Ind. Eng. Chem., Prod. Res. Dev., 15 (1976) 282-286.
2
H.H. Storch, N. Golumbic and R.B. Anderson, The FischerTropsch and Related Syntheses, Wiley, New York, 1951, p. 114. 3 R.B. Anderson, in P.H. Emmett (Ed.), CatalYsis, Reinhold, New York, 1956, Ch. 2, p. 29. 4 H.G. Stenger, H.E. Johnson and C.N. Satterfield, J. Catal.,
86 (1984) 477-480. 5 M. Boudart, J. Mol. Catal., 30 (1985) 27-38. 6 M. Boudart and M.A. McDonald, J. Phys. Chem., 88 (1984) 2185. 7 H.H. Nijs, P.A. Jacobs and J.B. Uytterhoeven, J. C. S. Chem. Commun., 4 (1979) 1095-1096. 8 H.H. Nijs and P.A. Jacobs, J. Catal., 65 (1980) 328-334. 9 P.A. Jacobs, in B. Imelik, C. Naccache, Y. Ben Taarit, J.C. Vedrine, G. Coudurier and H. Praliaud (Eds.), Catalysis by Zeolites, Elsevier, Amsterdam, 1980, pp. 293-308. 10 H.H. Nijs, P.A. Jacobs and J.B. Uytterhoeven, J. C. S. Chem. Commun., (1979) 180-181. 11 C.R. Robert and C.H. Bartholomew, J. Catal., 85 (1984) 78-88. 12 J. Zwart and R. Snel, J. Mol. Catal., 30 (1985) 305-352. 13 F. Liu and C.H. Bartholomew, J. Catal., 92 (1985) 376-387. 14 M.A. Vannice, Catal. Rev.-Sci. Eng., 14 (1976) 153-191. 15 G.A. Somorjai, Catal. Rev.-Sci. Eng., 23 (1981) 189-202. 16 A.T. Bell, Catal. Rev.-Sci. Eng., 23 (1981) 203-232. 17 P. Biloen and W.M.H. Sachtler, Adv. Catal. Rel. Subj., 30 (1981) 165-216. 18 G. Henrici-Olive and S. Olive, Angew. Chem. Int. Ed. Engl., 15 (1976) 136-142. 19 R.J. Madon, J. Catal., 57 (1979) 183-186. 20 C.N. Satterfield and G.A. HUff, Jr., J. Catal., 73 (1982) 187-197. 21 P.A. Jacobs and D.V. Wouwe, J. Mol. Catal., 17 (1982) 145-160. 22 R.A. Dictor and A.T. Bell, Ind. Eng. Chem. Process Res. Dev., 22 (1983) 681-684. 23 G.A. HUff, Jr. and C.N. Satterfield, J. Catal., 85 (1984) 370-379. 24 F. Liu, J.L. Rankin and C.H. Bartholomew, Cl Mol. Chem., (1986) 369-385. 25 C.H. Bartholomew and R.B. Pannell, J. Catal., 65 (1980) 390-401. 26 E.J. Erekson and C.H. Bartholomew, Appl. Catal., 5 (1983) 323. 27 J.T. Richardson and R.J. Dubus, J. Catal., 54 (1978) 207. 28 C.H. Bartholomew and M. Boudart, J. Catal., 25 (1972) 173. 29 R.C. Reuel and C.H. Bartholomew, J. Catal., 85 (1984) 63-77. 30 J.E. Crawford, G.A. Melson, L.E. Makovsky and F.R. Brown, J. Catal., 83 (1983) 454. 31 M.E. Dry, in J.R. Anderson and M. Boudart (Eds.), Catalysis, 32
Science and Technology, Vol. 1, Springer Verlag, Heidelberg, 1981, Ch. 4, pp. 211-222. M.E. Dry, J. Mol. Catal., 17 (1982) 133-144.
453
DISCUSSION P. CANESSON : Have you measured the porous repartition of your various catalysts, since it is possible to limit the chain growth in FT process by capillary condensation of hydrocarbons and a subsequent hydrogenolysis on metallic particles as recently proposed (D. Vanhove et al., Appl. Catal., 9 (1984) 327). Would you comment on this? LIU Fu : The data of this study are not consistent with the hypothesis of Vanhove et al. (Appl. Catal., 9 (1984) 327) involving condensation and hydrogenolysis of liquid hydrocarbons-in the catalysts pores. Six catalysts in this study were prepared from the same support (alumina) and hence involving a single distribution of pore sizes; yet very significant differences were observed in the product selectivity (chain growth probability and maximum length of hydrocarbon chain) of 3% Co/A1203 catalysts prepared by different techniques. The model of Vanhove and coworkers would predict the same product selectivity for all three catalysts. Moreover, the analysis of the data from this study leads to the conclusion that longer residence times of hydrocarbon (longer times for chain propagation) leads to longer chains; the model of Vanhove and co-workers would lead to the opposite conclusion, since a longer residence time would favor hydrogenolysis reactions. It is evident that the conclusion of Vanhove and co-workers is in conflict with the results that were obtained from most experiments and most catalysts. This point of view has been exposed by Liu Fu and C.H. Bartholomew (J. Catal., 92 (1985) 376-387). XU Xiaoding : All the Co sites are used for two purposes: (1) dissociation of H2 and (2) CO adsorption. The Cn surface species are also used for two purposes: (1) addition of one Cl to Cn+l and (2) hydrogenation to hydrocarbons (eventually). Would you comment on it, since you did not consider them in your model? (If the reaction is the chemical reaction determined). LIU Fu : This question can be answered through three points: 1/ adsorption and dissociation of CO should be mainly considered in this situation because the difference of capacity for CO+H2 conversion only results from the difference of CO adsorption and dissociation, not H2' as proposed by B.E. Nieuwenhuys (Surf. Sci., 126 (1983) 307); 2/ the product type is not important here, so the length of the hydrocarbon chain (carbon number contained in hydrocarbon product) is concerned only; 3/ Cn + Cl leads to chain propagation, and Cn desorption leads to chain termination (irrespective of the formation of paraffins or olefins, and through H2 addition into hydrocarbon chain or elimination from hydrocarbon chain). Hence, the competition of the rates between the chain propagation and the chain termination will determine the chain growth probability.
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B. Delmon, P. Grange, P.A.Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV
455
© 1987 Elsevier SciencePublishers a.v.. Amsterdam - Printed inThe Netherlands
SELECTIVE FISCHER-TROPSCH CATALYSTS CONTAINING IRON AND LANTHANIDE OXIDES B.G. BAKER and N.J. CLARK School of Physical Sciences, Flinders University, South Australia 5042 SUMMARY Catalysts with high selectivity for alkene production by the Fischer-Tropsch synthesis have been prepared. They consist of highly dispersed iron on an alumina support which has been impregnated with a lanthanide oxide. Selectivity towards light alkenes exceeds 90% and methane production is controlled at less than 5%. The essential steps in the catalyst preparation are (i) heat treatment of y-alumina to change the pore structure, acidity and surface area. (ii) impregnation of the alumina with a solution containing the lanthanide element in sufficient quantity to coat all of the support surface. (iii) impregnation with a ferric solution of controlled pH and containing an amount of iron comparable to the amount of lanthanide. INTRODUCTION The synthesis of hydrocarbons from carbon monoxide and hydrogen (the FischerTropsch synthesis) over a variety of transition metal catalysts is an established technology, however the product is generally far from optimum (refs. 1-4). Straight chain saturated hydrocarbons are principal products and methane and high molecular weight oils and waxes are undesirable products. The overall carbon number product distribution does not match the market requirements and the products of low intrinsic value inflate the refining costs to produce the desired range of liquid fuels. The present work has aimed to modify the properties of iron catalysts to achieve greater selectivity towards a narrower distribution of more valuable products.
A particular strategy is to form light alkenes which may be
converted to liquids by oligimerisation or alkylation (ref. 3,5).
High
selectivity to alkenes and suppression of the formation of methane are essential requirements of the catalyst. Such a catalyst must also have resistance to coking under conditions which favour the formation of hydrogen deficient products. These objectives have been achieved in the present work by highly dispersed iron catalysts containing lanthanide elements. CATALYST SUPPORT Catalysts were prepared on supports derived from a spheroidal y-AlzOs containing about 0.2% NazO (Merck). It has a surface area of - 78 m2/g, a porosity of - 0.7 mL g-l and an apparent surface pH ~ 8.0. The alumina was
456
sieved and only the fraction with particle diameters between for catalyst preparation.
was used
125-150~
It was found that a superior catalyst support was
produced by heat treating this y-Al z0 3 • The y-Al z0 3 was dried in an oven at - 150°C, preheated at 800°C for 15 minutes then tumbled in a furnace at high temperature for a specific time.
Typical times and temperatures used were
10-15 minutes at l250°C, 30 minutes at l200°C, 2-3 hrs at Il50 oC.
The tumbling
was particularly important when short heating times were being used so that the alumina was uniformly heated. of - 20
mZ
HT-alumina.
The product of this treatment had a surface area
g-l and an apparent surface pH = 10.
It is referred to as
Excess heating at high temperatures caused deterioration of the
alumina support with further loss in area and decrease in basicity of the support.
X-ray diffraction shows that the structure of HT-alumina is
predominantly a-Al z0 3 • The structural change, effected by heating, was studied by scanning electron microscopy.
Alumina heated to l150°C for only 15 min had no clearly resolved
pore structure.
Further heating at this temperature for 180 min resulted in
the development of a more open pore structure, (see Fig. 1). Measurement of the pore size distribution in HT-alumina by mercury porosimetry show radii 0.012 and 0.007
~m.
Pores in these size ranges account for
60% of the total surface area of the sample. The catalyst has a somewhat broader distribution of pore sizes.
The
characteristic small pores are smaller indicating that impregnated iron and praseodymium oxide have penetrated these pores. A number of commercial samples of y-A1 20 3 were tested.
Only those contain-
Fig. 1. Scanning electron micrograph of HT-alumina after heating at l150 oC, 3 hr. (scale ~m)
457
ing a sodium impurity, usually 0.2-0.3 percent, showed the structural and Only samples containing sodium showed the chemical changes reported here. basic surface reaction suggesting that a soluble compound of sodium was present at the surface after the heat treatment. It was recognised that this sodium could influence the impregnation stage of catalyst preparation and would be a Washings from the HT-alumina were analysed component of the final catalyst. by I.C.P. spectra and it was shown that 60% of the sodium impurity is water soluble and that almost 100% is soluble in acid at pH = 1, which is the approximate pH of the impregnating solutions used for catalyst preparation. CATALYST PREPARATION The catalyst is prepared by impregnation of the support first with a solution of the rare earth nitrate followed by microwave drying and calcining to decompose the nitrate. Secondly a solution of ferric nitrate of pH 1-1.2 is introduced by the incipient wetness technique; again followed by drying and calcining.
The catalyst is reduced at 400°C in hydrogen before proceeding to
the test of Fischer-Tropsch activity. Typical loadings were 3 percent by weight of the rare earth oxide and 2 percent by weight of iron. These quantities on HT-alumina result in a highly dispersed catalyst. The incipient wetness volume of HT-alumina is 0.65 mL/g. Solution concentrations were set to achieve desired loadings at a single impregnation. The pH of the iron solution was found to be a significant parameter. HTalumina is strongly basic and iron(III) nitrate solutions are strongly acidic. At pH>l significant time dependent polymerisation occurs in iron(III) solutions. After investigation of the variation of the pH of the iron impregnating solution on catalyst performance a standard impregnating method using freshly prepared iron(III) nitrate at pH by microwave heating.
I was adopted.
Drying of the catalyst was generally
Oven drying leads to inhomogeneous catalyst preparations
due to gross metal ion transport by a chromatographic process.
This can be
overcome by vacuum drying or microwave drying, the latter being much faster. REACTION TESTING Catalyst performance was tested in a microreactor. Typical conditions were; temperature 260-300°C, flow 900 hr- 1 , pressure 800 kPa. The ratio of carbon monoxide to hydrogen ranged from 2 to 0.5. Results are presented in tables 1-4. Analyses were by gas chromatography of samples taken by a heated line from a gas sampling valve. Extended tests on these catalysts have shown that the activity and selectivity are maintained over several weeks with remarkable freedom from coking. Highest yields of a1kenes are found with CO/H 2 > 1, a condition which generally leads to coking of iron catalysts.
458 TABLE 1 Effect of iron-loading, modifier and support on F-T product distributions (weight percent). Reaction pressure 800 kPa; flow 900 hr- 1 CO/Hz = 2 Catalyst Composition Wt %
Fe
6.0
6.0
1.6
1.6
CO converted to hydrocarbon (%)
-14.0
'c
300
Methane
Alkanes
0.3
800/ 120
1250/ 10
1250/ 10
0.3
10.4
14.7 280
300
280
280
280
30.4
22.6
21.9
25.8
14.1
7.0
7.8
16.7
14.0
17.5
13.4
7.4
17.6
22.9
18.4
15.9
17.1
25.6
C 4 C 5 C6
8.5
12.5
12.2
11. 7
12.9
20.7
2.6
6.9
6.0
3.3
11.8
6.0
1.3
2.7
2.7
1.8
8.4
3.9
C + 7
1.3
2.3
2.7
1.8
13.6
7.1
C2 C3 C 4 C 5 C 6 C + 7
13.9
8.5
16.4
24.8
3.7
5.6
5.9
2.3
2.8
3.4
2.5
1.5
2.8
1.0
1.9
2.5
2.0
1.2
0.9
0.5
0.9
0.7
2.1
1.4
0.4
0.2
0.4
0.4
0.4
0.2
0.4
0.4
1.4 2.3
1.2 2.1
C2 C3 Alkenes
2.4
2.0 2.0
Heat treatment of y-alumina 'C/min.
Reaction temperature
1.6
1.0
PrSOll
25
~ Alkanes
o
20 f-
-
Alkenes
>z w
I
j
u 15 0: w a, I-
:I:
10 f-
CJ
iii
3=
5
o
I~
r!l
3
m 4
... 5
PRODUCT
~
~ n
n
9 10 7 8 6 CARBON NUMBER
11
12
Fig. 2. Product distribution (weight percent) frOm F-T reaction on 2% Fe/3% PrSOll/HT-alumina catalyst. Reaction conditions as in table 1.
459
TABLE 2 Effect of modifier on F-T product distribution (weight percent) from a 2% iron Reaction temperature 280°C, pressure 800 kPa, flow
catalyst on HT-A1 20 3 • 900hr- 1 , CO!H 2 = 2 Catalyst Modifier
3% Pr6011
CO converted to hydrocarbon (%) Methane C 2 C 3 C 4 C 5 C 6 C + 7
Alkenes
C 2 C 3 C 4 C 5 C 6 C + 7
Alkanes
,, 50
3% La 20 3
1. 8% K20
3.9
2.5
0.6
7.5
5.1
6.0
27.6
13.8
16.2
9.4
12.2
11.8
19.5
22.2
16.7
17.8
17.4
17.8
19.2
13.8
16.3
9.0
14.1
13.3
13.0
16.0
6.6
5% Ceria (Tech.Grade)
3% Tech. Ceria (depleted)
5.4
2.6
4.8
10.2
8.5
8.6
10.0
5.1
17.6
12.9
14.4
11.5
6.4 5.9
0.6
0.2
4.7
4.1
0.8
0.3
3.0
2.3
1.2
0.5
0.3
3.2
1.2
3.1
0.3
0.1
3.0
1.1
2.8
0.1
0.1
2.0
0.7
2.3
3.3
0.8
1.0
,,
f-
Z
w
., ,
u
n:: w
10
W
5
0-
-'
a
::;: 0
-'
w >0.5 2
5 6 7 8 9 10 11 3 I, PRODUCT CARBON NUMBER
Fig. 3. Chain growth characteristic
data of fig. 2.
12
460 TABLE 3 Effect of reaction conditions on F-T product distribution (weight percent) from 1.6%Fe/2%Pr6011/HT-Alz03' CO/Hz = 2 Temperature 260 260 260 280 280 300 320 (OC) Flow (hr-1)
1800
1800
900
900
900
900
900
Pressure (kPa)
1240
800
800
800
800
800
800
2.9
2.5
5.2
16.8
12.5
22.6
37.2
CO converted to hydrocarbon (%) Methane
A1kenes
C2 C 3 C4 C5 C 6 C 7
Alkanes
C 2 C3 C 4 C 5 C6 C 7
+
+
7.3
8.6
7.3
8.0
5.0
6.0
8.6
7.8
9.2
7.7
7.0
15.0
14.8
16.0
13.1
15.3
14.5
15.4
23.8
23.8
24.4
12.0
13.2
13.2
13.9
17.5
16.1
15.2
11.1
10.8
12.4
12.1
12.0
11.3
10.0
5.7
8.0
7.1
6.0
15.3
14.8
15.1
11.7 1.3
7.9
7.5
26.7
22.5
8.0 22.1
3.4
4.0
3.1
2.9
0.9
0.6
2.1
2.3
2.2
1.8
1.0
1.4
2.0
2.0
2.2
2.0
1.9
0.7
1.0
1.2
1.6
1.1
2.2
2.3
0.6
0.8
0.9
1.1
0.8
1.4
1.1
0.5
0.7
1.0
3.8
2.5
4.0
2.8
0.9
1.4
2.0
Figure 2 illustrates the product distribution obtained from a catalyst containing 2 wt % Fe and 3 wt % Pr6011 on HT-alumina. as for data in Table 1.
Reaction conditions are
Alkenes comprise> 90% of the product, mainly between
Cz and Ca, and methane < 5%. The data is plotted logarithmically as mole percent versus C number in Figure 3. The departure from the straight line shows that the inherent disadvantage of a Schulz-Flory distribution has been overcome. The relatively sharp cut-off of chain growth, evident in both Figure 2 and Figure 3, is a most desirable feature of the distribution.
The a1kenes are
predominantly 1-a1kenes, i.e. the double bond is at the end of a straight chain. DISCUSSION The reaction results (table 1) show that an improvement in catalyst properties can be realised by impregnating the active component on an alumina support heat treated in accordance with the procedure described.
For example a
catalyst comprising Fe impregnated on a HT-alumina support exhibits higher
461 TABLE 4 The effect of CO/H 2 ratio on F-T product distribution (weight percent) from 1.6%Fe/1%PrS011/HT-A1203 Temperature (OC)
260
280
240
260
260
260
GHSV (hr -1)
900
900
900
900
900
900
Pressure (kPa)
800
800
800
800
800
800
2
2
0.5
0.5
2
0.5
5.1
15.6
4.8
12.8
3.6
17.4
6.9
7.4
10.2
13.0
6.6
15.6
9.4
6.9
8.0
6.1
9.8
4.9
15.0
15.3
14.3
16.7
14.9
16.6
13 .1
13.7
11.4
12.4
13 .5
12.3
12.7
13.5
11.8
11.7
14.2
10.1
8.8 24.0
9.9
7.4
7.2
11.6
6.5
19.6
17.2
9.2
21.9
11.1
CO/H 2 CO converted to hydrocarbon (%) Methane
Alkenes
C2 C 3 C 4 C5 C 6 C + 7
Alkanes
C2 C 3 C4 C 5 C 6 C + 7
2.8
2.3
4.0
8.0
3.0
7.1
1.9
3.0 2.6
3.5 3.4
1.7 1.1
3.5
1.5 1.1
1.8 2.1
3.3
2.2
3.1
3.5
0.5
3.1
0.8
1.7
2.1
2.4
0.3
2.2
2.3
3.3
4.9
3.0
0.9
3.7
2.83
1.27
Mole ratio CO 2/H 2O
activity and better selectivity than a catalyst comprising 1.6 wt % Fe impregnated on a conventional y-a1umina support: 0.3 to 10.4% CO converted to hydrocarbons.
the activity increases from
Furthermore, the catalyst including
the HT-a1umina support results in a substantial reduction in undesirable methane formed in the product and a significant increase in preferred alkenes having carbon numbers between C2 and Cs• A further improvement in properties is obtained by impregnating the HTalumina with both the active component and a modifier component such as praseodymium oxide.
For example a catalyst comprising 2% Fe, 2% PrS011 on
HT-a1umina exhibits higher activity and better selectivity than catalysts which do not include praseodymium oxide. Table 2 illustrates the effect of a modifier component in catalysts comprising 2% Fe on a HT-alumina support.
Catalysts prepared with lanthanide
modifier components exhibit better activity and selectivity properties than a
462
catalyst prepared with 1.8 wt % K20, a conventional modifier in F-T catalysts. In particular the catalysts which include the lanthanide modifier components Lanthanide elements found to be effective in form alkenes and less methane. suppressing the formation of methane and favouring light alkenes are praseodymium and lanthanum. Cerium was found ineffective. Technical grade ceria (mixed rare earths) was effective and could be improved by partial removal of cerium. Table 3 shows the effect of reaction conditions on product distribution for catalysts comprising 1.6 wt % Fe and 1.0 wt % Pr6011 on a HT-alumina support and catalysts comprising 1.6 wt % Fe and 2.0wt%Pr6011 on a HT-alumina support. As the reaction temperature increases there is a significant increase in activity with only a marginal deterioration in selectivity. Table 4 illustrates the effect of the CO/H 2 ratio on activity and product distribution. on HT-alumina.
The catalyst tested comprised 1.6 wt % Fe and 1.0 wt % Pr6011 The significant point evident from the table is that an excess
of hydrogen still results in the formation of alkenes as the major product. The selectivity of these catalysts is a function of the support, the high state of dispersion of iron and on the presence of a lanthanide oxide. Praseodymium appears to have the most favourable effect, possibly because it has accessible multiple valencies and changes of oxide composition are kinetically favourable (ref. 6).
The proximity of a semiconducting oxide to
active iron crystallites is suggested as the most likely explanation of the modified behaviour of iron as a Fischer-Tropsch catalyst. These catalysts have very long active life and appear to have an advantage A in both durability and selectivity over other alkene selective catalysts. catalyst containing manganese and zinc oxide (ref. 7) bears some comparison but produces much larger yields of ethylene. The HT-alumina, prepared as described here, has soluble sodium and aluminium These are important which are redeposited during the catalyst preparation. constituents and a further study of the effect of varying the aluminium content of the impregnating solutions has been undertaken (ref. 8): Patent applications for these catalysts have been filed (ref. 9). ACKNOWLEDGEMENTS This work was supported by the Australian National Energy Research Development and Demonstration Council. REFERENCES 1 H. Pichler, Gasoline synthesis from carbon monoxide and hydrogen. Advances in Catalysis, IV, (1952) 271-341. 2 M.A. Vannice, Synthesis of hydrocarbons from CO and H, Catal. Rev. - Sci. Eng. 14, (1976) 153-191.
463
3 4 5 6 7 8 9
D.L. King, J.A. Cusamano and R.L. Garten, A. Technological perspective for catalystic processes based on synthesis gas. Catal. Rev. - Sci. Eng. 23 (1981) 233-263. M.E. Dry, Advances in Fischer-Tropsch chemistry. Ind. Eng. Chern., Prod. Res. Dev., 15 (1976) 282-286. R.A. Sheldon, Chemicals from synthesis gas, Reidel, 1983, pp.7l-73. B.G. Hyde, D.J.M. Bevan and L. Eyring, On the praseodymium + oxygen system, Phil. Trans. Roy. Soc. London 259 (1966) 583-614. B. Bussemeier, C.D. Frohning, G.H. Horn and W. Kluy, German Offen. 2,518,964 and 2,536,488 (1976). J. Abbot, N.J. Clark and B.G. Baker, Effects of sodium, aluminium and manganese on the Fischer-Tropsch Synthesis over alumina-supported iron catalysts, submitted for publication in Applied Catalysis. B.G. Baker, N.J. Clark, H. McArthur and E. Summerville, International Patent Application PCT/AU83/00110.
464
DISCUSSION J. KIWI: It is not possible in a classical thermodynamic situation to form FeA1204 at the temperatures you hydrogenate CO in your reaction. Could you elaborate how this perovskite is formed on your catalyst? B.G. BAKER: The reduction of the catalyst in hydrogen at 400°C results first in the reduction of iron oxide to metallic iron. We have followed this process by XPS analysis and found that in the presence of alumina, reaction does occur and that, at the ratio of Al/Fe in our catalyst, iron is finally found to be in the +2 and +3 valence states. Because the sample does not charge in the electron spectrometer, we assume that both Fe and Al are in a semiconducting oxide phase and propose the solid solution spinel Fe304-FeA1204' The perovskite that I reffered to was PrFe03' The XPS evidence is that Pr is not present in the reduced sample as Pr6011' the oxide present after calcination. I think that the formation of these compounds at 400°C would be kinetically limited if bulk preparations were attempted but believe that reaction occurs readily in the very thin (only a few monolayers) overlayer on the catalyst support.
J. BARRAULT : You give in your paper some details about the negative effect of cerium oxide, but how do you explain the positive effect of lanthanum oxide on the acidity and selectivity of iron? B.G. BAKER: The rare earths which show a positive effect on selectivity all have access to the +3 valence state. Cerium is more likely in the +4 state whereas praseodymium readily assumes either state. Lanthanum could form LaFe03' L. GUCZI : Iron tends to agglomerate under reducing conditions on a support. Effect of oxide on this agglomeration is considered as this additive prevents this migration and very small particles are stabilized even if it is in the form of oxide. The outcome of small particles is twofold: i) inactive carbide formation is prevented and ii) weakly bonded hydrogen being responsible for the hydrogenation step is depleted (e.g. L. Guczi et al., Appl. Cat. 22). This selectivity for olefins increases. I wonder whether you have considered this possibility too? B.G. BAKER: I agree that non-reducible oxides could have this effect. It is possible that with the ratios of Al/Fe and Pr/Fe in our catalysts there might not be metallic crystals present after reduction. The selectivity we observed is not characteristic of metallic iron, however very small crystals in proximity to semi-conducting oxide could have quite different catalytic properties. A. VANNICE: Have you studied Fe dispersed on Pr6011 only? lytic behaviour was observed in the absence of A1203 ?
If so, what cata-
B.G. BAKER: We have prepared unsupported catalysts by homogeneous precipitation. However, these contained Fe, Al, Pro We have not studied the Fe-Pr system in the absence of Al. E.V.W. GRITZ: What is the stability of your catalyst with time on stream and what is the influence of higher synthesis gas pressures? B.G. BAKER: We have studied the reaction of CO/H? = 2 on our catalyst at pressures up to 12 atm for periods of - 3 weeks wlthout loss of activity or selectivity. The catalyst has been tested over similar periods of time at pressures up to 20 atm in another laboratory with similar results. The selective catalyst shows no evidence of coking under any condition investigated.
465
P. CHAUMETTE : Have you any idea on the influence of C02 added to syngas on the performance of the catalyst? B.G. BAKER reactivity product of relatively
: We have not conditioned catalyst in C02 and have not tested the to syngas containing added C02' This gas is however a major bythe reaction of CO with H2 and the catalyst must be exposed to high concentrations of C02 under normal reaction conditions.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers BV, Amsterdam - Printed in The Netherlands
467
TITANIA SUPPORTED IRON-RUI'HENIUM CATALYSTS FDR FISrnER-'IROPSCH SYN'IHFSIS
LIN LIwrt-, FRANK J. BERRy WANG
CHENGYJ-
2,
DU fDNGZHANG
l,
LIAOO OONGBAI
l,
TANG RENYUAN l,
and ZHANG suI.
lr:alian Institute of Chemical Physics, P.O. Box 100, 129 Street, Dal.i.an (people's Republic of China) 2nepartrrent of Chemistry, University of Birmingham, P.O. Box 363, Binningham BIS 2TI' (United Kingdom) SUMMARY
Iron-rutheniun catalysts supported on titaniaS'l differing surface areas have teen prepared and characterised by in-situ Fe l'tIssbauer spectroscopy and terrperature programred reduction techniques ('IPR). The results have shown that interactions tetween iron and titania supports are strongest when high surface area titania is used. fbN"ever, such interactions are weaker than those l:etween iron and silica and tetween iron and aluminia. The catalytic properties of these materials for Fischer-Tropsch have teen evaluated in a microreactor system at 23SoC and at 25 Kg/an2synthesis pressure. The tncorporataon of iron into titania-supported rufheni.im catalysts has a significant influence on selectivity. In contrast, the selectivities of less active silica-supported catalysts are insensitive to the iron concentration. INI'ROOOcrrON Significant attention has teen given in recent years to supported bimetallic catalysts for Fischer-Tropsch synthesis.
This interest is largely a result of
the promise which such materials have srown for :i.rrproving the selectivity of Fischer-Tropsch reactions.
Although iron-ruthenium catalysts have teen studied
in the past (1-3) the investigations have teen mainly perfonned when the biIretallic phase has teen supported on alunina or silica.
It is therefore
pertinent to note that titania has teen shown to exhibit strong metal-support interaction (s-tSI) with many transition metals (4,5) and that these s-tSI effects have teen found to influence the catalytic properties of such materials. Given that titania-supported iron-ruthenium catalysts have attracted only limited attention in the past (6,7) we have initiated an investigation of J::oth
iron- and iron-ruthenium catalysts on titania supports with differing surface areas.
In this paper
Y.e
report on our studies of these titania-supported
catalysts by in-situ Ml)ssbauer spectroscopy- and t.errperature programred reduction (TPR) techniques and corrpare the metal-support interactions with those in catalysts supported on alunina and silica.
we
also assess the influence of
such interactions on the catalytic selectivity of the catalysts for FischerTropsch synthesis.
468 EXPERIMENTAL
catalyst preparation Three types of titania with different surface areas were used as supports. Both Ti0 with a surface area of llm2/g and Ti0 with a surface area of 2[C] 2/g 2[A] 237m \\'ere obtained from conmercial suppliers whilst Ti0 with a surface 2[B] area of 5en?/g was prepared by hydrolysis of titanium(IV) chloride. T1E supports were irrpregnated with equeous solutions of iron (III) nitrate and ruthenium (III) chloride by the incipient wetness technique and dried in air at 60°C (12h) and at 110°C (lh) before calcination in air at 480° (4h).
8arrples
for examination by 57Fe Mt;ssbauer spectroscopy were prepared with 50% of the total iron content being CCfTPOsed of isotopically pure iron-57. tion of the alumina-
T1E prepara-
and silica-supported catalysts has teen deseriJ::ed pre-
viously (8). 'I'errperature prograrnred reduction and ~ssbauer
spectroscnpy
'IPR rreasurerrent.s \\ere perfonred using a 95: 5 argon-hydrogen gas mixture (3Qnl;1ni.n) at a prograrnning rate of 16
0C;1ni.n.
~~ssbauer
spectra were recorded
at 298K in a quartz in situ cell following treatment in flowing hydrogen at 235, 450, 500 and 600°C as descriJ::ed previously (8) . quoted relative to rretallic iron. Fischer Tropsch reactions
All chemical isarer shifts are
Carron rronoxide hydrogenation reactions \\ere performed in a flow-type mrcroreactor of 6 rrm Ld. and 220 nm length using 5 ml of catalyst of 20-40 standard rresh.
All catalysts \\ere prereduced in flowing hydrogen (240 ml/g catalyst/h) 2 at 9Kg/an pressure for 4 bours , Ruthenium and iron-ruthenium catalysts \\'ere
prereduced at 235°C whilst iron catalysts \\ere treated at 450°C.
catalytic
reactions were perfotmed according to proceedures described previously (8) at 2 235°C and 25 Kg/an with a hydrogen to carbon rronoxide ratio of 2:1 and GiSV of 250/h. RESULTS AND mSaJSSION
The nature of iron and ruthenium in the catalysts The iron-57 M:lssbauer spectra recorded from the titania-, alumina-, and
silica-, sUJ;P:>rted catalysts containing 1% iron following calcination in air srowed quadrupole split abso:rptions characteristic of iron (III) in srre Ll, particle supe:rparamagnetic a-Fe
T1E changes in cationic oxidation state 20 3. which occured as a result of treabnent in hydrogen at different terrperatures are sUl\llBrised in Table 1The results recorded fran the titania-supported catalysts
s~
that,
despite its initial resistance to reduction at 235°C, the catalystcorrposed of 1% iron on low surface area TiOiA] was most amenable to CO\T!Plete reduction to rretallic iron at 600°C.
Given that facile reducibility of the iron oorrponent
of the catalyst is a reflection of weak rretal-support interaction (8), the
469 results recorded from the catalysts corrposed of 1% iron on titania suggest that the extent of rretal-support interaction is dependent on the surface area of the support with the strongest rretal-support interactions occuring on higher surfare area titania.
Indeed the results indicate that the extent of rretal-
support interaction follows the order Ti0
2[C]
> Ti0
2[B]
> Ti0
2[A].
TABLE 1 Iron-57 Mtlssbauer pararreters recorded at 298K from iron catalysts suworted on titania, alumina and silica. catalyst
Treatment
l%Fe/I'iO/AJ
H 2, 23S~C H2, SOO C H 600° C 2,
1%Fe/I'iO2 [BJ
23SoC H 2, H2, 500°C ° H 2, 600C
l%Fe/I'iO2 [c]
H 2, 235~C H2, 500 C H2, 600°C
1%Fe/Si0
2
H 235°C 2, H 600°C 2,
1%Fe/Al
20 3
H2, 235 ° C ° H2, 600C
r
ctO.OS /rnns-l
MO.O
Fe (III) Fe (III) Fe (II) a-Fe a-Fe
0.3S 0.36 1.21 0.01 0.07
0.87 0.91 1.71
Fe (III) Fe (II) Fe (II) a-Fe Fe (II) a-Fe Fe (III) Fe (II) a-Fe Fe (II) a-Fe
0.39 0.84 LOS 0.08 1.01 0.09 0.35 0.83
0.82 2.29 1.96
Oxidation state of Iron
/nms-
H±10
Spectral
/kG Contribution (%) 100 8 6 86 100
329 329
329
81 19 16 84 8 92 100 44 56 17 83
320
94 6 56 23 21
327 327 0.81 1.70 329
-0.02
0.99 0.00
2.10
Fe (III) Fe (II) Fe (III) Fe (II) a-Fe
0.30 0.85 0.30 0.85 -0.01
0.90 1.90 0.80 1.90
Fe (III) Fe (II) Fe (III) Fe (II)
0.30 0.90 0.30 0.90
0.70 2.00 0.70 2.00
97 3 68 32
The 57Fe Me>ssbauer spectra recorded from higher metal loaded titania sup-
ported catalysts fomed by calcination in air at 48JoC showed eight line patterns which could be interpreted in terms of the superposition of a sextet characteristic of large particle magnetically ordered a-Fep3 on a doublet representative of srrall particle superparamagnetic a-Fe The changes in 20 3• the oxidation state of iron in a catalyst corrposed of S% iron supported on low surface area titania[A] following treatment in hydrogen are sl.llTllE.rised in Figure 1.
The results show that 72% of the iron content in the freshly pre-
pared catalyst fomed in air was present as large particle a-Fe
203
(denoted
470
as a-Fe 20 3) which underwent facile reduction in hydrogen to rretallic iron. In contrast the small particle iron oxide (denoted as Fe (III)) underwent initial reduction to iron (II) which was only arrenabl.e to reduction to rretallic iron at elevated terrperatures.
The results su;mest that the interaction J:etween srnal.l,
particle superparamagnetic a-Fe large particle a-Fe
203
and Ti0 is stronger than that J:etween 203 2[A] and titania[A].
/i a-Fe
!
I
I
235
450 500
600
'c
Temperature of treatment in hydrogen Figure 1 Variation of the oxidation state of iron in 5%Fe)riO Z[A] following treatment in hydrogen The data recorded fran similar experiments using a 5% iron - 1% ruthenium
catalyst sugJOrted on Ti0
are represented diagrarrrnatically in Figure 2. 2[A] Whilst the reduction of the large particle a-Fe appeared to follow the sane 203 route as ~bserved in 5%Fe/ri0 in that it underwent direct reduction to 2[A] rretallic iron the reduction of the small particle a -Fep3 was different and, instead of fanning an iron (II) species, gave an iron - ruthenium alloy.
The
results irrply that ruthenium has the capacity to interact with the small particle a-Fe
203 ruthenium alloy.
and Induce the reduction of iron and the fonnation of an iron -
471
The behaviour of a ruth:mium-rich catalyst conposed of 1%Fe-5%Ru/ri0
under similar reducing conditions is illustrated in Figure 3.
[AJ
2 The significantly
larger proportion of small particle a-Fep3 in the catalyst prepared l::ly calcination in air presumably reflects t.be enhanced dispersion of the smaller iron concentration.
In principle it would J:::e reasonable to expect that such
an enbancerrent; of the dispersion vould cause an increase in the interaction J:::etw=en t.be iron carponent and the titania support and thereby inhibit reduction rot, on the other hand, the higher ruthenium content could J:::e envisaged as exerting a greater acceleration of the reduction of the iron (III) species. The results \<.hich are surrmarised in Figure 3 show that the 1:alance J:::etw=en t.be se tv.o effects leads roth to the formation of an intenrediate iron (II)
species and to the tranformation of a larger fraction of tre total iron content to t.be iron - rutheniun alloy.
%
~~~-------I
80 rIe
\
I 60
\
r
40
c
Ql
~
20
I
I
\
I
\\
~
I~
l
a-Fe!
~~
.
t
fu-----n
Fe-Ru Alloy
Fe(III)
L----J-_ 235
450 500
600 DC
Temperature of treatment in hydrogen Figure 2 Variation of the oxidation state of iron in 5%Fe-l%Ru/ri~ rAJ following treat:rrent in hydrogen VIe v.ould also mention bere one striking observation during our studies of
low Iretal loaded catalysts which will be the subject of a further report (9). It v..e.s found that a O.2%Fe - 1%Ru/Ti02 catalyst was reduced at 400°C in hydrogen
472
to iron (II) and an iron - n.rt:henium alloy but that continued treatment in the reducing atIrosphere at 500°C was acconpanied by partial conversion of the iron (II) to iron (III).
This phenomenon was
observed in low iron content catalysts (O.2%Fe) on all three titania supports l:ut not in titania supported catalysts with iron loadings exceeding 1% nor
on low iron loaded catalysts supported on silica or alumina.
%
g
80
.~
~
Z
.~
a-Fe 60
~
~
C
o
re-s« Alloy Fe(II) 285
450 500
600°C
Temperature of treatment in hydrogen Figure 3 Variation of the oxidation state of iron in l%Fe-5%Ru/I'iOiA] follCMing treatment in hYdrogen catalyst nay be carpared to that of 2[c] 1%Fe/Sio 2 and a 1%Fe/Al203 catalysts where all three supports have surface 2/g. areas in the range 2(X)-):)() m The M~ssbauer spectroscopy results collected The reducibility of the 1%Fe/Ti0
in Table 1 sbow that all three catalysts undergo similar changes when treated
in hydrogen at 235°C. However, following continued treatment at
exi'c
the extent
of reduction follows the order 1%Fe/Ti02[C] > 1%Fe/Si0 > 1%Fe/Alp3' The 2 results infer that, of the three SUg;lorts examined here, alumina interacts most strongly with the iron oooponent of the catalysts.
These results are endorsed
by the data frcm the examination of three iron - ruthenium catalysts supported on titania, silica and alunina by 'IPR (Figure 4 ).
The peaks in the ION'
473 temperature region of the 'IPR profiles (2<X)-2?DoC) represent the reduction of pure ruthenium (10) whilst the nedium terrperature peaks (250-280° C) ccrrespond
to the formation of iron - ruthenium alloys and the high temperature peaks (3?D-4?DoC) reflect the reduction of isolated iron species.
The results depic-
ted in Figure 4 therefore show that iron - ruthenium alloy formation is most readily achieved on the titania support, where no isolated iron species can be . di cab.ve . f a strongerI~tal-me:tal . de tecte d , an d are t he reby m 0 Interactaon '-u::t1l.een azon and ruthenium on titania than is achieved on either alumina or silica. possible that the
~er
It is
interaction bet1.'.een the titania and the iron carponent
of the catalyst is an irrportant feature of this stronger interaction between iron and rutheniun.
100
200
800
400
600
BOO~
Figure 4 Terrperature proqrerrmed reduction profiles recorded fran (a) 1%Fe-l%Ru/I'i02[A]; (b) 1%Fe-l%Ru/A1 (c) 1%Fe-l%Ru!Si0 203; 2 Catalytic activity and selectivity of the titania sUpported iron - ruthenium catalysts Carron monoxide hydrogenation reaction results over sane titania supported rutheniun and iron - ruthenium catalysts, together with sane data recorded fran
474 silica-
and alumina--
sUIPJrted catalysts, are sumarised in Figure 5.
The
results show that although the 1%Ru/I'i0 [ c] catalyst has a lov.er activity than 2 its alumina - supported counterpart its activity is significantly higher than that of a silica - supported catalyst with a similar metal loading. also srow that the incorporation of iron into the Ru/Al
The results
catalyst is aco:m-
203 panied by a decrease in catalytic activity for carl:x:m monoxide conversion whilst the titania-
and silica-
supported catalysts are less sensitive to the in-
clusion of iron. The results are consistent with our previous studies (8) which have derronstrated the significantly higrer activity of RU/Al with RU/Si0
catalysts, as corrpared 203 catalysts with similar metal loadings, and that the addition of
2 iron to the rutheniun carponent of the alumina - supported catalyst results
in a significant decrease in activity. activity of the Ru/Al
W= have also
SOO\\ll1
that the higher
catalyst may be associated with the strong interaction
203 t-etv.een tre metal and the support and that the incorporation of iron which
interacts strongly with ruthenium causes a v.eakening of the interaction between ruthenium and tre alumina support.
SUch strong ruthenium - support interactions
are not a feature of silica supported catalysts. The product distributions tabulated in Figure 5 show that the alumina-
supported catalysts tend to give higher rrolecular weight hydrocarJ:xms whilst silica supported catalysts generally give tre gaseous hydrocarbons ,
It is
interesting to record that the selectivity of tre titania supported catalysts tend towards an intenrediate position and favour the fonnation of hydrocarrons in tre gasoline range.
It is pertinent to rrention that the incorporation of
iron into the Ru/AlP3 catalyst which causes a 'Weakening of the Iretal - support interaction as a result of a strong iron - rutheniun interactico, also gives rise to a shift in the product di stribution towards shorter chain hydrocarl:x>ns (8).
Henre it appears that tre incorporation of controlled arroizrts of iron
into tre Ru/I'i0
system gives rise to a change in the nature of the 2[C] ruthenium - support interaction and that the ensueing strong iron - ruthenium
interaction leads to tre fonnation of catalysts containing iron - ruthenium alloys with a high selectivity towards the gasoline fraction.
For exarrple a
O.2%Fe-l%Ru/I'iOic] catalyst has a selectivity of 75% of the total hydrocarron yield to the C -C fraction. The incorporation of further conrentrations of 5 15 iron causes the selectivity to change towards the fonnation of gaseous hydrocarbons similar to the products obtained from silica - supported ruthenium catalysts wh3re the Iretal support interaction is 'Weak. cma..USICNS One special feature of the freshly prepared titania supported iron and
iron - ruthenium catalysts is the presence of significant quantities of large particle o,-Fe 203 which canJ::e reduced directly to rretallic iron independently
475 catalyst
Activity
Product distribution
lTg/min. gRu
Carbon number
234 5 6-10 142.5
1 83.9
11-15
~~.
16+
~,,:.;-.:..:.-~
23
5
4
6-10
11-15
16+
~~=~ I 11-15
[
38.0
1 72.8
23
69.0
6-10
11-15
5
16+
!
~=~~~ 1 23 4
0.2%Fe-l%Ru/I'i02[C]
5
4
6-10
11-15 16+
~~_~~~~~~~~~ 1
2 3
4
5
6-10
11-15 16+
67.8
1
23
6-10
11-15 16+
27.2~E~I~
1%RU/Si0 2
1 0.2%Fe-l%RU/Si0 2
2
2 3
5
4
6-10 11-1516+
27.5
1 1%Fe-l%Ru!Si0
5
4
2 3
5
4
6-10
11-15 16+
28.5~4D 50
100%
Figure 5 catalytic activity and selectivity of SCIre iron - ruth:mium catalysts supported on TiO [C~,A1203' and Si02• 2 Reaction conditi&: Terrperature, 235°C; Pressure, 25Kg/an , 2/1; GHSV 250 h- l H 2/CO, of any iron - ruthenium alloy forrration.
Tre absence of this phenorrenon in
analogous alumina - supported catalysts is probably a reflection of the difference in t.be strengths of iron - titania and iron - alumina interactions. Our previous evidence for a strong interaction 1::Etv.een iron and alumina (8)
may reflect the existence of a spinel-type structure at the interface between iron and alunina which inhibits the reduction of the iron - containing species.
476 TI-e results presented here slow that the interaction between iron and titania
is rather weak, indeed, the formation of a spinel-type structure in the iron titania system is m.likely.
Hence the titania - supported large particle
a-Fep3 may be envisaged as m.dergoing direct reduction to rretallic iron even in the presence of ruthenium.
SUch a process does not appear to l::e allowed for
the alunina - supported com.terpart and this may be the rrost substantial difference l:etw=en the Fe-Ru;'ri02 and Fe-RU/Al 0 systems. 2 3 These principles may l:e inportant in explaining the unusual, oxidation of titania - supported iron (II), fonned by low terrperature hydrogen reduction, to iron (III) at higher terrperatures in reducing atlmspheres.
I t is possible that
at terrperatures of ca.500-600° ruthenium might exist in the so called SMSI state in the titania - supported catalysts and, given that the iron carponent interacts weakly with the titania support, any iron (II) species associated with the ruthenium atoms may transfer electrons to ruthenium as the latter forms alloys with the isolated iron particles.
In this connection we would also point
out that the Lncorporat.Ion of 0.2% iron into the ruthenium catalysts produced a rrore marked effect on the activity and selectivity of the titania supported catalyst than that supported on silica where the ruthenium interaction with the support; is weak.
This feature of l:ehaviour suggests that the iron species in
the titania - supported catalysts exist in a highly active state as a result of their close connection and interaction with the ruthenium corrponent., ACKNOwr..ED:;EME1\lTS
v.e acknowledge the support of The Chinese Academy of Sciences and The Royal Society. v.e thank Tioxide p.l.c for a gift of titanium dioxide. REFERENCES
1 2 3 4
5 6 7 8 9 10
L. Guczi, K. Mattusek, I. Manninger, J. Kiraly and M. Eszterle, in:
B. Delman, P. Grange, P. JaCDbs and G. Pon::elot (Ed.), Preparation of catalysts II, Elsevier, Amsterdam, 1979, pp. 391-400. 1. Dezsi, D.L. Nagy, M. Eszterle and L. Guczi, J. Physique (Paris) C 2, 40 (1979) 76. L.M. Tau, S. :eorcar, D. Bianchi and c.o. Bennett, J. catal., 87 (1984) 36. S.J. Tauster, s.c. Fm.g and R.L. Garten, J. Amer. Chern. Soc., 100 (1978) 170 J. santos, J. Phillips and J.A. Dumesic, J. catal., 81 (1983) 147. L. Guczi, Z. SChay and I.Bogyay, in: B. Delrron, P. Grange, P.A. Jacobs and G. Poncelet (Ed.), Preparation of catalysts III, Elsevier, Amsterdam 1983, pp. 451A.M. Van der Kraan, R.C.H. NOOnekens and J.W. Niemantsverdriet, in: Abstracts of Invited and Contributed Papers of International Cbnference on The Applications of the ~ssbauer Effect, Ieuven, Belgium, pp. 4.14. F.J. Berry, Lin Liwu, W9.ng Chengyu, Tang Renyuan, Zhang Su and Liang Ibngbai, J. Chern. Soc. Faraday Trans. 1. 81 (1985) 2293. F.J. Berry, Lin Liwu, Wang Chengyu, Du Hongzhang, Tang Renyuan, Zhang SU and Liang IbngJ:ai, Unpublished results. Tang Renyuan, Liang D:ngbai and Lin Liwu, Unpublished results.
477
DISCUSSION R.I. BICKLEY: I note that the calcination (480°C) of your specimens corresponds to conditions that are intermediate between those used by van Ommen et al. of 450 DC and those used by ourselves, Bickley et al., where a minimum calcination temperature of 500°C has been used. van Ommen has suggested that the iron remains exclusively at the surface of the impregnated particles whereas we infer that at 500 DC and greater a good bulk dispersion of iron has been achieved. Do you think that you have any evidence for the presence of bulk iron species within the Ti02 particles at your calcination temperature? F.J. BERRY: We have found no evidence for the presence of iron species within titania in the materials which we have examined. L. GUCZI : IS for a-iron is 0.0 mm/s in the Mossbauer spectra, whereas IS for FeRu alloy changes in the range between -0.2 to +0.06 mm/s. On what assignment could you separate a-Fe and FeRu from the spectral area? F.J. BERRY: The a-iron identified by us was magnetically ordered and was not superparamagnetic. Hence, the chemical isomer shifts for a-iron and the ironruthenium alloy (which gave singlet component to the spectrum) could readily be determined. J.G. van OMMEN: Can you explain why small Fe203 crystals on Ti02f in contrast with large Fe203 crystals cannot be reduced further than the Fe I state, while this reduction from Fell to Fe is possible in the case Ru is present. Or can you give a suggestion about what you think is happening? In our paper, we suggested in case of a monolayer of Fe203 that the formation (a reduction with H2) of a surface ilmenite prevents the iron to be reduced further than Fell state on Ti02. F.J. BERRY: Small iron(III) oxide crystals supported on titania are reduced to iron(II) which interacts strongly with the support and is not amenable to further reduction. In large particles of Fe203 some on the iron content does not interact strongly with the support and can therefore be reduced to metallic iron. Ruthenium renders small particle iron(lll) oxide more susceptible to reduction to iron(O) by allowing the formation of the iron-ruthenium alloy. This implies an intimate relationship between iron and ruthenium particles during the thermal reduction in hydrogen.
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479
B. Delmon, P. Grange. P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
PREPARATION OF CARBON-SUPPORTED K-Fe-Mn AND Fe-Mn CATALYSTS USING CARBONYL CLUSTERS J. VENTER1, M. KAMINSKy2, G. L. GEOFFROy2, and M. A. VANNICE' 1Department of Chemical Engineering, The Pennsylvania State University, University Park, PA 16802 (USA) 2Department of Chemistry, The Pennsylvania State University, University Park, PA 16802 (USA) SUMMARY Highly dispersed Fe-Mn bimetallic particles have been formed on a high surface area amorphous carbon black support using Fe-Mn and K-Fe-Mn carbonyl clusters containing different Fe/Mn ratios. These catalysts have been characterized by hydrogen adsorption and CO chemisorption at 195K and 300K, and by their kinetic behavior for CO hydrogenation at 1 atm. The K-promoted KMnFe and KMnFe 2 catalysts gave 85-90 wt% ethylene, propylene, and butene with CH. as the only other detectable hydrocarbon product. Comparison with other Mn-Fe catalysts reported in the literature shows that these catalysts not only have a much higher selectivity to light olefins but also have higher activity per gram iron. INTRODUCTION Various reports have noted that introduction of Mn into Fischer-Tropsch catalysts increases their selectivity to olefinic products (1,2).
A promising
method for obtaining supported Mn/Fe catalysts is the use of organometallic mixed-metal carbonyl clusters as the metal precursor. these clusters has a number of advantages such as:
The utilization of
(1) uniformly supported
clusters can be obtained with stoichiometric metal ratios, namely Mn 2 , Mn 2Fe, MnFe, MnFe 2 ; (2) promoters can be added to the MnFe, MnFe 2 , and Fe. clusters by preparing the potassium salts of the carbonyl clusters, Which also ensures initial intimate metal contact; (3) zero-valent metals can initially be obtained on appropriate supports by decarbonylation of the clusters at mild temperatures, which can also enhance dispersion. High surface area carbon has certain properties that make it an attractive support for Fischer-Tropsch catalysts.
The high surface area provided by the
small pores allows the distribution of the carbonyl clusters and stabilizes highly dispersed particles, presumably due to the presence of physical barriers which suppress surface migration and the resultant agglomeration. The carbon can be made virtually sulfur-free by a high temperature reduction in H2 which, in addition, also eliminates surface hydroxyl and carboxyl groups (3) •
In this study the preparation, characterization and catalytic behavior of a family of carbon-supported Fe, Mn, Fe/Mn and K/Fe/Mn cluster catalysts are described, with an emphasis on those catalysts which gave unusually high selectivities to light-weight olefins.
480
EXPERIMENTAL The amorphous carbon black used as a support in this study was CSX-203 from Cabot Corporation with a sulfur content of 0.13% after treating in H. at 1223K for 12 h and a BET surface area of 1400 m2 / g .
Before impregnation the carbon
was heated to 573K under dynamic vacuum of 10- 4 kPa for 8 h to remove any physisorbed water. The compounds Mn.(CO),O and Fe,(CO),. were purchased from Alfa-Ventron and the clusters Mn.Fe(CO)" dures (4-6).
and K[HFe,(CO)l,] were prepared by literature proce-
The remaining clusters were prepared by modifying established
routes to known salts to give the desired NEt,+ and K+ salts (5). K[Fe.Mn(CO),,]--K[Mn(CO)s] was prepared by reacting Mn.(CO),o with Na/K alloy in THF solution.
This was added to a refluxing THF solution of Fe.(CO).
under N. and the solution was refluxed for 12 h.
Subsequent removal of THF
was followed by addition of CH.CI. and then Et.O to dissolve all the residue. Addition of n-hexane and cooling to ooC for 12 h gave microcrystalline K[Fe.Mn(CO),,], identified by its IR spectrum. NEt,[Fe.Mn(CO),,]--Na[Mn(CO)s], prepared from the reaction of Mn.(CO),O with Na/Hg, was added to Fe.(CO). in THF followed by an 8 h reflux. Subsequent addition of NEt,Br in EtOH to this solution followed by cooling to ooC for 12 h gave microcrystalline NEt,[Fe.Mn(CO),s]' K[FeMn(CO).]--This cluster was prepared by irradiating a THF solution of Fe(CO). and K[Mn(CO)s] with a pyrex-filtered 450W Hanovia medium pressure Hg vapor lamp for 3 h with precautions taken to ensure that the solution did not warm above 35°C.
Removal of solvent gave microcrystalline K[FeMn(CO).] which
was washed with CH.CI. to remove residual Fe(CO)s' NEt,[FeMn(CO).]--This cluster was prepared from the above sample of K[FeMn(CO).] by metathesis with NEt,Br using a procedure similar to that for NEt,[Fe.Mn(CO),.]. The clusters were stored in a glove box and supported on the carbon by an incipient wetness impregnation technique under N. using standard Schlenck techniques and dry and degassed THF as solvent.
The Mn.(CO),o + Fes(CO),.
catalyst with Fe/Mn = 0.5 was prepared by sequential impregnation of the two clusters, using Mn.(CO),o first.
A coimpregnated catalyst was also prepared
using the same two carbonyls in a single solution with a metal ratio of Fe/Mn =
2.
Metals analyses for these catalysts are given in Table 1.
son, a more conventional catalyst was prepared from Fe(N0 3 )
.
For compari-
and Mn(N0 3 )
.
using an aqueous impregnation technique (3). A low temperature reduction (LTR) pretreatment consisted of heating to 473K under flowing He (40 cc/min) and reducing in H. at this temperature for 2-5 h until no solvent could be detected in the gas stream by gas chromatography.
A
481 TABLE 1 Percentage metal loadings of carbon-supported oatalysts \Imole Metal g Catalyst
wU Metal g Catalyst
Carbonyl Cluster
(Mn/Fe) Ratio
Mn
Fe
Total
Mn
Fe
Total
Hn2(CO)10
2.56
---
2.56
1166
0
466
High
Mn2Fe(CO),11
1.35
0.611
1.98
245
114
359
2.15
NEtll[MnFe(CO)9]
1.54
1.53
3.08
281
275
556
1.02
NEt4[MnFe2(CO)'3]
1.611
3.13
4.78
299
561
860
0.53
Fe3(CO),2
---
3.78
3.78
0
678
678
0
K[MnFe(CO)9]
1.33
1.16
2.50
236
203
439
1.16
K[MnFe2(CO)'3]
0.98
2.03
3.01
178
364
542
0.49
K[HFe3(CO)11 ]
---
11.38
11.38
--
786
786
0
2Mn/Fe; Carbonyls
3.06
1.52
4.59
557
273
830
2.04
2Fe/Mn; Nitrates
1. 37 2.09
3.116
250
377
627
0.66
2Fe/Mn; Carbonyls
1.25
3.75
228
1148
676
0.50
2.50
TABLE 2 Representation chemisorption measurements on fresh oatalysts
H2 0-200
CO 300K
(Fe + Hn) HTR CO H2 0-400 300K
Mn2(CO)1O
0.07
0.002
0.06
0.002
---
---
---
---
Hn2Fe(CO),4
0.008
0.006
0.011
0.28
0.03
0.02
0.13
0.90
NEt4[MnFe(CO)9]
0.02
0.02
0.06
0.30
0.04
0.04
0.11
0.60
NEt4[MnFe2(CO) 13]
0.05
0.02
0.08
0.27
0.08
0.02
0.13
0.42
Fe3(CO)12
0.07
0.53
0.09
0.90
0.01
0.53
0.09
0.90
K[MnFe(CO)9]
0.04
0.02
0.14
0.10
0.09
0.04
0.30
0.22
K[MnFe2(CO)13]
0.05
0.05
0.13
0.27
0.08
0.08
0.20
0.40
K[HFe3(CO),1 ]
0.04
0.30
0.06
0.41
0.04
0.30
0.06
0.41
2HnlFe; Carbonyls
0.03
0.02
0.05
0.09
0.10
0.06
0.18
0.26
0.12
0.69
---
---
0.20
1.15
0.07
0.35
0.09
0.50
0.10
0.53
~mole
2Fe/Mn; Nitrates 2Fe/Mn; Carbonyls
gas/~mole
LTR
Carbonyl Clusters
--- --0.06 ---
pmole gas/pmole Fe LTR HTR CO CO H2 H2 0-200 300K 0-400 300K
482
high temperature reduction (HTR) procedure consisted of heating to 673K under flowing H2 (40 cc/min) and reducing at 673K in H2 for 16 h. For chemisorption experiments on the fresh catalysts, the LTR pretreatment consisted of heating under dynamic vacuum (10- 4 kPa) to 473K and reducing at this temperature in H2 for 2 h.
The catalyst was evacuated at the reduction temperature for 1 h
prior to the CO uptakes. The H2 desorption technique was essentially that used by Amelse et al.(7). The sample was cooled under flowing H2 from the reduction temperature of either 473 or 673K to 273K, evacuated for 10 minutes, and then heated to the original reduction temperature in a closed, known volume.
The amount desorbed
was determined from the pressure, which was measured after both a 20-min and a 60-min period.
A dual isotherm method was used for the CO chemisorption at
both 195K and 300K, with irreversible CO being determined at a pressure of 200 Torr (26.7 kPa). All kinetic data were obtained at 1 atm (101 kPa) and a H2/CO ratio of 3 under differential reaction conditions in a glass plug-flow reactor. A PerkinElmer Sigma 3 gas chromatograph fitted with Chromosorb 102 columns and subambient temperature programming interfaced with a Hewlett-Packard integrator was used to obtain gas analyses.
Addition details have been given elsewhere (2).
The conversion of CO to hydrocarbons (HCs) was kept below 7%. RESULTS AND DISCUSSION Chemisorption These results are summarized in Tables 2 and 3.
An important pattern in
the data (Table 2) is the drastic increase in chemisorption values when using an HTR step rather than an LTR treatment.
Considering the CO uptake values at
300K, it can be seen that they increase by factors of 5 to 50 except for the catalysts containing no Mn--Fe.(CO) '2 and K[HFe.(CO) ,,]--which showed increases by factors of only 1.7 and 1.4, respectively.
When these chemisorp-
tion values are normalized to the amount of iron present, the variation among the apparent dispersion of the fresh samples after HTR is between 0.42 and 0.90 for the non-K-containing catalysts, and all samples after HTR are very well dispersed.
After an LTR step, however, the variation in CO (300K)/Fetotal
is between 0.02 and 0.53 with all the catalysts except Fe.(CO)'2 and K[HFe.(CO),,] having values less than 0.08, and the values are now approximately an order of magnitude lower than that obtained for the Fe.(CO)'2 catalyst. This indicates a decreased chemisorption capacity after the LTR step due either to an incomplete decarbonylation of the mixed clusters or to a significant Mn-Fe interaction which is overcome to a large extent after the HTR treatment.
The latter possibility implies metal segregation during the HTR treat-
483 TABLE 3 Chemisorption measurements after HTR on used catalysts (after kinetic studies) ~mole
Fe
gas/~mole
Carbonyl Clusters
COad
@
300
H2
COa 195K
CO 300K
Mn2(CO) 1Ob
0.07
0.03
0.00
0.0
Mn2Fe(CO)14
0.14
0.14
0.32
2.3
NEt4[MnFe(CO)9]
0.14
0.13
0.26
2.0
NEt4[MnFe2(CO)13]
0.08
0.08
0.16
2.0
Fe3(CO)12
0.08
0.07
0.10
1.5
K[MnFe(CO)9]
0.17
0.19
0.15
0.8
K[MnFe2(CO) 13]
0.12
0.15
0.12
0.8
K[HFe3(CO)11]
0.05
---
0.04
---
2Mn/Fe; Carbonyls
0.09
0.15
0.31
2.1
2Fe/Mn, Nitrates
0.11
0.15
0.97
6.4
2Fe/Mn; Carbonyls
0.09
0.20
0.35
1.8
COad @ 195
a Evacuated 60 minutes between isotherms. bBased on Mn-content; ~mole/~mole Mn. ment, in agreement with the work of Jensen and Massoth who suggested formation of Mn-oxide platelets covering reduced iron particles, thus separate phases
(8,9), where the Fe phase was formed by high-temperature treatments (9). Catalytic Behavior Activities for the catalysts are given in Table 4.
The turnover frequen-
cies (TOFs) were higher after LTR than after HTR for all Mn-containing catalysts although activities(per g) were lower after the LTR step.
This is a
consequence of the lower chemisorption uptakes obtained after a LTR pretreatment; however, it clearly demonstrates that active catalysts can be obtained without the use of an HTR step typically required for iron.
The activity of
the sequentially impregnated Fe,(CO)12/Mn2(CO)lO (2Mn/Fe) sample was substantially higher than that of the Mn 2Fe(CO)1. carbonyl cluster and nearer to that of the Fe-only sample, indicating that a greater extent of interaction between the Fe and Mn is achieved when stoichiometric clusters are used as precursors. Addition of Mn to the Fe clusters (Without K) diminished the activity, even when normalized to the iron present.
The addition of K to the MnFe and MnFe 2 clusters enhanced specific activity and TOF values for the former cluster, but
had less effect on the properties of the latter.
484 TABLE 4 Catalytic activities of carbon-supported clusters for CO hydrogenation after LTR and HTR: T s 548K. 1 atm. H./CO = 3.
Ini tial Clusters
(a) CO Conversion (%)
(a) Activity (HmO\CO) g ca ·s
LTR:
Activity (Hmoles Product 3) I1moies Fe.s x 10 Total
CO2
HC
(b) TOF (s-l xl03)
CH4
HC
CH4
Mn2(CO)10
0
0
0.004
0.004
--
--
0
0
Mn2Fe(CO)14
0.06
0.005
0.077
0.036
0.042
0.042
2.4
0.2
NEt4[MnFe(CO)9]
0.05
0.004
0.033
0.033
0.013
0.003
0.3
0.1
NEt4[MnFe2(CO)13]
1.30
0.14
0.51
0.27
0.24
0.059
10.8
2.7
Fe3(CO)12
2.7
0.58
1.61
0.76
0.86
0.16
1.6
0.3
K[MnFe(CO)9]
1.3
0.04
0.70
0.52
0.18
0.06
4.7
1.6
K[MnFe2(CO )13]
0.8
0.08
1.79
1.29
0.50
0.07
6.5
1.0
K[HFe3(CO)II]
2.2
0.34
1.55
1.11
0.44
0.05
1.6
0.2
2Mn/Fe; Carbonyls
0.8
0.04
0.32
0.18
0.15
0.05
2.7
0.9
2Fe/Mnl Nitrates
--
--
--
--
--
--
--
2Fe/Mn; Carbonyls
0.8
0.10
0.40
0.18
0.22
0.10
Mn2(CO)10
0
0
0.006
0.006
--
--
Mn2Fe(CO)14
0.5
0.007
0.12
0.06
0.06
0.02
0.07 0.02
NEt4[MnFe(CO)9]
0.4
0.05
0.55
0.38
0.17
0.03
0·3
NEt4[MnFe2(CO)13]
2.4
0.52
1,66
0.75
0.91
0.20
2.2
0.5
Fe3(CO)12
2.0
0.80
2.24
1.06
1.17
0.23
1.3
0.3
0.08
1,46
1.06
0.41
0.04
1.9
0.2
1.68
-
--
1.5 00.7 c
HTR: 0
0
0.05
K[MnFe(CO)9]
1.4
K[MnFe2(CO)13]
1.7
0.29
2.46
0.78
0.10
2.0
0.2
K[HFe3(CO)II]
6.1
0.60
1.90
1,15
0.77
0.12
1.9
0,3
2Mn/Fe; Carbonyls
1.8
0.90
0.69
0.36
0.34
0.09
1.3
0.3
2FelMn; Nitrates
1.9
0.20
1.08
0.56
0.52
0.13
0.5
0.1
2Fe/Mn; Carbonyls
4.1
0.22
0.98
0.49
0.50
0.10
0.9
0.2
:
(a) CO converted to hydrocarbons (HCs) only (~» Based on CO adsorption at 300K on fresh samples ( Based on CO (195K); CO/Fes • 1:2
A typical set of selectivity results at low CO conversion is reported in Table 5.
It is clear that the average molecular weight of the products
decreases as the Mn content increases, indicating that the chain growth probability decreases. Longer chain products are favored after HTR over all the Mn-containing catalysts indicating that the Fe-Mn catalysts become more iron-like after an HTR step, which is again consistent with the proposal that phase separation occurs after an HTR step.
The olefin-paraffin ratio (aPR)
485 TABLE 5 Selectivity of CO hydrogenation after LTR and HTR pretreatments of P = 10lkPa, H2/CO - 3.
carbon-supported FeMn and KFeMn catalysts. Temp (·C)
Catalyst LTR:
CO Conv. to HC ($)
Selectivity (HC Mole $) C, \a
J
C2
C2
C
C3
C4
C4
+
C5
l:
C2.3
~
l:
C2.3
(CO-HC)
---
High
1.6
1.8
300
0.06
100
-
-
-
306
0.21
37
29
24
10
NEt4[MnFe2(CO)'3]
275
1,3
43
26
7
19
-
Fe3(CO),2
249
2.7
40
11
12
15
8
K[MnFe(CO)9]
31
13 23
-
14
-
High
2.9
20
-
9
6
High
2.3
-
3.5
1.4
--
--
13
-
3.9
0.9
-
-
- - - - -
--
High
1.0
0.9
Mn2(CO),0
327
Mn2Fe(CO),4 NEt4[MnFe(CO )9]
0
302
1.3
38
K[MnFe2(CO)13]
250
0.8
31
33
K[HFe3 (CO)l1]
280
2.2
37
28
-
2Mn/Fe; Carbonyls
300
0.8
50
22
11
17
2Fe/Mn; Nitrates
-
-
-
-
-
-
2Fe/Mn; Carbonyls
284
1,0
61
18
8
- - - 5 - 5
-
6.4
0.9
9
1.3
·0.7
-
6
High
2.6
12
- - - -
0.7
!!!R: Mn2(CO),0
348
0
-
-
-
Mn2Fe(CO)14
314
0.52
50
17.
19
9
6
NEt4[MnFe(CO)9]
309
1,3
38
22
15
17
9
NEt4[MnFe2(CO)'3]
251
2.4
48
10
16
12
6
Fe3(CO),2
225
2.0
42
11
12
14
6
2
5
K[MnFe(CO)9]
290
1.4
25
35
30
-
11
K[MnFe2(CO)'3]
250
1.7
28
29
24
-
11
8
27
3
20
-
-
15
16
10
5
K[HFe3 (CO) 11]
285
1.9
41
2Mn/Fe; Carbonyls
300
1.8
47
-
- - 7
8
2Fe/Mn; Nitrates
279
2.2
43
10
18
10
7
-
2Fe/Mn; Carbonyls
260
2.2
46
9
19
9
7
-
-
1.6
1.8
1
1.0
0.8
8
1.4
0.5
-
High
2.4
High
2.4
3
15.7
1.7
4
3
1.2
1.2
7
4
0.8
1.0
6
5
0.7
1.0
(al A "-" indicates that no such products was detected.
for the C2 and C, products after the LTR step shows that very high selectivities can be obtained, even for certain unpromoted catalysts, especially that prepared from NEt,[Fe 2Mn(CO)l']'
The K-promoted clusters retain this high
selectivity after an HTR treatment. The most significant result of this study is the preparation of highly dispersed Fe-Mn and K-Fe-Mn catalysts which have selectivities to Cg-C. olefins as high as 85-90 wt% with CH. being the only other detectable hydrocarbon product.
For example, KFe 2Mn/C after LTR and KFeMn/C after HTR gave 11-14
wt% methane, 30-31 wt% ethylene, 31-39% wt% propylene, and 20-25 wt% butene.
486 These high olefin yields remained constant during long periods (20-40 h) on-stream (2).
Although certain Fe-Mn/C catalysts are quite active and have
high OPRs after an LTR pretreatment, in particular the Fe 2Mn catalyst, the selectivity of these olefin-producing catalysts can be significantly increased by promoting the catalysts with K by using it as a counter ion in the carbonyl clusters, and in certain cases an activity enhancement is achieved in addition.
1.2 >-.
.Iol
..-I
0.8
~
.Iol
s
0.4
o 1.0
o Mn 2Fe
MnFe 2
MnFe
KMnFe
Fe3
KMnFe
KFe 2
3 2Mn/Fe
(10)
(12) (11)
(14) (13)
(16) (15)
Fig. 1. Comparison of specific activity (mole CO reacted to hydrocarbons per mole Fe per second x 10- 3) and selectivity to olefins (fraction of olefins in C2 + C, hydrocarbons) of Fe-Mn catalysts with previous results. The numbers represent the appropriate references. The reaction conditions in each study were as follows: (10) - 1 atm, 280°C, Ha/CO = 1; (11) - 5 atm, 300°C, Ha/CO 1.5;(12) - 2 atm, 290°C, Ha/CO = 0.6; (13) - 1 atm, 240°C, Ha/CO = 1; (14) - 11 atm, 280°C, Ha/CO = 0.8; (15) - 10 atm, 250°C, Ha/CO z 1.5; (16) - 1 atm, 250°C, Ha/CO = 2; This study - 1 atm, 275°C, Ha/eO = 3.
487
These enhancements in activity and selectivity using potassium carry with them the price of increased CO. production; however, improvement may be obtained by operation at higher pressures--this has not yet been studied.
The optimum
choice between the less active and less selective unpromoted catalyst and the K-promoted catalysts is not known at this time. The catalytic behavior of the Fe-Mn catalysts described here can be compared to that reported by others.
Figure 1 presents a comparison of the
best catalysts from each of the referenced studies on a specific activity (mole CO-to-hydrocarbon products per second per mole iron) basis.
On this
basis it is clear that not only are the selectivities to olefins of our Fe-Mn/C catalysts comparable to or better than those reported by other groups, even though our H./CO ratio is much higher, but the average chain length is much shorter.
Also, the activities are much higher even though our operating
pressure was only 1 atmosphere. ACKNOWLEDGMENTS This study was funded by NSF Grant No. CPE-8205937. acknowledges SASOL for partial support.
One of us (JJV)
Some of the results contained in the
Tables and Figure 1 have been previously submitted for publication in the Journal of Catalysis.
REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15.
H. Kolbel, K. D. Tillmetz, Belgian Patent 837628. J. Venter, M. Kaminsky, G. L. Geoffroy, M. A. Vannice, J. Catal., submitted for publication, and references cited therein. H.-J. Jung, P. L. Walker, Jr. and M. A. Vannice, J. Catal. 75 (1982) 416. E. H. Schubert and R. K. Scheline, Z. Natureforschung, 20B (1965) 1306. J. K. Ruff, Inorg. Chem., 7 (1968) 1818. H. A. Hodali, C. Arcus and D. F. Schriver, Inorg. Synth., 20 (1980) 218. J. A. Amelse, L. H. Schwartz and J. B. Butt, J. Catal. 72 (1981) 95. K. B. Jensen and F. E. Massoth, J. Catal., 92 (1985) 98; J. Catal., 92 (1985) 109. J. M. Stencel, J. R. Diehl, R. A. Anderson, M. F. Zarochak, and M. F. Pennline, to be published. (a) J. Barrault, C. Forquy and V. Perri chon , Appl. Catal., 5 (1983) 119. (b) J. Barrault and C. Renard, Appl. Catal., 14 (1985) 133. L. Bruce, G. Hope and T. W. Turney, React. Kinet. Catal. Lett., 20 (1982) 175. W. D. Deckwer, H. J. Lehmann, M. Ralek and B. Schmidt, Chern. Ing. Techn., 53 (1981) 818. W. L. van Dijk, J. W. Niemantsverdriet, A. M. van der Kraan and H. S. van der Baan, Appl. Catal., 2 (1982) 273. (a) C. N. Satterfield and H. G. Stenger, I&EC Proc. Des. Dev., 23 (1984) 26. (b) c. N. Satterfield and G. A. HUff, J. Catal., 85 (1984) 370. (a) H. SchulZ and H. Gokcebay, 9th Conf. on Catal. of Org. Reac., Charleston, South Carolina, USA, April 1982. (b) H. SchulZ, 5th IntI. Conf. on Het. Catal., Varna, Bulgaria, October
4M
16.
3-4, 1983. (c) W. Benecke, R. Schulz, H. G. Feller and M. Ralek, Proc. 8th Intl. Congo on Catal., Vol. 4, Dechema, Frankfurt, Germany, 1985, p. 219. V. L. Kuznetsov, N. F. Danilyuk, I. E. KolosQva and Y. I. Yermakov, React. Kinet. Catal. Lett., 21 (1982) 249.
489
DISCUSSION M.V. TWIGG: 1/ Is it necessary to use "cluster" compounds in order to make catalysts having the specific properties you describe? 2/ In particular, what is the effect of using monomeric species such as Fe(CO)5 and HMn(CO)5 ? 3/ Could the anhydrous preparative conditions themselves be important? M.A. VANNICE: Although we have not used the two precursors you mention, we have found that optimum selectivity to olefins is achieved using stoichiometric clusters rather than mixtures of single-metal clusters giving the same Fe/Mn ratios. This can be seen in Table 5 by comparing the mixed-metal clusters to those with similar Fe/Mn ratios but prepared from mixtures of Fe3(CO)12 and Mn2(CO)10 or the nitrate salts of Fe and Mn (designated carbonyls and nitrates, respectively). Coprecipitation techniques have been used to prepare bulk Fe-Mn catalysts with high olefin selectivity; however, reported chain growth is much greater and activities per g iron are lower. The anhydrous preparative conditions are very important because these clusters are very sensitive to oxygen and water, particularly those containing Mn, and can readily oxidise to the individual oxide phases. V. PERRICHON : 1/ My question is relative to the state of manganese during the reaction. Since you start from zero-valent manganese complexes and you reduce the catalyst under H2' how can you involve Fe2Mn04 as the essential species for the activity? Can you comment on the state and the role of Mn ? 2/ Do you think that the iron active species are in a metallic state, or in a non-completely reduced state? M.A. VANNICE : 1/ Our hypothesis is that a mixed spinel with a stoichiometry very near that of Fe2Mn04 is responsible for the high selectivity to light olefins (J. Venter, M. Kaminsky, G.L. Geoffroy and M.A. Vannice, J. Catal., in press). This is based on Mossbauer studies in the literature Which have associated the presence of the spinel with the production of olefins (N.K. Jaggi, L.H. Schwartz, J.B. Butt, H. Papp and M. Baerns, Appl. Catal. 13 (1985) 347). The Mn will definitely exist in an oxidized state under reaction conditions, as discussed in our paper referenced above, and its precise role is not yet fully understood. 2/ We presume that most of the iron is in an incompletely reduced state. Mossbauer studies directed toward this question are now underway. J. BARRAULT : In the presented results, the selectivity is compared at very low CO conversion. What is the evolution of selectivity with CO transforma-" tion? It could be an important point in the comparison of your results with catalytic results reported in the literature. M.A. VANNICE: Based on previous studies, one might expect that the olefin/ paraffin ratio would decrease as CO conversions increase from below 1% to higher values; however, we have observed little change in selectivity (both carbon number and olefin/paraffin ratio) up to conversions of CO to hydrocarbons of 4%. These low conversions were purposely obtained to eliminate heat and mass transfer effects from the kinetic data. Further studies are planned in which much higher CO conversions will be achieved. D.E. RESASCO : Have you done any study about the energetics of adsorption of CO or H2 on these catalysts by either TPD or calorimetric measurements ? M.A. VANNICE : Not yet. We are planning to determine integral heats of adsorption on these samples using our modified DSC system. E. NEWSON : 1/ What was the ash content of your carbon since many commercial carbons have wt% ash contents and the major component is frequently iron?
490
2/ Have you done any deactivation experiments with S in the feed since Fe is so sensitive to S ? M.A. VANNICE: The ash content of this CSX 203 carbon is about 0.7% and a typical analysis ha5 been reported (M. Kaminsky, K.J. Yoon, G.L. Geoffroy and M.A. Vannice, J. Catal. 91 (1985) 338). Calcium is the major impurity (-0.2%) while iron content is very low at about 22 ppm. We have conducted no poisoning experiments with S and, in fact, we pretreat our carbon in HZ at 950°C prior to impregnation to remove sulfur impurities which can indeed pOlson the iron. L. GUCZI : The OH group concentration is negligible on carbon surfaces. Depositing transition metal carbonyl clusters over A1203 or Si02 there is an immediate decomposition of the cluster framework. It seems to me that here on a carbon support this is not the case; therefore, one can expect a CO chemisorption with a stoichiometry over 1 CO/M due to the lesser affected cluster framework. Did you find any such effect? It may point to the integrity of the cluster which is already indicated by the very high activity even at low reaction temperature. M.A. VANNICE : We have not yet observed COad/Fe ratios greater than unity for the clusters; however, although the clusters appear to be completely decarbonylated at 473 K in H2' we do not yet know the extent of agglomeration to form larger particles during this period. Although quite high dispersions are indicated in Table 2, the corresponding Fe crystallite sizes range from about 1 to 4 nm, indicating some agglomeration of clusters occurs within the pore structure of the carbon. We also have some preliminary results from Mossbauer spectroscopy and DRIFTS that Fe3(CO)12 can decompose to Fe(CO)5 at low temperatures on this carbon. J.W. GEUS : Carbon is a very suitable support for a promoted Fischer-Tropsch catalyst since it does not react with potassium ions as alumi2a and silica. However, a carbon support exhibiting a surface area of 1400 m /g is very liable to display the complication mentioned by Satterfield, viz., that the higher molecular weight products are retained in the catalyst and are thermally CH 4 is at 400°C at the cracked to olefins. Though the equilibrium C+2H2 lefthand side, part of the carbon may react to CH4 during the thermal pretreatment for 16 hours at 400°C, especially since Fe is strongly catalysing this reaction. Consequently, the difference in selectivity shown by the LTR and HTR catalysts may..be ascribed to the fact that after the high-temperature treatment the pores of the carbon support have been assuming a larger diameter. As a result, higher molecular weight products can more easily escape out of the catalyst which leads to an apparently different selectivity. Have you any evidence of an effect of the pretreatment at 400°C on the porous structure of the catalysts? My second question is concerned with hydrogen adsorption by small iron particles. It has been observed that hydrogen adsorption on small iron particles is activated. Accordingly, your procedure of admitting the hydrogen at high temperature is suitable for very small iron particles. Have you any observation about the adsorption of hydrogen by the small iron particles at room temperature ? M.A. VANNICE: We have been very aware of the possibility of gasification of the carbon support, and have studied this reaction in our Fe/carbon catalysts to obtain a standard pretreatment w~ich minimizes this possibility (H.-J. Jung, Ph.D. Thesis, Penn State Univ., 1981). Over the 16h reduction period at 400°C, less than 1% loss occurs with this CSX-203 carbon as no CH4 was detected in the effluent H2 stream, and BET measurements on the used catalysts indicate very similar apparent surface areas (which are probably somewhat too high due to capillary condensation). However, your point is important and it is possible that some small increase in pore size could result. Under CO hydrogenation conditions, no gasification of the carbon support occurs because of the lower temperature and the inhibitive effect of CO. In addition, we do not believe
491
that any chromatographic effect exists because of the low boiling points of these light olefins (and paraffins), the constant selectivities with times on-stream up to 100h, and relative invariance with CO conversion. In agreement with our earlier studies in which we found, as other have, that H2 chemisorption is quite activated on very small (1-4 nm) crystallites (Ref. 3, plus H.-J. Jung, M.A. Vannice, L.N. Mulay, R.M. Stanfield and W.N. Delgass, J. Catal. 76 (1982) 208), H2 adsorption on these small mixed-metal particles after reduction was very low (J.J. Venter, M.S. Thesis, Penn State Univ., 1985). It was for this reason we utilized the high-temperature cooldown/desorption technique described by Butt and coworkers (Ref. 7).
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
493
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
CATALYTIC ACTIVITY OF CARBON SUPPORTED CATALYSTS FOR CO-HYDROGENATION AND THEIR PREPARATION BY OXIDATIVE DECOMPOSITION OF Fe(CO)5 U. Peters, H. Greb, R. Jockers and J. Klein Bergbau-Forschung GmbH, Franz-Fischer-Weg 61, 0-4300 Essen 13, FRG
ABSTRACT Activated carbon supported Fe catalysts were prepared by oxidative decomposition of Fe (CO) from the gas phase. It could be shown that metal distribution and dispersitin are critically determined by the degree of activation of the supports used. Activity and selectivity of these catalysts in respect to CO hydrogenation were determined in a plug flow fixed bed reactor and both show distinct dependencies on dispersion and distribution of the metal phase in the catalyst pellets. INTRODUCTI ON Due to its specific properties activated carbon lends itself as a support for catalysts and has been examined for a great number of different catalysts viz. reactions (ref. 1-8). Of particular interest are the pore structure and the specific surface area which can be varied within a wide range. It should be possible therefore to obtain a high dispersion and, consequently, sufficiently high surface areas of the metal phase. It is generally known that the "history" of a catalyst influences the dispersion and distribution of the metal phase (ref. 9). The influence of the kind of support and of the preparation on the dispersion - and thus on the activity of these catalysts for CO hydrogenation - was observed also in the case of Fe/C catalysts (ref. 2,10). The catalysts presented in this paper were obtained by oxidative decomposition of Fe(CO)5 from the gas phase (ref. 11). Apart from the small number of treatment steps (decomposition of Fe(CO)5 and subsequent reduction) this method, unlike impregnation, offers the chance of attaining any desired Fe contents in the catalysts and of eVidencing the influence of the pore structure of the support on dispersion and distribution of the Fe phase. EXPERIMENTAL The cylindrically shaped activated carbons, serving as support, were made from hardcoal using partial H20 gasification as activation method (ref.12). The percentage of gasified carbon is defined as degree of activation or burn
494
off. Along with the texture and with a burn off diameter and an
increasing burn off there occurs a systematic change both in the specific surface of the support. We used activated carbons of 0 %, 21 %, 35 %, 51 %, and 64 % and grain size of 2 mm average length of 4 mm.
Preparation The catalysts were prepared by decomposition of Fe(CO)5 contained in a N2 gas stream (cFe(CO) = 5 x 10-4 molll) which was mixed with air and then flown through a con~inuOuslY revolved heap of activated carbon. The thus obtained catalysts contained 10 % Fe by wt. Characterization The catalysts were characterized by the following procedures: - The distributions of pore radii were determined by adsorption of methanol and Hg porosimetry. - The metal and metal oxide were characterized by Moessbauer spectroscopy at temperatures of 298 K, 77 K, and 4 K and X-ray diffraction using Cu-Ka radiation (Siemens, Typ K810). - Thermo-gravimetric analyses (TGA) (Mettler type TA 1) were carried out in H2 in the heat range between 298 K and 1170 K, applying a heating rate of 10 K/min. The gaseous products were identified by mass spectroscopy and continuously registered during the experiment. - The dispersion viz. mean particle size of the Fe phase were defined by chemisorption of CO on the reduced catalysts at 298 K (reduction conditions: vH = 100 ml H2/min . g cat; T = 675 K; t = 16 h). The adsorption isotherms we~e determined by the vOlumetric Carlo-Erba apparatus (Sorptomatic 1800). - The Fe distribution within the catalysts pellets was determined by microprobe analysis. CO Hydrogenation CO hydrogenation was carried out in integral conditions in a plug flow fixed bed reactor (10 = 30 rom) charged with 3 g of catalyst. The synthesis gas used was blended, without any further treatment, from the pure gases (CO = 99.997; H2 = 99.999; N2 = 99.996 % by vol) and consisted of 60 % by vol H2, 30% by vol CO and 10 %by vol N2. CO hydrogenation was performed at pressures of 0.1 MPa at a temperature of 550 K and of 1 MPa, 1.5 MPa, and 2 MPa resp. at a temperature of 525 K, with space velocities of 600 h- 1 at normal pressure and 3800 h- 1 at increased synthesis pressure.
495
The products were analyzed on-line by gas chromatography specified in table 1. TABLE 1: Analytical equipment for products of CO hydrogenation analyzed products
Hewlett-Packard Gaschromatograph 5880A, Capillary column (50 m chemical bonded SE 54 fused silica), gas samplig valve, FIO+MSD(He) Temp. program: O°C (20 min)- 4°C/min - 175°C (60 min) packed column (0.6 m Molecular sieve 13 X) gas sampling valve, TCD (Ar), Temp.: O°C 5840A, packed columns (1.2 m Molecular sieve 13 X, 1 m Porapac Q), gas sampling valve, TCD (He) Temp.: 60°C
Hydrocarbons C1 - CIS H CH 4 2, CO, CO 2, H20, CH4, N2
The fact that the integrator of the GC HP 5880 A can be programmed in BASIC permits adaptation of the hydrocarbon product distribution by means of the Schulz-Flory distribution law (e.g. ref. 13) immediately upon analysis. Hydrocarbons were identified by a GC/MS-equipment (HP MSD 5970). RESULTS AND DISCUSSION Characterization Fig. 1 shows the systematic increase of the pore volume along with increasing support burn off. The deposition of about 15 to 16 % by wt Fe203 corresponding to a Fe concentration of about 11 % by wt - has a significant influence on the pore structure (table 2) .
.....
64
Cl IS) IS)
,
100
51
E
,., ~ L..O
U
u
'"'
III
35
E :3
->
21
III
0
0
0
...c
:l
III
~
0
a.
1"
Fore Radius Cnm]
Figure 1: pore size distribution of supports
496
TABLE 2: Pore characteristics of supports and freshly prepared catalysts Support No. 0.12 0.12 0.11 0.12 0.13
0 21 35 51 64
4 470 690 910
0.02 0.20 0.26 0.34 0.37
1010
0/10 21/10 35/10 51/10 64/10
Catalyst Fe content Vm~cro3 Vm!cro wt% cm /cm cm /cm 3 11. 5 12 11 11 11 n.d.=
0.1 0.1 0.12 n.d. 0.12 0.2 0.25 0.12 0.27 0.12 not determined
The Moessbauer spectra measured at ambient temperature of the Fe/activated carbon catalysts freshly prepared are completely identical and independent of the support burn off. They are characteristic for superparamagnetic Fe203 (fig. 2) and do not show any signs of splitting even at 4 K. After reduction in H2, 6 hours of atmospheric CO hydrogenation and subsequent reoxidation with oxygen from air there will occur a significant change in the Moessbauer spectrum of the catalyst with a burn off of 0 % (fig. 3a). In this case the characteristic six-peak spectrum of elemental iron will appear. In the case of catalysts with higher support burn off (figures 3b and 3c), the reduced iron is almost completely re-oxidized into a very finely distributed Fe203, which can be concluded from the central doublet. Besides there will occur lines of X-carbide (catalyst 35/10) and of elemental iron (about 2 %, catalyst 64/10).
e
lIS
c:
III
"
c)
L
I-
-10 -S 0 5 10 Ve 1oc i ty / mm*s-\
Figure 2: Mossbauer spectra at 298 K of freshly prepared catalysts: a) 0/10; b) 35/10; c) 64/10
-10 -S 0 S 10 Ve 1oc i ty / mm*s-\
Figure 3: Mossbauer spectra at 298 K of reduced and reoxidized catalysts: a) 0/10; b) 35/10; c) 64/10
497
XRD did not give any diffraction peaks for samples of freshly prepared catalysts. After reduction and re-oxidation the intensive diffraction peaks of metallic Fe were only found for a sample of catalyst 0/10. The above results render obvious that certain differences occur between the deposited Fe203 during reduction and that such differences are due to the support burn off. In order to examine this mechanism more in detail we ran thermo-gravimetric tests in a H2 atmosphere. Fig. 4 shows the differential mass losses as a function of the temperature. There are two distinctive types of mass loss diagrams: - The catalysts 0/10 and 21/10 show distinctive maxima below 750 K caused by the release of CO, CO 2, and H20. - The mass loss for catalysts 35/10, 51/10, and 64/10 in the same temperature range is caused by the formation of H20 and CO 2, with no distinctive maximum observed. What strikes with these catalysts is the high mass loss in the temperature range between 720 K and 920 K which is attributed to the gasification of the support under release of CH 4. g ] zsm [m K*g
6T
1.5
1.0
0.5
64/10 .....--=;:=:::....-----7 5 1/10 \---=~-----...::~~=---/
~--'--;''-----r---'--'----.---r---r---(
1100
35/10 21/10 t 0/10 C'
st
,1i
As the above results show, the behaviour during reduction of the Fe203 as well as the interaction between the reduced iron and the support material are influenced by the support burn off. It was demonstrated by CO chemisorption measurements that these differences are mainly due to the Fe phase dispersion. Table 3 contains the amounts of irreversibly adsorbed CO for the different catalysts. The table gives furthermore the relevant values derived for dispersion and mean Fe particle diameters /10/ .
498
TABLE 3: Dispersion and mean Fe particle size determined by CO chemisorption* /10/ Catalyst
CO-Uptake (mmol/gcat)
0/10 21/10 35/10 51/10 64/10
0 0.05 1.27 0.72 0.50
Dispersion
0.025 0.64 0.37 0.25
Fe particle size (nm) Fe/CO= 1/1 Fe/CO= 2/1 30.0 1.2 2.0 3.0
15.0 0.6 1.0 1.5
* Adsorption at 298 K The results of CO chemisorption show that the subdivision of catalysts in two groups, as occurred during thermo-gravimetry, has to be attributed to the discrepant dispersions of the Fe phase. For the first group (b.o. < 35 %) the particles, lead to bulk reduction will, notwithstanding the very small Fe 203 Fe crystals which - as becomes evident by the Moessbauer spectrum of the 0/10 catalyst - do not reoxidize by access of oxygen from air during the lengths of times observed ( 6 month). For the second group of catalysts (b.o. ~ 35 %) there appears a very high Fe dispersion which is stabilized by the supports. From the above observation it can be assumed that Fe203 is deposited mainly in the micro-pores. These micro-pores contribute mainly to the total surface area which is rather small at low support burn offs. In the case of high burn offs (b.o. ~ 35 %) these micro-pores will prevent the metal crystallites from sintering, and stabilize them. In the case of low burn off values the available micro-pore volume will be insufficient for uptaking the quantity of Fe203 applied, thus inducing the formation of great bulk Fe-crystals during reduction. The activity of metallic support catalysts will, however, be critically influenced not only by the dispersion but also by the distribution of the metal within the particles of the support. For this reason we determined the Fe distribution in the catalyst pellets by micro-probe analyses (fig. 5). Up to a burn off of 35 % a homogeneous Fe profile is observed. Higher burn off levels lead to the formation of shell catalysts. As we found out recently (ref. 14) by more detailed examinations, this is due to the higher rate of decomposition of Fe(CO)5 along with increasing burn off.
499
Figure 5: Fe distribution across catalyst pellets a) 0/10; b) 35/10; c) 64/10 CO Hydrogenation Table 4 summarizes the results of CO hydrogenation obtained when using different catalysts at reaction pressures of 0.1 MPa and 2 MPa. TABLE 4: Typical results of CO-hydrogenation for Fe/C-catalysts
Catalyst
0/10 21/10 35/10 51/10 64/10
pressure: O'!lMPa, temperature: 550 K, t = 0.33 h S.v = 630 h chain growth Conversion (%) Alkene/Alkane ratio probability CO
o
37.2 76.4 71.1 42.0
o
17.7 23.6 19.7 19 6
0.022 0.004 0.014 0.045
pressure: 2 MPa, te~~erature: S.v. = 3800 + 200 h 0/10 21/10 35/10 51/10 64/10
o
5.1* 14.1 93.0 91.1
o
5* 13.8 61.6 56.8
* after 4.8 h time on stream
0.0583* 0.083 0.009 0.021
0.05 0.01 0.04 0.13
o 02 0.04 0.22
525 K, t
0.295* 0.472 0.052 0.159
0.45 0.50 0.44 0.46
0.11
= 20
0.503* 0.756 0.132 0.333
h
0.47* 0.57 0.45 0.45
500
As at atmospheric pressure the catalysts are subjected to rapid deactivation, the table gives the initial activity after a reaction time of 20 minutes. The distribution of hydrocarbons can even under these instationary conditions be expressed by the Schulz-Flory law. Independently of the type of catalysts a chain propagation probability of 0.45 has been determined. On the other hand the activity and olefin contents of the hydrocarbons show a strong dependence on the type of catalyst used. Catalyst 35/10, showing the highest Fe dispersion, turned out to be the most active, too, which was up to the expectations. The minimum of olefin contents may also be explained by high dispersion: the a.m. catalyst has at the same time the highest metal surface available for hYdrogenation of the olefins which are primary products of CO-hydrogenation. Besides the formation of linear and branched paraffins and olefins there will also occur formation of BTX-aromates under these reaction conditions.
100 r"1
~ 1-1
80
c 0
60
III L-
III
> 40 c 0
U I
0
20
u 5
Ifj me
[h)5
20
25
Figure 6: CO-conversions of different catalysts with time on stream p 2 MPa: v 21/10; 035/10; 051/10; '" 64/10 P = 0.1 MPa: -51/10 The deactivation of Fe catalysts during CO hydrogenation is known to be considerably reduced by application of increased synthesis pressures (ref. 15). This is represented on fig. 6 for the catalyst 51/10. Furthermore fig. 6 shows that, on the other hand, higher reaction pressure (2 MPa) will not induce all the catalysts to an increased and permanently high level of activity.
501
Catalysts 0/10 and 21/10 behave, just as during synthesis at atmospheric pressure, inactive or just slightly active which is explained by the low dispersion of the Fe-phase. Catalysts 51/10 and 64/10 show an almost constantly high activity over the entire period under review. Although Fe-dispersion is highest with catalyst 35/10 among all the catalysts concerned. a substantial decline in activity is observed to occur immediately after the reaction started. We believe that this behaviour has to be attributed to some blockage of the iron deposited within the catalyst pellets. by high molecular liquid reaction products. This hypothesis is supported by calculations done by other groups (ref. 16) and is furthermore revealed by the chain propagation probability of 0.57 derived from the Schulz-Flory distribution law. For catalysts 51/10 and 64/10, both of long-lasting activity, chain propagation probabilities amount to 0.45. For these catalysts blockage by long-chain reaction products is excluded due to the altered pore structure and the thin shell wherein the iron is deposited and which implies short diffusion paths. Run under a pressure of p = 2 MPa the reaction does not yield aromates. We suppose that at atmospheric pressure conditions the aromates are coke precursor or are produced by decomposition of graphitic deposites. The existence of graphitic carbon has been evidenced elsewhere (ref. 17) and may be responsible for the rapid deactivation of Fe catalysts at atmospheric pressure conditions. CONCLUSIONS Applying the described method and using appropriate activated carbons as support. it is possible to produce highly active catalysts. During catalysts preparation it turned out that the pore system exerts a strong influence on the dispersion and distribution of the deposited Fe phase. It will have to be clarified by further studies. yielding detailed kinetic data, in how far the activity of catalysts and the behaviour with time on stream are functions of transport mechanisms. ACKNOWLEDGMENTS We would like to thank Mr. B. Gatte and Prof. M. Philips from Pennsylvania State University and Mr. M. Deppe and Prof. M. Rosenberg, Ruhr-Universitat· Bochum. for measuring Mtissbauer spectra and helpful discussions.
502
REFERENCES 1 2 3 4 5 6 7
8 9
10 11 12 13 14 15 16 17
M. Kaminsky, K.Y. Yoon, G.L. Geoffroy, M.A. Vannice J. Catal. 21 (1985), 338 F. Rodriguez-Reinoso, J.D. Lopez-Gonzalez, C. Moreno-Castilla, A. Guerrero-Ruiz, J. Rodriguez-Ramos FUEL 63 (1984), 1089 E. Kikuchi, A. Koizumi, Y Aranishi, Y. Morita J. Japan Petrol. Inst. 25 (1982), 360 A.P.B. Sommen, F. Stoop, K. van der Wiele Appl. Catal. 14 (1985), 277 V.H.J. de Beer, F.J. Derbyshire, C.K. Groot, R. Prins, A.W. Scaroni, J.M. Solar FUEL §l (1984), 1095 A.W. Scaroni, R.G. Jenkins, P.L. Walker, Jr. Appl. Catal. 1i (1985), 173 F.F. Gadallah, R.M. Elofson, P. Mohammed, T. Painter Preparation of Catalysts III (Edited by G. Poncelet, P. Grange, P.A. Jacobs) Elsevier Science Publishers B.V., Amsterdam (1983), 409 J.L. Schmitt, Jr., P.L. Walker, Jr. Carbon 1Q (1972), 87 J.W. Geus Preparation of Catalysts III (Edited by G. Poncelet, P. Grange, P.A. Jacobs) Elsevier Science Publishers B.V., Amsterdam (1983), H.-J. Jung PhD Thesis, Pennsylvania State University (1981) H. Greb, K.-D. Henning, J. Klein, U. Peters DE-PS 33 30 621 H. JUntgen Carbon ~ (1968), 297 G. Henrici-Olive, S. Olive Angew. Chern. 88 (1976), 144 U. Peters, R. Jockers, J. Klein Paper presented at "Carbon '86" F. Fischer, H. Pichler Brennstoff-Chem. 20 (1939), 41 G.A. Huff, Jr., C.N. Satterfield Ind. Eng. Chern. Process Des. Dev. 24 (1985), 986 H.P. Bonzel, H.J. Krebs Surf. Sci. 21 (1980), 499
503
DISCUSSION T. HATTORI: You have mentioned that micro-pore is blocked by Fe203' But the difference in volume of micro-pore is not so large between carbon 21 and 35. What you mentioned is the blockage of pore mouth? R. JOCKERS : Yes, you are rigth. The comparison of available micropore volume of the support and the volume of deposited iron oxide shows that only a small part of the micropore volume can be occupied by Fe-oxide. The drastic decrease of the micropore volumes of the prepared catalysts is a strong indication for pore mouth blockage. J.W. GEUS : 1/ Your results are suggesting that your catalysts prepared from supports of a high burn-off are exhibiting diffusion limitation in the reaction. This can be established by measuring the activity on crushed catalyst tablets. Did you vary the size of your carbon support particles and did that affect the conversion ? 2/ We get a more homogeneous distribution of iron over the support during the decomposition of iron carbonyl by first exposing the support to the carbonyl containing gas at a temperature so low that the carbonyl did not decompose and subsequently rapidly raising the temperature. Since we used a fluidized bed, we could raise the temperature fast. Have you experience with addition of the carbonyl at low temperature and subsequent heating? R. JOCKERS : 1/ Diffusional limitation during reaction may be an explanation for the observed dependence of CO-conversion on burn-off of support. We made activity measurements with a crushed catalyst (burn-off = 35%, particle size = 0.08 - 0.5 mm) and found the same loss of activity with time on stream that we have observed with the uncrushed catalyst. Detailed kinetic measurements are necessary in order to clarify the role of diffusional limitation during CO-hydrogenation. Another explanation for the dependence of activity on burn-off is based on the influence of burn-off on mean iron particle size. Very small Fe-particles result on activated carbons which have 35% burn-off. Possibly these particles are converted to Fe-carbides very fast, and are blocked by inactive carbidic carbon layers in consecutive reactions. 2/ We performed Fe(CO)5-decomposition at various temperatures (273 K - 373 K) and found that the Fe-distribution across the support can be influenced by the decomposition temperature. By example with a burn-off = 35%. there is a change from a homogeneous iron distribution to an egg-shell catalyst, if the decomposition of Fe(CO)5 is carried out at 323 and 373 K, respectively. This fact could be explained by an acceleration of the decomposition rate with increasing temperature. We think that your observation can be explained by a volatilization of adsorbed Fe(CO)5 and a further diffusion of this compound to the centre of support particles during the raise of temperature.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
ON 1HE MECHANISM OF FORMATION OF COLWIDAL MONODISPERSE METAL BORIDE PARTICLES FROM REVERSED MICELLES COMPOSED OF CTAB l-HEXANOL- WATER
I. RA VET, 1 B.NAGY and E. G. DEROUANE Facultes Universitaires Notre Dame de la Paix., Centre de Recherches sur les Materiaux Avances, Laboratoire de Catalyse, 61 Rue de Bruxelles, B-5000 NAMUR.
SUMMARY A new method, proposed for the preparation of very small (d = 20-70 A) monodisperse particles of Ni 2B, Co 2B and Ni-Co-B, is examined from a mechanistic point of view. These catalysts are synthesized by reduction with NaBH4 of Ni 2+ and Co 2+ ions solubilized in a reversed micellar system composed ofCfAB (cetyltrimethylammonium bromide), l-hexanol and water. This work points out the role of micelles in obtaining colloidal catalysts as well as the essential parameters which control the size of particles formed in this system. The catalytic properties of the bimetallic particles are tested for the room temperature liquid phase hydrogenation of crotonaldehyde. INTRODUCflON In heterogeneous catalysis, the size of catalysts has a considerable importance. Much effort is devoted therefore to increase the quality of catalysts by increasing their specific area. Recently, a new method has been proposed to synthesize very small quasi monodisperse particles. The particularity of this mode of preparation is the use of a reversed micellar system as reaction medium. Metal salts are dissolved in the inner water cores of micelles and reduced with hydrogen,hydrazine or sodium borohydride. By this means, small particles ofPt (refs. 1-5), Pd,
Ir, Rh (refs. 1-3) ,Au (ref. 6), Ni boride (refs. 7-9), Fe boride (refs. 10-11) and Co boride (refs. 12-13) were obtained. In this paper, we study the mechanism of formation of nickel boride, cobalt boride and bimetallic Ni-Co boride by reduction, with NaBH4 ' of NiCl 2 and CoCl2 solubilized in the reversed micellar system composed of crAB (hexadecyltrirnethylammonium bromide), I-hexanol, water. These particles prepared either in ethanol or in water are known to be good catalysts for the hydrogenation of alkenes (refs. 13-23). Cobalt boride also catalyzes the hydrogen production from hydrogen peroxide (ref. 24) or from sodium borohydride (ref. 25) and the reduction of nitrate into ammonia by NaBH4 (ref. 24). On the other hand the selectivities of nickel and cobalt boride are very different for the hydrogenation of a,~-unsaturated
carbonyl groups (ref. 26).
Nickel boride favours the C=C double bond hydrogenation while cobalt boride hydrogenates
505
506
selectivelythe C=O double bond. Hence, the formation and the catalytic properties of the bimetalliccompound will be also compared. EXPERIMENTAL Materials The commercial products, l-hexanol (Merck, 98%), CTAB (Serva, 99%), NiClZ.6HzO (Merck, p.a.), CoClz.6HzO (Merck, p.a.), NaBH4 (Aldrich, 99%), crotonaldehyde (Merck, for synthesis), ethanol (Merck, p.a.) were used without further purification. Preparationof the particles The metal boride particles were synthesizedfrom several reversed micellar systemsof different compositions (see table 1) and, for each composition,the metal ion concentrationwas varied (see table 2) TABlE 1 Micellar compositionsused for the preparationof metal borideparticles Type
Composition (weight%) CTAB Hexanol
particlesprepared
HzO
1
2 3 4 5 6 7
12 16 20 12 15 18 20
18 24 30 39 38 37 36
70 60 50 49 47 45 44
Ni-B and Ni-Co-B Ni-B Ni-B Co-B Co-B Co-B Co-B
TABlE 2 Concentrationsof metal ions used for the preparationof Ni-B and Co-B Type
A B C D E
[Niz+] xl0Z (Molal) a
[Coz+] xlOZ (Molal)a
1.00 2.50 5.00 7.50 10.00
0.50 2.50 5.00 10.00 15.00
a) Molal concentrationversus total micellar solution. The bimetallicparticleswere obtainedfrom a constanttotal metal ion concentration(5 x 10'Z molal) varying the [Co2+]I([Ni2+]+[C02+]) ratio from 0 to 1 by step of 0.1.
507
The synthesis was carried out in a glove box under argon atmosphere to prevent the oxidation of the particles.Figure 1 shows the preparation scheme of the catalysts (ref. 8) (the compositions noted in Tables 1 and 2 are the final compositions). Electron microscopy The average size of the metallic particles was measured using a Philips EM 301 electron microscope in the transmission mode. For these measurements the particles were ultrasonically dispersed in butanol and deposited on grids covered with Formvar.
x y
714 H20 A molesofNiCI2 B molesof CoCl2
712
X eras Y Hexanol 3714 H20 A molesof NiCI2 B molesof CoCI2
714 H20 C molesofNaBH4 with C = 3(A + B)
x
crxs
Y
Hexanol H20 molesof NiClZ molesof CoClZ molesofNaBH4
Z A
B
C
Fig. 1. Preparation scheme of the particles Hydro~enation
tests
The crotonaldehyde hydrogenation tests were conducted at atmospheric pressure (760 mrnHg
± 10) and at room temperature (23 ·C ± 2) in a slurry type static reactor with continuous stirring by following volumetrically the consumption of hydrogen. The catalysts were synthesized in situ under Argon flow. For the activity measurements the reaction was carried out in an ethanol (90 weight %)-micelles (10 weight %) mixture, [crotonaldehyde] = 4 x lO-Z molal, [metal] = 5 x 10-3 molal. For the selectivity study, the a,p-unsaturated aldehyde was hydrogenated in a mixture composed of the micellar system used for the preparation of the catalysts, [crotonaldehyde] = 1 molal and [metal] = 5 x lO-Z molal.
508 IH-NMR After hydrogenation, the reaction mixtures were analyzed by IH-NMR to identify the products obtained. The spectra were recorded at room temperature on a Broker CXP-200 spectrometer at 200 MHz.The stabilization of the magnetic field was carried out using a deuterium lock ofD20 contained in the internal part of two concentric tubes, the external tube being filled with the solution to be characterized. RESULTS AND DISCUSSION Solubilization sites of ions and sizes of inner water cores of reversed micellar aggregates In ealier works (refs. 8, 12, 13), 13C-NMR measurements have shown that, in the reversed micellar system CTAB - 1-hexanol- water, the Ni2+ and C02+ ions are located in the inner water cores of the micelles quite close to the interface. Indeed, on average, one hexanol molecule is included in the first coordination shell of C 02+ ions, while one or more hexanol molecules participate to that of Ni 2+ ions. The size of the micellar water cores has been determined by an indirect method,based on 19p_NMR measurements of probe molecules (ref. 27). The average radii (refs. 8, 9, 12, 13) of the aggregates (rM) containing the precursor ions are important parameters in the understanding of the formation of colloidal particles. Size of monometallic particles Monodispersed colloidal nickel boride and cobalt boride particles are obtained by reducing, with NaBH4 , the precursor ions solubilized in the water cores of micelles. The composition of the catalysts has been determined by XPS (ref. 28) as being respectively Ni 2B and CazB. The size of the particles so prepared is always much smaller than that obtained by reduction of Ni2+ or Co2+ in ethanol or water (d=2500-4000 A). In the micellar system, the average diameter depends on the composition of the initial solution (see fig. 2). 70 70
.------,---r--....,
r~
CTAB 1HEXANOLI H:zO
30 30
18%170%/12%
o
5
10
- . 102 [NiH] (molal)
15
--+
200
5
10
15
-.10 2 [Co++] (molal) --+
Fig. 2. Average diameter of Ni 2B and Co 2B particles as a function of precursor ions .concentration for different micellar compositions.
509
We observe (see fig. 2) that the particle size as a functionof the precursorions concentration passes through a minimum. This phenomenom can be rationalizedif one analyzes the nucleation and the growth processes of the particles. Principles for the formation of the colloidal particles To form a stable nucleus a minimumnumber of atoms is required (refs.9, 29). The nucleationstep is always slower than the growth process. At the very beginingof the reduction, nucleationonly occurs in those watercores which contain enough ions to form a nucleus. At this moment,the micellaraggregatesact as "reactioncages" where the nuclei are formed. On the other hand, the micellar system being dynamic,rapidly the water cores rearrange. The other ions broughtinto contactwith the existingnuclei essentiallyparticipateto their growthprocess. This latter being faster than nucleation, no new nucleusis synthesizedat this moment. As all the nuclei are formed at the same time and grow at the same rate, monodisperse particlesare obtained. In summary,the particle size dependson the the number of nuclei formed at the very beginningof the reduction and this numberis a functionof the number of water cores, containing enough ions to form a stable nucleus,reached by the reducing agent before the rearrangementof the system. Quantitative modelfor the formation of the colloidal particles The first step in the calculationof the essentialparameterswhichcontrol the particlesize is the study of the distributionof the ions in the micellar water cores. Knowing the averge radii of the micellaraggregates(rM) (refs. 8, 9, 12, 13) and the total volumeof water (VT) per kilogramof micellarsolution(see table 1), one can calculatethe number of water cores per kilogramof reversedmicelles (N M ), neglectingthe solubilityof water in the l-hexanol organicphase: NM = -
-
-
-
(1)
This parameter(NM) and the initialconcentration in metal ions expressedin molality(see table 2) allow us to determine the averagenumberof ions per water core (nions):
[ions] x 6.023x1023 nions = - - - - - - -
(2)
The ions are statistically distributedin the aggregates.To calculatethis distribution, the statisticsof Poisson are perfectly adequate(ref. 30). They give the probability
510
ions per water core (k is an integer taking the values 0, 1,2,3, ... ), provided the average number of ions per water core ( A. = nions) is given:
(3) k!
The number of nuclei formed (Ng) when the ions solubilized in one kilogram of solution are reduced, is proportional to the number of aggregatescontaining enough ions for nucleation. If the minimum number of ions required to obtain a nucleus is i, Ng can be calculatedfrom the followingrelation: 00
(4) 00
00
:t Pk is the probabilityto have i ions or more per aggregate; hence NM :t Pk is the k=i
k=i
number of water cores containingi ions or more; F is a proportionalityfactor taking into account the proportionof aggregatesreachedbythe reducing agent before rearrangementof the system can occur. In this equation, we do not know the values of i and F but we can calculate all the other parameters. Indeed the number of nuclei (Ng) is the number of particles prepared and the latter is given by:
(5)
w
where Wt is the total weight of catalystsprepared per kilogram of micellar solution; w is the weight of one particle, and
[ions] ~ata ~=
~ x
where Mcata is the molecularweight of the catalyst;x is the number of metal atoms per moleculeof catalyst; w=
413 1t (dl2)3Mv cata
where d is the diameter of the particles measuredby electron microscopy (see fig. 2); Mv cata is the
volumic mass of the catalyst.
(7)
511
Calculation of the proportionality factorF For all the particlessynthesized, we have calculatedthe proportionality factorF, varying systematically the valueof the minimumnumberof ions requiredto form a nucleus (i), Only if i takes the valueof 2, is factorF reasonably constant (see table 3). The order of magnitudeof the factorF is always 10-3. That means that, at the very beginning of the reduction, i. e. when the nucleiare formed, only one aggregate per thousandis reached by the reducing agentNaBH4 . Anotherfact showsthat the nucleation occursat the very beginning of the reduction. Indeed, the averageradiiof the watercores usedfor thecalculation of the formation parameters of colloidalparticles,are measuredfor the systemcontaining only 3/4 of the total amountof water which is the compositionof the solutionbeforethe additionof the reducingagent (see fig.1). If the fmal compositionis used, however, no coherentresults can be obtained. The order of magnitudeof the factorF is constantbut its valuedecreases with the increaseof watercontent in the reversedmicellarsolution(see table 3). This phenomenon can be easily understood becausethe rearrangement rate of the micellarsystemincreases with the water amount (ref. 10)and hence,the number of aggregates reachedby the reducingagent before rearrangement decreases. As the numberof nucleiformed decreases, for a constant concentration of precursorions, the particlesize increases with the increaseof water in the system (see fig. 2). TABLE 3 Importantparameters for the formation of the colloidalparticles. Type [Ni2+]
LA I.B I.C I.D 2.A 2.B 2.C 2.D 2.E 3.A 3.B 3.C 3.D 3.E
rM
x102
(A)
(molal)
a
1.00 2.50 5.00 7.50 1.00 2.50 5.00 7.50 10.00 1.00 2.50 5.00 7.50 10.00
10.2 11.9 13.7 14.6 11.7 13.2 14.8 15.4 15.7 13.4 14.8 16.1 16.8 17.2
v.
NM nNi2+ d x10-22 (A) (g) a,b a c 2.11 0.29 1.33 1.13 0.87 3.45 0.72 6.24 1.86 0.32 1.29 1.17 0.92 3.27 0.81 5.58 0.77 7.82 1.54 0.39 1.16 1.30 0.90 3.35 0.79 5.72 0.74 8.14
44 36 32 37 45 42 40 40 51 67 49 45 46 49
w xlOI9 (g)d
0.64 3.52 1.60 1.93 3.21 1.35 4.81 2.09 0.64 3.77 1.60 3.06 3.21 2.65 4.81 2.65 6.41 5.49 0.64 12.44 1.60 4.87 3.21 3.77 4.81 4.03 6.41 4.87
Ng xlO-I8 b
NglNM
1.82 8.29 23.77 23.01 1.70 5.23 12.11 18.15 11.67 0.51 3.28 8.51 11.93 13.16
8.63.10'5 6.23.10-4 2.73.10"3 3.19.10-3 9.14.10'5 4.05.10-4 1.32.10-3 2.24.10-3 1.52.10-3 331.10- 5 2.83.10-4 9.45.10-4 1.51.10-3 1.78.10-3
00
LPk
F xl03
k.2
0.0347 0.3119 0.8587 0.9859 0.0415 0.3265 0.3265 0.9752 0.9964 0.0589 0.3732 0.8474 0.9780 0.9973
2.5 2.0 3.2 3.2 2.2 1.2 1.6 2.3 1.5 0.6 0.8 1.1 1.5 1.8
512
Type [Co2+] rM x102 (A) (molal) a
xlO-22 a,b
a
4.B 4.C 4.D 4.E 5.A 5.B 5.C 5.D 5.E 6.A 6.B 6.C 6.D 6.E 7.A 7.B 7.C
3.14 2.94 2.75 2.75 2.48 1.83 1.61 1.50 1.46 1.90 1.18 1.06 0.96 0.92 1.47 0.94 0.82
0.48 1.02 2.19 3.28 0.12 0.82 1.87 4.01 6.19 0.16 1.28 2.84 6.27 9.82 0.20 1.60 3.67
2.50 5.00 10.00 15.00 0.50 2.50 5.00 10.00 15.00 0.50 2.50 5.00 10.00 15.00 0.50 2.50 5.00
nCo2+
NM
8.9 9.1 9.3 9.3 10.4 11.5 12.0 12.3 12.4 12.1 14.2 14.7 15.2 15.4 13.7 15.9 16.6
d
v.
(A) (g) c
w
Ng xlO-1S
(g) d
b
31 1.61 1.26 28 3.22 0.93 28 6.43 0.93 30 9.65 1.14 59 0.32 8.71 49 1.61 4.99 41 3.22 2.92 34 6.43 1.67 38 9.65 2.33 67 0.32 12.76 52 1.61 5.96 46 3.22 4.13 41 6.43 2.92 46 9.65 4.13 71 0.32 15.18 59 1.61 8.71 48 3.22 4.69
F
NglN M
x10 19
x103
LPk k=2
4.07.10,4 1.18.10'3 2.51.10,3 3.08.10.3 1.49.10,5 1.76.10-4 6.25.10-4 2.57.10,3 2.84.10,3 1.32.10-5 2.29.10-4 7.36.10,4 2.29.10,3 2.54.10,3 1.43.10'5 1.97.10,4 8.38.10,4
12.78 34.62 69.14 84.65 0.37 3.23 11.03 38.50 41.42 0.25 2.70 7.80 22.02 23.37 0.21 1.85 6.87
4.8 4.3 3.9 3.7 2.3 0.9 1.2 2.8 2.9 1.2 0.6 1.0 2.3 2.5 0.8 0.4 1.0
0.0842 0.2716 0.6430 0.8389 0.0066 0.1984 0.5577 0.9091 0.9853 0.0115 0.3661 0.7756 0.9852 0.9994 0.0175 0.4751 0.8810
a) Values given for the system containing 3/4 of the total amount of water. b) Values given for 1 kilogram of solution. c) W t is calculated with MNi2B = 128.23 g/mole and M Co2B = 128.68 g/mo1e. d) w is calculated with M v Ni2B = 7.9 g/cm 3 and M v Co2B = 8.1 g/cm 3 (ref. 31)
0.75
1.0
r...
G) 0.75
j
0.50
@ 0.75
!I-l
0.50
0.25
G
J e-,
4
J
2
•
0.50 2
15 -k-+
~
.}
O.6
0.4
I
3J
0.8
Q"
0.2
0.25
0.00
Fig. 3.
0.0 0
6
9
12
0
246 ---lNi++j. 102 (moIa1)_
0
15-k-+
Fig. 4.
Fig. 3. Number of Ni 2+ ions per aggregate for the system CTAB 18%, hexanol 70%, H 20 12%. A [Ni2+] = 1 x 10-2 molal, C [Ni2+] = 5 x 10,2 molal, D [Ni2+] = 7.5 x 10-2 molal Fig. 4. Variation of the number of nuclei formed ~er aggregate and of the probability to have 2 or more ions per aggregate as a function of the Ni + concentration in the system CTAB 18%, hexanol 70%, H2O 12%
.
Z
513
The results of table 3 also allow us to explain the minimum in the particle size as a function of the concentration of ions (see fig. 2). For a constant micellar composition, at low ion concentration, only few water cores contain the minimum ions (2) required to form a nucleus, hence few nuclei are formed at the very beginning of the reduction and the size of the metal boride particles is relatively high. When the ion concentration increases, the distribution of precursor ions in the micelles is very different (see fig. 3) and the number of nuclei obtained by reduction increases faster than the total number of ions (see fig. 4). It results in a decrease of the catalyst particle size. When more than 80% of the water cores contain two or more ions (see fig.3 and 4), the number of nuclei formed remains quasi constant with increasing ion concentration. Hence, the size of the particles increases again. Study of the formation of the mixed nickel-cobalt boride Particles We have synthesized in the same micellar system (I.C) nickel boride and cobalt boride particles. In the first case, the value obtained for F is equal to 3.2 xlO- 3, in the second case, 17.4 xI0-3. As for these experiments the rearrangement rate of the micelles is constant in first approximation, the difference between the values ofF is probably due to the different solvation of the two types of ions at the interface (refs. 8, 12, 13). The Ni 2+ ions being multiply coordinated with hexanol at the interface, their mobility is lower and hence the probability of collision between the two reduced Ni atoms required to form a nucleus is also lower. In other words, the rate of nucleation is higher for cobalt boride than for nickel boride particles. From the same micellar system (1.C), bimetallic particles of Ni-Co-B have also been prepared. The average particle size and the width of the size distribution measured by electron microscopy is shown in figure 5 (e). No coherent values are obtained for the factor F, if the particles are considered to be homogeneous bimetallic catalysts. On the other hand, knowing the values of F for Ni 2B and COzB in this micellar system (l.C), we have calculated the expected sizes for the case where a mechanical mixture of separate particles of the monometallic borides is formed. These values are reported on figure 5 (0 for Ni 2B and 11 for Co 2B) as well as the weighted average sizes for these two types of particles (p). Indeed, only the latter can be compared with the experimental results which are average sizes. The averge sizes so calculated (q) are close to those measured experimentally (e). Nevertheless, in most of the cases, the
experimental size distributions are too narrow to correspond to a mechanical mixture of monometallic catalysts. Hence, the particles are probably bimetallic, but not completely homogeneous. Indeed, the nucleation rate is higher for Co2+ ions than for Ni 2+ ions(see above), and the nuclei are formed preferentially from Co2+ ions and the particles contain more nickel at the surface. This heterogeneous composition of the bimetallic particles is further confmned by the catalytic tests.
514
20
Fig. 5. Particle sizes as a function of the molar fraction in cobalt in the catalysts. • experimentalsizes. Hypothesisofa mechanicalmixture of Ni2B and C02B: 0 sizes calcutated for Ni2B, !! sizes calculated for C~B,cweighted
average sizes for Ni2B+C02B
Catalyticproperties of Ni-Co-B The colloidal catalystshave been tested for the liquid phase hydrogenation of crotonaldehyde at room temperature. On Ni2B, the C=C double bond is rapidly hydrogenated to form butanal but, in the reaction conditionsused, the reaction is not very selective. Only 20% of aldehyde is formed. On the other hand, on C~B,
the carbonyl function is hydrogenated very selectively but
the activity of the catalyst is lower. The only productdetected is the unsaturatedalcohol (butenol). On the mixed particles, the reaction products are the same as on Ni2B. The activity of the bimetallicparticles varieslinearlyas a functionof the amount of nickel in the catalyst (see fig. 6). When the particles contain both Ni and Co metal atoms,the reaction occurs essentiallyon nickel. This phenomenoncan be explained by the much lower activity of cobalt boride than that of nickel boride (ref. 26). In addition, the crAB molecules are strongly adsorbed on cobalt sites (ref. 32) and act as inhibitorfor the C=O hydrogenationcatalyzed by these sites. The activityof the bimetallicparticlesas a functionof their nickel content confrrmthe composition proposed for these catalysts. Indeed, as the reaction occurs on nickel sites only, the activity would drop to zero at Xe l (C02B). Fig.6 shows that the activity of the bimetallic catalyst is different from that expected for particlesof homogeneouscomposition (dottedline). The additionalactivityis explainedby the surfaceenrichmentin nickel in the colloidalmixed nickel-cobaltparticles. These results are in complete agreementwith the mechanismproposed for the formation of these particles, where a cobalt boride nucleus was supposed to be at the origin of
515
all particles. Nevertheless, the concentration gradient is much smaller than in big particles prepared in ethanol (ref. 26). In the latter case, the activity passes through a maximum as a function of nickel content. 1.00 r----r-,...----,r----r--r---,-"""T"'-,----,-..., HJC-CH=CH-CHO - . - ~C-CH2-CHiCHO 0.75
0.50 [crotonaldehyde J = 4 x 10 - Zmolal 0.25
PHz= 760 mmHg
T= 296 K 0.00 "'--......._...J...-~"'----'-_...I..----'_--'-_..&.-.-----'_'-U 0.0
0.1
0.2
0.3
'"
0.4
0.5
0.6
0.7
0.8
'"
0.9
1.0
•
x
Fig. 6. Catalytic activity of Ni-Co-B as a function of the molar fraction in cobalt in the particles for the liquid phase hydrogenation of crotonaldehyde at room temperature. CONCLUSIONS Knowing the solubilization sites of the precursor ions, the sizes of the micellar water cores, and the sizes of as synthesized metal boride particles, we have proposed a mechanism for the formation of the colloidal particles. The essential parameters have been quantified. We have shown that two metal atoms are required to form a stable nucleus andthat nucleation only occurs in those aggregates which are reached by the reducing agent before the rearrangement of the system can occur (one per thousand aggregates). This study also allowed us to have some informations about the composition of the bimetallic (Co-Ni-B) particles. Their catalytic activity has confirmed that the particles are rich in nickel at the surface because the nucleation rate is higher from CoZ+ions than from Ni Z+ ions. On the mixed catalysts, the crotonaldehyde is hydrogenated essentially on the nickel sites which are more active and less obstructed by adsorption of CT AB than the cobalt site. REFERENCES 1 2 3 4
M. Boutonnet, 1. Kizling, P. Stenius and G. Maire, Colloids and Surface, 5 (1982) 209. M. Boutonnet, Ph. D. thesis, universite Louis Pasteur, Strasbourg, (1980). P. Stenius, 1. Kizling and M Boutonnet, U.S. Pat. 4, 445, 261 (1984). T.H. Hsieh and H.S. Fogler, 5 th International Conference on Surface and Colloid Science, lecture 175, Potsdam, New York (1985). 5 A. Wathelet,Memoire de Licence, Facultes Universitaires, Namur, (1984).
516
6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32
K Kurihara, 1 Kizling, P. Stenius and J.H. Fendler, J. Am. Chern. Soc, 105 (1983) 2574. D. Rosier, Work for the end of studies, Institut Superieur Industriel, Anderlecht (1984). A. Gourgue, Ph. D. thesis, Facultes Universitaires, Narnur, (1981). 1 B.Nagy, A. Gourgue and E.G. Derouane, Stud. Surf. Sci. Catal.,16 (1983) 193. N.B. Lufirnpadio, Ph. D. thesis, Facultes Universitaires, Narnur, (1983). N.B. Lufirnpadio, 1 B.Nagy and E.G. Derouane,in KL. Mittal and B. Lindman (Eds.), Proc. Int.Syrnp. on Surfactants in Solution, Lund, Juny 27-July 2,1982, Plenum Press, New York, Vol. 3, 1983, pp. 1483. 1. Ravet, N.B. Lufirnpadio, A. Gourgue and 1 B.Nagy, Acta Chirn. Hung., 119 (1985) 155. 1. Ravet, A. Gourgue, Z. Gagelica and J. B.Nagy, Proc.8 th Int. Congress On Catalysis, Berlin West, July 2-6, 1984, Verlag Chernie, Weinheim-Basel, Vol. IV, 1984, pp. 871. D.G. Holah and S. Friberg, J. Am. Chern. Soc, 94 (1981) 12. Y. Okamoto, Y. Nitta, T. Irnanaka and S. Teranishi, J. Chern. Soc. Faraday Trans I, 75 (1979) 2027. RP. Buisson and N. Joseph, C. R. Acad. Sci. Paris, 232 (1951) 627. H.C. Brown, C.A. Brown, V.K Ahuja, J. Org. Chern., 85 (1963) 1005. C.A. Brown, V.K Ahuja, J. Org. Chern., 38 (1973) 2226. L.B. Luttinger, lOrg. Chern, 27 (1962), 1591. T. Irnanaka, Y. Nitta and S. Teranishi, Bull. Chern. Soc. Jpn., 46 (1973) 1134. RW. Mitchell, L.l Pandolfi and P.C. Maybury, 1 Chern.Soc. Chern. Cornrnun, (1976) 172. D.G. Ho1ah, lM. Hood1ess, A.N. Hughes and L. Sedor, J. Catal., 60 (1979), 148. Y. Okamoto, Y. Nitta, T. Irnanaka and S. Teranishi, 1 Catal., 64 (1980) 397. lM. Pratt and G. Swinden, 1 Chern.Soc. Chern. Cornrnun, (1969) 1321. H.I. Schelsinger, H.L. Brown, AD. Finholt, 1 Gilbreath, H.R Heekstra and EX Hyde, 1 Am. Chern. Soc., 75 (1953) 215. Y. Nitta, T. Irnanaka and S. Teranishi, Bull. Chern. Soc. Jpn., 53 (1980) 3154. T. Nguyen and H.H. Ghaffarie, J. Chern. Phys. 76 (1979) 513. 1. Ravet, J.Riga, 1 Verbist, E.G. Derouane and J. B.Nagy, in preparation. P.C. Hiernenz. Principles of colloid and surface chemistry, Marcel Dekker, New York, 1977, pp.233, 234. RD. VoId and MJ. VoId. Colloid and interface chemistry, Addison Wesley, London, 1983, pp. 181-186. G.V. Sarnsonov, 1.M. Vinitskii. Handbook of refractory compounds, IFIIPLENUM, New York, (1980), 96. 1. Ravet, A. Gourgue and 1 B.Nagy, in KL. Mittal and P. Bothorel (Eds), Proc.Int.Symp. on Surfactants in Solution, Bordeaux, July 9-13,1984, Plenum Press, New York, in press.
517
DISCUSSION P. CANESSON : During the preparation of the catalyst, you add an excess of NaBH4' What is the behaviour of this excess, and do Na+ ions play any role during catalysis ? I. RAVET : We add an excess of reducing agent because it is concomitantly hydrolysed. In addition, the hydrolysis reaction is catalysed by the particles obtained. Hence, after the formation of the metal boride, the elimination of this excess is very rapid and no more NaBH4 is present in the medium during the hydrogenation tests. Concerning the possible role of the Na+ ions, we have no indication but the hydrolysis of NaBH4 also introduces B02 in the solution. This component is adsorbed on the catalysts and causes a decrease of their activity. Fortunately, we can eliminate the B02 ions from the particles by neutralizing the metaborate with an acid solution and by subsequent washing of the particles. J.-P. PUTTEMANS : Have you measured the BET or metallic surface area of your catalysts? I. RAVET : Not yet. The BET measurements will be carried out in the near future. We shall compare the BET surface area before and after metaborate and CTAB elimination. J. SCHEVE: Did you observe an effect of stirring because this should influence the diameter of the water cores ? I. RAVET : The diameter of the water cores only depends on the composition of the micellar system. However, we observed an effect of stirring on the particle size. Indeed, the proportion of water cores reached by the reducing agent before the rearrangement of the micellar system can occur is influenced by the stirring and consequently the number of nuclei formed is different. This variation is reflected on the value of the F factor.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
519
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
PREPARATION OF NICKEL CATALYST FROM NICKEL CONTAINING CHRYSOTILE
Y. ONO, N. KIKUCHI and H. WATANABE Department of Chemical Engineering, Meguro-ku, Tokyo 152 (Japan)
Tokyo Institute of Technology,
Ookayama,
SUMMARY Nickel-containing chrysotiles (NixM93_x(OH)4Si205) were prepared under hydrothermal conditions. They showed the catalytic activity for ethylene dimerization after their treatment under vacuum at high temperatures, which leads to the dehydroxylation of the chrysotiles. The active sites for the dimerization are plausibly nickel cations exposed to the surface. Nickel metal catalysts were prepared by reducing the chrysotiles with hydrogen. The catalytic activity for ethylene hydrogenation at -50 0C increased with the reduction temperature up to 450 0C, and decreased by further increase of the reduction temperatures, though the extent of Ni (II) reduction monotonically increased with the temperature. Electron micrograph study revealed that the particle size of Ni metal increased with the reduction temperature, indicating that the lowering of the hydrogenation activity at higher reduction temperature can be ascribed to the lowering of the metal surface area due to the growth of the metal particles. Very uniform nickel particles can be prepared especially from the chrysotile of x = 0.6.
INTRODUCTION Chrysotile M93(OH)4Si05 is a layered magnesium silicate and can be synthesized under hydrothermal conditions (ref.
1).
Because of its defined crystal
structure together with its high surface area, chrysotile may be useful as a catalyst or as a support for a catalytically active species. Suzuki and Ono (ref. 2) studied the reaction of 2-propanol over chrysotile and discussed on the acid-base property of the surface. Catalytically active materials could also be obtained by incorporating metal ions like Ni(II), Co(II) and Al(III) in place of all or part of the Mg(II) ions in the structure. Such applications has been reviewed by Swift (ref. 3). Kibby et a L, (ref. 4) reported on the preparation of nickel metal catalysts by reducing nickel containing chrysotile. Jacobs et al. (ref. 5) studied the catalytic property of the reduced garnierite (a nickel-containing mineral with chrysotile structure). In this work, (1) chrysotiles in which part or all of the Mg(II) ions were replaced by Ni(II) ions were prepared and their catalytic activities for the dimerization of ethylene were studied, and (2) nickel metal catalysts were prepared by reducing Ni(II)containing chrysotiles with hydrogen and their catalytic activities for hydrogenation of ethylene and carbon monoxide were examined.
520 METHODS Synthesis of Ni-containing chrysotiles Nickel-chrysotile (Ni3(OH)4Si20S) was synthesized according to the method described by Noll (ref.
and Sholten et a L, (ref. 6). Sodium metasilicate
1)
(Na 2Si0
of 94.7 g and nickel sulfate (NiS0 46H 20) of 141 g was dissolved 39H20) 3 into 600 cm of water. To this solution, 19.1 g of sodium hydroxide was added and the mixture was stirred. The resulting gel of pH 11.8-12.0 with the composition of Ni/Si
loS and H20/Si ~ 118, was then placed in an autoclave. The temperature was raised to 295 0C in 4 h and kept at 295 0C for 8 h. After cooling, the insoluble product was washed and dried at l80 oC. The materials ~
with both Ni and Mg cations (NixM93_x(OH)4Si20S) were synthesized similarly except that the part of nickel sulfate was replaced by magnesium chloride
Catalytic reactions The treatment of the chrysotile samples and the catalytic reactions were carried out with a conventional gas-recirculation system. During the reduction of the chrysotiles with hydrogen, the water produced was trapped in aU-shape tube. The extent of the reduction was estimated from the pressure drop of hydrogen. The rate of ethylene hydrogenation or carbon monoxide was determined also from the pressure drop of the system. RESULTS AND DISCUSSION Structure of the synthesized materials X-ray diffraction patterns confirmed that the synthesized materials had the chrysotile structure. Fig. lea) shows the electron transmission micropraph of Ni-chrysotile. The hallow tube form of the chrysotile is clearly seen in Fig. lea). The length of the tubes ranges from 650 to 11000 A, mainly ZOOO-3000 A.The outer and inner diameter of the tubes are 250-400 A and 70-110 A, respectively. Fig. lea) also shows the layer structure of the nickel-chrysotile. The synthesized materials with Ni/Mg
~
1 and Ni/Mg
~
1/4 showed essentially identical X-
ray diffraction patterns to Ni-chrysotile and had similar physical forms. The BET surface areas of Ni-chrysotile and the materials with Ni/Mg 63.2, 59.5 and 47.6 m2g- l, respectively.
~
1 and 1/4 are
Heat treatment of nickel crysotile To know the thermal stability of the prepared nickel chrysotile, the thermal analysis was made. The weight loss of the sample with an endothermic peak was observed in the temperature range of 500-6000C. The weight loss corresponded with the loss of all the hydroxyl groups as water. The X-ray diffraction pattern was examined for the samples heated under vacuum at various tempera-
521
Fig. 1. Electron micrograph of nickel-chrysotile. (a) as prepared. (b) after heating under vacuum for 2 h at 700°C.
522 tures for 2 h. The sample treated at 400 0C gave similar pattern as the original one,
while the pattern due to chrysotile structure was completely lost for the
samples treated above 500 0C, Fig.
in conformity with result of the thermal analysis.
l(b) shows the electron micrograph of nickel chrysotile after heat treat-
ment at 700 0C for 2 h under vacuum. As mentioned above, this sample showed no X-ray diffraction lines due to the chrysotile structure. However, Fig. l(b) clearly shows that the tubular form of the material was mostly retained and even the layer structure of the material was also kept although it was more or less disordered. The samples with Ni/Mg
1 and Ni/Mg
1/4 behaved very similarly with
nickel chrysotile.
Dimerization of ethylene The reaction of ethylene over Ni-chrysotile was carried out at SOoC in a conventional gas-recirculation system with initial pressure of 47.6 kPa.
Prior
to the reaction, Ni-chrysotile was heated under vacuum for 2 h at varying temperatures. The main reaction is the dimerization of ethylene,
hexenes being
also formed over the samples heated over 450 0C. Among the butenes produced, trans-2-butene is most predominant over the samples treated at 300 and 400 0C. On the other hand, 450
0C,
I-butene was most predominant over the samples treated over
indicating that the mechanism of the dimerization depends on the heat-
treatment temperature of Ni-chrysotile. Over Ni-chrysotile treated at 500 0C, the butene distribution changed with reaction time. Thus, the fraction I-butene decreased gradually with reaction time and that of 2-butenes increased instead. The activation energy of the dimerization was estimated to be 20.0 kJ mol- l from the temperature dependence of the initial rate over the sample treated at 500 0C. Fraction of I-butene among butenes decreased with the reaction temperature, indicating the increase in the rate of butene isomerization with temperature. The initial rates of the reaction calculated from the pressure drop-time curves were plotted as a function of the temperature of the heat treatment in Fig. 2. The rate increased sharply in the temperature range of 450-600 0C. Since the surface area did not change appreciably by the heat treatment (Fig. 2),
the
change in the activity is not caused by the change in the surface area. The fact that the temperature range of the activity rise agrees with the temperature of the dehydroxylation of the chrysotile, indicates that the activity rise may be caused by the exposure of nickel cations with unsaturated coordination. It is well known that nickel species in zeolites or on silica surface are active catalysts for the dimerization of ethylene (refs.
7,S).
In order to obtain the information on the active centers for the dimerization,
several experiments were performed. The reaction of ethylene was carried
523 out at so?c with ethylene containing varying amount of carbon monoxide. The rate of dimerization decreased as the increase in the amount of carbon monoxide introduced into the reaction system. The activity for the dimerization was almost completely lost with the introduction of 3 x 10- 6 mol of carbon monoxide to 1 g of Ni-chrysotile. Provided that one molecule of carbon monoxide poisons one active center,
the ratio of the number of active centers to the number of
nickel atoms is 3 x 10- 4, indicating that the very small fraction of nickel cations located on the surface of the chrysotile can be active centers for the dimerization. As for the active species, Yashima et al. (ref. 7) reported that highly dispersed nickel metal gives active centers in nickel(II)-exchanged Y-zeolites, while Bonneviot et al. (ref. 8) reported that Ni(I) cations are responsible for the dimerization. The effect of the reduction of the catalyst was examined. Thus,
Ni -chrysotile pretreated at 600 0 e was exposed to hydrogen at 3000C for
varying periods. From the hydrogen consumed, the extent of the reduction of nickel(II) cations was estimated. Fig. 3 shows that the rate of ethylene dimerization at SOOC decreased only slightly with the extent of the reduction. On the other hand,
the rate of hydrogenation of ethylene at OOC increased sharply
with the extent of the reduction, indicating that the active centers for the dimerization and those for the hydrogenation are entirely different. Since the active centers for the hydrogenation are plausibly metallic nickel,
'0'1
.
15
/
-'
'e
E oE 10
!D' I
0 .-
-..... 5 (»
e 'c
....
f -o-o-i-o-t::o-.
,
0'1
100 r--r
-Em
•
(» L
50
•
ra.-
0
_I
the possi-
u
m =' If) L
0 300 400 500 600 700 Evacuation Temperature I °C
Fig. 2. Change in the initial rate of ethylene dimerization and the surface area of nickel-chrysotile with the temperature of the heat treatment.
524 bility of the participation of metallic nickel to the dimerization can be eliminated. The catalytic activities for ethylene dimerization over chrysotiles with Ni/Mg = 1 and 1/4 were also studied. The conversions of ethylene were 2.8, and 9.3% for chrysotiles of Ni/Mg
=
1 and Ni/Mg
=
5.5
1/4 and Ni-chrysotile,
respectively, after 1 h reaction at 50 0 e with the initial ethylene pressure of 40.7 kPa and 1 g of catalyst.
Reduction of nickel chrysotile Nickel chrysotile dehydroxylated at 600 0e was reduced at varying temperatures with hydrogen (initial pressure of 15.1 kPa) in a gas-recirculation system with a cold-trap,
which trapped water formed by the reduction. The rate
of the reduction was fast at the beginning and it slowed down at the later stage. In Fig. 4 was shown the temperature dependence of the degree of the reduction after 2 h-exposure to hydrogen. The higher the reduction temperature, the higher was the degree of the reduction. Thus, at 300 o e, the rate of the reduction was very slow. The degrees of the reduction at 400, 480 and 600
0e
were 20, 63 and 94%, respectively. The rate of the reduction depended very much on the content of nickel in the chrysotile structure. Thus, the rate of the Ni(II) reduction was much slower on Ni-Mg-chrysotile (Ni/Mg the degree of the reduction was 66.5% after
=
1/4),
for which 2 h-reduction at 600 o e. This
indicates that magnesium(II) cations in the chrysotile structure stabilize the nickel cations or the chrysotile structure and serve to inhibit the reduction of nickel (II) cations in the structure.
Hydrogenation of ethylene Hydrogenation of ethylene over reduced Ni-chrysotile (50 mg) was carried out at -50 o e with an equimolar mixture of hydrogen and ethylene (4.0 kPa each). Nichrysotile was reduced at various temperatures for 2 h with the initial hydrogen pressure of 15.1 kPa. The initial rate of the hydrogenation was plotted as a function of the temperature of the reduction in Fig. 4, where the degree of the reduction was also shown as described above. Though the degree of the reduction increased with the temperature of the reduction, the rate of ethylene hydrogenation went through a maximum around 4800 e of the reduction temperature, where the degree of the reduction was 63%. Fig. 5 shows the change in the catalytic activity per reduced nickel as a function of the reduction temperature. It decreased monotonically with the reduction temperature, indicating that the growth of nickel metal particles with the reduction temperature. Similar experiments were carried out over Ni-Mg-chrysotile (Ni/Mg = 1/4). this case, the activity for
hydrog~nation
In
increased with the reduction tempera-
ture. The activity per reduced nickel for Ni-Mg chrysotile was shown also in
525 Fig. 4. The activity per reduced nickel for Ni-Mg-chrysotile does not differ greatly from that for Ni-chrysotile, indicating that the state of the metal particles does not depend greatly on the starting material, but depends more on the temperature of the reduction. Hydrogenation of carbon monoxide Hydrogenation of carbon monoxide over reduced nickel chrysotile was carried out in a gas-recirculation system. Nickel chrysotile was reduced with hydrogen (initial pressure,
26.7 kPa) at varying temperatures. The reaction temperature
was 300 oC, and the initial pressures of carbon monoxide and hydrogen were 6.7 and 13.3 kPa, respectively. Water produced was removed from the system with a cold trap. In contrast to the hydrogenation of ethylene, the activity for carbon monoxide hydrogenation increased monotonically with the reduction temperature. The ratio of the activity per reduced nickel for the samples reduced at 600 0C and for those at 400 0C was 0.6, much larger than the ratio of 0.15 in ethylene hydrogenation. This indicates that the active centers for the hydrogenation of carbon monoxide was different from those for the hydrogenation of ethylene. It has been reported that the hydrogenation of ethylene is a structure-insensitive reaction and that the hydrogenation of carbon monoxide is structure-sensitive.
10
10'5
Ie
-- E4 ~, C
.
0(5
.- E
CUlO. 3
I-
......
~
~
2
0
'0 ~
m 0 0:: ~
I-
8
e_
~o
4
g~
1J
>-
0
0
--o t:E.
6 ~§ Ci·
I
I
b C
of
Nb
E
_.~o
->: -==t-
0 ......
::r:
2 '0 ~ ~
I
4 1 2 3 Ni (II) Reduced I %
5
m 0 0::
Fig. 3. The effect of nickel(II) reduction on the catalytic activities of nickel-chrysotile for the dimerization and the hydrogenation of ethylene. Dimerization: SOoC, C2H4 = 40.7 kPa. Hydrogenation: OOC, C2H4 = H2 = 13.3 kPa.
526
~
o
4
100 §
E
M'
'a
--~
2
......
Fig. 4. The dependence of the initial rate of ethylene hydrogenation and the degree of Ni(II) reduction on the reduction temperature of Ni-chrysotile. Reduction: initial hydrogen pressure 15.1 kPa, 2 h. Hydrogenation: -SOoC, ini tial pressures, C2H4 = H = 4.0 kPa. 2
o
-
-
I-
OL-------l...------JL..-.--~
300
400
500
600
Reduction Temperature foe Fig. 5. The dependence of the initial rate of ethylene hydrogenation per reduced Ni(II) ions on the reduction temperature of Ni-chrysotile. Reduction and reaction conditions are described in Fig. 4.
527
Fig. 6. Electron micrographs of (a) Ni-chrysotile and (b) Ni-Mg-chrysotile (Ni/Mg = 1/4) reduced at 500 0C for 2 h.
528
Electron Micrograph of Nickel Particles The nickel-containing chrysotile samples after reduction was examined by their electron micrograph. Though the nickel-chrysotile reduced at 3000C showed the catalytic activity for ethylene hydrogenation (Fig. 3), nickel-metal particles were not observed in the electron micrograph. The nickel-chrysotile after reducion at 350°C for 2 h gave the nickel particles with the diameter of 15-20 A ,
the number of particles being very small. The size of nickel particles and
the number of particles increased with increasing reduction temperature. Thus, the average particle sizes of nickel metals formed after 2 h reduction were 34, 36 and 82 A for the reduction temperature of 400, 470 and 600 0C, respectively. Fig. 6(a) shows the electron micrograph of nickel chrysotile after 2 h-reduction at 500 0C. While a large number of nickel particles were observed, it is also clear that the physical form of the original chrysotile was not completely retained. Ni-Mg chrysotile (Ni/Mg = 1/4) after reduction at 500°C gave the nickel-particle size similar to Ni-chrysotile reduced at the same temperature (Fig. 6(b)). Here, however, the particle size distribution is more uniform than that of nickel-chrysotile reduced at 500 oC,
and the physical form of the
chrysotile was maintained. Thus, it is concluded that very uniform metal particles can be prepared on the support with a very uniform physical shape. It is no doubt that the size of nickel particles can be varied by changing the content of nickel in the chrysotile structure and the conditions of conducting reduction of nickel cations.
REFERENCES 1 2 3 4 5 6 7 8
W. Noll, H. Hirscher and W. Syberts, Kolloid Z., 157 (1958) 1-11. S. Suzuki and Y. On o , Applied Cata1., 10 (1984) 361-368. H. E. Swift, in J. J. Burton and R. L. Garten (Eds.), Advanced Materials in Catalysis, Academic Press, New York and London, 1977, Ch. 7, p.230. L. L. Kibby, F. E. Massoth and H. E. Swift, J. Cata1., 42 (1976) 350-359. P. A. Jacobs and H. H. Nijs, J. Catal., 64 (1980) 251-259. J. J. F. Scholten, A. M. Beers and A. M. Kie1, J. Catal., 36 (1975) 23-29. T. Yashima, Y. Ushida, M. Ebisawa and N. Hara, J. Catal., 36 (1975) 320-326. L. Bonneviot, D. Olivier and M. Che, J. Mol. Catal., 21 (1983) 415-430.
529
DISCUSSION J.A. SCHWARZ: You have used reduction temperature to correlate your activity data for ethylene hydrogenation. Most of your interpretation has been directed toward the structural changes that occur during the heat treatment. Could you comment on the possibility of compound formation during H2 treatment at elevated temperatures ? Y. ONO : The change in the catalytic activity for ethylene hydrogenation seems to be well correlated to the dispersion of the nickel particles formed. Since the starting material of our catalysts is nickel silicate, we simply presume that unreduced nickel cations still remain in some form of nickel silicates. G. BELLUSSI : During the reduction of Ni-chrysoti1e at a temperature higher than 400°C, probably a phase transformation or a loss of crystallinity of the Ni-chrysoti1e takes place. Have you tried to correlate these variations with the data of activity in the hydrogenation reaction? Y. ONO : Before the reduction, the Ni-chrysoti1e was pretreated at 600°C. At this stage, X-ray diffraction of the chrysotile was completely lost, although the shape of the original crystals was still retained (Fig. l(b)). By reducing the material at temperature above 500°C, the original shape was also lost as seen in Fig. 6(a). It seems, however, that the activity can be well correlated with the change in the particle size and the extent of Ni(II) reduction. We cannot find any positive evidence for correlating the structural change in the support material with the change in the catalytic activity for hydrogenation. D. REINALDA : By reduction of Ni, are there acid sites created in the lattice? Y. ONO : Though we have not measured the acidity directly, the data for the catalytic activity for isopropy1alcohol dehydration suggest that the material is not so acidic.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
531
PARAMETERS INFLUENCING THE PREPARATION AND CHARACTERIZATION OF SODIUM ON ZEOLITE CATALYSTS L.R.M. MARTENS. W.J.M. VERMEIREN. P.J. GROBET and P.A. JACOBS Laboratorium voor Oppervlaktechemie. Katholieke Universiteit Leuven. Kardinaal Mercierlaan 92, B-3030 Leuven (Heverlee) (Belgium)
SUMMARY Optimization of a sodium on support catalyst for room temperature isomerization of butenes is reported. In situ ESR spectroscopy of the working catalysts allows to discriminate between the different sodium species present. Sodium azide impregnation using a methanolic slurry on zeolites L. X. Y and alumina followed by slow thermal decomposition of the adsorbed az tde , turned out to be tne preferred method. Thermoanalytic, ESR and IR characterization of the catalyst showed that neutral sodium clusters "glued" to the wall ot the support through a chemical interaction between residual nitrogen and support oxygen.
INTRODUC I ION In basic catalysis the reaction intermediates are of the carbanion-type. the stability of which increases as follows tertiary
<
secondary
<
primary
It is evident that in this way different selectivities can be reached compared to acid catalysis. As an example oligomerization and codimerization of olefins is relatively easy and side-chain alkylation of alkylaromatics is possible (ref. 1). For olefin isomerizations. double bond shift is much faster than cis-trans isomerization (ref. 2). Hydrogenation of unsaturated compounds is possible as well (ref. 3). The catalyst activity is very much sUbstrate dependent. Molecules containing allylic or benzylic hydrogen atoms are very acidic and undergo rapid isomerization or are excellent substrates for addition reactions (ref. 4). Oxides like MgO, ZnO. CaO or their mixed oxides clearly show basic behaviour (ref. 5). Supported alkali metals have been used as heterogeneous basic catalysts as well (ref. 2). Alkali metals on alumina and zeolite supports can be prepared by a metal vapour adsorption method or by impregnation in liquid ammonia. Both methods are industrially unattractive and suffer from major draWbacks such as instability during transport and easy sublimation of the excess alkali metal. We reported on the impregnation of
532
sodium azide from a methanol slurry on a dry and non-acidic zeolite (ref. 6). The azide is thermally decomposed in situ in the catalytic reactor in order to overcome the above mentioned drawbacks. These sodium-on-zeolite catalysts exhibit typical base catalyzed selectivities as derived from butene isomerization (ref. 7). In such catalysts ESR spectroscopy distinguishes between three kinds of sodium an extra-lattice metal phase (Na~), intracrystalline Na~ clusters and finally ionic clusters which in case of (ref. 7). A positive correlation was found between Y-type zeolites are Na~+ the concentration of the intra-crystalline sodium clusters and the activity for double bond shift and cis-trans isomerization (ref. 7). It is also known that the concentration of the respective sodium species is dependent on the rate of the temperature increase during the azide decomposition and on the structure of the zeolite. On zeolite L and on y-alumina mainly neutral clusters are found in the support pores, while on NaY zeolites ionic clusters are formed as well, located in the sOdalite cages (ref. 7). After slow heating during the azide decomposition step the neutral intracrystalline species are highest in concentration, while during fast heating the formation of charged clusters is favoured. The chemistry involved in the formation of charged clusters is clear and the arguments advanced in recent studies are consistent. In contrast to early interpretations (ref. 8), it is now clear that charged clusters are only formed in zeolites containing sodalite cages (ref. 6,7,9). In zeolites were no other species are present also no basic catalytic activity is found (ref. 7), which is in line with their location 1n hidden sites of the lattice. In Y-type zeolites Na~+ clusters are present (ref. 8) which according to computer graphics (ref. 13) fit very well in the sodalite cages. They are formed by charge transfer from sodium metal to the 4 Na+ cations already present in these small cages (ref. 9) : ---~~
3+ + Na4 + Na
(1)
In the present work, it was aimed to determine systematically all parameters which can influence the concentration of the respective sodium species and to compare the azide loading method with other more classical preparation methods. The active species on an optimized catalyst was also characterized as far as possible.
533
METHODS Materials The catalyst supports used were the zeal ites NaX, NaY and L from Union Carbide. y-Alumina was Spheralite SAP350 from Rhone-Poulenc with an average pore size of 2.5 nm and a specific surface of 350 m2g- 1 • All supports were contacted at 343 K with an excess of an aqueous sodium dithionite solution to remove iron impurities which obscure the ESR spectra. The samples were dried at 343 K and calcined at 723 K in case of the zeol ites and at 873 K for y-alumina. To the dry supports was then added a slurry of reagent grade sodium azide in methanol so as to fill completely the pore volume of the supports. This material is dried at room temperature and stored in absence of moisture. Azide decomposition is carried out in situ in the catalytic reactor. First all traces of the solvent methanol are removed at 523 K under a purge of dry helium. Then the temperature is raised up to 673 K at a rate of 3 K per minute, which is close to the slow azide decomposition rate reported earlier (ref. 7).
Technigues The catalytic reactor consists of a quartz tubular reactor, the exit of which is connected to a gaschromatograph. This tubular reactor fits in the cavity of an ESR spectrometer. The description of this particular set-up has been reported earlier (ref. 7). The ESR measurements (g 6HZ) were made at room temperature unless mentioned otherwise. Infra-red (IR) measurements have been made on a PE 580B grating instrument connected to a data station and computer. Adsorption of gases was also carried out in-situ using an IR cell mounted in the spectrometer. The gases and vapours used were of highest purity. They were further passed through an oxygen and water filter to depress the impurity content down to 0.5 ppm. Thermoanalytic experiments were performed on a Mettler thermoanalyzer connected to a computer. The heating rate was 4 K per minute while the balance was being flushed with ultra-pure helium at a flow of 10 liter per hour. Inert alumina was used as DTA reference. The amount of sample used was 25 mg.
RESULTS AND DIS~USSION Comparison of the azide method with classical preparation technigues Obvious methods with whiCh sodium metal can be loaded on supports such as alumina and non-acidic zeolites are the fol lowing:
534
i. vapour deposition by heating of the bulk metal (ref. 8,9); ii. impregnation with a solution of alkali metal in liquid ammonia (ref. 2); iii. thermal activation of adsorbed sodium azide (ref. 7). ESR spectra of the different Na-species present on NaY after using these three methods, are shown in Fig. 1.
9 =2.078
A
19=2.003 X40 _
1 B
-1
C
500G
t---l
t
X40
Xl
Fig. 1. Comparison of the ESR spectra of Na-species on NaY zeol ite obtained after heating at 673 K of a catalyst made by A, impregnation of Na in liquid ammonia; B, vapour deposition and C, decomposition of adsorbed sodium azide. In each case, 7 % by weight of Na per 9 of dry zeolite was used.
On the three catalysts the common ESR signals reported for Na on NaY zeolites (ref. 6-9) are observed. Na~ particles are mainly formed on the catalyst made in liquid ammonia, while via vapour deposition all three species are obtained. With the azide method few extra-lattice particles can be observed and the intensity of mainly Na~, the active clusters for e.g. butene isomerization, is enhanced considerably. Moreover, the first two methods give a sodium mirror in the quartz reactor, which is not the case when the azide method is used.
Optimization of the impregnation of sodium azide In order to optimize the concentration of the intracrystall ine neutral clusters impregnation of an azide slurry in water or methanol was compared with a method consisting in grinding together the dry solids lsodium azide and zeolite NaY). Representative ESK spectra are shown in Fig. 2.
535
9 =2.078 A
9=2.003
X40
8 X 40
c
Xl
100 G I-----l
Fig. 2. ESR spectra ot Na species after calcination at 673 K of 3.80 mmol sodium azide on 1 g of NaY zeolite obtained by: A, impregnation with water and C, methanol as solvent, and B, grinding together dry NaN 3 and NaY.
Using the impregnation with methanol the intensity of the catalytically interesting Na~ signal is at least 5U times more intense than in the other two cases. The water impregnation method gives the lowest concentration of sodium species, which is possibly the result of a hydrolysis reaction of the azide during impregnation and/or activation. The co-grinding gives an unusually high concentration of charged clusters hidden in the sodalite cages. It is evident that the presence of water and subsequent hydrolysi s of the azide has to be avoided during the preparation of the catalyst.
Decomposition of adsorbed azide Fejes et a1. (ref. 10-12) have introduced the use of sodium azide in zeal ite science to neutral ize residual acidity after its thermal decomposition. In all cases either physical mixing of the sol ids or impregnation 1n water was used (ref. 10). The same authors (ref. 12) also advance certain arguments from which they conclude the presence of at least two residual azide ions after high temperature decomposition. Thermoanalytica I curves of the decomposition of NaN 3 on alumina and zeolite NaY is shown in Fig. 3. The exothermic decomposition of the azide is clearly visible at 685 K for alumina. For NaY the decomposition is more complex but occurs in the same temperature range. Comparable thermoanalytic curves have been obtained by Kiricsi et al ,
536
(ref. 12) for the two already mentioned methods of impregnation.
10,....-----,.------------,-------------...., A a
B
Cl
E
A
b
B
<, 1Il 1Il
o
...I
5
I:I:
e
W
~
EXO
c
c
ENDO
700
300
TEMPERATURE/ K
Fig. 3. Thermogravimetric (A), DTG (B) and OTA (C) curve of 3.85 mmol NaN per g of alumina (a) or zeol ite NaY (b), prepared according to the met~anol impregnation method.
The figure also shows the clear separation between the methanol desorption and the weight loss associated with release of nitrogen from the azide. In this way the amount of nitrogen left in the system can be determined accurately (Table 1).
TABLE 1 Thermoanalytic characteristics of NaN 3 decomposition. Support
Position of DTA exotherm/K
% of original
alumina a NaY(49)a NaY(56)a NaX(8b)
685 661 665 671 659 673
33.4 O.U 18.0 21.3 31.9 32.3
L
Cayb
Residual N
~ Values in brackets refer to the number of Al atoms per unit cell. 2+ NaY(56) exchanged in reflux conditions, containing, containing Ca for 90 %
of the exchange capacity.
The data show that depending on the nature of the support the amount of residual nitrogen left in the catalyst varies considerably. On alumina,
537
zeolites L and CaY approximately one third of the nitrogen is not released. This corresponds closely to the stoichiometric requirements of the following reaction liT
(2)
This stoichiometry is closely followed when no sodalite units are present in the substrate as with zeolite L and alumina or when these units are devoid of cations such as in CaY (ref. 14) and thus when no charged cl usters can be formed. It is indeed found experimentally that on zeolite L or alumina (ref. 7) as well as on CaY the hyperfine 13 lines ESR spectrum is never obtained. The lower amount of N retained in NaX and Y zeolites can be understood in terms of the preferential formation of charged clusters according to equation 1. The decrease of the residual N content with decreasing aluminium content of the faujasite type zeolites X and Y is clearly established in Table 1. In Table 2 the line widths for the Na~ cluster signal are given for the various supports used.
TABLE 2 Support dependence of the line width of the ESR signal with g = 2.003 for azide loadings of 3.85 mmol/g. support
Signal width/g
NaX(85) NaY(56) NaY(49)
16.00 6.35a 5.60 4.25 3.90
L
alumina a At 110 K.
The table shows a significant variation of the signal width and consequently ot the cluster size with the nature of the support. A size effect on the conduction ESR resonance of small sodium particles in solid NaN 3 has been reported (ref. 15.16). The experimental line width was found to increase with decreasing particle size from 100 to 1.0 nm (ref. 15). This suggests that for faujasites the size of the supercage clusters increases with decreasing Al content of the framework. On zeolite L and alumina this cluster size sti 11 increases further. The change of the residual N content (Table 1) is in no way
538
related to the size of the neutral clusters lTable 2). Therefore, the change in the former values for faujasites must be exclusively related to the change in the number of charged clusters. In this context, this number has to increase with decreasing Al content. Experimenta'lly it is not obvious to confirm this using ESR spectroscopy in a quantitative way. From all this it has to be concluded that the overall composition of the catalytically actlve Na species is identical and Na and N are present in equal atomic amounts. The occlusion of un reacted azide in the sodalite cages as suggested earlier (ref. 12) is not very probable since on every support investigated the same ratio ot Na to N is found.
Isomerization activity The base catalyzed isomerization of cis-2-butene into its normal isomers should increase with the line width of the signal at g = 2.003, provided all previous hypotheses are true. At constant 1ine width, which is found for a given support, the activity has to increase with the azide loading. This is indeed what is experimentally is observed (Fig. 4).
'#-
100
<,
e o
.
80
III
G)
>
60
o
40
e o
20
O ......._...J....._--L_ _. L - _........
_~
o
2
3
Si/AI
4
5
Fig. 4. Conversion of c-2-butene at 296 Kover Na-faujasites with different Al contents loaded with a, 3.85; b, 3.07; c, 2.31 and d, 1.54 rnmol NaN 3 per g of zeolite.
At any azide loaQing, NaX zeolite is the most suitable support. Its activity progressively increases with increasing azide contents. The activity observed for zeolite L and alumina corresponds to that expected from the sequence of line widths (Table 2).
539
Interaction of the neutral clusters with the support After thermal activation of NaN 3 on NaY and CaNaY (ref. 12) it was found that IR absorption still occurs in the region where N-N stretching of the azide is expected. The cited work shows that the original stretching mode of azide (at 2090 cm-1 ) gradually disappears while two new peaks reappear (at 2170 and 2030 em-I). The authors tentatively attributed this to occluded azide which is thermally stable and is accommodated in different locations of the structure. The 2030 cm- 1 band is attributed to supercage azide and the 2090
a
b
723 K
t ~
....... r:: 0 III
10 "10
E
1
o
III
r::
..
l'Il
~
2112
2400
2000
2400
2000
Wavenumber / cm- 1 Fig. 5. Room temperature IR spectra of NaN 3 on different supports during thermal activation.
cm- 1 band to extra zeol itic azide. The 2170 cm-1 band represents then azide occluded in the sodalite cages. The IR spectra of the present catalysts at two stages of their thermal activation are shown in Fig. o. Essentially the same features are obtained for three supports, alumina, NaY and CaY. The i ntens i ty of the bands decreases after advanced outgass i ng but the residual bands are not of negligible intensity. Since on alumina and zeolites the same phenomenon is observed, the attribution of the bands to specific
540
locations of undecomposed azide as advanced for zeolite Y (ref. 12) has to be rejected. A more plausible explanation is to assign them to the N-O vibration of either 1inear (2181 cm- 1) (species 1) or bridged (2 low frequency bands) (species II) N-O formed through chemical interaction of residual N with oxygen from the support :
Al - 0 - N - Na(Na n) II
In this picture the neutral cluster is "glued" to the support surface giving it an unusual thermal stability.
CONCLUSIONS Out of three methods which can be used to prepare sodium-on-support catalysts, the azide impregnation method after slow thermal decomposition of the adsorbed azide gives the highest concentration of neutral sodium clusters, located in the supercages and characterized by an ESR signal at g = 2.003. The method can be optimized when a methanolic slurry of azide is added to the dry zeolite, thus avoiding any azide hydrolysis prior to its thermal decomposition. Using the latter method, thermoanalysis shows that in case of alumina and zeolite supports, whose structure does not contain sodalite cages and consequently on which no charged sodium clusters can be formed, one third of the nitrogen of the azide is not released. When sodalite units are present, less nitrogen is kept in the system as a consequence of the formation of charged clusters in which only charge transfer from the metal to the hidden cations but no nitrogen from the azide is involved. The ESR line/width of the signal at g = 2.003 is substrate depending. The more aluminium is in the faujasite-type zeolite, the larger is the line width and consequently the smaller are the corresponding clusters. A positive correlation is further obtained between the base catalyzed butene isomerization activity and the width of the ESR signal. Infra-red spectra suggest that residual nitrogen associated with the neutral clusters interacts chemicallY with oxygen of the support, thus glueing the neutral clusters to the support wall.
ACKNOWLEuliMENT This work is performed in the frame of a concerted action on shape selective catalysis, sponsored by the Belgian Ministry of Science Policy.
541
L. Martens acknowledge further a research grant and P. Grobet and P. Jacobs a research position from N.F.W.O.-F.N.R.S. (Belgium).
REFERENCES 1 H. Pines, in: H. Pines (Ed.), The Chemistry of Catalytic Hydrocarbon Conversions, Academic Press, New York, 1981, pp. 135-151. 2 H. Pines and W.M. Stalick, in: H. Pines and W.M. Stalick (Eds.), Base-catalyzed Reactions of Hydrocarbons and Related Compounds, Academic Press, New York, 1977, pp. 25-79. 3 K. Tanabe, in: B. Imelik, C. Naccache, G. Coudurier, Y. Ben Taarit and J.C. Vedrine (Eds.), Catalysis by Acids and Bases, Elsevier, Amsterdam, 1984, pp. 1-14. 4 D.J. Cram, in: D.J. Cram (Ed.), Fundamentals of Carbanion Chemistry, Academic Press, New York, 1965, pp. 1-47. 5 W.O. Haag and H. Pines, J. Amer. Chern. Soc., 82 (1960) 387-391. 6 L.R.M. Martens, P.J. Grobet and P.A. Jacobs, Nature, 315 (1985) 568-570. 7 L.R.M. Martens, P.J. Grobet, W.J.M. Vermeiren and P.A. Jacobs, in: Proc. 7th Int. Zeolite Conf., in press. 8 J.A. Rabo, C.L. Angell, P.H. Kasai and V. Schomacker, Disc. Faraday Soc., 41 (1966) 328-349. 9 M.R. Harrison, P.P. Edwards, J. K1inowski and J.M. Thomas, J. Sol. St. Chern., 54 (1984) 330-341. 10 P. Fejes, I. Hannus, I. Kiricsi and K. Varga, Act. Univ. Szegediensis, Act. Phys. Chern., 24 1978) 119-130. 11 P. Fejes,!. Kiricsi, 1. Hannus, T. Tihanyi and A. Kiss, in: B. lme1ik, C. Naccache, Y. Ben Taarit, J.C. Vedrine, G. Coudurier and H. Praliaud (Eds.), Catalysis by Zeolites, Elsevier, Amsterdam, 1980, pp. 135-140. 12 I. Kiricsi, I. Hannus, A. Kiss and P. Fejes, Zeolites, 2 (1982) 247-251. 13 S. Ramdas and J.M. Thomas, Chemistry in Britain, January 1985 (1985) 49-52. 14 W.J. Mortier, in: Butterworth Co (Publ.) Ltd., Compilation of Extra Framework Sites in Zeolites, GUildford, 1982, pp. 20-25. 15 M.A. Smithard, Sol. St. Comm., 14 (1974) 411-417. 16 D.A. Gordon, Phys. Rev. B, 13 (1976) 3738-3746. 17 P.A. Jacobs, F.H. Van Cauwelaert, E.F. Vansant and J.B. Uytterhoeven, J. Chern. Soc., Faraday Trans. I, 69 (1973) 1056-1068. 18 P.A. Jacobs, F.H. Van Cauwelaert and E.F. Vansant, J. Chern. Soc., Faraday Trans. I, 69 (1973) 2130-2139.
542
DISCUSSION B. NOTARI : I assume that the catalyst you have described must be used only in the absence of water and oxygen forming compounds like oxygenated products. Under these anhydrous conditions have you examined the thermal stability of the catalysts? This could be important for applications where themodynamics require operating at high temperatures. L.R.M. MARTENS: The activated Na-on-Y catalyst is very sensitive to H20, 02 and CO 2 in the feed resulting in the formation of NaOH, Na20 and Na2C03, respectively. In absence of those poisons, the catalytically active . Na O particles are stable up to 673 Kand retain their basic properties. At h{gher temperatures sodium is evaporated from the support. No evidence for sintering inside the zeolites cages is available. A. LA GINESTRA : Can you give more details about the influence on the catalytic activity in the butene isomerization by changing the Na-support or the thermal treatment ? L.R.M. MARTENS: The butene isomerization is catalyzed by metallic intrazeolitic sodium. The amount and the size of those particles is support dependent, e.g. on NaX a larger amount of smaller particles is formed in comparison to NaY. Reaction selectivity, however, is equal on all supports and is governed by the carbanion chemistry. The overall activity depends upon the number of active particles. Slow activation (2K/min) of the catalyst gives the highest amount of intracrystalline sodium particles, while for rapid ac (30K/min) more catalytically inactive extra-lattice Na particles and Na 4+ clusters are formed. P. RAJARAM : 1/ What is the life of the sodium azide loaded catalysts for the isomerization reactions? 2/ What is the effect of sodium azide loading on the mechanical properties of zeolite pellets?
3ivation
L.R.M. MARTENS: 1/ Deactivation of the Na-on-support catalyst for the isomerization of n-butenes at room temperature is only dependent upon the amount of poison (02, C02' H20) in the feed. The activity remains constant during at least 48 hours using feed gases with impurity content lower than 5 ppm. 2/ Sodium azide impregnation was performed on zeolite agglomerates prepared by pressing the zeolite powder in a laboratory press into flakes, followed by crushing and sieving of these flakes. The 0.25 - 0.50 mm agglomerates used in the bench-scale reactor (or in-situ in the ESR spectrometer) are resistant to this impregnation and activation procedure of the impregnated sodium azide. It is not yet known how zeolite extrudates would behave during this loading and activation operation. S. CARTLIDGE: 1/ Can you give some information on the distribution of sodium species between the super-cages and sodalite cages of zeolite Y? Which of the sites are catalytically active? 2/ How clear is the butene isomerization reaction in respect to coke formation and by-products ? L.R.M. MARTENS: 1/ At this moment, we are no~ able t8 give quantitative information on the amount of sodium used to form Na~+ or Nay clusters in the sodalite and supercages, respectively. With in§reaslng NaN30loading there is a more rapid saturation of the ESR signals of Na a+ than of Na~'OBy in situ ESR-spectroscopy of a working catalyst we found tnat only the Nay particles are catalytically active. 2/ The butene isomerization reaction at room temperature and atmospheric pressure is a clean reaction. Coke nor side-products like isobutene or oligomers are formed.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
543
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
PREPARATION OF BIMODAL ALUMINA AND OTHER REFRACTORY INORGANIC OXIDES -SUITABLE SUPPORTS FOR HYDROTREATING CATALYSTS K. ONUJIA
Research Center, Mitsubishi Chemical Industries Ltd. 1000 Kamoshida-cho, Midori-ku, Yokohama 227 (Japan)
ABSTRACT A novel method that can produce bimodal aluminas and other refractory inorganic OXides was developed. A selected carbon black which has a chained and reticulated structure was incorporated into, for example, boehmite gel. At the final stage of the preparation the carbon black was burned out in the presence of an oxygen-containing gas. The active alumina thus obtained was revealed to have bimodal pore distribution. The pore distribution could be varied according to the kind and the amount of carbon black. Further studies also proved that this method is applicable to other inorganic oxides. A catalyst prepared by using the bimodal alumina as support was shown to have high activity for hydrotreating of heavy oils, especially, for hydrodesulfurization and hydrodemetallation. The catalyst was also found to have a good durability. INTRODUCTION Activity, selectivity and durability are three major elements of catalyst performance. Though there are so many factors that influence them including active metal component itself, size and shape of metal, surface structure, interactions between metal and support and so on, pore distribution is one of the most important factors. The reason is that the pore gives not only paths for reacting and produced molecules but also sites for active metal's deposition. In addition, contaminated impurities and produced carbonaceous materials also deposit on the pore, which crucially affects the catalyst life. For catalyst producers, therefore, it is a great concern how to control the pore distribution of a catalyst. Methods of pore control may be devided into three categories. The first is to regulate particle sizes, for pore distribution is primarily related to interparticle spacings. A number of methods are proposed for it in the step of production of precursor OXides or hydroxides (ref. 1) , in the step of gelation (ref.2) and in the step of calcination. The second is to widen interparticle distances by decreasing surface tension(ref.3), by utiliZing vapor pressure of water when heated(ref.4) or by gasification of added materials(ref.5). The third is to give pores by
burni~g
out or by leaching
added organic
materials(ref.6). In the last case pore distributions are intrinsically bimodal,
544
for the pore derived from interparticle spacings is also present. It is not so easy to obtain such bimodal pore distribution in practice. One reason of it is that there are not appropriate additives which can give pore at a desired position. Another reason lies in the difficulty of making shaped and calcined body without changing other physical properties of the oXides and without weakening the mechanical strength of the body. Presumably the reasons mentioned above explain why the pores of the conventional so called "bimodal" aluminas are either broadly distributed or too large. Here a new preparative method of bimodal alumina and other refractory inorganic oxides is described. Characteristics of the method lie in the position and the sharpness of pore distribution, the capability of varying pore position and pore volume and the strength of the shaped body. Application of this bimodal alumina for catalyst support is also presented. EXPERIMENTAL Preparation of bimodal alumina Preparation of alumina is presented as a typical example of preparative method. Boehmite gel powder Pural SB (Condea Chemie) was mixed with carbon black. The mixture was kneaded with 4.3% aqueous nitric acid for 30 minutes in a batch type kneader. The kneading was continued for another 30 minutes after addition of 2.1% aqueous ammonium hydroxide. The resulting pasty batch was then extruded by a vacuum extrUder. The extrudates thus obtained were dried and calcined in an electric furnace at about 600'c in air. Care was taken as to keep away from abrupt increase of temperature by burning carbon black. Experimental procedures of hydro treating reactions Loading cobalt and molybdenum to alumina was achieved using a conventional impregnation technique. Hydrotreating reactions were performed either by an autoclave reactor or by a fixed bed flow reactor. The condition for autoclave reaction was; reactor volume 200 ml, catalyst 6.0g, oil 60g, pressure of H 2 2G, 160kg/cm temperature 360 or 390'c, reaction time 3hr. -1
The condition for flow reaction was; catalyst 100ml, LHSV 0.7 hr ,H 2/oil 3 2 850 Nm /kl, pressure of H 150 kg/cm G. In order to avoid back mixing and 2 channeling, the catalyst was mixed with equal volume of carborundum. The catalyst was presulfided before the reaction. Temperature was kept as to maintain 1.0% of sulfur in product oil. Used feedstocks were atmospheric distillation bottoms of !rabian light or Arabian heavy crudes. Concentrations of impurities were;
sulf~
3.0%, vanadium
31 ppm, nickel 6 ppm for Arabian light and sulfur 3.9%, vanadium 94 ppm, nickel 30 ppm for Arabian heavy.
545
RESULTS AND DISCUSSIONS Preparation of bimodal alumina Boehmite gel is well-known as an appropriate source of active alumina. A commonly used method of preparation of alumina is to add acid or acid and base to the gel for peptization. Then the peptized mixture is shaped. dried and calcined. The resulting alumina has pores with less than 10 nm radius (micropores). These micropores are essentially derived from interparticle spacings of alumina crystallites. The object of this research was to develop pores in the range of 10 to 100 nm radius (mesopores). After extensive studies it was found that a selected carbon black with specified size and structure is just suitable for the purpose (ref.7). The carbon black was incorporated into the gel at the first stage of formulation and finally burned out to leave voids within the alumina solid. A typical example of a pore distribution of the alumina thus obtained is drawn in Fig.l. Other than micropores. there is seen a large quantity of
TABLE 1 Physical properties of carbon black. Particle size
Carbon black
(nm)
A B
70 22 31 40 70
c
D E
DBP absorption (ml/l00g) 125 130 130 120 80
narrowly distributed mesoSurface ar~a
(m /s) 20 110 80 65 20
Particle size: measured from transmission electron microscope. DBP: dibutyl phthalate.
pores. A highly porous character of the alumina should be also noted. The pore distribution and the pore volume could be varied by the selection of kinds of carbon black and by controlling the quantity of carbon black. Table 1
D 1.5 <, oJ
E ..... CD
E
::::J 0
oJ
1.0
>
CD
L.
0
a, CD
0.5
>
+= 0 oJ
::::J
E
::::J
u
a
100
10
Pore
Radius (nm)
Fig. 1. Pore distribution of bimodal alumina.
1000
546
TABLE 2 Physical properties of alumina. Carbon black A B C D
E
Surface area 2/g) (m
Pore volume (ml/g)
274 277 274 268 272
Pore peaks
total
micro
meso
0.965 1.024 0.995 0.976 0.985
0.728 0.610 0.618 0.628 0.800
0.233 0.408 0.369 0.334 0.177
(nm) 6.4 5.4 6.0 5.5 5.7
20 13 15 15 12.5
Pore volume: total=3.75-7500 nm radius, micro=3.75-10 nm radius, meso=10-100 nm radius, Pore peaks: radius of pore peak in pore distribution curve, Amount of added carbon black is 30 wt%. shows the physical TABLE 3 Pore distribution of alumina. Carbon black (wU) 30 40 50 70
properties of several kindS of
Pore volume (ml/g) total 0.965 1.153 1.100 1.091
micro 0.728 0.705 0.640 0.497
Cabon black A is used.
Pore peaks
carbon black used.
(nm)
Pore distributions
meso 0.233 0.435 0.449 0.587
6.4 6.3 5.4 4.6
20 28 30 31
and other properties of the prepared aluminas are in Table 2. It is seen that the peak
positions of mesopore reflect the particle size of the used carbon black (except carbon black E). Amount of carbon black could also change the pore distribution and pore volume, as is shown in Table 3. By increasing the quantity of carbon black, the mesopore volume increases and the position of mesopore peak moves toward a larger radius. Further studies on the effects of carbon black were performed using Co-Mo/alumina catalysts. It was possible because loading cobalt and molybdenum to this bimodal alumina did not change the bimodal character basically, though the pore volume was decreased and the pore position was shifted to larger radius to some extent. Fig. 2 and Fig. 3 show the relations between amounts of carbon black and pore structures of catalysts. The same tendency as Table 3 was obtained. But it is worthwhile to note that more than 20-30 wt% of carbon black was required to obtain mesopore as distinguishable one. Taking into account the fact that micropore volume was maximUm at around 10 to 20 wt% addition of carbon black, it was concluded that when added small amount of carbon black almost all particles are present inside of alumina solid. Only when adequate amount of carbon black is added,it begins to appear on the surface of the solid. Fig. 4 shows the effects of carbon black particle size, which is in accord with the results of Table 2.
40 ,.... 20
,....
-'
E
<,
CI ....,
.8
.:.0: l:l CD
CD
c
CD
L-
a
6
a..
.4
T
T
1.2 ,....
.a
CI
<,
Micropore Volume
-'
§
.....CI
CD
<,
E
E ...,
:::l
-'
a
>
CD
E
CD
:::l
L-
a
-'
a
a..
>
-'
CD
0
+-
a
~
.... Mesopore Peak
8
a..
c:
lD
t
,
E
"iii
.:.0: :::l
.. ......
tr'
c: ....,
>+-
-'
547
L-
a
a..
.4
r 0
20
40
60
80
0
Amount of Carbon Blacl< (wt,,) Fig. 2. Relation between amount of carbon black and physical properties of catalyst.
20
40
60
80
Amount of Carbon Black (wt,,) Fig. 3. Relation between amount of carbon black and pore distribution of catalyst.
Discussions on the effect of carbon black Carbon black is a kind of ultra-fine particles, which seems to be only one available material that meets our object. The uniformity and fineness of the carbon black enables us to acquire expected narrowly distributed intermediate pore openings (mesopore). Carbon black particle sizes are variable over a wide range (10 to 300 nm) by a skillful technique of carbon black production. An important character of the carbon black is that particles link together at the edge of respective particle to form a highly reticulated chained structure. A photograph of the structure is shown in Fig. 5. The extent of growth of the chained structure can be roughly estimated from oil absorbancy (Table 1). The fact that the mesopore volume and the position are rather small when used carbon black E (see Table 2 and Fig. 4) can be interpreted by the poorly chained structure of the carbon black E. This chained structure also explains why the shaped body with such high porosity keeps its mechanical strength.
548
40 ,.., E c 20
....
..... ....................
:lit. 0 CD
a. CD
S
L-
--00
'"r"Ii~;%
8
0
0
a.
.,* -= ~~
Mesopore Peak
Mlcropore Peak
6
T
4f,
~
Micropore Volume CD
§
.4
-oJ
o
> ~
o
~VOl~';'"
Mesopore
,......-"""500nm
Fig. 5. (TEM).
Photograph of carbon black
•
.2
a.
o
20
40
60
80
Carbon Black Particle Size (nm)
Fig. 4. Relation between carbon black particle size and pore distribution of catalyst. o DBP abs.: 120-130 ml!100g D DBP abs.: 80 ml!100g added carbon black is 30 wt%.
From the SEM image of the surface
Fig. 6. Photograph of surface of bimodal alumina (SEM).
of alumina (Fig. 6), presence of mesopores which are formed by burning chained carbon black is seen. Consideration on the mechanism of appearance of mesopore will be helpful to comprehend the advantages of using such carbon black. As has been stated already (see Fig. 3), the mesopores appear only when over 20 wt% of carbon black is added and the micropore volume is maximized at around 10 to 20 wt% addition. This indicates that remnant void spaces surely exist but when added small amount of carbon black, they are recognized only as micropores (a extreme case of ink-bottle pore). If additives had only separated spherical shapes, inkbottle pores would remain even if a larger amount of them was added. This is illustrated in Pig. 7A. Addition of much larger amount would make it possible to obtain mesopores at the surface but it would give a detrimental effect to
549
the shaped body. On the other hand, in the case carbon black with highly chained structure used, probability of appearance of an end of the structure on the surface is considerably increased. This is shown in Pig. 7B and seems to explain the reason why the addition of relatively small amount of carbon black can produce desired mesopores. Another merit of using carbon black should be mentioned, which has substantial
B
importance in commercial production. It is combustibility of carbon black. Additives must have mild combustibility, for requirement of high temperature for combustion or high exothermicity of burning is very harmful to resulting bodies. In this point, carbon black has
Fig. 7. Image of pore structure.
an excellent nature.
Preparation of other inorganic oxides In the preparation of aluminas, raw sources are not limited to boehmite gel. It is possible to utilize, for example, l-alumina, f- or X-alumina, gibbsite and so on. In these cases, preparative methods are not necessarily the same as in the case of boehmite gel and,_of course, pore distributions are different. But the bimodal character of resulting aluminas is always retained. It is also
.... Q
1.5
"oJ
E ...... CD
E 1.0
:J 0
oJ
>
CD 0
L-
n, 0.5 CD
>
+=0 oJ
:J
E
:J
u
0
10
Pore Pig. 8.
100
1000
RadiUS (nm)
Pore distribution of trimodal silica-alumina.
550
possible to obtain 9- or$-alumina with bimodal pore distribution by calcination at higher temperature. This technique has been proved to be applicable to other inorganic oxides. Thus bimodal titania, silica, silica-alumina, zeolite-alumina, alumina-clays and so forth were obtained successfully. If two components of oXides were mixed, even a trimodal pore distribution could be obtained in some cases. Fig. 8 shows the trimodal pore distribution of silica-alumina. The smallest micropores are attributable to the intercrystallite spacings of silica. Hydrotreatment of heavy oils A potential effectiveness of bimodal catalyst has been recognized in the field of hydrotreatment of heavy oils. As has been pointed out, loading active components to the bimodal alumina does not change the bimodal character. Therefore a set of Co-Mo/alumina catalysts was prepared and examined by using an autoclave reactor or a fixed bed flow reactor. Aging behavior of the catalysts was also followed. Activity tests by autoclave.
Results of hydrotreating reactions are shown
in Figs. 9 and 10. Loaded catalytic components were 4 wt% of CoO and 13 wt% of Mo0 • The other reaction conditions are written in experimental section. 3
.....
100
'#.
0
ot-
.!c 90
a a
a 80 ot-
...J
.oJ
+-
CD
E
! T
......
70 T
72
0 :j: 0
:j: 0
70
L.
°C
72
N
....
....
:3
:3 It)
CD
0
0
C 0
.!::!
...J
a
80
l( +~
+-
0
CD
E CD 0
~ ..... c
0
0
+-
0 :j:
...J ...J
'#.
100
'#.
~ ..... c 90
0 :j: 0
......
......
0
D
0
:3
...J
68 T
:3 It)
CD
20 40 60 0 80 Amount of Carbon Black (wt,,)
Fig. 9.
Relation between amount of
carbon black and activity.
0
68
r
60 0 20 80 40 Carbon Black Particle Size (nm)
Fig. 10.
Relation between carbon black
particle size and activity.
551
From these Figs., it is seen that a high level of desulfurization activity is kept over a wide range examined. On the other hand demetallation activity rises up abruptly when added more than 20-30 wt% of carbon black. This clearly corresponds to the appearance of mesopores (see Fig. 3). Thus mesopores are obviously important for demetallation. Micropores and mesopores of the bimodal catalyst seem to share their roles in the reaction, that is, micropores for desulfurization and mesopores for demetallation, since metal containing molecules are relatively large. The reactions in Figs. 9 and 10 were performed by using Arabian light residue as feedstock. The use of Arabian heavy residue gave the same tendency. As to confirm this remarkable effici-
Reactions by fixed bed flow reactor.
ency of the catalysts (ref. 8), consecutive reactions were carried out using a fixed bed flow reactor. The high performance of the catalyst is evident from Fig. 11, where metal contaminants are effectively reduced while desulfurization activity keeps high level. In order to clarify the effect of mesopore position, further studies were undertaken using catalysts which had different pore distributions. To assure the reliability of results, other properties of the catalysts such as total pore volume and micropore peak position were kept constant. As is seen in
Cat.R D
D
0
00 0
D
0
0
0
o
D
D
D
0
0
D
0
0
0
100 200 On Stream Time (hr)
Fig. 11.
o C
o
0
o Cat.A 300
Typical run of flow test using bimodal catalyst. Cat. A (bimodal catalyst) Cat. R (reference catalyst (monomodal»
552
Fig. 12, activities for removal of
~,.
impurities (metals, asphaltene,
fE
nitrogen and residual carbon) are increased with increase of mesopore
d390~
peak radius. On the contrary,
w
~ ~ 370
fRi ::J*
'----------
interpreted by coke deposition as is generally accepted.
lib 1-'=
CD
+= 1: ::J
Co
These results teach us that the selection of pore distribution is
80
very important for this hydrotreating
~
60
.§
....0
Asphalten
40
~
.oJ
~---_
C
> 0
e 20
CD
a: "t#.
decreased. Initial decrease of the desulfurization activity may be
+- .... CCD LCI) CDC Co CD
CI)
desulfurization activity was
0
0
,
0
20
reactions. At the same time, the
... ..
CCR (.3)
.... -----°0
30
-
Nitrogen 0
properties of feedstock oils should be considered • Evaluation of aged catalysts. For practical use, not only initial activity but also catalyst
40
Mesopore Peal< Radius (nm)
life must be examined. For this purpose, several kinds of catalysts were loaded in a commercial reactor,
Fig. 12. Relation between mesopore peak radius and activity. (*1) Temperature required for maintaining 1% sulfur concentration after 200 hr on stream. (*2) Temperature increased during 100 hr to 200 hr on stream. (*3) Conradson carbon residue.
withdrawn from the reactor after one year and evaluated by an autoclave. The relation between used carbon black particle size and residual desu1furization activity is shown in Fig. 13 together with amount of coke and metal deposits. It is
noticed that in spite of an appreciably large amount of accumulated deposits, decrease in activity of these bimodal catalysts is rather small. A comparison between bimodal and monomodal catalyst is made in Fig. 14. It is striking that the bimodal catalyst can hold approximately twice amount of metal deposits to keep the same level of activity as that of monomoda1 one. Applications of the bimodal alumina and other inorganic oxides are, of course, not limited to this field of reaction. Use of this bimodal alumina to coal liquefaction (ref. 9) and to selective hydrogenation (ref. 10) is reported. Addition of carbon black is also effective for the production of zeo1itic hydrocracking catalyst (ref. 11).
553
'"
....III 1.0 ....a :go •8
8-
....
it .... .!
eO
+-
'iii
(.J
~\
G)
G)
a
..-e
....
....~--
so E
«
....---...-
CoKe
0&-..------
~
s ~ ....
\
\~
.'"'"
.4
\
G)
c:: o e
160
.6
\
Cat. R ... \
.2
\
1; c::
o
II)
20
40
60
Amount of Metal Deposit
'1:120
80 (wt~)
....e
g
Pig. 14. Relation between amount of metal deposit and reaction rate.
(.J G)
1; eO
Cat. A : bimodal catalyst. Cat. R : reference catalyst (monomodal).
c::
1
o
20
40
60
80
Carbon BlacK Particle Size (nm)
Ratio of reaction rate constant= k(aged)/k(initial).
Pig. 13. Effects of carbon black particle size on residual activity and metals and coke deposits.
CONCLUSIONS Addition and final burning of some kinds of selected carbon black give mesopores ranging from 10 nm to 100 nm radius to alumina and other refractory inorganic oxides. Characteristics of the metod are: 1) the position of mesopore peak and the volume of mesopore are variable by changing the properties and the amount of carbon black, 2) in spite of high porosity, shaped body keeps a gOOd mechanical strength, 3) the method is applicable to a wide range of inorganic oxides. Thus the preparative method is very useful for obtaining suitable catalysts and catalyst supports. Bimodal Co-No/alumina is an excellent catalyst for hydro treatment of heavy oils. The catalyst has a high hydrotlesulfurization and hYdrodemeta1lation activity and also has a gOOd durability which assures a prolonged catalyst life.
554 ACKNOWLEDGEMENT This article was partly reproduced from Sekiyu-Gakkaishi (Jurnal of The Japan Petroleum Institute) by courtesy of The Japan Petroleum Institute. The author would like to express his grateful acknowledgement to the Institute.
REFERENCES 1
A.L.Dicks, R.L.Ensll, T.R.Phillips, A.K.Szczepura, M. Thorley, A. Williams and R.D. Wragg, Jurnal of Catalysis, 72 (1981) 266. 2 T. Ono, Y. Ohguchi and O. Togari, Proceedings of Preparation of Catalysts III, Louvain-la-Neuve, Belgium, 1983, 631 pp. 3 Japan Kokai 1979-25907. 4 J.H. DeBoer, A. Van der Heuvel and B.G. Linsen, Jurnal of Catalysis, 3 (1964)
268. 5 R.F. Vogel, G. Marcelin and W.L. Kehl, Applied Catalysis, 12 (1984) 237. 6 D. Basmadjian, G.N. Fulford and B.I. Parsons, Jurnal of Catalysis, 1 (1962) 547 ; R.E. Tischer, ibid., 72 (1981) 255 ; R.E. Tischer, N.K. Narain, G.J. Stiegel and D.L. Cillo, ibid., 95 (1985) 406. 7 US 4,508,841. 8 US 4,448,896. 9 Japan Kokai 1984-213791. 10 Japan Kokai 1984-123539. 11 EP 138783.
555
DISCUSSION E.B.M. DOES BURG : What is the thermal stability of the bimodal pore distribution in, for instance, alumina? K. ONUMA : For alumina, bimodal pore distribution can be retained even at about -1,lOO°C. Temperatures higher than _1,100°C will change the crystal l tne phase to a-alumina and bimodal pore will disappear. J.D. CARRUTHERS: Could you quantify the statement that these materials have good mechanical strength? What would be the side crush strength of an extrudate that contains 30-40% carbon for example? K. ONUMA : Minimum value of crush strength for cylinder type alumina (1.2 mm ¢, 4 mm 1) is 2 kg/piece and for polylobe, 3 kg/piece. Efforts to strengthen the shaped material are in progress and an average value of 4 kg/piece has been realised for polylobe.
J. SCHEVE: Are you able to burn off the carbon black completely? K. ONUMA : Yes. Combustibility of the used carbon black is very good and carbon black can be completely burned off without damaging the material's physical properties. BIRKE : How do the mechanical properties change if you add carbon black till 80 wt %? Is it possible to extrudate mixtures with such a high content of carbon black? K. ONUMA : Too high addition of carbon black is unfavourable from the view point of mechanical strength and economy.
This page intentionally left blank
557
B. Delmon. P. Grange. P.A. Jacobs and G. Ponce let (Editors). Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V" Amsterdam - Printed in The Netherlands
HYDROTREATING NiMo/SEPIOLITE CATALYSTS: INFLUENCE OF CATALYST
PREP~
RATION ON ACTIVITY FOR HDS, HYDROGENATION AND CHAIN ISOMERIZATION REACTIONS 1 1 1 2 F.V. MELO , E. SANZ , A. CORMA , A. MIFSUD lInstituto de Catalisis y Petroleoquimica, C.S.I.C., Serrano, 119, 28006-Madrid (Spain) 2Instituto de Fisico-Quimica Mineral, C.S.I.C., Serrano, 115, 28006-Madrid (Spain)
SUMMARY A series of NiMo catalysts supported on a sepiolite: a) in its natural state, b) modified by acid leaching, and c) modified by cation exchange, have been prepared. The preparation variables studied were: Method of metal deposition, amount of active phase, sepiolite pretreatment, and temperature and time of sulfurization. The cataly tic activity for HDS, hydrogenation, and cracking-isomerization has been studied by feeding a thiophene-cyclohexene -cyclohexane mixture and carrying out the reaction in the following conditions: 300° and 400°C reaction temperature, 20 Kg.cm- 2 total pressure, and 3 to 1 molar ratio of H2 to hydrocarbons. An optimum for HDS and hydrogenation activity was found for a 12% wt Mo03, and 5% wt NiO, prepared by simultaneous impregnation by the pore volume method at Ph = 5.0. The optimum conditions with these catalysts are 400°C and 3 hours of sulfurization. An increase in the acidity of the support produces a decrease of HDS and hydrogenation and an increase of the cracking-isomerization activities. A good correlation between HDS and the concentration of an XNiO. Mo03 phase is found. The XNiO.Mo0 3 phase is completely sulfurized to a modified MoS while NiMo0 and Mo0 are only slightly sulfurized. 4 3 2, INTRODUCTION The hydrotreating of the heavy fractions of petroleum is becoming more and more a necessity. This is due not only to economical reasons which impose a better use of the botton Of the barrel, but also to the ecological necessity of decreasing SOx and NO
x
emi-
ssions. The upgrading of heavy oils together with a deep
hydrodesulfur~
zation can be carried out either by catalytic or non catalytic procedures. The catalytic hydrotreating of heavy oils and residues, although quite efficient, is still limited by a fast, irreversible deactivation of the catalyst due to an irreversible adsorption of metallic contaminants, mainly nickel and vanadium (refs.1-7). On the other hand, the non catalytic procedures (refs. 8-12), in prin-
558 ciple less atractive from the point of view of conversion and selectivity, are in many cases competitive with the catalytic processes due to the absence of deactivation. One alternative to improve the efficiency of catalytic hydrotreating is to prepare and test non expensive catalysts, e.g.
nat~
rally occurring materials (refs. 13-18), which decreases the cost of the irreversible deactivation. The testing of new catalysts can be carried out using commercial feeds, or more frequently with a
l~
bora tory test reaction. To choose a test reaction one has to take into account that the processes involved in hydrotreating heavy oils, viz. HOM, HOS and cracking-isomerization, involve hydrogenation, HOS, and acid functions in the catalyst (ref. 19), and
ther~
fore the test reaction should give information about these three functions. In this work, a naturally occurring magnesium silicate (sepiolite), instead of classical supports such as A1 20 Si0 or Si0 3, 2 2-A1 20 3, has been used to prepare hydrotreating Ni-Mo catalysts. Their catalytic properties for hydrogenation, HOS, and crackingisomerization have been measured by means of a model feed composed by cyclohexene, cyclohexane and thiophene. Finally, a correlation has been observed between the oxide and sulfide phases on the surface and catalytic activity and selectivity. EXPERIMENTAL Catalyst preparation The sepiolite used as support proceeds from Vallecas (Spain). It was used. 1) in its natural form, 2) treated with HCl
(1-3N so-
lutions) to leach different quantities of magnesium (ref. 20), and 3) with part of magnesium exchanged by aluminum (ref. 21). An A1 20 3 Girdler T-126 was used as reference carrier. For the Mo and Ni impregnations, aqueous solutions of (NH4)6M07024.4H20 (PROBUS A.R.) and Ni(N0 3)2 (PROBUS A.R.) of adequate concentrations were used. The impregnation was carried out either with excess of impregnation solution with vaporization by means of a rotary evaporator at 60°C and reduced pressure (RE samples), or with the exact solution volume to fill the pores (PV
sa~
ples. The two impregnations were either simultaneous (SIM-samples) and successive (SUC-samples). In the last case, molybdenum was always impregnated first and a calcination step was carried out be fore the second impregnation. was:
The
calcination treatment
From ambient temperature to 550°C with a linear
used
559
heating rate of
~
1 C.min-1,
with intermediate isotherm
steps at 110°C (5 hours), 150°C (2 hours), 300°C (2 hours) and 550°C (4 hours). The calcined catalysts were sulfurated in the catalytic reactor before the activity test with a mixed flow of H2 and H2S (20% by volume). The temperature and time of sulfuration were varied in the ranges of 300°-500°C and 0-5 hours,respectively. Catalytic activity test The catalytic experiments were made in a fixed bed reactor. The experimental conditions were: pressure, 20 Kg.cm -2 ; hydrogen flow, 9.5 dm3.h- 1; liquid flow, 1.5 cm 3.h- 1; catalyst weight, 1-2 g; tern perature, 300°c for 3 hours and then 400°C for 4 hours; liquid feed composition, 30% wt cyclohexene, 68% wt cyclohexane, and 2% wt thiophene. The gas and liquid products were analysed by gas chromatography using a Porapak Q plus silica gel column and 15% carbowax 1500 on chromosorb WAW column, respectively. Conversion to thiophene is referred to thiophene introduced, while the rest of the products are referred to cyclohexene introduced. A pseudofirst order kinetic model has been used to calculate the activity of the catalysts. Characterization techniques X-ray diffractograms of the samples were obtained in a Philips PW 1710 with CuKa radiation to identify the possible compounds of Mo and Ni formed in the calcined and sulfided catalysts, and a con ventional TEM, Philips 300, was used to estimate the dispersion of the active phases. For the catalysts supported on sepiolite it was possible to identify by X-ray diffraction several oxides and sulfi des, while for the catalysts supported on Al 203 no crystalline species were observed. To estimate the surface acidity of the catalysts pyridine was adsorbed at room temperature, and the pyridine retained after desorption at 150°C and 10- 4 Torr was studied by i.r. spectroscopy. RESULTS AND DISCUSSION Usually, when the authors try to compare the influence of different supports on HDS catalysts they use the same preparation conditions and the same metal loading (generally the optimum found for Al 20 3) with the different supports. This can be quite misleading, since different supports interact differently with the Ni and Mo compounds, so that the optimum preparation conditions differ
560
for different carriers. Therefore, a systematic study of the
diff~
rent variables involved in the preparation of the catalyst has to be carried out if the real possibilities of a new catalyst are to be found. Firstly, the optimum temperature and time of sulfurization were studied (Fig. 1). At the temperature of 400°C during 3 hours
Q
4
...s:
b
40 ~
I
s:
~
~2 :t:
-'£
;;'2 0 :t:
-'£
0
0 10
10
~
~
I
s:
0 >- 5 :t: -'£
0
r
.c
~
05 >-
:t: -'£
0 O[
h
Figure 1. HDS and hydrogenation constants for 12% wt M00 and 5%wt 3 NiO SIM-PV-catalysts supported on natural sepiolite. a. Effect of sulfurization temperature. b. Effect of sulfurization time. Reaction temperature: 0, 300°C;
~,
400°C.
the HDS and hydrogenation activity were at a maximum, while no differences in isomerization were observed with temperature and time of sulfurization. When the pH of the impregnating solution was changed no influence of the pH of the solution on the isomerization was observed. On the other hand maxima for both HDS and hydrogenation activities were observed at pH = 5.0 (Fig. 2). Therefore these were the conditions used for the preparation of the different NiMo/sepiolite catalysts described below. Influence of the metal deposition procedure The catalytic activity for the reactions under consideration obtained with a series of catalysts, all of them with 5% wt NiO and 12% wt M00
but with different procedures of incorporation of 3 the oxides, is given in Fig. 3. Four facts clearly emerge:
561
10 -
40 ~
I
'j5-
s: VI
~20 x
o
?==: I
I
I
I
I
2
4
6
8
10
pH solution
Figure 2.HOS and hydrogenation constants of 12% wt !'4003 and 5% wtNiO SIM-PV-catalysts supported on natural
of impregna-
sepiolite~.
ting solution. Reaction temperature: 0, 300°C1 &, 400°C.
25 20
k
15
';;10 5
o
1
2
3
sol. excess sue. impreg.
4
1
2
3
sol. excess sim. impreg.
234 pore volume sue. impreg.
1
234
pore volume sim. impreg.
Figure 3. Activity constants for catalysts with 12% wt M00 and 5% 3 wt NiO on natural sepiolite prepared by different methods. 1,kHOS1 2, k HV01 3, k I SOM1 and 4, kOEHYO' Reaction temperature:~, 300°C and . , 400°C.
1) Cracking products (C n' n~5) are not observed, 2) the activities for isomerization and dehydrogenation are not sensibly affected by the preparation procedure, 3) a maximum in activity for HOS and drogenation is found with a sample prepared by simultaneous
h~
impre~
nation of Mo and Ni using the pore volume procedure, and, finally, 4)
there is a direct correlation between activity for HOS and hy-
drogenation. It is not surprising that cracking products are not observed, since there are no Br6nsted sites of medium and strong acidity,
re~
ponsible for cracking of olefins and alkanes, as shown by the absence of the pyridium band at 1545 cm- 1 after pyridine desorption 4 at 150°C 10- torr, in all of the four samples. The low isomerization activity to methylcyclopentane and methylcyclopentenes could
562
then take place on the molybdenum and nickel sulfides or any mixed sulfide, or also on relatively weak acid sites. Indeed, it is known that nickel and molybdenum sulfides are active for branching
isom~
rization of alkanes and cycloalkanes (ref. 22). The possibility of relatively weak acid sites being active for isomerization, preferently of cylcohexene, has to be considered since it is observed that in the isomerized fraction methylcyclopentene is the majoris£ merized product on catalysts with low hydrogenation activity. On the other hand, it is clear, that the isomerization of cyclohexene via carbonium ions is an easy reaction which does not need strong acidity. The differences in activity for HDS and hydrogenation observed with the samples prepared by different methods can be due to
diff~
rences in the dispersion of the oxides or to differences in the phases formed during calcination. An attempt to measure the disp.eE sion has been done by means of TEM in all of the catalysts oxides and sulfides were as big
the
nrn). On the other
crystals(~200
hand the following phases have been identified on the calcined
sa~
pIes by X-ray diffraction: Sepiolite, NiO (file nQ 22-1189 (ref. 23», i"100
(5-508), NiM00 4 (16-291), NiM00 (18-879) and XNiO.M00 3 4 3 (12-348), this being the most important of the two last species.
The same species were detected by Laine and Pratt (ref. 24) in NiMo/AI
catalysts with high molybdenum loadings (>15%). In a 20 3 first approximation, the activity for HDS has been related here
with the concentration of each of these species (measured as the area of the most representative diffraction peak), a direct correlation being only observed for the XNiO.M00
3
phase (Table 1).
TABLE 1 Comparison of HDS activity at 300°C and area of XNiO.M00 at 3.34 ~. Catalyst
kHDS,h
SUC-RE SIM-RE SUC-PV SIM-PV
3.2 6.5 6.6 10.2
-1
3
species
XNiO.MOO area 3 26 29 30 44
During sulfuration, the XNiO.M00 phase disappears, the two NiMo0 3 4 phases are not affected, and the area of M00 sligthly decreases. 3 A large amount of MoS2 is formed, which is higher the higher the
563
amount of XNiO.M00
present in the precursor. It must also be 3 pointed out that the (002) peak of the MoS phase is always quite 2 broad. This can be a consequence of the low crystallinity of the MoS 2 phase, or also of the presence of small crystallites. A third possibility which can explain the diffractogram of the MOS
and 2, which has to be considered in the light of work by Candia et al. (ref. 25), is the incorporation of Ni atoms occupying molybdenum positions in the MoS
2 phase. When a NiMO/AI 20 catalyst was prepared by simultaneous 3
impre~
nation using the pore volume procedure with 5 and 12% wt NiO and M00 , respectively, no peaks were obtained in the diffractogram, 3 and a much higher dispersion of the phases than in the NiMo/sepiolite was observed by TEM. However, the HDS activity of the NiMo/
3 was lower than that of the NiMo/sepiolite (Fig. 5). Up to this point the results show that a mixed oxide of the
/A1 20
type XNiO.M00 which is easily sulfurated to a modified MOS spe2 3, cies, could be the precursor or the active phase for HOS. Furthermore, the dispersion of the oxides, considered globally, is not ne cessarily related with the HOS activity. Effect of Ni content The effect of Ni content on the activity for thiophene HDS and cyclohexene hydrogenation is shown in Fig. 4. A maximum in activity is observed for a Ni/Mo atomic ratio of
~0.8,
pretty much in
the range found with Al 203 (ref. 25-28). In this series, the hydrogenation activity does not follow, the same behaviour as HOS, but increases monotonically with increasing Ni content. This result, together with the hydrogenation behaviour observed with the cata-
20
8 ~
I
I ~
~10 ~ ~
~
~4 ~ ~
Figure 4. HOS and hydrogenation constants for SUC-RE-catalysts supported on natural sepiolite and 12% wt Mo0 3 . Reaction temperature: 0, 300°C, ~, 400°C.
564
lysts prepared by different procedures, indicates that, besides the sites which are active for both HDS and hydrogenation, there are others probably related with NiS or with mixed phases rich in nickel. Effect of sepiolite treatments As said above, sepiolite is a hydrated magnesium silicate with a crystalline structure (ref. 29) consisting of laths joined together at their corners, with channels running the whole length of the fibre-like crystals. The octahedral sheet of the sepiolite is formed by magnesium ions, which can be leached by acid treatment leaving a siliceous skeleton (ref. 30) or a mixture of silicagel and unattacked sepiolite (ref. 20). In this way, a sample of
sepi~
lite with 20% of the octahedral magnesium extracted by acid leaching was used to prepare a catalyst by pore volume impregnation and 5 and 12% wt of NiO and M00
respectively. The activity 3, results are given in Fig. 5, which shows practically no differen-
ces in behaviour, with respect to the natural sepiolite. It has been recently shown that by exchanging the magnesium 3 cations,
ions located at the edges of the octahedral sheets by Al+
it is possible to prepare a sepiolite retaining its original
stru~
ture, but with a surface acidity comparable to that of amorphous silica-aluminas (ref. 21). A catalyst was prepared in the same way 3 as above, but using an Al+ exchanged sepiolite. In Fig. 5 it can be seen that this catalyst shows a lower HDS and hydrogenation activity, but higher isomerization and cracking (15% relative to cyclohexene + cyclohexane) activities than any of the other
previou~
SO 40
'.it;':...20 e 30 10
o
L~ 234
A~03
1 234 untreated sepiolite
2 3 4 acid treated sepioli te
~ 1 2 3
4
Al exchanged sepiolite
Figure 5. Activity constants for SIM-PV-catalysts with 12% wt M00 3 and 5% wt NiO on different supports. 1, k HDS; 2, k Hy D; 3, k I SOM; and 4, kDEMYD' Reaction temperatures: ~ ,300°C a n d . ,400°C.
565
ly prepared catalysts. In Fig. 6 it can be seen, from changes in the 1545 cm- 1 band, that the Al+ 3 exchanged sepiolite has medium and strong acid sites, 1545
NiMo/Al ssp.
and that part of them disappear after impregnation and calcination, a fact which has also been observed in zeolites (ref. 31). In any case, u l.t.houqh a good catalyst for hydrocraeking, this sample does not
LU
present an interesting activi-
U
Z <{
ty for HDS.
a:l
c:: o
Finally, to check if the co
V1 aJ
rrelation observed in Table
-c
between HDS activity and the
Al sep.
concentration of the xNiO.M00 3 phase could be generalized, in Fig. 7 the HDS activity vs XNiO.M00
concentration has 3 been plotted for the catalysts. 1700
1500
'V, em"
1
1300
A direct correlation is observed in practically all
cases.
Figure 6. i.r. spectra of pyridine retained after evacuation at 150°C 4 and 10- torr.
101-
'" 0" Q
VI
I
Cl
Z
""
I
51-
I
I
<:/0 I
9
Figure 7. Correlation between
o
HDS activity at 300°C and the I
I
30 40 area of the peaknt 3.34A,lllI.
area of the major peak of xNiO.Mo0 3 phase.
566
REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27
28 29 30 31
W.I. Beaton, Chern. Process. Eng., 11 (1977) 1. W.K.T. Glein, Dev. Pet. Sci., 7 (1978) 243. E. Furimsky, Erdoel Kohle Erdgas Petrochem., 32 (1979) 383. A. Nielsen, B.H. Cooper and A.C. Jacobsen, Prepr. Div. Pet. Chern., A.C.S., 26 (1981) 440. P.W. Tamm, H.D. Harnsberger and A.G. Bridge, Ind. Eng. Chern. Process Des. Dev., 20 (1981) 262. S.T. Sie, in B. Delmon and G.F. Froment (Eds.), Catalyst Deactivation, Elsevier, Amsterdam, 1980, p. 545. B.G. Silbernagel and K.L. Riley, in B. Delmon and G.F. Froment (Eds.) Catalysts Deactivation, Elsevier, Amsterdam, 1980,p.313. S. aka, J. Kimoto, H. Matsui, H. Seki and M. Kaneko, Prepr. Div. Pet. Chern. A.C.S., 24 (1979) 678. T.Y. Yan, Prepr. Div. Pet. Chern. A.C.S., 28 (1983) 691. Y. Jacquin, H. Toulhoat, A. Quignard and J.F. Le Page, Revue de l'I.F.P., 38 (1983) 371Y. Fukui and Y. Shiroto, Chern. Econ. Eng. Rev., 15 (3) No. 166 (1983) 1Heavy Oil Proceesing Handbook, The Institute of Heavy Oil Proce ssing (Eds.), Chemical Daily Co., Ltd., Tokyo, 1982. Y. Shiroto, T. Higashi and T. Ono, U.S. Patent 4,367,167 (1983). S. Asaoka, T. Ono and Y. Shiroto, U.S. Patent 4,439, 312 (1984). C.D. Chang and A.J. Silvestri, Ind. Eng. Chern. Process Des. Dev., 15 (1976) 161. S. Marengo, A. Iannibello and A. Girelli, in Preparation of Catalysts III, G. Poncelet, P. Grange and P.A. Jacobs (Eds.), Elsevier, Amsterdam, 1983, p. 359. G. Sep~lveda, M.L. Cubeiro, C.M. LOpez, M. Micarelli, L. Uva, R. Torres, V. Galindo, L. Flores and M. P~rez, VIII lb. Symp. cat., Spain, C-18 (1982) 671. C. Dodet, F. Noville, M. Crine, P. Marchot and J.P. Pirard, Appl. Catal., 11 (1984) 251R.A. Ware and J. Wei, J. Catal., 93 (1985) 100,122 and 135. A. Corma, J. P~rez-Pariente and A. Mifsud, Clay Minerals, in press (1986). A. Corma, V. Forn~s, A. Mifsud and J. P~rez-Pariente, Clay Min~ rals, 19 (1984) 673. O. Weisser and S. Landa, Sulphide Catalysts. Their Properties and Applications, Pergamon Press, New York, 1973. Powder Diffraction File, Joint Comittee on Powder Diffraction Standards, (1976). J. Laine and X.C. Pratt, Ind. Eng. Chern. Fundam., 20 (1981) 1. R. Candia, B.S. Clausen, J. Bartholdy, N.-Y. Tops¢e, B.Lengeler and M. Tops¢e, Proc. 8th. Int. Congress on Catalysis, Berlin, 1984, Vol II, p. 375. S.P. Ahuja, M.L. Derrien and J.-F. Le Page, Ind. Eng. Chern. Prod. Res. Dev., 9 (1 970) 272. B. Delman, in Proceedinqs of the Climax Third International Con ference on Chemistry and Uses of Molybdenum, H.F. Barry and P. C.H. Mitchell (Eds.), Climax Molybdenum Co., Ann. Arbor. Michigan, 1979 p. 73. N. Sharma and A.K. Kar, Indian J. Technol., 21 (1983) 510. K. Brauner and J. Pressinger, Tschermats. Min. Petro Mitt., 6 (1956) 120. L. Gonz~lez, L.M. Ibarra, A. Rodriguez, J.S. Moya and S.J. Valle, Clay Minerals, 19 (1984) 93. V. Forn~s, M.I. V~zquez and A. Corma, Zeolites, 6 (1986) 125.
567
DISCUSSION M. TOKARZ: 1/ In what form was the extrudates 1 2/ What is the average pore size of 3/ Do you have any evidence showing sed on the carrier 1 I believe this of Mo- and Ni-phases were big enough
catalyst support prepared, pellets or both support and ready catalyst 1 how Mo and Ni are distributed and disperis important due to the fact that crystals to be detected by XRD.
F.V. MELO : 1/ The catalyst was pelletized, crushed and sieved and the 0.250.42 mm size was used for the reactions. 2/ The natural sepiolite has an average pore radius of 38 Xand the catalyst with 5% wt NiO and 12% wt Mo03 rrepared by simultaneous impregnation has an average pore radius of 88 ~. 3/ With electron microscopy we have observed different types of crystals distributed on the external surface of the sepiolite fibers and there are no preferential sites. We are trying to identify the composition of these crystals by microdiffraction. B. DELMON : You correlate HDS activity with the quantity of nickel molybdate phase (xNiO.Mo03). In NiMo/A1203, activity seems to correlate better with the perfection of the NiMoO x bilayer adhering on alumina than with NiMo04' 1/ Can you exclude that, in addition to xNiO.Mo03 that you detect, there be a large quantity of bilayer adhering on sepiolite 1 2/ If all Ni is in the form of xNiO.Mo03, how could you explain the correlation xNiO.Mo03 vs. activity, remembering that, on the contrary, CoMo0 formation corresponds to less active HDS CoMo/A1203 catalysts 1 Is this due to 4the solid state chemistry-or-the reduction-sulfidation of xNiO.Mo0 3 1 (compared with that of CoMo0 4 1) F.V. MELO : 1/ We cannot completely exclude the existence of bilayer adhering on sepiolite, but with electron microscopy we can conclude that the extent of bilayer, if there is, must be small. 2/ The sulphidation of xNiO.Mo03 phase is much faster that the su1phidation of the other phases present in the catalyst. These phases are not considerably sulphided in our sulphidation conditions. The xNiO.MoD3 can be the precursor of the active phase. A. SCHUTZ: You prepare modified sepiolites by "exchanging" some Mg by A1 tn view of preparing a supported Ni-Mo catalyst. Did you compare some activity results of this support with attapulgite which is a similar mineral with non-negligible amounts of aluminum? F.V. MELD : We have tested the acti vity of supported NiMo catalysts on attapulgite (palygorskite) and the activity of these catalysts is lower than the activity of the catalysts supported on Al+3-exchanged sepiolite. In the A13+exchanged sepiolite, the a1umi nium is located in octahedra1positi ons at the border of the channels while in the case of the attapulgite the A13+ present in the naturally occurring samples is located within the octahedral sheet. The effect of this cation can be different for the two supports. A.J. van HENGSTUM : You have presented results on the effect of removal part of the Mg from the sepiolite structure or exchanging it for Al. However, all catalysts investigated were prepared according to a standard route, which was more or less optimized only for catalysts based on the natural sepiolite. Have you also looked at the influence of the preparation conditions (e.g. the pH of the impregnation solution) for the other support materials 1 F.V. MELD: Indeed, in all cases the preparation conditions have an influence on the activity of the catalysts. We are following the work in this direction to optimise HDS activity and to modify the selectivities of the reactions : HDS, hydrogenation, isomerization and cracking.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
569
NEW URANIUM-BASED HYDROTREATMENT CATALYSTS. G. AGOSTINI 1, M.J. LEDOUX Z, L. HILAIRE and G. MAIRE Laboratoire de Catalyse et Chimie des Surfaces, UA4Z3-CNRS, Universite L. Pasteur, 4 rue Blaise Pascal, 67000 Strasbourg (France).
SUMMARY This article deals with the preparation, the physical characterization (EXAFS, XPS) and the reactivity of NiU; Al Z0 3 and NiMoU;Al Z0 3. The uranium is impregnated from a solution of uranyl acetylacetonate by chemical reaction with the hydroxyl groups of the support. HDS of thiophene and HDN of pyridine at atmospheric pressure are tested. Such a preparation of the uranium blocks the formation of aluminate (Ni), increases strongly the dispersion of the active phase and favorizes the hydrogenation of the reaction products. INTRODUCTION Uranium has already been added as doping agent on hydrotreatment catalysts in order to prevent rapid coking and metallation of the surface during the process of residues (ref. 1). The impregnation steps were found to be unimportant, uranium could be added after or before Ni, Mo whatever the initial salt : uranyl nitrate, acetate, sulfate or halogenides. Other authors have also used uranium with Co or Ni and/or Mo either to hydrodesulfurize light gas, naphta or residues (ref. Z) or to reduce HZS by the Claus process (ref. 3). We found (ref. 4) that the conditions of impregnation and the use of a specific uranyl salt, the acetylacetonate (acac), which has not been published up to now, were determining for the specific action of uranium, in terms of dispersion and antialuminate functions (ref. 5). PREPARATION Uranyl acetylacetonate synthesis. A stochiometric yield synthesis of UOZ(acac)Z has been developed following the equation 40°C UO Z(N0 3)Z + Z acacH
1present address
---f
K2 C0 3
t UOZ(acac)Z + ZHN0 3
--,.
~
HZO
+ ZKN03 + CO Z
Good-Year International Tyre Technical Center, Colmar-Berg (Luxembourg). ZTo whom all correspondence should be addressed.
570
These reactions occur in water solution; the yellow precipitate must be protected from any trace of acidic vapour or liquid. It is not air sensitive. It is slightly soluble in water and hexane and very soluble in ethanol. Its composition has been checked by elementar analysis (%C 25.6-found 26.2, %H 3.00-found 2.95). Impregnation. The y-alumina (Ketjen CK300 or Procatalyse) is pretreated with a 1:1 water ethanol solution (volume of water
=
porous volume of alumina) and dried at
120°C in order to provide a maximum and a reproducible amount of hydroxyl groups on the surface. A 10% in excess of porous volume of an ethanolic solution of uranyl acac is contacted with this alumina and placed at 120°C for 12 h in order to perform the following reaction :
o OH
OH
Al
Al
?
~
,......,U .........
o
"
(0)
/
I
° I
Al
+ U0 2(acac)2
0
Al "
/
+
t
2acacH
(0)
(acacH boiling point is 118°C). With such a method, one can obtain about 16% in weight of uranium if after completion of the reaction, the solid is washed to eliminate the excess of unreacted complex. This solid is then dried and grilled at 500°C for 2 h in a rotative furnace, in a flow of dry, clean air. The subsequent impregnations, Mo from paramolybdate or Ni (Co) from nitrate, were performed by following the usual incipient wetness method. PHYSICAL CHARACTERIZATION EXAFS The compounds have been studied by EXAFS at the LURE (Orsay , France) using the beam from the DCI ring storage. The experimental details and the calculation procedure have been reported elsewhere (ref. 5). The oxide U30S has been used as reference in order to extract the measured phase function for U-o because of the inaccuracy of the Teo and Lee approximation (ref. 6) for this couple. This approximation was accurate enough for the other couples. The nature, the number (amplitude) and the distance of the neighbors of the U or the Mo atom for different compounds are reported in table 1.
571
The expected values for the nature and the number were found on the starting complex. The two double bond U=O distances, as the two single bond U-O distances, were observed on the basic catalyst U/A1 203 ex-acac, meaning that the chemical reaction between the hydroxyl groups of the alumina and the starting complex occurred as expected. These distances are retained after the double impregnation of Mo and Ni although the number (amplitude) was less accurate because of a less good spectrum. Mo as neighbor, could not be detected because of the too low concentration of this metal. Mo edge has not been yet exploited on this oxidic precursor. TABLE 1 Geometrical parameters extracted from EXAFS measurements. Compounds Complex U0 2(acac)2 U edge
nature
number
Oxygen Oxygen
2 ± 0.5 4 ± 0.5
Carbon U/A1 203 ex-acac (precursor) %U = 13% U edge
Oxygen Oxygen
1. 75 2.33
Sulfur
2.1 ± 1 1.3 ± 1
2.19 2.94
Sulfur
3.3 ± 0.5
2.45 (2.41)a
Molybdenum
2.7 ± 0.5
2.78 (3.15)a
Uranium
2.1 ± 0.5
2.98
Oxygen
Idem, in situ sulfided
Oxygen
Mo edge
2.795
1.72 2.33
%Mo = 2%, %Ni = 0.5% U edge
Idem, in situ sulfided
2 ± 0.5 2 ± 0.5
1.68 2.22
2.3 ± 1 3.9 ± 2
NiMoU/A1 203 ex-acac (precursor) Oxygen
U edge
7.7 ± 2
distance (1\)
aDistances in MoS 2 (ref. 7 and 8). In order to study the variation of the surface structure after sulfidation (simulating the reaction conditions), this precursor has been sulfided by H2S, in situ, and the spectrum on U edge has been recorded. The double bound oxygen atoms have been replaced by sulfur atoms while the two single bound oxygen atoms binding U to Al were not substituted. The distance U-S is typical of the distance in US 2' the U-O distance is slightly shortened (2.19 vs 2.33) (see
572
Figure l a), The EXAFS on the Mo edge gave a Mo-Mo distance rather different from the typical distance obtained by EXAFS on conventionnal catalysts (2~8 vs 3.15 A) where this distance corresponds to the MoS 2 structure (ref. 7 and 8), while the Mo-S distance is close to the usual one (2.45 vs 2.41 K). On Figure lb, a Mo -
S
XS
/
SX
\ /
Mo • - - •• Mo - - ./••
/S'/ S
u
/ \
u 0
0
\ .. '·'~'sl
(a)
S
/ \
Al
/
\
o
/
\
X
u
\
/ (b)
/ \
0
o
I
I
Al
Al
Al
(0)
Al
I's S
\
u Al
-l-w.
s
\ /
/
S
0
H or Mo
Fig. 1. Structural models of the sulfided catalyst. a) For U edge. b) For Mo edge. structure is proposed to illustrate these findings. Each Mo atom is bound via two sulfur atoms to two uranium atoms; a third sulfur atom binds together two Mo atoms. Such a model takes into account all the parameters determined from the EXAFS on the U and Mo edges. It is interesting to note that the Mo-Mo distance, much shorter than in MoS 2, is close to the Mo-Mo distance found in organometallic complexes, L,e, 2.8 1\., with M<>"S'Mo structure (ref. 9). This anchoring of molybdenum via sulfur atoms to two uranium atoms, these uranium atoms being themselves monolayered-bound to the alumina support in a pseudo-epitaxial structure, can retard the formation of large MoS 2 pellets during the sulfidation pretreatment of heavily Mo charged catalysts (ref. 8) and, then, authorizes the preparation of catalysts containing larger 8IIIounts of active, accessible Mo than the usual optimal concentration of 7-8 % in weight metal. This dispersion effect of uranium ex-acac has been proved for many other metals, such as the 3d transition series (ref. 5). X.P.S. In order to confirm the role of the impregnation procedure in the formation of this pseudo-epitaxial structure of uranium oxide on alumina, an XPS
573
study has been made on NiU/AIZ03 ex-acac and ex-nitrate oxidic precursor, HZ reduced precursor and HZ/HZS (Z% vol.) sulfided catalysts, assuming that this structure exists whatever the nature of the metal (Ni, Mo and/or NiMo) accompanying uranium. The XPS spectra have been recorded on a VG ESCA III; the samples, when pretreated, have been transferred under argon from the reactor to the analysis chamber. The Al ZP3/Z peak at 119.6 eV has been used as reference for the supported catalysts and the C Is peak at 284.8 eV for the pure oxides (U02 and U308). The corrected values for the binding energies are reported in Table 2. TABLE 2 Binding energies from XPS measurements. Energies for Ni (eV)
Compounds
2P3/2
satel. 2Pl/Z
856.9
861.7
873.3
Idem
(HZ reduced)
856.3 853.6
862.2
873.0
Idem
(H2/HZS sulfided)
853.7
NiU/AI203 ex-acac b (precursor) Idem (H2 reduced) Idem (H2/HZS sulfided)
857.2
861.8
856.2
861.6
NiU/A1 203 ex-nitr.a(precursor)
4f 7/ Z 4f 5/ 2 satel. 402.2
381.6 380.3
392.2 391.3
398.8
380.8
391.8
398.8
NO
38Z.4
393.1
380.2
391.0 398.2
381.0
392.0
399.6
U308 (ref. Is C)
380.4 381.0
391.1 391.7
398.7
U02
379.9
390.8 397.6
a%Ni
873.8 873.Z
for U
853.7
(ref. Is C) 2.1%, %U
= 15.8%.
b%Ni
2.4%, %U
= 16.2%
In order to determine the nature of the uranium oxide present on the precursors' surface, it is not wise to compare binding energy values found on supported and bulk oxide because the correction made for the charge effect does not use the same reference. However, one can observe on NiU ex-acac a shift of 0.9 eV towards higher binding energies when compared to NiU ex-nitr., meaning a higher degree of oxidation for uranium in the ex-acac precursor and corresponding to the shift observed between U0 2 and U03 (ref. 10 and 11). It is more convenient to examine the shape of the peaks and the presence or the absence of the shake-up satellite structure associated to the U4f
574 levels.
a
!
!
'---_--r_ _
C
i !
/
ex-acac
eV
380
400
eV ex-nitrate
400
d eV
......_inr-----""'II'IiIP---~
Fig. 2. XPS spectra. 2 acac) NiUI Al 203 ex-acac, 2 nitr) NiUI A1203• a) oxidic precursor. b) H2 reduced. c) H2/H2S sulfided. d) ex-nitro U 4f 7/ 2' high resolution. On Fig. 2, are reported the different spectra corresponding to the U 4f 7/2 and 5/2 on the two catalysts. The spectra of the oxidic precursors (a) show, at low resolution (checked at high resolution) the absence of the shakeup satellite on NiU ex-acac while this satellite is present on NiU ex-nitro According to a preceding study (ref. 12), U03 is the only uranium oxide without a shake-up satellite associated to the 4f peaks; the 2p to U Sf excitation
°
would be the only possible shake-up process occurring in the final state which follows U 4f ionization. Then, only U02 and U30S (which can be written U0 2.2U03) will show the satellite. The high resolution spectrum reported on Fig. 2d shows that on NiU exnitr., the U 4f peaks are not doubled as we observed for U30S and could corres-
575
pond to the singulet of UOZ' (See ref. 13). Then, from the XPS studies of the precursors one can deduce : - the uranium oxide on NiU ex-acac is neither U30 g nor UO Z but presents a spectrum similar to U0 3• The superficial pseudo-epitaxial structure found by EXAFS is not in contradiction to this observation. - the uranium oxide on NiU ex-nitro is different from the preceding one;
the different values observed could cor-
respond to cristallites of UOZ' The binding energies found on the H2 reduced catalysts are very similar concerning the uranium oxide (probably U02_x) while almost 25% of the Ni is reduced on NiU ex-nitro and no reduction is observed on NiU ex-acac. A strong interaction between Ni and U (uranate) inducing a large dispersion could account for the Ni ex-acac behavior. On NiU ex-nitr., some Ni 20 3 cristallites, easily reducible, are probably present together with Ni aluminate and Ni uranate. The binding energy study of the sulfided catalysts did not give much information about the surface composition except the possible presence on the surface of metallic Ni, because one cannot distinguish between Ni 3S2 and Ni. (Table Z : 853.6 eV for Ni metal and 853.7 eV for Ni 3 S Z' i f this sulfide exists !). The surface concentration of the different metals, estimated from the XPS analysis, is very instructive. Table 3 shows these relative concentrations compared to the bulk composition. TABLE 3 Ni/U surface atomic ratio from XPS analysis. Catalysts
bulk composition
oxidic precursor
H2 reduced
ex-nitrate
0.52
0.3 ± 0.2
0.3 ± 0.2
ex-acac
0.55
1.8 ± 0.2
0.3 ± 0.2
The sulfided values are not accurate because of the Ni peak shape. The relative amount of Ni on the surface of NiU ex-acac is four times higher than the bulk composition while on the NiU ex-nitro the two values are very close. This enrichment is also consistent with the formation of a superficial highly dispersed nickel uranate which will be the precursor of a very highly dispersed sulfide as observed with Mo by EXAFS. The same value found on the H2 reduced precursors shows that the ex-acac catalyst must not be reduced before sulfidation because it will lose most of its peculiar properties.
576
The XPS study has shown the great importance of the impregnation procedure on the dispersion role of the uranium. CATALYTIC TESTS HDS Five series of NiMoU/A120 3 ex-acac catalysts have been tested with different loading of Mo (1%, 3%, 7%, 12% and 18%) and different Ni/Ni + Mo atomic ratio (from 0 to 1). On Fig. 3 is illustrated the activity of the three last series for the thiophene HDS expressed in mo1e/s.g(cata1yst). The reaction conditions were: 20 torr of thiophene, 760 torr total pressure, temperature 227°C, on a fixed bed flow differential reactor on presu1fided catalysts. The rates have been determined at very low conversion by plotting the conversion versus the inverse of the space velocity. The activity, in the same conditions, of the PROCATALYSE NiMoHR346 is also reported for comparison.
Act.HDS
15
10
5
Fig. 3. HDS activity of NiMoU/A1 203 catalysts. The full study on these catalysts will be published later (ref. 14) but one can see on this figure that the consequences on the activity of the dispersant role of uranium are obvious.
577
The pyridine HDN activity of two series of presulfided Ni catalysts are compared, NiU/Al Z03 ex-acac and Ni/Al Z03• The reaction has been performed at 450°C, under 4.25 torr of pyridine and a total pressure of 760 torr. On Fig. 4, one can see a very different behavior of the two series. In the presence of uranium,
the activity increases when the Ni concentration diminishes (higher
dispersion) while in the absence of uranium this activity decreases with the Ni concentration (formation of inactive aluminate). One must also observe that the activity of the NiU catalysts is always higher than the activity of the Ni catalysts whatever the Ni concentration. The study of the selectivity in products and of the total surface acidity has led us to conclude (ref. 5) that the differences between the two sets of catalysts could be understood if i) uranium blocks the formation of nickel aluminate because of the formation of a nickel
uranate, allowing also a very high dispersion of the Ni, and if ii) one assumes, with the fraction of uranium which is not involved in this uranate, the formation of a superficial uranium sulfide having a strong hydrogenating capacity. This last property, very well proved by the amount of hydrogenated pro-
ducts even in the absence of Ni or Mo, compared to the products obtained without uranium, is independent of the impregnation process and could account for the anticoking effect observed precedingly (ref. 1 and 2). (see Fig. 5).
Act. HDN
5 4
~
3
«
NiU
--
Ni
2
1
.
5
1~
%Ni
Fig. 4. HDN activity of U/A1 20 3 c, NiU/AI Z03., A1203- and Ni/A1 Z03& catalysts.
578
% Cs sat/ total Cs 20 r - = - - - - - - - - - - - - - - - - - - - - - - ,
NiU
15
10
5
(
•
Ni
&
5
0
10
% Ni
Fig. 5. Percentage of saturated C5 hydrocarbons in the total C5 fraction for the reactions reported on Fig. 4. ACKNOWLEDGMENT The ANVAR-Alsace is acknowledged for its financial support (nOA8302008). REFERENCES I 2 3 4 5 6 7 8 9 10 11 12 13 14
E.E. Davies and C.R. Pout, French Pat. 2205363. T. Nicklin, J. Clack, B.H. Holland, R.J. Whittaker and F. Farrington, Brit. Pat. 1 221 05!. T. Nicklin, Canad. Pat. 975534. M.J. Ledoux, G. Maire, R. Benazouz and G. Agostini, Eur , Pat. 84.401508.1. G. Agostini, These de Docteur-Ingenieur, Strasbourg, (1986). B.K. Teo and P.A. Lee, J. Amer. Chem. Soc., (1979), 2815. B.S. Clausen, H. T(lipsoe, R. Candia and B. Lengeler, Pr oc , of an International Conf., Exafs and Near Edge Structure, Stanford 1984, (Ed. SpringerVerlag), p , 18!. N.S. Chiu, S.H. Bauer and M.F.L. Johnson, J. Catal., 98, (1986), 32. W.-H. Pan, M.E. Leonowicz and E.I. Stiefel, Inorg. Chem., 22, (1983), 672 and T.R. Halbert, S.A. Cohen and E.!. Stiefel, Organomet., 4, (1985), 1689. J. Verbist, J. Riga, J.J. Pireaux and R. Candano, J. Electron. Spectrosc., 5, (1974), 193. R. Delobel, H. Baussart, J.M. Leroy, J. Grimblot and L. Gengembre, J. Chem, Soc ,; Faraday Trans.I, 79, (1983), 879. J. Weber and V.A. Gubanov, J. Lnor g, Nucl. Chem., 41, (1979), 693. J .L.G. Fierro, E. Salazar and J .A. Legarreta, Surf. and Interface Anal., 7, (1985), 97. M.J. Ledoux and G. Agostini, to be published.
579
DISCUSSION J.G. van OMMEN: Your preparation of uranium oxide monolayers is very much the same as our preparation of iron on vanadium oxide monolayers. If you start from UOZ (AcAc)Z' you get a stabilized monolayer of UZ03-like surface structure of uranium oxide, while starting from UOZ(N03)Z you get a U30g-like oxide. What is the reason why the UZ03-like structure is stabilized on the y-A1 Z03 ? M.J. LEDOUX: The monolayer will have a U03-like structure, not UZ03. The best explanation we have is that the pseudo-epitaxial structure blocks the uranium in its highest degree of oxidation (+VI) and does not allow the partial reduction which occurs generally with the metastable U03 bulk oxide (formation of UO Z + ZU0 3 = U30S)' B. DELMON : Your results indicate that, when sulfided, your catalysts (NiMoU/ A1Z03) possess a sort of U-Mo-S double layer on A1Z03 (Fig. 1). This is quite different from the small crystallites observed in other HDS-HYDTT catalysts, and this does not seem to leave space for the Ni promoter on MOS Z edges. Where do you believe the Ni promoter sits? M.J. LEDOUX: At low concentrations of Mo « 7% weight metal), we believe that Ni (or Co) is more probably bound to uranium (because of the strong interaction) than to Mo and this is why we observed a very poor synergetic effect (see Ref. 5). At high concentration of Mo. when most of the uranium monolayer is covered by Mo, the only possibility for Ni is to stay on the top of Mo (at least in the oxidic form of the precursor). LI Da-Dong : My question concerns impregnation. You said: "When the y-alumina is pretreated with an ethanol solution, this can provide a maximum amount of hydroxyl groups on the surface". Can you tell us more details about the amount of hydroxyl groups on the surface before and after pretreatment ? M.J. LEDOUX: In order to have both the maximum amount of OH groups on the surface (because of the chemical process of impregnation) and reproducible results (same amount of OH groups), we adopted the following procedure of pretreatment of the alumina. After a Zh calcination in pure air at 550°C, the y-alumina was impregnated with a 1:1 water-ethanol solution (the volume of water being equal to the pore volume of the alumina) for lZh, then left lZh at lZ0°C in dry air, and after rapid cooling, contacted with the acetylacetonate solution. No quantitative measurement of the density of the OH groups on the surface has been performed. J.R.H. ROSS: 1/ In addition to the work reported in your references Z and 3, Nicklin and his coworkers developed a successful series of high-temperature steam reforming catalysts based on Ni-UOz-aAlz03' The main feature of these was the formation of a nickel uranate phase. Is there any evidence for the formation of such a phase here ? Z/ One of the reasons for the withdrawal of the NUA catalysts from the market has been the problem of the (depleted) uranium in the catalysts. Do you think that this will be a problem with your materials? M.J. LEDOUX: 1/ The very high dispersion of the NiU catalysts (monolayer) made useless the XRD analysis and did not allow us to have pattern of nickel uranate. However, XPS analysis showed that the NiU ex-acac oxidic precursor had a very different nature than the NiU ex-nitrate and could be due to a pseudo-nickel uranate. Z/ Psychological behaviours are outside my competence! The lack of information of the public on depleted uranium is a problem that all the scientific community has to tackle because many uranium-based catalysts (oxidation, ammoxidation, steam reforming, hydrotreatment, CO/HZ' etc .•. ) are very efficient and selective.
580
G.M. PAJONK : Your method to prevent the migration of Ni 2+ species into alumina where it builds up the spinel is very interesting. Is there a lower limit of the U oxide layer than the full monolayer to obtain this effect 7 Beside is this barrier for Ni 2+ ions a chemical or just a physical one 7 J.M. LEDOUX: 14 to 16% of uranium fweight metal) corresponds to the optimum concentration on aluminas of -200 m /g. But 9-10% w.m. is already very efficient. Concerning physical or chemical barrier. a physical barrier at the level of a monolayer is also a chemical barrier. R. PRADA SILVY : My question concerns the preparation of the catalyst. What should be the influence of the impregnation procedure : a) consecutive U + Mo impregnation b) simultaneous Mo + U impregnation and c) consecutive Mo + U impregnation on the catalytic properties of Mo-U/A1203 system? How should be the Ni sulfided species distribution after impregnation and activation of each one of the catalysts prepared about and their implication on the total catalytic activity 7 J.M. LEDOUX: The aim of this work was to use uranium as a support additive (promoter 7) and not as an active phase. But we have tested some catalysts where Mo has been impregnated before. or together with uranium. These catalysts did not present any interesting results for HDS reaction. For the second question. see the answer given to Prof. B. Delmon.
B. Delmon, P. Grange, P.A. Jacobs and G. Ponce let (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
581
ON THE ADDITION OF VARIOUS METALS AS INORGANIC SALTS OR ORGANOMETALLIC COMPLEXES TO A MoS2-~AI203 HYDROPROCESSING CATALYST PREPARATION, CHARACTERIZATION AND HYDROGENATION ACTIVITY A. WAMBEKE 1, H. TOULHOAT 2, J.P. BOUTROIS 2, J. GRIMBLOT 1,3 and J.P. BONNELLE 1 et Homogene, U.A. C.N.R.S. 402, 'Laboratoire de Catalyse H~t~rogene 59655 Villeneuve d'Ascq Cedex (France). 2Institut Fran~ais
du P~trole,
B.P. 311, 92506 Rueil-Malmaison Cedex (France).
3To whom correspondance should be adressed. SUMMARY The samples prepared by addition of a metal salt or complex to the MoSYA120~ based catalyst are all active in toluene hydrogenation in presence 6f sulphor. Ni appears to be the more effective promoter whereas Cu has no detectable effect. Co and Fe promoter effects are intermediate. The experimental results indicate that the procedure used to prepare the catalysts is, in general, adequate to progressively decorate the supported MaS, slabs when the ratio Promoter/Mo increases. The 0 chemisorption measuremen~ on Fe and Ni promoted samples confirm that this mo~ecule reacts with all the coordinatively unsaturated sites. INTRODUCTION The active phase of hydroprocessing catalysts is made of MoS 2 crystallites, generally well dispersed as small single slabs, on aoy -A1 203 support. Usually, the catalytic activities involved in the hydroprocessing reaction, namely the C-S or C-N bond hydrogenolysis and the hydrocarbon hydrogenation are greatly enhanced by the presence of a second metal added to the MoS 2 phase. Several models have been proposed to account for the synergy (called also the promoter effect) between the two metals present on Al 203 in a sulphided form. Up to now, the most convincing description is that of the Topsee Group who claimed the presence, on the supported and unsupported catalysts, of a Co-Mo-S phase in which the cobalt (promoter) ions are located as individual species around the small MoS 2 crystall ites (1-2). A structural modell ing of such a phase composed of supported single MoS 2 slabs has pointed out the crucial role of the edges (highly unsaturated in the working conditions) at which the promoter species are located (5-6). From theoretical calculations, it was deduced that the second meta I (P) acts rea lly as a promoter if an electronic transfer occurs from P to Mo through the sulphur bonds (7-8). Genera lly the promoter ion (Co, Ni. .. ) is added to the catalyst by impregnation of a nitrate solution of the desired cation before. during or after the molybdate impregnation. The drying and calcination steps lead to the oxide
582
or commercial form of the catalyst. Acti vation transforms it into supported sulphides. An alternative to this very classical preparation procedure is to take benefit of the high react i vi ty of the edges of the supported MoS 2 phase to "decorate" them with the promoter by using a molecule or an inorganic salt which is decomposed in the subsequent step of activation. Obviously such a procedure, which will be described in this paper, implies that the supported MoS 2 phase is formed "in situ" to avoid any oxygen adsorption which occurs also at the edge positions. The catalysts prepared according to this method have been tested in toluene hydrogenation and the effect of the various promoters used has been compared. The 02 chemisorption capacity of the MoS 2Al 203 catalysts promoted by Ni or Fe have also been reported. EXPERIMENTAL Catalyst preparation The starting material we used was a Mo0 3 (14 wt%)-~AI203 catalyst calcined at 773K and moderately sulphided by an organic molecule which contains polysulphur groups according to the SULFICAT (r)process described in Ref .9. The promoter salt or the organometallic complex in solution (water or organic solvent) was added to this presulphided solid by a pore filling method. The concentration of the impregnating solution was chosen to give the final catalyst with the desired P (promoter)/Mo atomic ratio. After maturation and drying under vacuum at 423K, the solids were activated in the catalytic reactor with H2 at 573K. This treatment decomposes the promoter compound and during the same time the sulphidation into MoS 2 is completed. We tested several promoters -Fe,Ni,Co or Cu- initially present in different salts such as nitrate, acetate, octoate and stearate or in metalocenes (nickelocene and ferrocene). A Ni acetylacetonate complex was also used. Mo and Promoter contents were determined by X-ray fluorescence analysis. Oxygen chemisorption The oxygen uptake measurements were performed at 333K (10) in a pulse system on the catalyst previously activated under H2 at 573K and purged with an argon flow. Catalytic activity measurements The toluene hydrogenation reaction was carried out at 623K under hydrogen (60 atm.). The charge consisted of a mixture of cyclohexane (78 wt%), toluene (20 wt%) and 8000 ppm of sulphur present as thiophene, a molecule which is totally converted into H2S under these experimental conditions. The H2/hydro-
583
carbon ratio was 450 with a space velocity of 2 hr- 1 The products have been collected at intervals of 1 hr and analyzed by gas chromatography. We reported the steady state activities A calculated according to A = log (1/1-X) in which X is the toluene conversion calculated as X = E saturated hydrocarbons/Esaturated hydrocarbon + toluene. RESULTS AND DISCUSSION In Figure 1, the toluene hydrogenation activity A is plotted as a function of the atomic ratio a = P/Mo for the four series of samples. The nature of the promoter compound used for the preparation is noted in the symbol list. For comparison, the results obtained on the Mo0 3- Al 203 precursor and on the NiO-Mo0 3- A1 203 and CoO-Mo0 3- A1 203 industrial catalysts HR 346 and HR 306 from Procatalyse prepared according to classical methods (successive steps of impregnation, drying and calci nation) are al so reported. Prior use, these catalysts were sulphided with DMDS (dimethyldisulphur) spiked feedstock according to recommanded industrial procedure. By comparison with the catalyst without promoter, the activity linearly increases (except with the cobalt octoate), as a function of the atomic ratio P/Mo until a value a M which is close to 0.45 for the Ni series of samples prepared with the nickel salts. This aM value is simi lar to that previously determined on optimized Co-Mo or Ni-Mo industrial catalysts. In the modelling of supported MoS 2 catalysts (5,6), it was suggested that the promoter addition is optimized when all the edge sites of the MoS 2 single patches are occupied (decorated) by the promoter ions. The number of edge sites, re lati ve to the total number of Mo atoms in one single slab is a function of both size and shape of the MoS 2 cristallite. In the present investigation, all the samples are prepared with the same presulphided precursor and activated under the same conditions so that we suggest that the MoS 2 supported patches have an average size and shape which can accomodate 0.45 promoter ion per molybdenum. This saturation limit aM corresponds to the ratio r = P/P+Mo = 0.31 which is the value at which the maximum of catalytic synergy is generally observed for such catalysts (11). The linear increase of the catalytic activity when a
584
A 0.7
P=Ni
0.5 P=Co
0.3
• Co
OJ
--------. P=Cu
~:::::::::=------~~--------,*--
OJ
0.2
03
OA
0.5
0.6
*:
Fig. 1. Steady state hydrogenation activity A versus the atomic ratio a(P/Mo) 0: MaS, ; e : MaS? (sulph. DMDS) ; 6. : nitrate; acetate;. : metallocene ; • NiO-Mo0 3octoate , IJ : stearate ; ~ : acetylacetonate ; Al 203 and CoO-Mo0 3- Al 203 industrial catalysts sulphided with DMDS.
*
during the preparation steps may be invoked to explain such a behaviour. When a > aM (series prepared with Ni acetate), formation of bulk Ni sulphide is expected as the MoS 2 slabs are completely decorated and activity variations are small. If a part of the nickel ions at the MoS 2 edge position are concerned with the bulk Ni sulphide formation by phase segregation, then a loss of activity, as previously predicted (5-6), can be observed. This is the case for the Ni-Mo sample of highest For each series of samples, we can define (5) a promotion factor Op as being the ratio A(a = aM)/A(a= 0) between the activity of the completely promoted catalyst over the activity of the unpromoted catalyst. The Op values, reported in Table 1, confirm that nickel is the best promoter for this hydrogenation reaction.
585
TABLE 1
Promotion factor as deduced from figure 13.2
15.9C
3.1
Nib
Co
Fe
'"
1
Cu
tal NIckel salts (b) Nickel complexes (c) Calculated from the result on the industrial catalyst It can also be observed in figure 1 that the activity A begins to increase linearly even for low ~ values. This implies that all the added promoter ions are concerned with the MoS 2 decoration and the partial loss of the promoter ions into the support during the calcination step of classically prepared catalysts is avoided (5). In figure 2, a comparison is established between the promotion factor Qp we deduced from our results with those reported by Chianelli and Harris (8) about the hydrodesulphurization HDS of dibenzothiophene on various promoted-MoS 2
4
20
3
2
600
10
400 1
-*
200
1 Fe
Co
Ni
Cu
Fig. 2. Comparison between Qp. activity results on hydrodesulphurization of dibenzothiophene catalysts. Ahds (from ref. 8) and the theoretical activity parameter At h (defined in ref. 8) with the nature of the promoter P.
586
catalysts (classical preparation). In addition the theoretical estimation of the activity (activity parameter) is also reported. This theoretical approach predicts that a metal P acts as a promoter if an electronic transfer from P to Mo through p ....... S ...... Mo bonds occurs. The simi larity between the experimental results with the theoretical predictions seems to show that the promotion origin is mainly electronic for either aromatic hydrocarbon hydrogenation and HDS. Note only that Chianelli and Harris suggested a small poisoning effect of Cu in HDS whereas we observed no detectable effect in hydrogenation. In the literature on MoS catalysts, a relatively large number of 2-based papers is concerned with the use of probe molecules to better define the nature and number of the catalytic sites considered to be coordinatively unsaturated (C.U.S.). Oxygen, one of these probe molecule is reactive, even at low temperature. with MoS but for the moment the correlation between the 2 catalytic activity (HDS or hydrogenation reaction) with the amount of chemisorbed 02 is not always decisive (10,12). This is still more complicated when a promoter is present with MoS as the stoichiometry of chemisorption may be 2 modified. Despite these difficulties we have performed 02 chemisorption measurements on two series of promoted MoS
(P = Ni and Fe) samples and 2 plotted these data against the toluene hydrogenation activities (fig. 3).
A Q7
* •
as
*
0.3
_-
P.Fe
0.1
.... 1.0
15
....
2.0
-* N0 2
Fig. 3. Toluene hydrogenation acti yi ty A as a fu~ction of the amount of chemisorbed O? -NO - (in mol. 02 g- catalyst x 10 ). The same symbols are used in figure- 1. 2
587
Obviously these two independent measurements are fairly well correlated when considering a series of samples of a given promoter with different P/Mo ratios but the straight I ines have quite different slopes when comparing Ni and Fe samples. In the Ni series, two samples are outside the straight line. With nickelocene as the promoter precursor, some difficulties have been already mentionned concerning the catalytic activity and the result will not be further considered. The linear relationship we observe suggests that the number of 02 chemisorption sites is related to the number of catalytic sites. In the a = 0.41 Ni sample (from Ni acetate) almost all the edge sites of MO:~ are decorated with the Ni ions and the amount of 02 chemisorbed is 1.5.10 mol 02/ g catalyst. By comparison with the amount of 02 chemisorbed on MoS (1.0 x 10- 4 mol.0 2/g 2 catalyst), we can deduce that the stoichiometry O/Ni -edge position- is slightly higher to 1.5, a result quite simi lar to that previously reported (5). On the other hand, on the Fe-MoS samples large amounts of 02 are chemi2 sorbed whereas the activity variations are very small. Consequently, adsorption of the 02 molecule is not sensitive to the electronic factor which can predict promoter performances. 02 seems to be related to the number but not to the quality of the sites. CONCLUSION Impregnation of a presuIphided Mo0
starting material with several 3-rA1 203 salt or complex solution followed by an activation step which destroys the added compound and completely sulphides the supported phase leads to catalysts which are as active as the reference MoS2~A1203 system in toluene hydrogenation. The promotion factor decreases in the series : Ni> Co» Fe» Cu (With Cu no promoting or poisoning effect was detected). The fact that the activity variations are linear in general as a function of the ratio promoter/Mo up to "'0.45 implies that the MoS slabs are progressively decorate by the second 2 metal. Let us note that the industrial Ni-Mo-AI catalysts was optimized in 203 this way. The linear correlation between 02 uptake with hydrogenation activity but with two different slopes depending on the nature of P (Ni or Fe) leads to the conclusion that this molecule probes the "CUS" but this chemisorption is not sensitive to local electronic changes.
588
REFERENCES
2 3 4 5 6 7 8 9 10 11
12
H. Tops0e, B.S. Clausen, R. Candia, C. Wivel and S. M0rup, J. Catal. , 68 (1981) 433-452. C. Wivel, R. Candia, B.S. Clausen, S. M0rup and H. Tops0e, J. Catal .• 68 (1981) 453-463. H. Tops0e in J.P. Bonnelle et al. (Eds.), Surface Properties and Catalysis by Non-Metals, Reidel, Dordrecht, 1983, pp.329-360. H. Tops0e, R. Candia. N.Y. Tops0e and B.S. Clausen, Bull. Soc. Chim. Belg., 93 (1984) 783-806. S. Kaszfelan, H. Toulhoat, J. Grimblot and J.P. Bonnelle, Appl. Catal., 13 (1984) 127-159. S. Kasztelan, H. Toulhoat, J. Grimblot and J.P. Bonnelle, Bull. Soc. Chim. Belg .• 93 (1984) 807-811. S. Harris ana-R.R. Chianelli, J. Catal., 86 (1984) 400-412. S. Harris and R.R. Chianelli, J. Catal., gg (1986) 17-31. French Patent F 84-16540 assigned to EUR~T. J. Bachelier, M.J. Tilliette. J.C. Duchet and D. Cornet, J. Catal., 76 (1982) 300-315. J. Bachelier, J.C. Duchet and D. Cornet, J. Catal., 87 (1984) 283-291. P. Ratnasamy and S. Sivasanker, Catal. Rev. Sci. Eng~ 22 (1980) 401-429. P. Grange, Catal. Rev. Sci. Eng., 21 (1980) 135. B. Delmon, Proceedings "Climax Third International Conference on Chemistry and Uses of Molybdenum", H.F. Barry and P.C.H. Mitchell eds., Climax Molybdenum Company, Ann Arbor, Michigan, 1979, pp. 73-84. J.F. Le Page et al . "Catalyse de Contact" Editions Technip, Paris, 1978, p 236. S.J. Tauster, T.A. Pecoraro and R.R. Chianelli, J. Catal., 63 (1980) 515-519. R. Burch and A. Collins, Proceedings "Climax Fourth International Conference on Chemistry and Uses of Molybdenum", H.F. Bary and P.C.H. Mitchell eds., Climax Molybdenum Company, Ann Arbor, Michigan, 1982, pp. 379-383.
589
DISCUSSION B. DELMON : You present nice new approaches to preparation of HDS catalysts. Because of their novelty. this makes difficult a direct analogy with catalysts prepared in conventional ways. In such catalysts. R. Candia reports spectacular variations of the correlations between the amount of group VIII atom "decorating" the MoS2-like phase (so-called Co-Mo-S phase) and activity, according to preparation conditions. Compared to conventional catalysts, your method of preparation corresponds to major changes in the solid state chemistry transformations leading to the active sulfided species. and. hence. presumably. to major changes in the structure and texture of the active phase or phases. It is not unlikely that previously found correlations would fail in your case. You assert that the decorated system (Co-Mo-S type) is kept in your sample. What physico-chemical evidence have you? J.P. BONNELLE : We have no physico-chemical evidence that the decorated system (Co-Mo-S) is kept in our samples. Some Co or Ni salts give results very similar to the conventional catalysts with a linear increase of the activity versus the promoter concentrations up to the same optimum activity. In other cases particularly for promoter complexes (metallocene). some discrepancies appear which are probably due to incomplete decomposition or to changes in the struture and texture of the active phase or phases as mentioned in your question. A. LYCOURGHIOTIS : Why do you expect that your method of preparation decreases the interaction of the support with the promoter and thus enhances the decora~ tion of the MoS2 by the promoter ions ? J.P. BONNELLE : Our feeling is that the edge sites of the MoS2 slabs are reduced by hydrogen (creation of S vacancies) and thus very reactive to adsorb the promoter compounds. E.K. POELS : The proposal of Harris and Chianelli is. if I remember correctly. that the heat of formation of the Co-Mo-S and Ni-Mo-S mixed sulphides is particularly fit for the hydrodesulfurization reaction (due to reversible creation of S-vacancies). For the hydrogenation of toluene. I cannot see the necessity of such a requirement for the active site; couldn't it be the case that the Mo-A1203 system is merely a good support for Ni (or Co) dispersion limiting three-dimensional cubic nickel sulphide formation thus resulting in relatively high toluene hydrogenation activity. J.P. BONNELLE : The last proposal of S. Harris and R. Chianelli is a donor effect of nickel towards molybdenum ions. This effect can optimize the heat of formation of Mo-S bonds and so the reversible creation of S-vacancies. For the hydrogenation of toluene. the molecule adsorption takes place on sulfur vacancies and the donor effect is understood in terms of molybdenum back-donation to the aromatic ring which destabilizes the toluene molecule and increases the hydrogenation rate (J. Catal. 98. 17 and 229 (1986)). J.R.H. ROSS: Is there any evidence of migration of the metals into the alumina during use of the catalysts ? J.P. BONNELLE : No. The surface species. their repartition and relative amounts are the same before and after catalysis as seen by XPS and ISS.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
591
CONTROL OF CONCENTRATION PROFILES BY RATIONAL PREPARATION OF PELLETED HYDRO DESULFURIZATION CATALYSTS 1 P. GRANGE and B. DELMON J.L.G. FIERRO, Universite Catholique de Louvain, Groupe de Physico-Chimie Minerale et de Catalyse, Unite de Catalyse et Chimie des Materiaux Divises, Place Croix du Sud 1, 1348 Louvain-la-Neuve (BELGIUM)
1. On a sabbatical leave from Instituto de Catalisis y Petroleoquimica, C. S.I.C., 28006 Madrid, Spain ABSTRACT The objective of this work was to obtain Co and Mo concentration profiles into y-Al 0 extrudates by careful control of the impregnation parameters. It was shbw~ that the solute concentration, preparation procedure, and pH of the solution influence the shape and the extent of ion incorporation. Flat profiles were obtained bv the adsorption from solution procedure. at low pH and lona soakina time or additional preparative steps, while the incipient wetness impregnation procedure rendered always egg-shell (or degenerated egg-shell) profiles. We studied also the influence of the shape of the concentration on the HDS and HYD reactions of dibenzothiophene. INTRODUCTION Transition metals as oxides, sulfided, or in the reduced form, supported on alumina constitute the most important class of industrial catalysts. Most of them are used in the form of beads (approximately spherical) or pellets (cylinders) usually of dimensions of between 1 and 10 mm. Hydrodesulfurization (HDS), and more generally hydrotreating (HDN, HOM, etc.) catalysts very often correspond to the last category. Hydrotreating catalysts are commonly prepared by impregnation of dry or prewetted A1 203 pellets with ammonium heptamolybdate (AHM) (or, sometimes, some tungstate) and cobalt (or nickel) nitrate solutions. During impregnation and subsequent drying and calcination small crystallites or a molecular thick layer of the active ingredients are deposited on the internal surface of the A1 203 pellets (the so deposited oxides have to be reduced-sulfided for being catalytically active). In principle, these preparation steps are mass and/or heat transfer limited processes. Usually, impregnation equilibrium is not reached. This results in non-uniform concentration profiles along the radius of the pellet. Two excellent general reviews covering the problem of impregnation have been pUblished recently (1,2). All hydrotreating catalysts possess differenct catalytic functions: hydrogenolysis of the carbon-heteroatom bonds, hydrogenation, acid catalysed reactions,
592
etc. The reactions must be extremely well balanced for the maximum yield in useful products to be attained. The Group VIII/Group VI metal ratio is critical for the selectivity of the catalyst (hydrogenolysis/hydrogenation) as well as for the extent of the interaction of the Group VIII metal with alumina (formation of cobalt or nickel aluminate). On the other hand, various complicated diffusional processes take place during hydrotreatment, usually involving both a gaseous and a liquid phase. It ensures that manufacture of catalysts finely tuned to all the variants of hydrotreatment presently used in the petroleum and refining industry demands achieves precisely defined profiles of the active phase inside the pellets. Although relatively little attention has been given to the independent determination of the Group VIII and Group VI metal profiles, the variation of hydrogeno1ysis/hydrogenation activity with Group VIII/Group VI suggests that further increased performances could be achieved by controlling the Group VIII/Group VI ratio vs. radius profile. Our communication deals with this problem. An additional incentive in controlling independently the Group VIII and Group VI metals profiles comes from the remote control hypothesis (3,4) put forward for explaining the synergy between the metals of the two groups. If the hypothesis further strengthens, the nature of the active sites (hydrogeno1ysis and hydrogenation) would depend not only on the active phase composition, but also on the local H2/Sulfur ratio. As both concentrations depend on diffusion, the optima1ization of the wanted reactions would require an adequate Group VIII/ Group VI ratio vs. radius profile. Theoretical calculations indeed show that an adjustment of those profiles in pellets could lead to a remarkable change of the hydrodesulfurization/hydrogenation (HDS/HYD) selectivity (5). A careful control of the internal concentration profile of the active phase in pe11eted catalysts has been achieved since long in the case of Pt/A1 203 catalysts (6-8). A theoretical analysis of the concentration profiles arising during impregnation was given by Vincent and Merrill (9) and confirmed for Ni/ A1 20j catalysts. Cervel10 et al. (10) studied the impregnation of dry or wet A1 203 pellets with nickel solutions and concluded that short contact time renders uniform or degenerated "shell" catalysts, while longer impregnation time increased the surface concentration of the adsorbed nickel and produced a more uniform catalyst distribution. More recently, Komiyama and Merrill (11) have emphasized the influence of solution pH on the adsorption behaviour of Ni 2+ ions onto A1 203, and studied the control of impregnation profiles by the use of acids in the preparation of Ni/A1 203 catalysts. Attention has also been given to the preparation and behaviour of pel1eted alumina-supported chromia catalysts (12). Chen and Anderson used an electron probe microana1yser to determine concentration profiles in porous y-A1 203 spheres impregnated with aqueous solutions of chromium nitrate or chromic acid, and concluded that, with nitrate, the Cr was
593
deposited near the outside of the sphere, whereas with chromic acid the Cr penetrated deeper. The first comparable studies on Mo0 3/A1 203 catalysts were reported by Srinavasan et al. (13), who succeeded in obtaining a shell catalyst of Mo/A1 203 by chemical reaction of AHM, in the outer pores of pelleted alumina, with a reducing agent such as hydrazine. Shell catalysts so prepared have shown movement of the shell boundaries toward the pellet center on sintering at 750°C, while at 875°C molybdena was lost by vaporization (14). In the work presented in this communication, we attempted to prepare quite different types of molybdenum and cobalt distributions within the support pellet. For that we conducted a systematic investigation, based on the principles of surface and colloidal chemistry, capillarity and transport phenomena. EXPERIMENTAL PROCEDURES The catalysts were prepared by impregnating porous y-A1 203 extrudates (2.62 mm diameter. 10.6 nm average pore diameter) with aqueous solutions of AHM and Co(N03)2.6H20. Mo and Co were incorporated to the carrier following different procedures. One of these was the simUltaneous impregnation of Mo and Co according to the incipient wetness impregnation procedure. For this purpose, dry A1 203 extrudates were immersed in a solution containing both Mo and Co, whose volume exceeded by only 10% the overall pore volume of the alumina. The other preparations were carried out by two-step impregnation: Mo first and Co afterwards. A1 203 extrudates were immersed in a continuously-stirred solution of 0.018 (or 0.055) MAHM (volume of solid/volume of solution = 6.25) at p~3 and allowed to equilibrate for 1 (or 72) hour(s). The solution was then decanted and the extrudates kept at ambient temperature for 12h. Drying was at 100°C for 4h in a vacuum oven. The ultimate step was calcination. in two steps: 350°C for 2h and 500°C for 4h. Cobalt was subsequently incorporated by the adsorption from solution or by the incipient wetness procedure. Drying and calcination were the same as for the Mo preparations. A few extrudates for each preparation were imbedded in Eurepox resin and one half of the pellet was ground off along a plane perpendicular to the axis of the cylinder in order to reveal the cross section. Polishing was done with a se-; quence of diamond powders (sizes: 30-0.3 ~m) on a gyratory polishing plate. The polished section coated with a thick gold layer on a Balzers vacuum evaporation system. Profile analysis were obtained with a JEOL TEMSCAM 100 CX electron microscope equipped with a Kevex energy dispersive analyser. Point counting was performed along the whole diameter of the circular cross section at 50 ~m intervals in order to check the symmetry of the metal distribution. The results are presented either as Mo/Al or Co/Al vs. radius data. or. for the results concerning the Group VIII/Group VI variation, as the change of the atomic ratio r = Co/(Co + Mo) across the diameter.
594
RESULTS Impregnation of Molybdenum The adsorption of Mo on the A1 203 surfaces from aqueous solutions of AHM precursor is a complexe process governed by several factors. Most of them have been carefully analysed for the cases of particles of very small sizes (d < 0.2 mm) of the A1 203: in those cases flat concentration profiles along the diameter of the cross-section were usually obtained. In the light of these studies, we adopted several rules which satisfy the criterion of equal access, namely, all the points of the catalyst surface located at the same distance form the axis should have equal access to the same solute concentration. These were: i) the molybdena loading was controlled by varying the pH of the AHM solution; ii) the preparations were shaken for a constant time (lor 72h); iii) a relatively large volume of a dilute solution of AHM was used with a fixed quantity of A1 203 carrier (volume of solution/volume of A1 203 = 6.25) (15); and iv) the amount of Mo adsorbed must be greater than the Mo left in the pore volume of the carrier (dilute concentrations). Under the above experimental conditions, the Mo concentration is constant at equal distances from the axis of the pellets, and a symmetric distribution of Mo must be obtained. If the impregnation time is shorter than the one required to attain an uniform Mo-profile across the diameter of the pellet, non-uniform Mo profiles will be found. The same reasoning is also valid for the impregnation of cobalt. Screening of the Carrier Size Non-flat concentration profiles can, in principle, be obtained for any pellet diameter. However, the diameter should be large enough in order to let the preparation variables, viz., time of impregnation, Mo concentration, pH, presence of acids, etc., to act on the profile. For example, it was extremely difficult to prepare Mo-profiles other than flat when using A1 203 extrudates of 1.1 mm diameter (Fig. 1). Catalysts with egg-shell type of active phase profiles are very easily obtained using extrudates of greater size. In literature examples of metal oxides deposited on y-A1 203 concern CuO and Cr203 (12,16) using 5.5 rom diameter spheres, NiO (10,17) using 5 mm diameter pellets or 4 mm diameter spheres (11), and Mo03 using 4.8 mm diameter pellets (13), 3.9 x 3.0 mm ellipsoidal spheroids (14) or 4-4.2 mm diameter spheres (18). The sizes are quite representative of cases where qUite unequal distributions were obtained. Figure 1 shows a plot of typical Mo concentrations vs. radial position for a section through three different 1.10, 1.60 and 2.62 mm diameter cylindrical extrudates of calcined Mo0 3/A1 203 catalysts. Traverses were made along a diameter. As can be seen, even with constant preparation conditions, the shape of Moprofile depends markedly on the diameter of the pellet. From the results, we selected the 2.62 mm diameter A1203extrudates as the carrier to be used in
595
further studies.
d=t10mm c( <,
o
d=1.60mm
2
.10
d= 2.62mm
.05
~
o
d
d
J ~
o
d
Fig. 1. Influence of the size of A1 203 extrudates on the Mo profiles after impregnation with a 0.018 MAHM solution for 1h. Effect of pH Sonnemans and Mars (19) demonstrated that the Mo-loading could be varied by careful control of pH. They also suggested that lowering the pH induced a change of the solute concentration from the monomeric to polymeric ions. These differences can be understood on the basis of the various phenomena (including colloidal) which take place during the interaction between a solution of a polymerizable ion and a potentially charged surface. Several year ago Parfitt (20) and Brunelle (21) pointed out that the surface of an oxide should behave towards adsorption of ions in a similar way to colloids. A1 203, with an isoelectric point (IEP) in the pH range 6.0-8.0 (22), wull be positively or negatively charged when contacted with AHM solutions of pH lower or higher than the IEP, respectively. For these solutions, the relative concentrations of the anionic species are given by the ionic equilibrium, (1 )
which indicates that the adsorption of molybdate ions should be favored at high
596
pH. whereas at low pH the polymolybdate ions should dominate. From electrophoretic migration measurements. it was found that the IEP of the A1 203 used in these studies is 7.9. For obtaining AHM solutions with pH below and above the IEP. we added H2C204 or NH40H. respectively. Figure 2 shows the Mo profiles obtained by adsorption of a 0.018 MAHM solution for 1h at pH of 2.8. 6.2 and 10.8 respectively. As stated above. at pH = 10.8. a rather low Mo-loading and virtually flat profile inside the pellet is observed. whereas higher Mo-loading and degenerated Mo shell profile resulted at pH = 2.8.
-08
°1~-----7-----~
o
°1~----~------!
r/R
Fig. 2. Influence of the pH of the AHM solution on the Mo profiles in pelleted catalysts: 0 • 2.8; D , 6.2; and A , 10.8.
o
r/R
Fig. 3. Effect of the consecutive impregnations on the Mo profiles in pelleted catalysts :0, 1st impregn.; D, 2nd ir.lpregn .; A, 3rd i mpregn .
Consecutive Impregnation Consecutive impregnation with AHM can be used to enhance the catalyst load or uniformity. Figure 3 shows the Mo-profiles after one, two and three impregnations with a 0.018 MAHM solution at pH~3 for Ih. After each Mo impregnation, the precursor was dried and calcined to ensure that Me would not be redisolved in the subsequent impregnation. As can be observed, the Mo-loading increases with increasing number of impregnations. and the Mo-profiles tend to flatten out Furthermore, the difference of Mo-loading between two consecutive impregnations
597
decreases with increasing number of impregnations. This behaviour was expected since the zero point charge tends to decrease with increasing Mo-loading (15,23) thus decreasing the driving force across the length of the pores. Effect of Impregnation Time When the A1 203 is contacted with AHM solution, the liquid phase is imbibed into the pores due to capillarity, and simultaneously diffusion and adsorption start. This latter process, so-called diffusional impregnation accounts for the Mo-transfer from the exterior to the interior by diffusional relaxation, namely desorption of ions adsorbed near the mouth in the first place, diffusion, and readsorption. If the diffusional impregnation of Mo is performed as a step following the capillary impregnation, it may be accounted for by Fick's laws. According to the quasi-homogenous porous body-model (24), the relaxation time of Mo adsorption must be a function of Mo-loading on the carrier. It seems, therefore, appropriate to analyse several Mo-profiles obtained either at short or at long impregnation times. Figure 4a shows Mo impregnation profiles after the A1 203 extrudates have been contacted with a 0.018 MAHM solution for 1 and 72h, respectively, while Figure 4b shows the Mo profiles obtained after impregnation with a 0.055 M(three times more concentrated) AHM solution for the same lengths of time. As can be observed, whichever solute concentration, the Moprofiles almost levelled off at the longer impregnation time (72h), whereas a shorter impregnation time (lh) produced Mo-profiles which decreased toward the interior of the pellet, becoming sharper with increasing solute concentration. These characteristics, when combined with the effect of pH (cf. Fig. 2), provide an excellent basis for preparing Mo-profiles with the desired shape and Moloading (see below). Combined Profiles of Mo and Co The preparation of CoMo/A1 203 catalysts containing both Co and Mo with independently determined profiles is much more complexe than with Mo alone. Although the parameters considered above for the preparation of Mo-profiles are equally valid for the second component impregnation, some additional considerations must be taken into account. i) the cobalt incorporation is normally carried out by impregnation with a cobalt nitrate solution; ii) the extent of cobalt adsorption onto A1 203 increases with increasing pH, reaching its maximum value at a pH near the rEP of A1 203, but below that of the precipitation of cobalt hydroxide (15); and iii) in order to minimize the losses of inactive cobalt during intermediate calcinations of the precursors, e.g., surface and subsurface Co A1 204 phase, cobalt incorporation must be the last impregnation. It seems therefore evident (cf. Fig. 2) that the extent of Co and Mo incorporation follow an antagonistic dependence relative to the influence of the pH. Our objective was to combine
598
the Co and Mo profiles so that different r radius curves were obtained.
= Co/(Co
+ Mo) ratio values vs.
~4 0 ~
+
0
~.z
@
8 0 o
o
0 0
0
III
.« .... o
.~ ,2
'iii c
L
o
\'0'"
~
s
c
>~ I
><
:;;:
®
i
.'
°1L------:-.---------:::
°
0'
r/R
Fig 4. Combined effects of the AHM concentration : a) 0.018 Mand b) 0.055M, end the impregnation time: 0, lh; 0, 72h, on the Mo profiles of pelleted catalysts.
O'-------~----~
1
o
r/R
Fig 5. Cobalt and molybdenum impregnation profiles a) and their atomic ratios b).
Decreasing r Profile This profile was attained by the two-step impregnation procedures. For this purpose, the Mo/A1 203 precursor with an almost flat Mo-profile (Fig. 5) was impregnated with a cobalt nitrate solution by the incipient wetness impregnation procedure. The cobalt profile so obtained is of the egg-shell degenerated type (Fig. 5a). Because of the flat Mo-profile, the atomic ratio, r, decreased much more slowly inwards (Fig. 5b). Flat r Profile The most important characteristic of this profile is that the radial distribution of Co and Mo during the impregnation step must follow the same trend. It seemed of practical interest to incorporate both Co and Mo ions in one single step. This incorporation was carried out by simultaneous impregnation according to the incipient wetness procedure. The solution penetrated inside the pores by capillary forces, which produced degenerated shell profiles for both Co and Mo
599
~4
go
0
0
00
o 00
+
.3
0
u
3.2
0:
000-0-
'0
r = constant
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~
+
·2
I
0
~
°
<3
·1
°1
°
.2
·2
III
'"
III
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Vi
'"
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.'
i
~
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~
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1
Fig. 6 and 7 ratios b).
0
r/R
0
r/R
Cobalt and molybdenum impregnation profiles a) and their atomic
atoms (Fig. 6a). The Co-profile was below the Mo-profile since the solute concentration was lower than that of Mo. Increasing r profile Once Mo was incorporated to A1 203 extrudates by three consecutive steps (cf. Fig 3), cobalt was incorporated in a further step by adsorption at pH near 8.2. Since Co also tends to become more concentrated in the periphery of the extrudates, the impregnate was washed with methanol to decrease the Co level. In these conditions, the Co profile decreased slightly from the axis of the extrudates to their periphery (Fig. 7a), but, conversely, r increased (Fig. 7b). DISCUSSION The incorporation of Co and Mo onto dry A1 203 extrudates to produce extreme profiles of the active ingredients is a complex process which includes the penetration of solution by capillary forces, the adsorption of solute (Co and/or Mo) on the pore walls. and the transport of solute by diffusion in the pore
600
solution. Furthermore, during the drying step, the evaporation of water near the periphery of the particles concentrates the pore solution, producing a concentration gradient which drives more solute towards the center of the extrudates. If the evaporation of water is fast enough the concentration of solution within the pores would exceed the solubility, and the solute would precipitate, thus rendering a poorly disperse phase. As dry A1 203 extrudates are impregnated with an excess amount of AHM solution, the solution penetrates the particles to fill the pores by capillary forces. A concentration gradient of M070~4 (or MOO~-) ions in the pore solution is produced between the periphery and the center, and more molybdate ions enter a pore, part of these ions adsorbs on the pore walls and part remains in solution. The fraction of adsorbed molybdate ions, which is quantitatively described by the adsorption isotherm, is dependent among other factors, on the solute concentration and the pH of the solution. It is worth considering the effect of pH of the AHM solution. At low pH, the surface hydroxyl groups of A1 203 are protonated, r~o 05Al-OH + H+ _ Al-OH+ 7 24. (2) 2 The stability of the surface anion pairing seems to be enhanced by increasing the charge and size of the adsorbed anion. In acid solutions (H 2C204 at pH 3), the polymolybdate ions are preferentially adsorbed. However, the presence of a parallel pairing of C20~ions presumably occurs, since the Cl s photopeak in the X-ray photoelectron spectrum of this catalyst showed an overlaping peak whose binding energy is characteristic of carbonate species produced from the remaining oxalate ions during the calcination step (25). Factors other than charge may have some effect on the electrostatic adsorption. As D'Aniello (26) pointed out, the ionic charge density and its effect on the size of the hydration sphere around an ion in solution should have an important effect on the adsorption process. The above statement is also valid for the Co-impregnation. When Co incorporates (last step) at pH 8.2, the surface hydroxyl groups are ionized (Al-OH + OH- --. Al-O- + H20), hence the pairing of Co 2+ ions will be favored. An alternative method to incorporate Co and Mo in a single step is by impregnation with a solution of AHM and a negative Co-complex at PH 3. We were successful in preparing pelleted CoMo/A1 203 catalysts by impregnation of the A1 203 extrudates in a single step with M070~4 (or MOO~-) and CO(C204)~ions. However, the preparation is very sensitive to the charges of the individual ions in solution. Segregation of ions of different charges may also occur during adsorption or, if particularly high Mo- and Co- loadings are desired, the most highly charged (Mo70~4 ) ions can be preferentially adsorbed. Ideally, the Co-
601
complex should bear the same charge as M070~4 ions, and be of the same ligandtype. We emphasize that using a volume of impregnation solution equal to the pore volume is the simplest procedure. The method of adsorption from solution requires substantial soaking time and additional preparative steps. Most, but not all, of the effects of this method can be achieved with the incipient wetness impregnation by changing concentrations of solution or multiple impregnations. In short, by a careful control of the parameters which control the incorporation of the active ingredients into A1 203 interface and applying the ideas from colloidal chemistry, quite extreme Co- and Mo- impregnation profiles can be achieved. ACKNOWLEDGEMENT Financial support of this work was done from "les Services de Programmation de la Po1itique Scientifique" (Belgium). We wish to thank Dr D. Pirotte for making the electron microscope analyses. J.L.G.F. thanks CSIC (Spain) for a sabbatical leave. The authors are indebted to Dr L. Moscou, AKZO Chemie (The Netherlands) for supplying alumina carrier. REFERENtES 1. S.Y. Lee and R. Aris, Catal. Rev.-Sci. Eng. 27 (1985) 207 2. M. Kamiyama, Catal. Rev.-Sci. Eng. 27 (1985)~41 3. B. Delmon, Bull. Soc. Chim. Belg. 88:(1979) 979 4. B. Delmon, in "Proceedings 3rd Int:-Climax Conf. on the Chemistry and Uses of Molybdenum" (H.F. Barry and P.C.H. Mitchell, eds ) , The Climax Molybdenum Company, Ann Arbor, Michigan, 1980, p 33 5. J.M. Asua and B. Delmon, Ind. Eng. Chem. Fund. (in press) 6. R.W. Maatman and C.D. Prater, Ind. Eng. Chem. 49 (1957) 253 7. J.F. Roth and T.E. Reichard, J. Res. Inst. Cata~, Hokkaido Univ. 20 (1972) 85 8. G.H. van den Berg and H. Th. Rijnten, in "Preparation of Catalysts II" (B. De1mon, P. Grange, P. Jacobs and G. Poncelet, eds), Elsevier, Amsterdam, 1979, p 265 9. R.C. Vincent and R.P. Merrill, J. Catal. 35 (1974) 206 10. J. Cerve110, E. Hermana, J.F. Jimenez andlr. Me10, in "Preparation of Catalysts I" (B. De1mon, P.A. Jacobs and G. Ponce1et, eds), Elsevier, Amsterdam, 1976, p 251 11. M. Komiyama and R.P. Merrill, Bull. Chern. Soc. Japan 57 (1984) 1169 12. H.C. Chen and R.B. Anderson, Ind. Eng. Chem., Prod. Res. Dev. 12 (1973) 122 13. R. Srinivasan, H.C. Liu and S.W. Weller, J. Catal. 57 (1979) 8~ 14. P.R. Duncombe and S.W. Weller, AIChE Journal 31 (19~) 410 15. B. Arias, P. Grange and B. Delmon, in "ProceeQTngs of the 9th Iberoamerican Symp. on Catalysis", lisboa, 1984, p 1064 16. H.C. Chen and R.B. Anderson, J. Catal. 43 (1976) 200 17. J. Cervel10, J.F.G. de la Banda, E. Hermana and J.F. Jimenez, Chem. Ing. Tech. 48 (1976) 520 18. C.V. Caceres, M.N. Blanco and H.J. Thomas, in "Preparation of Catalysts III" (G. Poncelet, P. Grange and P.A. Jacobs, eds), Elsevier, Amsterdam, 1983, p 333
602
19. J. Sonnemans and P. Mars, J. Catal. 31 (1973) 20. G.D. Parfitt, Pure Appl. Chern. 48 (1~6) 415 21. J.P. Brunelle, in "Preparation of Catalysts II" (B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet, eds), Elsevier, Amsterdam, 1979, p 211 22. C.A. Parks, Chern. Rev. 65 (1965) 177 23. F.J. Gill-Llambias and ~M. Escudey-Castro, Chern. Comm. (1982) 24. P.B. Weisz, Trans. Faraday Soc. 63 (1967) 1801 25. J.J.G. Fierro, P. Grange and B. Delmon, in preparation 26. M.J. D'Aniello, J. Catal. 69 (1981) 9
603
DISCUSSION A. LYCOURGHIOTIS : I would like to make a comment and a suggestion to your very nice work. It is preferable to use the ZPC instead of IEP because the first parameter is a real surface parameter whereas IEP refers to the shear plane of the double layer. Moreover, I would like to suggest a slightly alternative explanation concerning the influence of pH on Mo profile: at pH < ZPC, the concentration of the adsorption sites for negative species (A10H2) is relatively high, inducing thus a rapid adsorption of the negative Mo species and therefore an enrichment of Mo in the outer part of the pellet. J.L.G. FIERRO: We thank you for your comment on the preferred use of ZPC instead of IEP due to the fact that this later parameter is directly related to the shear plane of the double layer. Our explanation on the pH effect on the Mo profiles is slightly different. At pH below the ZPC the protonated hydroxyl groups of A1203 surface increase, therefore, the adsorption of M070~4 ions is favoured. This explains the higher Mo concentration at low pH values, but the shape of the profile must be certainly controlled by mass transfer resistance of these ions from the periphery to the center of the pellet. A.J. van HENGSTUM": Upon covering A1203 with Mo03 the ZPC will decrease. Repeated Mo impregnations will therefore result in a modification of the interaction between the species in the solution and the surface of the support. Can you then explain the fact that Mo-profiles after the first, second, and the third impregnation have similar shapes? J.L.G: FIERRO: After the first Mo-impregnation the ZPC of the sample is lowered by about two pH units. It is, therefore, expected that at the second Mo-impregnation the interaction between Mo species and the surface should be weakened, and hence a lower Mo content incorporated at the second impregnation, and so on. But this is certainly an extremely idealized picture. During the second impregnation with AHM the Mo profile tends to be levelled off in the presence of NH! ions, incorporating more Mo in the inner part of the pellet with the subsequent Mo adsorption in the periphery of the pellet. Parallel to this phenomenon, the Mo present in solution in the pore volume of the pellets may precipitate during the drying step. Both effects would give rise to an "anomalous" high Mo content in the second and further impregnations. B. ANDERSSON : Can the pH effect on the concen profile be explained by transfer resistance since the larger M0 7024 ions will diffuse slower than Mo0 4-? The equilibrium between both ions is only effected by pH and not by adsorption on the surface.
mas~
6ration
J.L.G. FIERRO: There are two factors to be considered: the Mo concentration and the shape of ~~ profiles. The Mo concentration depends on the pH of the impregnation solution. At pH < ZPC the adsorption of M070~4 ions on positively charged OH groups of alumina surface is favoured, and hehce the Mo level must be higher than at pH values above ZPC. However, the shape of these profiles is influenced by the diffusion of polymolybdate ions from the periphery to the center the pellet. This diffusion does not seem to be different from that of Mo0 20f 4- ions since the Mo level reached at the axis of the pellet is lower when Mo impregnation is conducted in basic medium.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
605
INFLUENCE OF THE ACTIVATION PROCEDURE ON THE NATURE AND CONCENTRATION OF THE ACTIVE PHASE IN HDS CATALYSTS R. PRADA SILVY, J.L.G. FIERRO, P. GRANGE and B. DELMON Groupe de Physico-chimie Minera1e et de Cata1yse, Unite de Cata1yse et Chimie des Materiaux Divises, Universite Catho1ique de Louvain, Place Croix du Sud 1, B-1348 Louvain-la-Neuve (Belgium) SUMMARY The influence of two different activation procedures (reduction-su1fidation by consecutive or simultaneous reactions) on the catalytic properties and the structural changes of a commercial hydrodesulfurization catalyst in function of the temperature (400-800°C) was studied. It was concluded that the activation by consecutive reduction-sulfidation reactions at temperatures above 400°C produce a strong catalytic activity loss. Physico-chemical analysis carried out on samples reduced at temperatures above 400°C indicated several structural changes: i) CoA1204 decrease ii) metallic Co increase and iii) simultaneous MoO? and Al?(Mo04)~ formation. IR measurements of adsorbed NO and the sulfur content anarysis indicated that the sulfidation of the MoO species takes place at temperatures above 530°C. The XPS analysis carried ouf on catalysts reduced at te~peratures above 530°C and then sulfided at 400°C evidenced two types of Mo+ species: one is easily reducible or sulfidab1e at moderate temperatures and the other, difficult to reduce or sulfide at temperatures below 600°C. The activation by simultaneous reduction-su1fidation reactions produce catalysts with catalytic activity optimal in an extended temperature range (400-650°C). INTRODUCTI ON The activation of cobalt-molybdenum hydrodesulfurization (HDS) catalysts (namely reduction-sulfidation) is generally carried out at moderate temperatures (300-450°C), under pressures between 2-50 bar of sulfiding and reducing agents (1, 2). For a same batch of catalysts, the effectiveness of these catalysts depends on several parameters: i) nature of the activating molecules ii) sulfidation temperature and iii) activation procedure. We have previously studied how the nature of the sulfiding molecules (DMDS, thiophene, CS 2, butanethiol, H2S) and the su1fidation temperature (using a H2S-H 2 mixture) modify the activity and selectivity of a commercial HDS catalyst (3, 4). However, very little attention has been paid in the literature on the activation procedures so far. Three different procedures can be used to activate a HDS catalyst: 1(1) reduction (H 2) -+ su1fidation (H 2S\ Oxide precursor _(2) reduction + sulfidation (H 2S+H _ sulfided catalyst 2) ~(3) sulfidation (H 2S)-reduction (H 2
l
606
Recently it has been shown that, at 400°C, procedures (1) and (2) allow an optimal activation of the catalyst. The physico-chemical analysis indicated differences in structure and composition of the supported phases. The activity of the samples activated according to procedure (3) is poor (5). This communication is part of a research program aimed at a systematic investigation of the activation procedure on the nature and concentration of the active phase in a HDS catalyst. Because of the differences observed previously, we study 3 catalyst series which have been activated starting from the activation procedures (1) and (2). Correlations between the catalytic activity and the physico-chemical analysis by XPS, IR spectroscopy of adsorbed NO and sulfur content are reported. EXPERIMENTAL Catalyst and Activation procedures A commercial Co-Mo-/A1 203 catalyst (Procatalyse HR-306) with 3% CoO, 14% Mo0 3 and BET surface area of 195 m2g-1 was employed. Three different catalyst series have been obtained activating the precursor oxide by consecutive or simultaneous reduction-sulfidation reactions (starting from the activation procedures (1) and (2) mentioned above) at temperatures between 400 - 800°C. These catalyst series have been designated as (lA), (lB) and (2). The experimental activation conditions are described as follows: ~~r!~~_{!~l : two grams-samples (0.1 - 0.2 mm grain size) were reduced in H2 (6.L.h- l) at 400, 530, 650 or 800°C, for 4 h and after that, sulfided with a H2S-H 2 mixture (15% V H2S, 6.L.h- l, at the same temperature, at atmospheric pressure for 4 h. ~~r!~~_{!~l : the samples have been reduced in the same conditions than for the previous series but sulfided at 400°C for 4 h. ~~r!~~_igl : the oxide precursor was sulfided directly with the H2S-H 2 mixture (15% V H2S, 6Lh- l) at the different temperatures for 4 h. The samples activated at temperatures above 400°C, were previously sulfided at 400°C and then heated progressively (15°C min-l) up to the final activation temperature (530, 650 or 800°C) under the H2S-H 2 flow. The reactor was maintained isothermically during 4 h at the final temperature. After activation, the samples, were cooled down in argon and stored under isooctane avoiding air contact for physico-chemical characterization (6). Catalytic testing After introduction in the high pressure catalytic reactor, 2 grams of 42 Lh- l) the samples were re-sulfided "in situ" with a H2S-H 2 mixture (4% H 2S, at 400°C, and 30 bar, for 2 h. Their activities for the HDS of thiophene and
607
hydrogenation (HYD) of cyclohexene were measured at 237°C, at a total pressure of 30 bar, using a feed containing 0.5% thiophene, 30% cyclohexene and 69.5% cyclohexane (7). Physico-chemical characterization XPS : The X-ray photoelectron spectra (XPS) of the catalysts activated by the different procedures, were recorded on a Vacuum Generators [SCA 3 spectrometer equipped with a Tracor Northern TN-17l0 signal averager which allowed improvement of the signal-to noise ratio. The samples were pressed into the sample holders under iso-octane and kept protected from exposure to air by a meniscus of this liquid which was removed during pumping of the preparation chamber of the spectrometer. The Cl s' AL 2p' Mo 3d, M0 3p' S2p and c02p lines were recorded for each catalyst. All binding energies were referenced to Cl s at 285 eV. IR spectroscopy: The infrared measurements of adsorbed NO, were carried out in a cell assembled with greaseless stopcocks and NaCl windows which allowed thermal treatments under vacuum or under a H2 and H2S-H2 mixture flow. The reduced catalysts were re-reduced "in situ" at 400°C for 2 h. For the sulfided samples the re-sulfidation was carried out with a H2S-H 2 mixture (15% H2S) using the same conditions as described above. Before contacting with NO (99.6 %V), the catalysts were outgassed at 400°C for 2 h. Infrared spectra were obtained with a Perkin-Elmer 580-B spectrophotometer, interfaced to a data system. Dynamic NO chemisorption: Dynamic NO chemisorption (DNOC) at room temperature was measured using a pulse technique. The different catalyst series were re-sulfided using a H2S-H 2 mixture (15% H2S) at 400°C in a U-tube pyrex reacto~ at atmospheric pressure for 2 h and flushed in argon for 0.5 h at this temperature. The catalysts were maintained in a flow of He at room temperature and thereafter, 2 ml pulses of 5% NO in He were injected until the equilibrium was reached. Results are expressed as ~mol of gas adsorbed per gram of catalyst. Sulfur content: The catalysts sulfided by the different procedures, were burnt at 1100°C under a stream of oxygen. The evolved sulfur gases (S02 and S03) were absorbed in a solution of hydrogen peroxide and then the sulfuric acid formed was titrated with a standard solution of sodium tetraborate (6, 8).
608
RESULTS Catalytic activity The HDS and HYD conversions are presented in figures la and lb respectively as a function of the temperature and the activation procedure.
50 '0 III
t:l
::t 30
. -;,t
20 10
rb I
25
20 t:l
),.
::t 15
. -;,t
10 5
'00
500
600
700
PRETREA TMENT TEMPERA TURE
800 ('CI
Fig. 1. Conversion in HDS and HYD for the catalyst series (lA),(lB) and (2) as a function of the pretreatment temperature For the catalyst series (lA) and (18), the conversions for both reactions decrease progressively with increasing temperature. For the catalyst series (lB) the loss of the catalytic activity is much more pronounced than for the series (lA). For the catalyst series (2), the catalytic activity remains constant at sulfidation temperatures between 400 and 650°C and decreases for the sample activated at 800°C. Physico-chemical characterization XPS : Figures 2 and 3 show the c02p3/2 and M03p3/2 energy levels, respectively corresponding to H2-reduced catalysts (a) and the catalyst series (lA)(b), (lB) (c) and (2) (d) as a function of the temperature. As it can be seen, both binding energies (BE) and peak intensities are substantially altered by the activation conditions.
609
... ~
.'"
l!!
.....: l!!
....
...'"
e,
.. ....
"''''
-:
~~
e-,
.... "'''' .... ..... o:~
......
~
BE I,V)
'" ...-: '" 0: ..; ... l!! ~ e-,
Fig. 2. C02p3/2 energy levels corresponding to : a) reduced catalysts b) series (lA). c) series (lB) and d) (2) as a function of the temperature.
.... ...'" 'Ii .;
..
'" '"
'" '"
~'
'" ",'
e,
.. .'" '" ~
'
'" '"
..
BEI,v)
'" '" '" '" ~
.;
Fig. 3. M03p3/2 energy levels corresponding to : a) reduced catalyts b) series (la), c) series (lb) and d) (2) as a function of the temperautre.
610
CO 2 : For the catalysts reduced at temperatures between 400 and 800°C -p (figure 2A), the main and the satellite peaks (782.3 and 787.7 eV respectively) corresponding to C0 2+ ions in oxidized state are observed (presumably coA1 204like species). At reduction temperatures above 530°C, two new peaks appear at 778.7 and 793.7 eV. Their intensities increase with the reduction temperature. They can be assigned to C02p3/2,1/2 levels corresponding to cobalt atoms in the metallic state (9). The c02p3/2 spectra for the catalyst series (lA), (lB), and (2) given in the figure 2b, 2c and 2d, respectively, show the peaks corresponding to the CoA1 204 species. An additional signal situated at 779,5 eV can be attributed to C0 2+ ions in sulfided state (C0 9S8), CoS l+x' sulfided Co in CoMoS-like species, etc) (10,11,12). Comparing the spectra of the catalysts sulfided at 400°C, the highe~concentration of Co sulfided can be observed in the catalyst series (lA). At sulfidation temperatures above 400°C, (series (2)) the signals corresponding to the CoA1 204 decrease progressively and disappear for the sample sulfided at 800°C. But, a broad signal at about 786 eV is observed for the catalyst series (lA) and (lB) at temperatures above 400°C. This signal increases progressively with increasing temperature. This last signal is much more intense for the catalyst series (lA). M0 3p : Figure 3a shows the M03p3/2 energy levels corresponding to H2-reduced catalysts as a function of the temperature. The signal at 398.3 eV for the samples reduced at 400°C indicates that the Mo+ 6 ions are the dominant species. At reduction temperatures above 530°C, the lines are broadened (contribution of Mo+ 4, Mo+ 5 and Mo+ 6 ions) and their positions are slightly shifted towards lower BE. The peak at 395.2 eV corresponds to the Mo+ 4 species in the oxidized state (Mo0 2-like species) (13). The M03p3/2 energy level corresponding to the catalyst series (lA), (lB) and (2) (figure 3b, 3c and 3d respectively) show a peak at 395.0 eV which can be attributed to Mo+ 4 ions in sulfided state (MoS 2 like-species) (10,11). At 40QoC or above, the MoS 2-like are the prevalent species of the catalyst series (lA) and (2). For the catalyst series (lB), the intensity of the signal at 398,3 eV corresponding to the Mo+ 6 species is observed for the samples pre-reduced at temperature above 400°C. The concentration of these species is much higher in the reduced catalysts at 800°C. Infrared of adsorbed NO Typical IR spectra of NO chemisorbed on catalysts reduced at temperatures between 400 and 800°C are given in figure 4a. The three main bands between 1700 and 1900 cm- l represent NO chemisorbed on both Co and Mo exposed on the surface. The band near 1877 cm- l is characteristic on Co ions (symmetrical
611
stretching band), while that near 1700 cm- l is characteristic of NO chemisorbed on Mo+ o ions in reduced state (asymmetrical stretching band). The central o band (1795 cm- l) contains contributions from NO on both Co and Mo+ ions (asymmetrical and symmetrical stretching bands, respectively) (14). Comparison of the intensities of the bands at 1877 and 1700 cm- l, for the samples reduced at temperature above 530°C shows that the surface concentration of the Co and Mo ions decrease strongly. The IR spectra corresponding to catalyst series (lA), 18), and (2) as a function of the temperature are presented in figures 4b, 4c and 4d respectivel~ Three principal bands similar to those obtained for the spectra of the reduced catalysts are observed and shifted downward. The signal at 1852 cm- 1, is attributed to NO chemisorbed on Co+ 2 ions in the sulfided state (symmetrical stretching band). The band at 1680 cm- l represents the NO adsorbed on su1fided Mo species (asymmetrical stretching band). The central band (1790 cm- l) contains both contributions of NO chemisorbed on Co and Mo sulfided species (asymmetrical and symmetrical stretching, respectively) (14). At sulfidation temperatures between 400 and 530°C, for the catalyst series (la) and (18) (figure 4b and 4c respectively) the spectra show an additional band situated at 1705 cm- l, with corresponds to NO chemisorbed on Mo+ o ions in the reduced state. At temperatures above 530°C three simultaneous effects are produced: i) the asymmetric band corresponding to NO chemisorbed on Mo+ o ions in the reduced state (1705 cm- l) disappears. ii) The asymmetric band intensity of NO adsorbed on MO+ o ions in su1fided state (1680 cm- 1) decreases. iii) The intensity of the symmetrical band corresponding to NO chemisorbed on Co+ 2 ions (1852 cm- 1) increases strongly. The two last effects, can be observed for the catalyst series (2) in the same temperature range. Dynamic NO chemisorption Table 1 shows the influence of the temperature and the activation procedures on the NO uptake at room temperature. For the three activation procedures, the NO uptake varies in a similar manner to the catalytic activity. However, there are differences between the total amount of NO chemisorbed on the different catalyst series which varies as follows: (18) > (lA) > (2). Sulfur content The varia~on of the sulfur content in the catalysts as a function of the temperature and the activation procedure are shown in table 1. For the catalyst series (lA) and (2), the sulfur content remains constant between 400 and lDJ"C;
612
while, for the catalyst series (lB) the sulfur with increasing reduction temperature.
2OllO
1eoo
1100
2OllO
leoo
1100
ZCOII
1100
amount decreases progressively
1100
ZCOII
1100
1100
Wav.number (em"')
Fig. 4. IR spectra of adsorbed NO corresponding to : a) reduced catalysts b) series (lA), c) series (lB) and d) (2) as a function of the temperature. TABLE 1
Temperature (OC) 400 530 650 800
NO uptake (mol.g -1 ) cata1yst seri es (lB) (2) (lA) 100 117 117 101 138 115 99 109 124 79 97 100
Sulfur content catalyst series (lA) (lB) (2) 7.0 7.0 6.9 6.0 7.2 4.6 7.0 7.1 4.0 6.8
NO uptake and sulfur content analysis corresponding to the catalysts series (lA), (18), and (2). DISCUSSION In other to explain the evolution of the catalytic activity, it is convenient to study the physico chemical changes induced by the pre-reduction and the sulfidation conditions of the catalysts as a function of the temperature.
613
The XPS analysis indicate that H2-reduction temperatures above 530°C produce structural and composition changes: i) CoA1 204 decreases ii) Coo increases and iii) Mo0 2-like species crystallizes. The increase of Coo exposed on the surface seems to be due to the reduction of CoA1 204, XPS, XRD (15) and TPR analysis (16) carried out on pure CoA1 204 reduced at temperatures above 600°C have confirmed these changes. The presence of Mo0 2-like species in samples reduced at temperatures above 400°C, explains satisfactorily the loss of catalytic activity observed for the catalysts series (lA) and (lB). It is known that these species are inactive towards the HDS reaction and difficult to reduce or sulfide at moderate temperatures(17,18). The IR spectra of adsorbed NO for these two catalyst series demonstrated this fact. The presence of the band corresponding to NO chemisorbed on Mo+~ ions in the oxidized state (1705 cm- l) in addition to the band corresponding to the sulfided state (1680 cm- l), for the sulfided samples at 400 and 530°C shows the high stability of the Mo0 2 species towards sulfidation, Another evidences is the following: the sulfur content for the samples reduced at 650 and 800°C sulfided at 400°C (table 1) decreases sharply. If the sulfidation of the H2-reduced catalyst series is carried out at temperatures above 530°C (series (lA)), the sulfur content in the samples increase strongly. These results are in agreement with those obtained by Zabala (17) and Arnoldy et al (18) which indicate that the sulfidation of unsupported Mo0 2 species or supported Mo/A1 203 catalysts takes place at temperatures above 600°C. In addition to the XPS analysis of the reduced catalysts, in all the spectra the signal corresponding to Mo+ 6 species (398.3 eV) was observed. For the catalyst series (lB) this signal was also present and much more intensive in the sample reduced at 800°C (figure 3c). Several researchers have proposed different Mo+ 6 species in Mo/A1 203 and Co-Mo/A1 203 catalysts. Arnoldy et al (18) and Thomas et al (19) have studied by TPS and TPR the reactivity and stability of the Mo+ 6 species in supported Mo/A1 203 catalysts. They have observed two types of Mo+ 6 species: One is easily reducible or sulfidable at lower temperatures (bi-multilayer Mo+ 6 species in octahedrally (0) coordination). The other is difficult to reduce or to sulfide at lower temperatures (monolayer of Mo+ 6 species in tetrahedrally (T) or octahedrally coordination which have a strong interaction with the A1 203 support). Ratnasamy et al (20), propose three different Mo+ 6 species in a Co-Mo/A1 203 catalyst: a non reducible phase (A) of MO(T) which is A12(Mo04)3-like; an easily reducible phase (B) of bulk Mo0 3; and an intermediate reducible phase (C), probably Mo(O)' whose reducibility is increased with Co. In agreement with these studies, our XPS results seem to indicate that during the reduction of the catalysts at higher temperatures,
614
species like A1 2(Mo04)3 could be formed by two different ways: i) by a decrease of the coordination state of the Mo monolayer species fromoctahedrical to tetrahedrical by oxygen remotion during reduction or ii) by solid state reaction between the Mo species with the alumina support. As A1 2(Mo04)3 species arevery stable and difficult to reduce or sulfide at low temperature. The formation of these species could easily explain the presence of the peak at 39B.3 eV corresponding to Mo+ 6 observed for the sample reduced at BOOoe and sulfided at 400 0 e. This peak disappear if this sample is sulfided at BOOoe (series lA). This hypothesis is also supported by the sulfur content analysis which is lower for the catalyst reduced at BOOoe and then sulfided at 400 e (series 18) than the catalyst sulfided at BOOoe (series lA). For the catalyst series (2) the catalytic activity did not change at temperatures between 400 and 650 0 e. However, several changes in structure and composition of the supported phases have been observed by Mossbauer Emission Spectroscopy (MES) on aCo-MolAl 203 laboratory catalyst activated in the same temperature range (21). This catalyst showed a similar catalytic behaviour than the commercial one, and it is assumed that the structural changes observed would be comparable. The following effects were observed at sulfidation temperatures above 400°C : i) The support (A1 203) and the supported phases tend to sinterize together. ii) The concentration of C0 9SB and COS l+x increase. iii) The concentration of CoA1 204 and CoMoS decrease. The increase of the symmetric stretching band intensity corresponding to the NO chemisorbed on C0 2+ ions observed for the samples activated at 650° and BOO°C, could be explained by two antagonic processes. In one hand, the unsaturation number associated with the Co species decreases during the sintering at higher sulfidation temperatures (under these conditions, the adsorption mechanism of the NO molecules could change from dimers (22) or dinitrosyls (23) to monomeric species). On the other hand, the presence of COS l +x species (observed by MES) in these catalysts, could be responsible for the observed spectroscopic changes. The higher concentration of these species could be due either to the sulfidation of CoA1 204 (confirmed by XPS and XRD analysis on pure CoA1 204 sulfided at 650° and BOO°C) or to decomposition of the CoMoS phase. This last hypothesis seems to be reasonable if we compare our results with those obtained on NO chemisorbed on Ni/A1 203 catalysts. Topsde et al (24) observed that the NO adsorption on sulfided Ni/A1 203 catalysts give one strong adsorption band attributed to Ni 3S2 or NiS-line species. As COS l +x and Ni sulfided species have similar structure (hexagonal symmetry), the analogy between both Co and Ni systems may be support our hypothesis. The NO chemisorption on Mo/A1 203 and Co-Mo/A1 203 catalysts, has been correlated 0
615
with the coordinative unsaturation developed during the reduction or sulfidation (25). Results of table 1 indicate that for the three catalyst series the NO uptakes vary in a similar manner to catalytic activity. However, a direct correlation between the total amount of NO chemisorbed and the catalytic activity was not observed although both depends on : i) Dispersion state of the supported phases. ii) Vacancy number associated to the Mo species. iii) Reduction-sulfidation degree of the catalyst. iv) Nature and concentration of the promote atoms. Okamoto et al (26) have studied the adsorption of NO on reduced and sulfided Mo/A1 203 catalysts. They have observed that the adsorption of NO on sulfided catalysts is as strong as that on reduced catalysts. In our case, the reduction of the catalysts at temperatures above 400°C not only increases the number of unsaturation on the Mo species but also could break the interactions between the Co and Mo from the bidimentional bi-layer (CoMoO x) taking place on separated phases. The subsequently sulfidation of these catalysts would produce an increase of the sulfided Co concentration (species type C0 9S8, CoS l+x). These species should in part be responsible for the increases of the NO uptake and the greater XPS signal corresponding to the Co sulfided (779.5eV) observed for the catalysts series (1A) and (lB) (figure 2b and 2c, respectively). This suggestion could explain the differences between the total NO uptake observed for the different catalyst series and the catalytic activity results. CONCLUSION · We have shown that the activation of a Co-Mo/A1 203 HDS catalyst by consecutive reduction-sulfidation reaction at temperatures above 400°C produces a strong catalytic activity loss. Whereas if the reduction-sulfidation reaction is carried out simultaneously the activity remains stable in an extended temperature range (400-650°C). · The pre-reduction stage at temperature above 400°C produce several physicochemical changes : i) Co metallic increase. ii) CoA1 204 decrease. iii) simultaneous Mo0 2 and Al 2(Mo04)3 formation with a consequent decrease in catalytic activity. · IR of adsorbed NO measurements and the sulfur content analysis shown that the sulfidation of the Mo0 2 species takes place at temperatures above 530°C. · Two different Mo+ 6 species have been evidenced by XPS ; one is easily to reducible or sulfidable at moderate temperatures and the other difficult
616
to reduce or sulfide at temperatures below 650°C (species type A1 2(Mo04 )3) ' . There is not a direct correlation between the NO chemisorption and the catalytic activity although both depends on i) Dispersion of the supported phases. ii) Unsaturation number of the Mo species. iii) Degree of sulfidation of the catalyst. iv) Nature and concentration of the promote atoms. ACKNOWLEDGEMENTS The financial support of the "Services de Programmation de la Politique Scientifique (SPPS)" (Belgium) is gratefully acknowledged. R. Prada Silvy thanks the "Centro de Formation y Adiestramento Petrolero y Petroquim;co (CEPET)" (Venezuela) for a grant. We thank also to Mrs. M.E. Prada and Mrs. J. Liagre-Dieux for illustrations and typing this communication. REFERENCES 1 H. Hallie, Ketjen Catalysts Symposium, p. 49, Amsterdam, 1982. 2 B. Gates, J. Katzer, G.C.A. Schuit, Chemistry of Catalytic Processes, Mac Graw-Hill, Chap 5, 1979. 3 R. Prada Silvy, P. Grange, F. Delannay, B. Delmon, IXth Symposium Iberoamericano de Catalisis, 1141, Lisbon, 1984. 4 R. Prada Silvy, P. Grange, F. Delannay, B. Delmon, Polyhedron, V5, Nl/2, 195, 1986. 5 R. Prada Silvy, P. Grange, B. Delmon, to be presented in Xth Symposium Ibero-americano de Catalisis, 6-11 july 1986, Merida, Venezuela. 6 R. Prada Silvy, M. Sc thesis, Louvain-la-Neuve, 1984. 7 R. Prada Silvy, J.M. Beuken, P. Bertrand, B.K. Hodnett, F. Delannay, B. Delmon, Bull. Soc. Chern. Belg., 93, 775, 1984. 8 R.P. Jone~, P. Gale, P. Hopkins, C.N. Powell J. of Iron and Stell Institute, 505, May 1966. 9 J. Bonnelle, J. Grimblot, A. D'Huysser, J. Elect., Spect. and Rel., Phenom, 7, 151, 1975. 10 P. Gajardo, P. Grange, B. Delmon, Surf. Int. Anal., 3, 206, 1981. 11 R.F. Declerch-Grimme, P. Canesson, R.M. Friedman, J.J. Fripiat, J. Phys. Chern., 82, 885, 1978. 12 I. Alstrup, I. Chorkendorff, R. Candia, B. Clausen, H. Tops0e, J. Catal., 77, 397, 1982. 13 T. Patterson, J. Carver, D. Leyden, D. Hercules, J. Phys. Chern., 80, 15, p. 1700, 1976. 14 J.B. Peri, Preprints AM Chern. Soc., Div. Petrol. Chern., 29, 889, 1984. 15 R. Prada Silvy, P. Grange, B. Delmon, in preparation. 16 P. Arnoldy, J.A. Moulijn, J. Catal., 93, 38, 1985. 17 J.M. Zabala, Ph.D Thesis, Louvain-la-Neuve, 1976. 18 P. Arnold, Ph.D Thesis, Amsterdam, 1985. 19 R. Thomas, E.M. van Oers, V.H.J. de Beer, J.A. Moulijn, J. Catal.,84,275,l983 20 P. Ratnasamy, A. Ramaswamy, K. Banerjee, D. Sharma, N. Ray, J. Catal. 38, 19, 1975. 21 R. Prada Silvy, J. Ladriere, P. Grange, B. Delmon, in preparation. 22 H.C. Yao, W.G. Rothschild, IVth Internation, Climax, Conference on the Chemistry and Uses of Mo, Michigan, 31, 1982.
617
23 24 25 26
J.8. Pery, J. Phys. Chern., 86, 1615, 1982. N.Y. Tops~e, H. Tops~e, J. Cata1., 84, 386, 1983. K. Segawa, W.K. Hall, J. Catal., 77, 221, 1982. Y. Okamoto, Y. Katoh, Y. Mori, T. Imanaka, S. Teranishi, J. Cata1., 70, 445, 1981.
618
DISCUSSION D.M. HERCULES: I am surprised that you see Mo+6 on the surface of your catalYsts that have been treated with H? at 400°C. Are you certain that it is not Mo +? We have always observed reductions of Mo to +4 and +5 at temperatures of 400°C. Also, are you certain that you have A12(Mo04)3 present? We have observed it by Raman spectroscopy but not at loadings below 16% Mo03' R. PRADA SILVY : The XPS results corresponding to catalysts reduced at temperatures between 400°C and BOO°C showed the presence of a signal at about 236 eV which corresponds accurately to Mo+6 species in oxidized state (the presence of Mo+5 species in these samples is not exc~uded). When these samples were sulfided in H2S-H2 (15%V H2S) at 400°C, the Mo+ signal (M03d level) was not observed (perhaps by an overlap effect between the signals Mo+4, S2s and those of Mo+6). However, in the M03 spectra a small peak at about 399 eV was observed. This peak was more intensg for a sample reduced at BOO°C. We have proposed that the presence of this signal could be attributed to the formation of A12(Mo04)3- like species by two different ways: on the one hand, by a solid-state reaction between molybdenum species with the alumina and on the other hand, by a decrease of the coordination state of the Mo monolayer species (from octahedral to tetrahedral) by oxygen removal during reduction. It is well known that A12(Mo04)3-like species are difficult to reduce or sulfide even at high temperatures and they can be easily formed during the calcination step at temperatures above 600°C. We don't know the calcination temperature of our catalyst because it is a commercial HDS catalyst. The molybdenum content in this catalyst is 14% Mo03' It could be thought, in accordance with your observations, that these species could be present in the oxide precursor. We are studying, at the present time, these samples by Raman spectroscopy to confirm our results.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
619
PALLADIUM CATALYSTS FOR SELECTIVE GAS-PHASE HYDROGENATION OF PHENOL TO CYCLOHEXANONE J.R. GONZALEZ-VELAsC0 1, J.I. GUTIERREZ-ORTIZ 1, M.A. GUTIERREZ-ORTIZ 1, 2 2 M.A. MARTIN, S. MENDIOROZ , J.A.2 PAJAREs and M.A. FOLGADO 2 1 Dept. Qu;mica Tecnt cs , Univ. del Pais Vasco, Apdo. 644, 48080 Bilbao (Spain) 2 Instituto de Cata1isis y Petro1eoqu;mica, C.s.I.C., Serrano 119, 28006 Madrid (Spain)
!\BsTRACT Various Pd catalysts supported on different materi a1 s were prepared by Gifferent methods for use in the gas-phase hydrogenation of phenol to cycl ohexanone , The activity and deactivation of the catalysts are explained in terms of metal crystal size and catalyst acidity. Some important conclusions with respect to the effects of different stages of the preparation on the catalyst performance have been drawn and some general considerations on catalyst preparation have been established. INTRODUCTION Cyclohexanone is a critical intermediate in adipic acid and e-capro1actam production, the latter being used in the manufacture of fibres such as nylon 6 and polyamide resins (ref.1). In the Allied Chemical and Vickers-Zimmer processes (ref.2, 3), phenol undergoes 1i quid-phase hydrogenation on supported metals of the Pt group. As an alternative, in recent years research has been focused on the selective gas-phase hydrogenation of phenol to cyc10hexanone over supported pa11 adi urn catalysts with a Pd content between 0.5 and 5% and at temperatures between 350 and 525 K. Data on these catalysts are covered mainly by the patent literature (ref.4-8). It seems that the addition of 2-60% of alkali or alkaline earth metals (usually Cal hydroxides as a promotor has a beneficial effect on the behaviour of the catalysts, although this effect is not yet well understood. Depending on the type of catalyst used and the reaction conditions, the selectivity of the hydrogenation differs and the products may be cyclohexanol, benzene, cyclohexene or cyc10hexanone (ref.9). High activity and selectivity are catalysts as cyclohexanone is difficult to separate from required from ,;~he 1arge amounts of phenol because of the formati on of phenol-cyc1ohexanone and pheno1-cyclohexanol azeotropes. In this study, the activity of a Pd-CaO-A1 203 catalyst prepared by the conventional incipient wetness method optimized among a series of similar
620
preparations of Pd on silica,
alumina and activated carbon
(ref. 10, 11) was
compared with that of a new, strictly controlled series of Pd-Si0 Pd-sepiolite
catalysts,
prepared
by
adsorption
from
2, solution.
Pd-A1 203 and From this
comparison some conclusions on the influence of the methods of preparation on results obtained can be drawn.
~he
EXPERIMENTAL Supports A series of catalysts were prepared using different materials and methods, as follows. HARSHAW AL-3945 alumina, extruded gamma alumina, essentially neutral, 99% pure with Na
(100 ppm), Fe20~ 20 impurities. Its SBET was 240 m2g-
(100 ppm) and Si0 2 (100 ppm) as the main 3g- l• and its pore volume 0.59 cm
KETJEN CK-300 .alumina, impurity content below 0.5%, particle size 4.5 /Lm, SBET 188 m2g -1, pore volume 0.52 cm 3g -1 and isoelectric point 7.6. It showed partial crystallinity on X-ray diffraction. DAVIDSON G-952 silica, impurity content about 0.5%, mainly A1 (1500 ppm) 203 2g- l and Na (Na 0, 1000 ppm), particle size 2 /Lm, SBET 307 m and pore volume 1.65 cm 3g- 12• X-ray diffraction studies showed the amorphous character of the material, as expected from its high surface area. Its isoelectric point (3.8) is slightly higher
than those reported in
the literature (about 2), which
is
probably due to the high impurity content. Sepiolite, supplied by TOLSA (Madrid). This is a fibrous silicate with the theoretical formula Si12M98{OH)4030(H20)4.4H20, and with mineralogical characteri sti cs reported el sewhere (ref. 12). The textural parameters of the 2g-l, original material are particle size 0.8-2 /Lm, SBET 210 m mainly external, 3gl pore volume 0.61 cm and isoelectric point about 2. After thermal treatment at 723 Kin ai r , SBET decreases to 130 m2g -1 as a result of the structure "folding" (ref.13) caused by the loss of constitutional water. In
order
supports
to
ensure
that
the textural
di d not change duri ng the vari ous
characteristics of the different steps invol ved in
the catalyst
preparation, the original samples were treated at 973 K for 4 h for 5i0 A1
203
2
and
and at 723 K for 2 h for sepiolite prior to their utilization.
Catalyst preparation The catalyst on Harshaw AL-3945 alumina was prepared by the conventional incipient wetness (CI) method: the support was impregnated first with a 50 g 1- 1 so lut i on and then with a 0.2 M PdC1 solution in 0.1 M HCl (l cm3 per 2 2 gram of support). The amount of palladium was calculated so as to give
CaC1
a catalyst with the required metal content. Both impregnations were performed in Rotavapor at 10 mmHg and 305 K until evaporation of the solvent to dryness
621
was attained. The precursor obtained was dried at 393 K for 16 h and calcined in air at 973 K for 4 h. The final Ca content, measured as CaO, was 8.7%. Taking into account the isoelectric points of the supports, the remainder catalysts were prepared in a basic medium (ref.14). Additionally, when working on ,1\1 with an isoelectric point of 7.6, an acidic medium and corresponding 203 anionic species were used. The method ina bas i c cedi um (CE, from cat i oni c exchange) was to add to a weighed amount of support 0.84 M ammonia solution, the mixture being stirred continuously for 5 h at 343 K. After this, the desired amount of 10- 2 M [Pd(NH 3)4] (OH)2 solution, previously obtained by addition of excess of ammonia to PdC1 2, heating at 353-363 K for 15 min and drying, was added dropwise. Conti nuous st i rri ng was maintai ned for a further 5 h at the pre-treatment temperature (343 K) .In one additional series , lower contact times, 2 h, in both stages of preparation, were used. In acidic medium (AE, from anionic exchange), [PdC1 was employed as the 4]2active species. The following method was used: 150 cm 3 of 0.84 M HCl (pH 0-1) were added to 10 g of alumina and the mixture was stirred continuously for 5 h at 343 K. The corresponding acidic solution of the original PdC1 in excess of 2 acid, with a total complex content of 10- 2 M, was added and stirred for a further 5 h. The precursors were filtered TABLE 1 and thoroughly washed with Characteristics of the catalysts: palladium dilute ammonia solution or contents and particle size. HC1, in accordance with the previous medium used. Finally, Preparation support % Pd ~ method they were dri ed at 393 K for 16 h and calcined at 673 K for A1 0.93 CI 22 203(AL-3945) 4 h in air after gradually A1 AE 0.47 35 203(CK-300) i ncreasi ng the temperature at CE 1. 26 18 1.6 K mi n-1. 23 CES 1.26 CES 1.26 37 All the catalysts were AE 5.26 33 reduced in a flow of H2 (100 CE(2/2) 0.42 37 0.52 17 cm 3 min- l) at 523 K for 2 hand CE 1.03 CE 44 at 723 K for 1 h. Previously, CE(2/2) 1.17 76 CE 4.03 20 they were dried at 523 K in a Sepiolite CE 0.74 20 flow of helium and cooled to CES 0.74 24 room temperature under the CES 0~74 28 50 CES same conditions. The temperature 0.74 CES 0.74 70 during the whole treatment was always gradually increased,
a
622
taking 1 h up to 523 K and a further 1 h to 723 K. Finally, the catalysts were allowed to cool to room temperature in a flow of hydrogen. Various series with Pd contents of 0.5-5% were prepared. From these, different cata lysts were chosen; thei r metal contents, determined by atomi c absorption spectrometry, are shown in Table 1. In order to obtain different particle sizes..::he original alumina -and sepiolitesupported catalysts were sintered. The catalysts obtained are denoted by S in Table 1. Sjntering was carried out for 2 h in a flow of H2 (l00 cm 3min- l) at 823 and 973 K for alumina and at 823, 873, 9Z3 and 973 K for sepiolite-supported catalysts. The respective particle sizes of the sintered Pd, determined by H2 chemisorption as described below, were 23 and 37 ~ on A1 203 and 24, 28, 50 and 70 ~ on sepi 01 i te. The general characteri st i cs of the catalysts used are gi ven in Table 1. equipment and methods Characterization of the original support and catalysts was accomplished using conventional methods and equipment. Particle size was determined in a Model ZM Coulter Counter from Coulter Electronics Ltd. The surface area and porosity were measured by N2 adsorption-desorption in a 21 DOD instrument from Micromeritics and mercury intrusion in a 60000 p..s, i. porosimeter from Aminco. Isoelectric points were obtained by microelectrophoresis using a Shimadzu-Kalnew 2235 optical microscope and a Northrop-KUnitz electrophoretic cell. Elemental analysis for Pd was accomplished by atomic absorption spectrometry on dried samples disgregated by a suitable acid combination depending on the nature of the support. An AA 3030 spectrometer from Perkin-Elmer under standard operating conditions was used. Metal particle size was determined by hydrogen chemisorption in a high-vacuum volumetric apparatus with an MKS transducer for pressure measurements (accuracy 10-Z mmHg). Prior to an adsorption experiment the sample was outgassed at 573 K for 16 h. A monolayer of H2 was obtained by back-extrapolation of the isotherms obtained at pressures up to 100 mmHg and room temperature. The Pd crosssectional area was 0.078 nm 2 (ref.15). During the experiments, oxidation and re-reduction of one sample was performed in order to study deviations of the tabulated particle size. The conditions for oxidation were temperature 573 K, pressure 100 mmHg and time 2 h; re-reduction was completed under the same conditions as for the original reduction; prior to both treatments the samples were outgassed under a 10-6 mrnHg vacuum for 16 h at room temperature. Subsequent ly, Hz chemi sorpt i on was carri ed out under the standard conditions as indicated above. Ammon; a adsorpti on, not used as a standard characteri zation method, was applied to the determination of the acidity of some of the catalysts. A Perkin-
623
Elmer TGS-II thermogravimetric apparatus and a 3600 OS Data Station were used. A 100 cm 3min- l stream of ammonia previously dried in an on-line P205 cartridge was passed with helium as carrier gas (1:1) at room temperature on a previously He dried sample for 1 h and the change in weight was recorded. The ammonia flow was cut off and physisorbed ammonia was removed from the surface by the remaining He stream. The difference between the values gave the amount of retained ammonia and, in a semi-quantitative way, the. number of acidic centres at the temperature studied. Catalytic activity was determined in a fixed-bed 13 mm 1.0. reactor of the isothermal and plug flow type (ref. 10). The kinetic experiments were carried out at atmospheric pressure, 523 K, linear velocity of gases 6 cm s-l and space time 5 g cat. h mol-l. Because of the difficulty of operating with pure phenol, the feed was phenol-toluene (60:40, w/w) (inert, not giving detectable products during hydrogenation). The hydrogen to phenol molar ratio was 3:1. Under these conditions, inter- and intrapartic1e diffusional gradients in the catalyst were negligible (ref.16, 17). Hence, temperature differences between the bulk fluid and catalyst surface and gradients inside catalyst particles could be disregarded. Hydrogen was introduced into the pre-heater of the reactor through a flow controller. Phenol and toluene were also introduced into the pre-heater at a constant rate, using a high-precision device. The mixture leaving the reactor was cooled and passed through a gas-l i qutd separator. The analysi s of the products was performed on a Perkin-Elmer Sigma 38 gas chromatograph with N2 as the carrier gas and flame ionization detection. The column was 20% 2-diethy1hexy1 sebacate on Chromosorb WAW DHCS (80-100 mesh), 1 m long. RESULTS AND DISCUSSION Table 2 gives initial and 2 h conversion values in terms of turnover frequency (TOF, 0 and 2 h, respectively). The tabulated data were obtained from conversion measurements such as that in Fig. 1, corresponding to the Pd-CaOA1 203 catalyst, as an example. As can be seen, the initial TOF values are in a relatively narrow range for all the catalysts, increasing roughly with increasing metal particle size. Deactivation follows a more complicated pattern. If a reaction mechanism is accepted in which the irreversible change of the aromatic ring to a cycloalkene (reL18) is the rate-controll ing step, activity may be understood in terms of particle dispersion on the support; thus, apparently, the method of preparation, whatever it may be, has no influence on activity except through the metal crystal size attained; the supposed beneficial i nfl uence of Ca in Pd-CaO-A1203 cata 1tyst is then exerted only through a dispersion mechanism involving the metal phase.
624
However, from a strict point of view some di fference in activity can be detected TOF, 0 h TOF, 2 h for different supports and Catalyst min -1 mi n- l methods of preparat i on and these differences are more 22 29 Cl* 0.93/alumina-CaO tmpor-t.ant Tn catalyst deac4 63 AE* 0.47/alumina 22 15 1.26/alumina tivation. 24 1.26/alumina (823)** 11 When the different crystal 7 78 1.26/alumina (973)** 10 8 AE* 5.26/alumina sizes are obtained through 133 12 0.42/silica si nteri ng, as with the Pd 35 0.52/silica 59 supported on a1umi na (CE) 52 1.03/silica 8 104 6 lo17/silica and sepiolite, the activity 8 8 4.03/sil ica increases with i ncreasi ng 0.74/sepiolite 23 5 crystal size (Fig. 2). 0.74/sepiolite (823)** 34 6 0.74/sepiolite (873)** 31 9 Deactivation does not follow 0.74/sepiolite (923)** 44 5 the same pattern and to 0.74/sepiolite (973)** 53 o account for it some other * Method of preparation other than cationic assumptions should be made, exchange (Cl, conventional impregnation; related to the mechanism AE, anionic exchange). involved in the reaction, **Sintering temperatures (K). in conjuncti on with the peculiarities provided by the method of preparation. As indicated before, sintering is produced at temperatures above 773 K for 2 h ina 100 cm 3mi n-1 flow of hydrogen. Under such conditions some effects on TABLE 2 Activity of catalysts.
0.1 +
+
+
1
2
3 Time, t Ih
Fig. 1 Activity and selectivity of the Pd-CaO-A1203' ., Total; 0, cyclohexanone; +, cyc1ohexano1; 0, benzene.
625
"'i 60 ...
-
c: 'j!
LL
o~ 40 ...
..
OL...-_...I-_....... I I _--~_"""'--_ 10 40 30 20
I
SO
......._ I _L I -I 70 60
Fig. 2 Initial TOFs for various Pd - support catalysts ( 0, Pd-alumina; x, Pdsepi01 ite).
d/A the catalyst other than sintering of metal particles can be produced. Hence supported metals of the Pt group often contain residual H2 (ref.19l, which can cause an apparent loss of hydrogen chemisorption capacity and alter the catal t tyc activity and selectivity of catalysts (ref.20). These effects have been. detected, as an example, in n-pentane and n-hexane hydrogenolisis at 573-673 K (ref. 21,22l . Hydrogen chemi sorpti on at such high temperatures, probably affecting sub-surface layers of Pd (Ptl, may turn inactive a fraction of surface Pd (Ptl atoms by forming a non-stoichiometric hydride (ref.24l that cannot react with H2 or O2 during the room temperature chemisorption. The effect on chemisorption experiments would be to lower the amount of adsorbed gas, thus leading to underestimation of the surface area of the supported metal and consequently "enlargement" of the Pd crystal size. The tabulated diameters for Pd-alumina (eE) and Pd-sepiolite must be higher by a factor that depends on the sintering temperatures and degasification conditions (ref.23l. The second effect that can be produced on sintering at high temperature and hydrogen pressure is H2 spi 11 over from the metal surface to the support, as dissociated atoms or ions affecting the activity and selectivity of the catalysts and the reaction mechanism (ref. 25). This possibility, would imply the existence of an easy electron transfer mechanism between the metal and the support, which has been postul ated as the mai n feature Thus the spilled-over of strong metal-support interaction, S MS I • hydrogen may be incorporated on the support, enhancing the acidic character of the surface residual OH groups (ref.26-29l, giving sense to a reaction mechanism than could be produced through di ssoci ati ve adsorption of hydrogen, adsorpti on of phenol, reaction between adsorbed phenol and adsorbed or spilled-over
626
hydrogen and fi na1 desorption of produced cycl ohexanone. With a deacti vati on mechanism
based on coke
formation,
as
is
the case,
acidity must play an
important additional role in the whole process. We have tried to verify indirectly both, Hz-metal retention and HZ-support incorporation effects,
being aware of the difficulties involved in measuring
them by conventional techniques. Thus, a reoxidation experiment sepiolite (50 ~)
on a 0.74% Pd-
sample, accomplished at an intermediate temperature of 673 K
followed by re-reduction (ref.Zl) under standard conditions and chemisorption at room temperatures, as indicated above, produced a reduction in the Pd crystal size from 51 to 40 ~. In the same way, using ammonia adsorption, the acidic character of various sinterea sarples was also tested. Working on 0.74% Pd-sepiolite samples treated at roan temperature and 973 K as an example, a higher acidity (Z5 mequiv. per 100 g) on the sintered sample has been demonstrated, which confirms the assumptions made. Acidity may also be the cause of the activity decay of Pd-SiO (A E) . catalysts,
and Pd-A1 Z Z03 but as no possibi 1ity of spi llover is apparent, it must have
been generated during some of the steps in the preparation methods. Working with silica (CE), catalysts were prepared using different contact and pre-treatment times. Leaving aside the fact that higher contact times result in more dispersed catalysts because a redistribution of the adsorbed complex ions on the polarized sites with a competitive ion in solution is produced (ref.30), such as is the case here, higher pre-treatment times give rise to a more important neutralization of the surface acidity of the support which may be reflected in catalyst activity losses. With alumina (AE), the catalysts have been prepared after a prolonged
acid
treatment of the support (5 h with 0.84 MHC1), and 5 h contact time with the aci di c active phase solution, thus some important changes must have been produced on the support especially with regard to total acidity through hydrocloric acid adsorption and partial dissolution of alumina until equilibrium with excess acid is reached. Working with
CI 0.93% Pd-alumina,
which was prepared after previous Ca
incorporation in order to prevent acidity, the activity decay is much lower, although
a small
effect
is
also
present.
Through
ammonia adsorption
this
deactivation has been also explained in terms of residual acidity induced on the support by the 0.1 M HCl solution used in the active phase incorporation. When ammonia adsorption was carried out on the support alone, no acidity was detected on CaO incorporated alumina, but after Pd incorporation, both catalysts showed acidity, more for the untreated than for the Ca-treated support. CONCLUSIONS From these results, some conclusions can be drawn concerning the preparation of supported Pd catalysts for use in the gas phase hydrogenation of phenol to
627
cyclohexanone, which can be applied in a more general sense to catalyst preparation. The initial catalyst activity increases roughly in proportion to the metal crysta 1 si ze up to 70 thi s means that the method of preparation, whatever it may be, has no i nfl uence on acti vity except through the metal crystal si ze attained. However, on deactivation, the nature of the supports in relation to the different preparation stages does play an important role, as the modifications introduced may have a deleterious effect on the total conversion. With a deacti vat ion mechani sm based on coke formati on, as is the case here, aci dity must play an important role in the whole process. Hence one must be aware of the steps in the preparation scheme that can enhance it, and to try to prevent them if good catalyst performance is to be expected. The acidity has been enhanced in different ways on different supports. Thus, on alumina, this is achieved through acidic activation prior to active phase incorporation via adsorption from solution, and also through the acidic character of the PdC1 impregnant solution, which almost nullified the beneficial influence of the previously incorporated CaO. Acidity on silica was partially preserved through short contact times with the basic solutions of the active phase ammonia complexes. In sintered Pd samples on sepiolite and alumina, acidity was enhanced through H spilled over to the support under prolonged metal reduction 2 or sintering treatments at high temperature and H2 pressure. This spillover may induce not only SMSI between the finelly divided metal particles and the support, but also acidity, as has been shown, via residual OH bonds always present in the support. In a general sense, the different preparation steps may induce different modifications to different materials, and a decision regarding the optimal preparation method may require an exhaustive knowledge of the variables involved and thei r i nfl uence on the textural and structural parameters of the supports to be used.
t
Z
REFERENCES 1 H.J. Naumann, H. Schaefer, H. Oberender, D. Timm, H. Meye and G. Pohl; Chem. Tech., 29, 38 (1977). 2 M. Taverna; Hydrocarbon Processing, 49, 137 (1970). 3 K. Kahr; Ullmans Encyklopadie der Technischen Chemie, 96 (1975). 4 Pat. Brit. 1,257,609 (1971). 5 Pat. Brit. 1,332,211 (1973). 6 Pat. USA 4,092,360 (1978). 7 Pat. USA 4,203,923 (1980). 8 Pat. Can. 1,127,139 (1982). 9 H.A. Smith; "Catalysis", 5, edited by P.H. Emmett (Reinhold Publishing Corp .) p. 175, New York, 1957.
628
10 J.I. Gutierrez-Ortiz; Ph.D. Thesis, Universidad del Pais Vasco, 1984. 11 J.R. Gonzalez-Velasco, J.I. Gutierrez-Ortiz and A. Romero; Afinidad, 39, 543 (1982). 12 M.A. Fo1gado; Ph.D. Thesis, Universidad Complutense, Madrid 1983 13 K. Brauner and A. Preisinger; Mineral Petrogr. Mitt., 6, 120 (1956). 14 J.R. Brunelle; Pure Appl. Chern., 50, 1211 (1978). 15 J.R. Anderson; "Structure of Metallic Catalysts". Academic Press., New York 1975 16 J. I. Gutierrez-Ortiz, J.R. Gonzalez-Vel asco , A. Romero and J.A. GonzalezMarcos; Afinidad, 43, 80 (1986). 17 J. I. Gutierrez-Ortiz, J.R. Gonzalez-Vel asco, A. Romero and J.A. GonzalezMarcos; Afinidad, 43, 173 (1986). 18 T. Mathe, J. Petro, A. Tungler, Z. CsUros and K. Lugosi; Ac. Chim. Acad. Sci. Hun., 103 (3), 241 (1980). 19 P.G. Menon and G.F. Froment; "Metal-Support and Metal Additive Effects in Catalysis", B. Imelik et al. Eds., Elsevier, Amsterdam 1982. 20 Z. Paa1 and P.G. Menon; Cat. Rev., 25, 2 (1982). 21 P.G. Menon and G.F. Froment; Ac. Chim. Acad. Sci. Hun., 111 (4), 631 (1982). 22 P.G. Menon and G.F. Froment; J. Catal., 59, 138 (1979). 23 J. Sanz, J.M. Rojo, P. Malet, G. Munuera, M.T. Blasco, J.C. Conesa and J. Soria; J. Phys. Chern., 89, 5427 (1985). 24 P.A. Sermon and G.C. Bond; Cat. Rev., 8, 2 (1973). 25 D. Bianchi, M. Lacroix, G. Pajonk and S.J. Teichner; J. Catal., 59, 467 (1979) • 26 J.M. Parera, N.S. Figoli, E.L. Jablonski, M.R. Sad and J.N. Be1tramini in Catalysts Deactivation, B. Delmon and G.F. Froment Eds. Elsevier, Amsterdam 1980. 27 C. Blanco, J. Herrero, S. Mendioroz and J.A. Pajares; Proc. XXI Reunion Bienal de la R.S.E.Q.,Santiago de Composte1a, 1986. 28 P. Gajardo, T.M. Apple and C. Dybowski; Chern. Phys. Let., 74, 306 (1980). 29 Wm.C. Conner, J.F. Cevallos, N. Shah and V. Haensel; "Spillover of Adsorbed Species", J.M. Pajorik, S.J. Teichner and J.E. Germain Eds. Elsevier S.P. Amsterdam 1983. 30 T.A. Dorling, B.W. Lynch and R.L. Moss; J. Cata1., 7, 378 (1967).
629
DISCUSSION You have used hydrogen chemisorption to determine Pd dispersion in your catalysts. You know that hydrogen gives some problems in such measurements, due to absorption phenomena. How did you avoid such problems in your experiments and which evidence do you have that you detect only chemisorbed hydrogen? N.PERNICON~
M.A. MARTIN: Following current literature reports (Aben, Sermon) a temperature of 343 K and low pressures «20 mmHg) have been always used. Previously, in order to verify the suitability of the chosen conditions for the studied systems, SOr:1e additional exper tment at ron has been carried out. Working with Pd/si1ica catalysts, a series of isotherms have been completed at pressures up to 250~ and 343 K. Indeed the ,8 -phase has always appeared at pressures above 100 r.lr.lHg and higher, depend i nq on the Pd crystal size (the higher the size, the lower the pressure). Therefore, a working pressure lower than 20 mmHg was chosen. Wi th respect to temper-ature, the isobar at pressure ""0 correspondi n9 to vari ous hydrogen isotherms ranging between 343 and 423 K has been obtained; differences in H? uptake less than 5% have been found. This confirms the suitability of the recommended temperature (343 K) for the chemisorption experiments. The same cond it i ons apply to Pd/alumina catalysts. With sepiolite-supported catalysts the pressure to be used is more critical since formation of Pd-hydride (or at least a slope increase in the corresponding isotherms) was detected qat pressures as low as 50 mHg, especially wi'ch the higher crystal sizes (""801-1). TAUSZICK: It is very likely that the catalyst is deactivated by heavy organi c compounds, deri ved from phenol and not by carbon. Have you tri ed any characterization of these organic deposits? G.~.
M.A. MARTIN : A cause for deact i vat i on of catalysts duri ng operation is the formation of carbonaceous deposits which are strongly adsorbed on the surface, somehow blocking the active sites. Coke, with a composition C1H n (n between 0.5 especially and 1) or tq lnates by degradation of reactants and/or proC!ucts, aromatics, suffering succesive condensations. For the Pd/CaO/A1 203 catalyst we established the relationship between the coke content in the catalyst (determined by relative weight difference between the coked catalyst and itself after burning) and the operating conditions, namely space time, temperature and J.I., PhD Thesis, Universidad del Pais Vasco, operation time (Gutienez Or"~iz, 1984). From the results obtained it can be deduced that deactivation occurs in series with the main reaction, coke being derived mainly from cyclohexane. G.J. ANTOS: Did you note a difference in catalyst selectivity dependent on support u~ilised? What catalyst function do you think can affect the selectivity? Targets for selectivity are usually 95%. What can be adjusted to improve selectivity? M.A. MARTIN: Palladium is a very selective metal for hydrogenation of phenol to cycl ohexanone. Thi sis the reason why there has been a tendency in recent years to replace the two stage method (phenol - cyc1ohexano1 - cyc1ohexanone) over nickel by this economical single-stage method using palladium catalysts. Indeed, all prepared catalysts showed a high selectivity, with conversions to cyclohexanol less than 5%, and practically no other by-products being formed. Although some difference in selectivity was noted for the different catalysts, probably due "~O the support acidity and SMSI, as has been discussed for deactivation, they are too small to allow definite conclusions.
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B. Delmon. P. Grange. P.A. Jacobs and G. Ponce let (Editors). Preparation of Cstelvsts IV © 1987 Elsevier Science Publishers 8. V.• Amsterdam - Printed in The Netherlands
631
ALUMINUM-OXIDE-PILLARED MONTMORILLONITE: EFFECT OF HYDROTHERMAL TREATMENT OF PILLARING SOLUTION ON THE PRODUCT STRUCTURE J.P. STERTE and J-E. OTTERSTEDT Department of Engineering Chemistry, School of Chemical Engineering, Chalmers University of Technology, S-412 96 Goteborg (Sweden) SUMMARY A1203-montmorillonite complexes were prepared by interacting hydrothermally treated aluminum chlorohydrate (ACH) solutions with Na-montmorillonite in aqueous dispersion. The resulting materials showed surface areas in the range 150-400 m2jg and pore volumes in the range 0.2-0.45 cm 3jg. The surface areas decreased with increasing temperature of hydrothermal treatment of the ACHsolutions while the pore volumes increased with an increase of this temperature. Pore size distributions calculated from the desorption branches of N2-adsorption-desorption isotherms as well as X-ray powder diffraction analysis showed an increase in pore sizes with increasing temperature of hydrothermal treatment of the ACH-solutions. The uptake of aluminum oxide by the montmorillonite increased with increasing temperature of hydrothermal treatment. An optimal amount of Al added in the preparation giving a maximal surface area was found for each solution. This optimum was shifted towards larger amounts as the temperature of hydrothermal treatment of the ACH-solution increased. INTRODUCTION The process of smectite pillaring has recently attracted considerable attention. Due to their high surface area (200-500 m2jg), their microporous structure, and their acidic properties, pillared smectites have a potential interest as catalysts and as adsorbents. Smectite clay minerals have layer lattice structures tn which two-dimensional oxyanions are separated by layers of exchangeable, hydrated cations. Pillared smectites are generally prepared by cation exchange of a smectite with poly- or oligomeric cations, large enough to prop open the smectite layers, leaving part of the interlayer open and available for adsorption and for catalytic reactions. A number of such oligomeric cations are formed upon base-hydrolysis of metal salts. Most work on preparation of pillared smectites has involved pillaring with oligomeric aluminum-oxide-hydroxide cations (refs. 1-3). Solutions containing such species can be prepared either by base hydrolysis of aluminum chloride in aqueous solution or by reacting aluminum metal powder with an aqueous solution of aluminum chloride. The later type of solution is usually referred to as aluminum chlorohydrate (ACH). Pinnavaia (ref. 4) compared products prepared from these different types of solutions but found little
632
difference in aluminum uptake, basal spacing or catalytic properties of the products. 27Al NMR studies of the solutions indicated, however, that the pillaring species in the ACH-solutions were further aggregated than those in the base hydrolyzed aluminum chloride solutions. In line with Vaughan (ref. 5), Pinnavaia suggested that the basic structure of the pillaring cation in both cases was that of (A1 1304(OH)24(H 20)12)7+ proposed by Johansson (ref. 6). A high resolution solid state NMR study (27Al and 285i) of pillared smectites by Plee et al. (ref. 7) supported this structure. This study also showed that, upon calcining, there was ~ reaction between the pillars and the clay surface in tetrahedrally substituted smectites. This reaction did not occur in pillared smectites without tetrahedral substitution. Van Damme and Fripiat (ref. 8) showed, by a fractal analysis of adsorption processes by pillared clays, that the distribution of pillars in aluminum-oxide-hydroxide pillared smectites was close to homogeneity. Furthermore they found that the molecular sieving properties of this pillared clay were close to those of faujasite. Using the technique of cation exchange of the smectite with hydrolyzed oligomeric species, pillared smectites have been prepared with hydroxy-Zr (ref. 9), hydroxy-Cr (refs. 10-11) and hydroxy-Ti (ref. 12) pillars. The properties of hydroxy-Zr pillared smectites are similar to those of aluminumoxide-hydroxide pillared ones, i.e. surface areas in the range 200-400 m2/g, basal spacings of about t9 A and pore volumes of about 0.2 cm 3/g. Recently, Burch and Warburton (ref. 13) found that hydroxy-Zr pillared smectites can be prepared more thermally stable compared with aluminum-oxide-hydroxide pillared ones. Hydroxy-Cr pillared smectites were first reported by Brindley and Yamanaka (ref. 11). These materials exhibit basal spacings of about 18 A when freeze dried and about 14 A when dehydrated at temperatures above 250 0C. Pinnavaia et al. (ref. 14) prepared a hydroxy-Cr montmorillonite with a basal spacing of about 20.5 A after heating at 500 0C. This larger basal spacing was obtained by preparing the hydroxy-Cr oligocations at 95 0C, thus obtaining larger species compared with those obtained by hydrolysis at room temperature. Hydroxy-Ti pillared smectites with basal spacings of about 28 A, corresponding to an interlayer distance of about 19 A, have been prepared in this laboratory (ref. 12). These materials exhibit surface areas in the range 200-400 m2/g and a thermal stability comparable to that of aluminum-oxide-hydroxide and Zrhydroxide pillared smectites. As the number of metals which form poly- or oligomeric cations upon hydrolysis is limited, an alternative method for the preparation of pillared smectites, involving in situ hydrolysis of complex cations between the smectite layers, has also been used. For example, the preparation of 5i0 2-pillared smectites by Endo et al. (ref. 15) demonstrated the usefulness of this synthetic
633
approach. In their materials the layer separation was limited to one monolayer of Si02, corresponding to a basal spacing of about 13 A. Using a related method, Yamanaka et al. (ref. 16) prepared Fe203-pillared montmorillonite by the hydrolysis of (Fe3(OCOCH3)70H)+ in the interlayer region of the clay. This product was characterized by a basal spacing of about 17 A and a surface area of about 300 m2/g after heating at 500 0C. Recently, Lewis et al. (ref. 17) developed techniques for preparing smectites pillared with three-dimensional Si02-oligomers. These materials were prepared by reacting the smectite with oligomeric Si02-species containing organic substituents, for example 2-(2-silylethyl)pyridine oligosilsesquioxane pillars. Thus Si02-pillared smectites having basal spacings of about 19 A and surface areas of approximately 400 m2/g after heating at 500 0C were prepared. A major objective of our work is to prepare pillared smectites suitable as catalysts for cracking of heavy oil fractions. The use of previously reported pillared clays for this purpose is limited by the lack of thermal and hydrothermal stability of these materials. At the thermal and hydrothermal conditions in the regenerator of a fluid catalytic cracker (FCC), these materials rapidly lose most of their surface area and their catalytic activity. Furthermore a larger basal spacing (i.e. a larger pore size) would be desirable for this application. Vaughan et al. (ref. 18) found that, due to further polymerization of the pi 11 ari ng cations, hydrotherma 1 treatment of the ACH-so1ut ions (refl ux conditions) improved the thermal stability of the products. Tokarz and Shabtai (ref. 19) conducted a similar study using base hydrolyzed A1C1 3 solutions. They found that refluxing of the solutions for 6-48 hrs was sufficient to produce pillared smectites with higher thermal stability and higher porosity as compared with those prepared from solutions aged at room temperature for two weeks. Hydrothermal treatment of ACH-solutions at temperatures above about 120 0C is known to yield positively charged, fibrillar boehmite crystallites in colloidal solution (refs. 20-22). The size of these crystallites increases with increasing temperature and time of hydrothermal treatment. As part of our work on preparation of pillared smectites, we have investigated the pillaring of montmori l lont te with hydrothermally treated ACH-sol uti ons. This paper reports a study of the preparation and properties of these materials. The effects of temperature of hydrothermal treatment of ACH-solutions, and amount of pillaring solution added in the preparation were investigated. Products were characterized by N2-adsorption-desorption measurements, by X-ray powder diffraction, and by elemental analysis.
634
EXPERIMENTAL Aluminium Chlorohydrate (ACH) A stock solution of aluminium chlorohydrate was prepared by addition of aluminium metal powder to an aqueous solution of aluminium chloride hexahydrate. The aluminium powder was added in small portions over a time of approximately one week and the solution was kept at 70-800C during this addition. Thus a solution with an Al concentration of 5.55 M, having an A1 203/Cl molar ratio of 0.93 and a pH of 3.9 was obtained. Prior to the hydrothermal treatments of the ACH-solution, parts of the stock solution were diluted to an Al concentration of 0.176 M(1.8% A1 203, by weight) Hydrothermal treatment The hydrothermal treatment was carried out by heating diluted ACH-solutions in a stainless steel autoclave kept in a polyethylene glycol bath. Temperatures in the range 110-160 0C were used and a synthesis time of 24 hrs was employed in all treatments. After cooling, the product solutions were removed and deionized. The deionization procedure consisted of raising the pH to about 9, thus gelling the ACH, and washing the gel until chloride free (as determined by AgN0 3). Montmorillonite A Wyoming Na+-Ca 2+-montmorillonite (commercial designation Volclay Spy 200) was obtained from American Colloid Company. Quartz, present as impurity was removed by fractionation, using conventional sedimentation techniques. The <2~ fraction, which was practically free from impurities, as determined by X-ray diffraction analysis was used as starting material. The cation exchange capacity (CEC) of this montmorillonite fraction was determined to be 89 meq/ 100 g. Homoionic Na+-montmorillonite was prepared by subjecting this fraction to repetetive (4 times) ion exchange with aqueous 1.0 MNaCl at 250C. The exchanged clay was washed with distilled water until chloride free and dried at 600C. Preparation of A1 203-pillared montmorillonite Deionized and washed gel containing 1.8 g of A1 203 was dispersed in 300 ml of distilled water and the pH was adjusted to 3.9 (stable for 1 min) by adding 1.0 Macetic acid. The amounts of acetic acid required to reach this pH are given in table 1, for the different preparations. A decrease in this amount with increasing temperature of hydrothermal treatment of the ACHsolution was observed. The acidified solution was diluted to 600 ml and stirred for several hours prior to its use.
635
TABLE 1 Amount of acetic acid required to obtain pH 3.9 in solutions of deionized and washed ACH, hydrothermally treated at different temperatures. Temperatureoof hydrothermal treatment ( C) ACH b 110 120 130 140 160
mmoles acetic acid addeda per g of A1 203 34.4 22.2 14.4 4.7 3.9 3.1
solutions contained 1.8 g of A1 203 in 300 ml of distilled water Untreated ACH-solution.
~All
Na+-montmorillonite dispersions were prepared by prolonged stirring of 1.0 g of the montmorillonite in 2.0 1 of distilled water. The pillared products were obtained by adding calculated volumes of the hydrothermally treated, deionized and reacidified ACH-solutions to the montmorillonite dispersion to obtain A1203/montmorillonite ratios in the range 0.15-1.2 gig. The products were left in contact with the solution overnight, separated by centrifugation, washed four times with distilled water and dried at 60 oC. Characterization of products N2-adsorption-desorption isotherms were determined by a Digisorb 2600 surface area-pore volume analyzer (Micromeretics Instrument Corporation). The samples were first outgassed at 250 0C for 3 hrs and the isotherms were recorded at liquid nitrogen temperature. Surface areas were calculated using the BET equation and pore volumes were estimated as the volume adsorbed at a relative pressure of 0.995. Pore volume distributions were calculated from the desorption branches of the isotherms, using parallel plates as a geometrical model (ref. 23). The Halsey equation (ref. 24) was used for calculating the nitrogen monolayer thickness at a particular relative pressure. X-ray diffraction (XRD) analysis of A1 203-montmorillonite samples was performed on powder samples. The XRD patterns were obtained with a Philips powder diffractometer, using Ni-filtered. fine focus, CuKcr-radiation. Elemental analysis of products was carried out by atomic absorption spectroscopy (AAS) employing LiB0 2-fusion (ref. 25). Si02 was separately determined colorimetrically using the beta-silicomolybdate method (ref. 26). Analysis of calcium in magnesium chloride extracts of calcium saturated samples
636
by AAS was used to determine cation exchange capacities. RESULTS AND DISCUSSION Effects of hydrothermal treatment on the aluminum containing species in the ACH-solution Upon hydrothermal treatment of ACH-solutions, at temperatures above 120 0C, fibrillar boehmite is formed (refs. 20-22). X-ray powder diffraction patterns of the deionized and freeze dried ACH used in this study are shown in Fig. 1. The peaks characteristic of boehmite begin to appear for the sample treated at 120 0C and gradually grow more intense with increasing temperature of hydrothermal treatment. The fibrillar shape of the boehmite crystallites was confirmed for the samples treated at 140 0 and 160 0C by transmission electron mi croscopy. Dependence of surface area, pore size and elemental composition upon temperature of hydrothermal treatment of ACH-solution In Fig. 2, BET surface areas and pore vOlumes of samples prepared from the
160°C
~ ~
~
130°C 120°C
50
45
40
35
llOce ACH 30
25
~
20
~ 15
10
5
Degrees 29
Fig. 1. X-ray diffraction patterns of aluminum-oxide-hydroxide products obtained by hydrothermal treatment of ACH-solutions at given temperatures, fnl l owed by deionization and freeze dryi ng.
637
0- - 0.40 01
300
E
N......
u
E eCIJ
.
...0
0.30 200
~
~---
0.20
u
....0... ~
OJ'"
E
CIJ
(J)
01 ~
0
>
...0 CIJ
100
0.10
ACH
110
120
130
140
150
a..
160
Iernpercture.t'c )
Fig. 2. Surface areas and pore volumes of aluminum-oxide-montmorillonite complexes prepared from hydrothermally treated ACH-solutions. Amounts of solution corresponding to 0.6 g of A1 203 per g of montmorillonite were used in all preparations.
0«
ai 40
u
c
Cl
OIl
-0
...
30
CIJ
e>-
...c
20
Gl
~--
> 10
t:l ~
0-
w
ACH
110
120
130
140
150
160
Temperature, (C)
Fig. 3. Equivalent distances between the montmorillonite layers in aluminumoxide-montmorillonite complexes, calculated from the surface areas and pore volumes in Fig. 2.
638
starting ACH-solution and from ACH-solutions heat treated at 110 o-160 oC for 24 hrs are shown. All the ACH-solutions were deionized and reacidified as described under Experimental, before the pillaring procedure. For easier comparison, BET surface areas are given for all samples although the samples prepared from solutions treated at lower temperatures, due to their microporous nature, show better correlations with the Langmuir equation. The surface areas are decreasing with increasing temperature of hydrothermal treatment while the pore volumes show a continuous increase with increasing temperature. Equivalent layer distances were calculated from the pore volumes and surface areas, using parallel plates as a geometrical model. The values are plotted in Fig. 3 and indicate an increase in pore size with increasing temperature of hydrothermal treatment of the ACH-solution. Complete adsorption-desorption isotherms were determined for the A1 203montmorillonite samples and the isotherms are shown in Fig. 4. The isotherms recorded for the samples prepared from ACH treated at lower temperatures «130 oC) are of type I in the classification of Brunauer, Deming and Teller.
200 100 0
~ 200 ME $
-af
100
.Q
5III
"1J
0
CII
100
o
E :::l "'5 >
0 100 ,-
ACH
0 100
0'----1._--'-_-'-_...1..----'
o
Q2
0.4
Relative
0.6
0.8
1.0
pressure, PIPo
Fig. 4. N isotherms for aluminum-oxide-montmorillonite complexes.2-adsorption-desorption Amounts of Al-solution corresponding to 0.6 g and 1.2 g of A1 203 C preparation and the per g of montmorillonite were used in the~130 >130 oC preparations respectively.
639
This type of isotherm is characteristic for microporous materials in which multilayer adsorption of nitrogen is sterically hindered. The isotherms recorded for the samples prepared from ACH treated at higher temperatures (130-160 0C) are of type II. This type of isotherm is characteristic for mesoporous materials. The hysteresis loops are, for all samples, of type B in the classification of de Boer (ref. 23). Among the materials showing this type of hysteresis are those in which the pore structure is built up of parallel plates. Desorption pore volume distributions were calculated from the isotherms of the A1 203-montmorillonites, prepared from the untreated ACH-solution and from ACH-solutions treated at 120 0 , 130 0 , 140 0 and 160 0C using 0.6 g of A1 203 per g of montmorillonite for the first three samples and 1.2 gig for the last two. The pore volume distributions are shown in Fig. 5. For the samples prepared from solutions treated at lower temperatures only a fraction of the total pore volume of the materials is given in these plots, as part of the pore volume is found in pores smaller than those measurable by this method. It is, however, clear that a significant fraction of the pore volume of all samples investigated is found in the range covered. This fraction is increasing with increasing temperature of hydrothermal treatment of the ACH-solution at
o
140 C
o
40
800
40
800
40
800
40
80
0
40
80
't: (A) Fig. 5. Desorption pore volume distributions of aluminum-oxide-montmorillonite complexes, calculated from the isotherms shown in Fig. 4.
640
the expense of the fraction found in micropores. In the sample prepared from the solution treated at 130 0C, most of the total pore volume is found in pores with an interlayer distance of 16-24 A. The corresponding ranges for the samples prepared from solutions treated at 140°C and 160°C are 16-35 A and 4080 A respectively, indicating a broadening of the pore size interval with increasing temperature of hydrothermal treatment of the ACH solution. The sizes of the boehmite crystallites in the ACH-solutions treated at 140°C and 160°C were estimated from the X-ray diffraction patterns shown in Fig. 1. Using the full breadth at half maximum of the d(020)-peaks, the crystallite sizes were calculated to 51 A and 75 A for the 140°C-sample and for the 160 0Csample respectively. The corresponding values calculated from the d(120)-peaks are 33 A and 108 A, respectively. Taken into account that the model chosen to calculate the pore volume distribution probably underestimates the pore sizes to some extent, the correlation between the measured pore sizes and the ones expected from the boehmite crystallite sizes obtained is quite good. X-ray diffraction patterns for the samples prepared from the starting ACHsolution and from the ACH-solutions treated at 120°, 130 0 and 140 0C are shown in Fig. 6. The sample prepared from the untreated solution shows three orders of basal reflections, corresponding to a basal spacing of 18.6 A (or an interlayer distance of about 9 A) and is very similar to previously published X-ray diffraction patterns for pillared smectites (refs. 3-4). In the diffractogram recorded for the sample prepared from the solution treated at 120 0C this peak is shifted to 17.5 A and a second peak appears at 28 = 2.7 0 corresponding to a spacing of about 33 A (or to an interlayer spacing of about 23 A). This latter peak gets more intensive and is shifted to a somewhat lower angle for the sample prepared from the ACH-solution treated at 130 0C, while the first peak is shifted to 28 = 5.9° corresponding to a basal spacing of 15.5 A. For the sample prepared from the solution treated at 140°C, the high-angle peak is shifted to a somewhat higher angle while the low-angle peak disappears (or is shifted to 28 < 2°). The d(020)-peak of boehmite (28 = 14.5°) is seen in the diffractograms of the samples prepared from ACH-solutions treated at temperatures above 130 0C. Both the pore volume distributions and the X-ray diffractograms indicate steplike changes in interlayer distance with a step width of about 10 A, with changes in temperature of hydrothermal treatment of the ACH-solution. The pillaring species in ACH-solutions are generally believed to be the oligomeric (A1 1304(OH)24(H zO)1Z)7+ cation. The size of this cation is compatible with an interlayer distance of 8-10 A in the pillared smectite. The appearance of a 20 A spacing in the materials prepared from hydrothermally treated ACHsolutions could be explained either by a partial double layer of these cations, by growth of the original cations to larger complexes or by a combination of
641
130·C 120·C
ACH 16
14
12
10 8 6 Degrees 29
4
2
Fig. 6. X-ray powder diffraction patterns of aluminum-oxide-montmorillonite complexes prepared trom the untreated ACH-solution and from solutions treated at 120, 130 and 140 C. Amounts of Al-solution correspond3ng to 0.6 g and 1.2 g of A1 203 perog of montmorillonite were used in the~ 130 C preparation and 140 C preparation respectively. the of these factors. The hydrothermal treatment of the ACH-solutions undoubtedly leads to a further polymerization of the Al-species but also to a decrease of the positive charge per Al in these species. Due to this lower charge, more Al is needed to saturate the ion exchange capacity of the montmorillonite which, in turn, makes the formation of a double layer or of a partial double layer of Al-species between the clay layers more likely. The wider pore volume distributions of the samples prepared from ACH-solutions treated at 140 0 C and 1600 C indicate a wide variation in thickness of the boehmite crystallites, believed to be responsible for the pillaring in these products. It should also be noted that the lengths of the boehmite fibrils in these ACH-solutions (as estimated from electron micrographs) were of the same order of magnitude as the side length of the montmorillonite sheets. In Table 2, the elemental analysis of A1 203-montmorillonite complexes prepared from the untreated ACH-solution and from solutions treated at 1200 and
TABLE 2. Elemental analysis of the starting montmorillonite and of A1 203-montmorillonite products samples prepared from ACH-solutions Na+-montmo- a rillonite untreated b treated at 120 0C c treated at 1400C c amount A1203 f added (gig)
-
0.15 0.3
0.6
1.2
0.3
0.6
1.2
0.3
0.6
1.2
49.9 30.6 3.56 2.26 0.10 0.32 0.15 13.7
48.7 30.4 3.53 2.06 0.20 0.36 0.24 15.1
44.4 35.9 3.36 1.95 0.08 0.32 0.19 13.9
39.7 40.1 3.05 1.85 0.08 0.23 0.21 15.1
43.3 38.0 3.27 2.01 0.06 0.27 0.20 14.9
40.5 39.0 3.35 1.77 0.13 0.25 0.18 13.2
32.1 44.7 2.66 1.50 0.12 0.17 0.14 20.8
25.3 54.5 2.17 1.13 0.24 0.14 0.16 17.3
wt%o Si0 2 A1 203 Fe203 MgO CaO K20 Na 20 H oe 2
60.4 20.5 4.29 2.94 0.43 0.40 3.20 6.8
50.8 26.4 3.65 2.03 0.19 0.46 0.21 14.9
50.7 27.8 3.69 2.04 0.13 0.46 0.22 13.2
total:
99.0
98.6
98.2 100.6 100.6 100.1 100.3 102.0
-
98.4 102.2 100.9
~Na-exchanged montmorillonite used as starting material in all preparations. cDeionized and reacidified ACH-solution. dSolutions treated at gives temperatures for 24 hrs. eSamples were dried at 1100C overnigHt prior to analysis. flgnition loss between 110 and 1000 C. Amount of solution added in the preparation step, calculated as 9 of A1 203/g of montmorillonite.
s
643
140 0C are shown together with the corresponding analysis of the starting montmorillonite. The uptake of aluminum oxide by the montmorillonite increased with increasing temperature of hydrothermal treatment of the ACH-solution. This was expected as the charge per aluminum in the pillaring species decreases with increasing degree of polymerization which is also reflected by the amounts of acetic acid required to obtain pH 3.9 in the solutions prior to the pillaring procedure (see Table 1). Dependence of surface area, pore size and elemental composition upon the amount of Al added in the pillaring procedure Fig. 7 shows the surface areas of the samples prepared from the untreated ACH-solution and from solutions treated at 120 0 and 140 0C, as a function of the relative amount of Al (calculated as A1 Z03) added in the preparation step. Theoo-value represents the surface area of the deionized and freeze dried ACHsolution and the dashed lines represent the surface areas expected from physical mixtures of the montmorillonite and the given amounts of deionized and freeze dried ACH-solution. The amount of Al required to obtain a maximal surface area increases with increasing temperature of hydrothermal treatment of the ACH-solution. While only about 0.3 g of A1 203 per 9 of montmorillonite needed to reach this maximum for the sample prepared from the starting ACH-
A ACH
300
D
120·C
0
140·C
01 N-
E e
.... 200
e
-... U
e
::l
en 100
0.0
0.3
G6
1.2
00
Amount AlP3 added Ig/g)
Fig. 7. Surface areas of aluminum-oxide-montmorillonite complexes as a function of tbe amount of Al (calculated as A1 203) added in the preparation step.
644
solution, about twice this amount is required for the sample prepared from the solution treated at 120 0C. An interesting observation is the decrease in surface area of samples prepared from ACH-solutions treated at lower temperatures « 130 0C) when excess Al is added in the preparation step. Some workers have used a large excess in their preparations in order to ensure complete saturation of the montmorillonite with the pillaring cations (ref. 4). The surface areas of their samples may therefore be lower than the maximal ones. It should, however, be noted that the basal spacing of the samples prepared from the starting ACH-solution is continuously increasing with increasing amount of A1 203 added in the preparation. The basal spacing of the samples prepared from this solution and shown in Fig. 7, increases from 18.2 A for the sample prepared using 0.15 g of A1 203 per g of montmorillonite to 19.1 A for the 1.2 g/g-sampl~. The elemental analysis of A1 203-montmorillonites prepared from the untreated CH-solution and from solutions treated at 120 0C and 140 0C, using amounts of solution corresponding to 0.15-1.2 g of A1 203/ g of montmorillonite is shown in Table 2. The amount of A1 203 required to saturate the montmorillonite increased with increasing temperature of hydrothermal treatment of the ACH. This is consistent with the lower charge per Al in the solutions subjected to hydrothermal treatment compared with that in the untreated ACH-solution. Structure of aluminum-oxide-montmorillonite Pinnavaia et al. (ref. 4) proposed a structure of pillared smectites in which the pillaring species were all of the A1 13-type and evenly distributed over the smectite layer surface. According to this model, pillared smectites would have a truly zeolitic structure with only one major pore size. The results presented in this study partly contradicts this model as a significant part of the pores, even in the products prepared from the starting ACH-solution, seem to be larger than what would be expected assuming a partial monolayer of Al 13-species between the montmorillonite layers. Pore volume distributions similar to those shown in this paper for samples prepared from the starting ACH-solution have been measured in this laboratory for samples prepared from and from ZrOC1 2-solutions. The widening of the base hydrolyzed A1C1 3-solutions pore size distribution with increasing temperature of hydrothermal treatment of the ACH-solutions is at least partly due to a wider distribution in the degree of polymerization of the Al-species in these solutions. Furthermore, a lower ordering of these species between the montmorillonite sheets can be expected as the length of the boehmite fibrils approaches the side length of the montmorillonite sheets. It is by no means certain that all of the boehmite crystals in the products prepared from ACH-solutions treated at higher
645
temperatures (>130 0C) are situated between the clay layers. Part of the high surface area of these materials could be explained by a physical mixture of the starting materials. Conclusions Hydrothermal treatment of ACH-solutions before preparation of pillared smectites provides a new route for preparation of A1 203-smectite complexes with larger pores than those in conventional pillared smectites. The relatively high ordering of pore sizes in the conventional pillared smectites is, however, lost when the ACH-solutions are subjected to hydrothermal treatment prior to the pillaring procedure. The materials prepared from hydrothermally treated ACHsolutions may be interesting as catalysts in cracking of heavy oil fractions and of biomass oils, provided that the thermal stability is sufficient. The thermal stability, acidic properties and cracking activity of these materials are now under investigation. ACKNOWLEDGEMENTS The authors wish to thank the Swedish Board for Technical Development (STU) for financial support of this project. Helpful advice from P.G. Menon in interpreting the results and writing this paper is greatly appreciated. REFERENCES 1 G.W. Brindley and R.E Semples, Preparation and properties of some hydroxyaluminum beidellites, Clay Miner., 12 (1977) 229-236. 2 D.E.W. Vaughan, R. Lussier and J. Magee, Pillared interlayered clay materials useful as catalysts and sorbents, U.S. Patent No.4, 175, 090 (1979). 3 N. Lahav, U. Shani and J. Shabtai, Cross-linked smectites I. Synthesis and properties of hydroxy-aluminum montmorillonite, Clays &Clay Minerals, 26 (1978) 107-115. 4 T.J. Pinnavaia, M-S Tzou, S.D. Landau and H.R. Raythatha, On the pillaring and delamination of smectite clay catalysts by polyoxo cations of aluminum, J. Mol. Catal., 27 (1984) 195-212. 5 D.E.W. Vaughan, R.J. Lussier and J.S. Magee, Pillared interlayer clay products, U.S. Patent No. 4,271,043 (1981). 6 G. Johansson, On the crystal structures of some basic aluminum salts, Acta Chem. Scand., 14 (1960) 771-773. 7 D. Plee, F. Borg, L. Gatineau and J.J. Fripiat, High-resolution solidstate 27Al and 29Si nuclear magnetic resonance study of pillared clays, J. Am. Chem. Soc., 107 (1985) 2362-2369. 8 H. Van Damme and J.J. Fripiat, A fractal analysis of adsorption processes by pillared swelling clays, J. Chem. Phys., 82(6) (1985) 2785-2789. 9 S. Yamanaka and G.M. Brindley, High surface area solids obtained by reaction of montmorillonite with zirconyl chloride, Clays & Clay Minerals 27 (1979) 119-124. 10 J. Shabtai and N. Lahav, Cross-linked montmorillonite molecular sieves, U.S. Patent No. 4,216,188 (1980). 11 G.W. Brindley and S. Yamanaka, A study of hydroxy-chromium montmorillonites and the form of the hydroxy chromium polymers, Amer. Mineral, 64 (1979) 830-835.
646
12 J. Sterte, Synthesis and properties of titanium oxide cross-linked montmorillonite, submitted to Clays & Clay minerals. 13 R. Burch and C.J. Warburton, Zr-containing pillared interlayer clays I. Preparation and structural characterization, J. Catal. 97 (1986) 503-510. 14 T.J. Pinnavaia, M-S Tzou and S.D. Landau, New Chromia pillared clay catalysts, J. Am. Chem. Soc. 107 (1985) 4783-4785. 15 T. Endo, M.M. Mortland and T.J. Pinnavaia, Intercalation of silica in smectites, Clays & clay minerals 28 (1980) 105-110. 16 S. Yamanaka, T. Doi, S. Sako and M. Hattori, High surface area solids obtained by intercalation of iron oxide pillars in montmorillonite, Mat. Res. Bull. 19 (1984) 161-168. 17 R.M. Lewis, K.C. Ott and R.A. Van Santen, Silica-clay complexes, U.S. Patent No. 4,510,257 (1985). 18 D.E.W. Vaughan, R.J. Lussier and J.S. Magee, Jr. Stabilized pillared interlayer clays, U.S. Patent No. 4,248,739 (1981). 19 M. Tokarz and J. Shabtai, Cross-linked smectites IV. Preparation and properties of hydroxyaluminum pillared Ce- and La-montmorillonites and fluorinated NH4+montmorillonites, Clays & Clay minerals, 33 (1985) 89-98. 20 J. Bugosh, Fibrous alumina monohydrate and its production, U.S. Patent No. 2,915,475 (1959). 21 J. Bugosh, R.L. Brown, J.R. McWhorter, G.W. Sears and R.J. Sippel, A novel fine alumina powder, fibrillar boehmite, Ind. Eng. Chem. Prod. R&D 1 (1962) 157-161. 22 J. Sterte and J-E Otterstedt, A study on the preparation and properties of fibrillar boehmite, in preparation. 23 J.M. Thomas and W.J. Thomas, Introduction to the principles of heterogeneous catalysis, Academic Press, London, 1967. 24 G. Halsey, Physical adsorption on non-uniform surfaces, J. Chem. Phys. 16 (1948) 931-937. 25 J.H. Medlin, N.H. Suhr and J.B. Bodkin, Atomic absorption analysis of silicates employing LiB02 fusion, At. Absorpt. Newsl. 8 (1969) 25-29. 26 R.K. Iler, The Chemistry of silica, Wiley, New York, 1979.
647
DISCUSSION J.E. OTTERSTEDT : The negative charge on the surfaces of montmorillonite sheets may be reversed by running a clay slurry in~o a solution containing positively charged polycations, e.g. Al1304(OH)24{H20Ji2' Negatively charged pillars can then be used to crosslink the smectite. R. SCHOONHEYDT : 1/ Do your pillared montmorillonites have cation exchange capacity and how much? 2/ What is their thermal stability? J. STERTE : The products described in this presentation do have some cation exchange capacity. The CEC's of the products are in the range 0-20 meq/lOOg depending on the amount of A1203 used in the preparation and on the temperature of hydrothermal treatment of the solution prior to the preparation. It is however possible to regain some of the CEC by using the method described by Vaughan (ref. 5). The thermal and hydrothermal stabilities of these materials are under investigation. Preliminary results indicate that hydrothermal treatment of the ACH-solution prior to the preparation of pillared products results in an increase in their thermal stabilities. G.M. PAJONK : In your talk, you quoted that type II isotherms are specific of mesoporous. This point is not clear because it is well known that type II isotherms are representative of non-porous (or macroporous) material while it is type IV which belongs to mesoporous solids. It seems to me that by looking at Fig. 4, there is a trend (more of less strong) from type I to type IV isotherm (at 160°C). In other words, up to the temperature of treatment of 140°C, it seems to me that your A1203-montmorillonite belongs to the super microporous materials as defined recently by Dubinin and then, at 160°C, to the mesoporous category of solids. J. STERTE : no reply. L. GUCZI : If you had to decide between using pillared clay instead of zeolite, what would be your selection?
J. STERTE : The question can only be answered by assuming that you refer to a choice between pillared clays and zeolites as catalysts for catalytic cracking. In a commercial heavy oil cracking operation, the choice must be a zeolite cracking catalyst since there are no pillared clay cracking catalysts available. In a laboratory investigation of heavy oil cracking, the choice would be a pillared smectite since the limitations of zeolites to crack heavy ends are well known. Based on our present knowledge, we would choose a montmorillonite pillared with an ACH-solution treated at 140°C, but we may have changed our mind in about 6 months when we will have completed a comparative study of pillared smectites as cracking catalysts. M.G. HOWDEN: The thermal stability of a montmorillonite varies as indicated by its dehydroxylation DTA peak, which ranges from 500°C up to 750°C. At what temperature did your basic material dehydroxylate? What influence does this variation in dehydroxylation temperature have on the stability of the pillared product?
J. STERTE : The montmorillonite used in this study dehydroxy1ates gradually between 500 and 700°C. I do not know what effect the temperature of dehydroxy1ation of the montmorillonite has on the thermal stability of the pillared product. J. SCHEVE: 1/ Can you comment on the change of conversion and gasoline yield with time on stream? 2/ Till what weight percent could you burn off the high amount of coke you found?
648
J. STERTE : At this point, I am not able to answer these questions. This information cannot be extracted from MAT runs but requires pilot-plant cracking experiments. We have not yet made such runs in our pilot-plant cracking unit using pillared smectites as catalysts. J.W. JOHNSON: When you get the aged Al solution, then redissolve it with acid, does it form a homogeneous solution or suspension readily? Would you comme~t on the benefits of this treatment compared to the usual use of solutions of Al that have not been precipitated and redissolved? Transmission electron microscopy should prove useful in determining what fraction of the boehmite crystallites are located between the clay layers and which are separated from the layers. J. STERTE : Yes, they form a homogeneous solution readily. The hydrothermal treatment results in a polymerization of the A1-species involving a release of protons. The product solutions are very acidic and unstable at room temperature. If left for some time, Al is redissolved from the boehmite fibrils. By precipitation and redisso1ution at pH 3.9, a stable solution is obtained. I agree that it would be a good idea to investigate these materials using TEM. M. TOKARZ : 1/ I am not sure that I can agree that the evidence you have presented, namely the results of the XRD measurements, really indicate that the boehmite crystals are present in between the smectite layers. I would like to suggest you to study your samples with the use of TEM technique. This should give the direct answer if the boehmite particles are where you believe they are. 2/ The second question is connected with the fact that the boehmite needles are equally large as the montmorillonite platelets. Therefore, may be, they block the inter1ayer space preventing the organic molecules from entering in between the smectite layers. Could you comment it on? J. STERTE : 1/ See answer to preceding question. 2/ The boehmite fibrils are not equally large as the montmorillonite platelets. The length of the boehmite fibrils in the 140 and 160°C solutions approaches the side length of the montmorillonite sheets. The high surface areas of these materials indicate that they are not blocking the interlayer space of the products.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
649
SYNTHESIS AND PROPERTIES OF CROSS-LINKED HYDROXY-TITANIUM BENTONITE
SUN GUIDAl, YAN FUSHANl, ZOO OOIOOA2 and LIU ZHONGHUI lDepartment of Petro-Chemical Engineering, Fushun Petroleun Institute, Fushun, Liaoning, P. R. China 2Research Institute of Petroleum Processing, Beijing, P. R. China
ABSTRACT Cross-linked hydroxy-Ti bentonite was synthesized by exchange reaction of Na-bentonite with Ti(S04)2 solution under given conditions, which include mainly reaction terrperature, Ti4 + concentration in mixed slurry, Tilbentonite ratio and the order of dropwise mixing of bentonite slurry with Ti(S04)2 solution. The product has a basal spacing l6.2A and a surface area 204 W.lg. At lower terrperature it contains both Bronsted and Lewis type acidity, whereas at higher tenperature it has mainly Lewis type acidity. The cracking reactions of hydrocarbons over it proceed via carbonium-ion reaction mechanisms. ~thylnaphthalene on it undergoes a disproportionation reaction and the selectivity for naphthalene is higher than that of RE-Y zeolite catalyst. The catalyst cooprising 003 and NiO supported on it shows considerable pyridine hydrodenitrogenation activity.
INTRODUCTION Cross-linked smectites are a new class of IOOlecular sieve-like catalytic materials, which bring to attention of many research workers. The published papers about the studies in this field increase gradually in recent years, but many studies were concerned with cross-linked hydroxy-AI smectites with a large pore size, considerable acidity and intrinsic catalytic activity (ref. 1, 2, 3). Moreover, the catalyst carriers containing Ti02 showed significant influences on the performance of the catalysts, so studies of catalysts comprising Ti02 become one of the subjects which some researchers exert themshelves to investigate (ref. 4, 5). It is noteworthy that nearly no published paper about cross-linked
hydroxy-Ti smectites has been looked up in the meantime. The possible reason is that there are some difficulties in preparation of cross-linked hydroxy-Ti product. One of the difficulties is that in Ti(S04)2 solution the Ti02 precipitate tation
is easily
and rapidly formed under certain conditions. This precipi-
process includes the successional formations of binuclear, tetranuclear
and polynuclear titanium complexes (ref. 6). Because of this property of titanium salts, cross-linked hydroxy-Ti smectites can hardly be prepared by conventional direct polymeric cation exchange method.
650
In the present work a feasible preparation rrethod of cross-linked hydroxy-Ti bentonite (Ti-CLB) was found and serre interesting results were obtained on the basis of a systematic investigation of preparation and properties of Ti-CLB. EXPERIMENTAL
Montmorillonite The clay material used was the Heishan natural Ca-bentonite with a cation exchange capacity of 0.73 rreq/g produced in Lianong Province, China. The water slurry of homoionic Na-bentonite (Na-B) with particle s ize z Z pn used for the preparation of Ti-cLB was obtained from the water dispersed clay by ionexchange with sodium fonn cation resin and gravity sedirrentation rrethod. Preparation of Ti-cLB The Na-B slurry heated to 55
0C
was dropwise added to a Ti(S04)2 solution
heated to 55°C with occasional shaking at 55
0C
for IOOre than 10 hours, in
which the concentration of Ti(S04)2 was controlled in 0.06 M, then filtered and washed with distilled water until sulphatic ion free. The resulting solid
was dried at nooc. The Ti-CLBs by a titration rrethod were obtained by dropwise addition of a NaOH solution to a mixed slurry of Ti(S04)2 solution with Na-B slurry. The reference sanple, cross-linked hydroxy-A! bentonite (Al-cLB), was prepared by the sane rrethod as the references 1 and 2. Preparation of Ti-cLB supported NiO-W03 catalyst (NiO-W03/Ti-cLB) The Ti-CLB used as carrier was crashed to pieces and Irrpreqnated with a
Ni(N03)2
and (NH4)2 W04 solution containing 2% of NiO and 20% of W03 for 4
hours at room terrperature, then removed from the mot her liquor by filtration, 0C oC. dried for 2 hours at 120 and calcined for 4 hours at 450 Characterization techniques X-ray diffraction analysis was carried out with a Geigerflex D-9C X-ray diffractometer, by which the basal spacings (dOOl) of the products were determined. BET surface area and pore vo.lume distribution were obtained using a Carld Erba sorptomatLc instrurrent of series 1800, and the detennined sarrples were pretreated at 300
0C
and 0.04 torr for 1 hour. Differential thermal
analysis was carried out with a Model DA-2A instrurrent. An atomic adsorption spectrerreter was used for the determinatiml of Na20 and
CaO contents of the
serrp.les , Surface acitity of the sarrples was detennined by an amoonia adsorption differential thermal method (ADT) (ref. 7) and a n-butylamine titration method (BT). Lewis and Bronsted acid sites present on Ti-CLBs were estimated from IR measurements of pyridine adsorbed using a Perkin-Elmer 580B IR
651 spectrometer. The cracking activities of the samples were measured for n-hexane, n-bexadecane and «-methylnaphthalene cracking in a pulse microreactor at 400 oC, and catalyst sample of 0.1 g and feed amount of 0.3}l1 were used each tirre. Nitrogen was employed as the carrier gas. The hydrodenitrogenation activities of the NiO-W03/Ti-CLB catalysts were determined for pyridine hydrodenitrogenation in a continuous flow microreactor under the following conditions: oC, terrperature of 250 pressure of 40 Kg/em 2, space velocity of 3.6 ml/g.hr, hydrogen rate of 400 ml/min. and used catalyst of 1.5 g with particle size of 40-60 mesh, and the catalysts were presulfided with CS2 for 2 hours at 300
0C
prior to use. RESULTS AND DISCUSSION I t is not easy to prepare hydroxy-Ti polymer solution which can be used in
preparation of cross-linked smectites by conventional methods. We had tried to prepare it by reference to the preparation methods of the hydroxy-Al polymer solution (ref. 8) and hydroxy-Zr polymer solution (ref. 1), but faild. The addition of NaOH solution to Ti(S04)2 solution resulted in immediate formation of a white precipitate.
When Na2C03 solution was dropwise added to Ti(S04)2
solution, at first the white precipitate
formed could be dissolved rapidly,
but after a certain amount of Na2C03 solution (C032- /Ti > 1) was added,the precipitate became unsolvable. The Ti(S04)2 solution with C032-/Ti=1 which did not contain Ti02 precipitate was not able to undergo carplete crosslinking reaction with the clay. Moreover, when a certain amount of water was added to Ti{S04)2 solution, a large quantity of precipitate was formed, and when Ti(S04)2 solution was treated at a certain terrperature, at first the 4 rate of Ti + hydrolysis was slow, but after a tirre the rates of hydrolysis and precipitate formation became quicker and quicker (ref. 6). This shows that the Ti(S04)2 solution with a lower extent of hydrolysis in which there is no precipitate is not suitable for the CCIlplete cross-linking process, and when the hydrolysis extent is too high, the precipitate is formed rapidly, thus the solution is not also able to be used as cross-linking agent. According to the above considerations and the property of strong hydrolysis of Ti{S04)2 solution, we had investigated the preparation of Ti-cLB by direct reaction of Ti(S04)2 solution with Na-B slurry and found some suitable preparation 4 conditions. In the cross-linking process, the Ti + concentration in the mixed slurry and the order of dropwise mixing of the clay slurry with Ti( 804 )2 solution must be controlled strictly, so as to prevent the formation of precipitate, and a proper reaction terrperature and amount of titanium salt must be chosen, so as to increase the extent of cross-linking reaction, other-
652 wise the cross-linking reaction will be undergone incompletely and the product will not be separated from the water slurry by filtration. For exarrple, when Ti (S04 )2 solution reacted with the clay slurry at room terrperature, an incomplete cross-linking reaction was observed and a little amount of crosslinked product with doOl of 15.8
A was
separated by filtration merely~
Based
on the experimental results, the suitable cross-linking conditions for 0
Ti-cLB are mainly chosen as reaction terrperature of 55 C , reaction time of more than 10 hours, Ti4 + concentration in the mixed slurry of 0.06 M and Ti/bentonite (Ti/B) ratio of 4.0. The product prepared under the above conditions was spearated from the slurry by filtration basically completely. The basal spacing, surface area, acidity, catalytic activity, differential thermal analysis curve and Na20 and Cao contents of the obtained product show that it is a stoichianetric cross-linked bentonite.
TABLE 1 Properties of Ti-CLB, Al-CLB and RE-Y zeolite catalyst. Sarrple
doOl Surface Pore vol. distribution( %) Pore vol. (A) area(M2/g) 100A 7050- 30- 20- (ml/g) 0C 300 500 10011. 70A 50A 30A
Ti-CLB 16.2 204 Al-CLB 17.3 190 RE-Y cat. - 540
156 150
CaO
(wt%)
0.8 1.3 11.7 63.0 22.7 0.287 0.06 0.005 33.7 20.5 15.1 13.7 17.0 0.205 0.06 0.005 Average pore diameter 27A 0.700 -
Sane properties of Ti-cLB are listed in Table 1 and its X-ray powder diffraction pattern is shown in Fig. 1. The basal spacing measurement is based on the 001 diffraction which is shown as sharp reflection. The Na-bentonite treated with Ti(S04)2 solution gives a basal spacing of 16.2
A,
which is
slightly larger than that of the products by the titration method with a broad 001 diffraction peak (see Fig. 2). The diffraction peak at 29 of 19.9° of Ti-cLB in Figure 1 is the characteristic diffraction peak of the layer structure of the Heishan bentonite (020 diffraction), which shows that the essential structure of the bentonite layers was not decomposed,',eventhough.theNa..:.oontonite was treated in the acidic Ti(S04)2 solution for a long time.
653
5
10
15
20
2
4
6
8
10
12
14
28
28 Fig. 1. X-ray powder diffraction pattern of Ti-CLB dried at nooe.
Fig. 2. X-ray powder diffraction patterns of Ti-CLBs by titration method.
The Barrett-Joyner-Halenda method (ref. 9) was used to calculate pore volume
distribution from nitrogen adsorption isothenns. The results in Table 1 indicate that the pore volume distribution for Ti-CLB is mainly in the smaller pores «
5011.) and the average pore diameter is slightly larger than that in
RE-Y zeolite. Sane nitrogen sorption data for the Ti-cLB treated at 500 0 e were plotted according to the Langmuir and BET isothenns, which assume respectively monolayer and multilayer adsorption on open surface, by using the values of adsorbed gas, V, at 5 relative pressures PIP0 in the range O. 02 ~ PIP0 ' 0.32.
Figure 3 shows that only the Langmuir isotherm gave a linear plot. The surface area 204 M2 / g of Ti-cLB, which is slightly larger than that of Al-cLB, was calculated from the above results. This indicates that the high surface area is due to the microporous structure of the interlayered spaces fomed by stable hydroxy-Ti pillars. Therefore, the decrease of surface area of the Ti-cLB with increasing temperature shows that the interlayering hydroxy-Ti pillars vary with temperatures. The differential thermal analysis curve of Ti-CLB in Figure 0-GOO o 4 shows that in the treatment temperature range of 300 e, for the TiCLB an endothermic reaction corresponding to water and hydroxyl loss fran the interlayering pllars took place. This is consistent with the results given by Yamanaka et al (ref. 10) and Kodama et al (ref. 11), who indicated that cross-linked srnectites undergoes dehydration of water molecules strongly associated with hydroxy-metal cations and dehydroxylation of interlayering hydroxy-metal pillars between 200
0
and 500
0
e.
654
940
0.1
200
0.3
400
600
800
1000
Temperature (OC) Fig. 3. Plots of PIP - PIP V (L) and PIP - P/V(P _P)o(BET) ofor Ti-cLB 0 0
Fig. 4. Differential thennal analysis curve of Ti-cLB
Chemisorption of ammonia was used to assess acidity of Ti-CLB. The acidity
was expressed relatively in temperature difference potential (}JV) r which corresponded to the exotherm of the ammonia adsorption differential thennal analysis curve. The}JV value of total ammonia adsorption was used to stand for relative total acid amount, and that left after desorption for 30 min. for stronger acid amount. The acidity data obtained by Artr in Table 2 show that the Ti-cLB gave appreciable ammonia adsorption, about 1/3 of that of the RE-Y catalyst and slightly larger than the Al-cLB on a weight basis; on a surface area basis, the Ti-cLB adsorption is comparable to the RE-Y catalyst. Acidity of Ti-CLB decreases with increasing temperature. For exanple, the 0C
total acid amount of Ti-CLB treated at 300 is 912 }JV, 711}JV at 4000C and oC. 507 }JV at 500 The acidity data by BT indicates that the Ti-CLB is a weakly acidic rraterial, and has a srrall amount of stronger acidity of pKa" -5.6. Figure 5 shows that the Ti-cLB has both Bronst.ed and Lewis acidity at lower temperatures (below 300 oC) and rrainly Lewis acidity at higher temperatures, and that the acidities decrease with increasing tenperature. The acidity of Ti-cLB is consistent with the results in our previous study of surface acidity of cross-linked srrectites (ref. 3).
655 TABLE 2
Acidities of Ti--eLB, Al--eLB and RE-Y catalyst treated at 500°C. SanI:>le
Acidity by ADT (pV) Total
Ti-CLB Al-CLB RE-Y cat.
507 425 1385
Stronger 160 108 510
Acidity by BT (meq/g) pKa+4.8
+3.3
-3.0
-5.6
1.2 1.3 1.2
0.7 0.45 1.0
0.3 0.2 0.8
0.1 0 0.5
A
t
1491
-
B
1600
1500
1400
Wavenumber (cm- 1) F'ig. 5. Infrared spectra of pyridine adsorbed on Ti-CLB 0(B), o(C) at l500bA), 250 350 and 450 C(D).
Column temperature
Fig. 6. Prograrmed tenperature gas chromatogram of products in hexadecane (C16) (A) and c(-methylnaphthalene (Cll) (B) cracking over Ti--eLB.
Ti--eLB was tested for n-hexane(C6), n-hexadecane(C16) and c(-methylnaphthalene(Cll) cracking at 400°C. The experimental results are given in Table 3. The hexane cracking activity of Ti--eLB is much lower than that of RE-Y zeolite
catalyst, but slightly higher than that of Al--eLB. The hexadecane cracking conversion over Ti--eLB is lower than those over Al-CLB and RE-Y
656 catalyst and for Ti-CLB a good correlation was obtained with total acid amoilllt by MYr. A sinple corrparison of the product distributions obtained in cracking of hexadecane with Ti-CLB, Al-CLB and RE-Y catalyst can be seen in Table 3. The Ti-cLB gave more products in the C3 to C4 range and fewer in the Cl to C2
range ((Cl+C2)/(C3+C4) -c 0.02) and appreciable errourrt of isorrers (see Fig. 6A). This product distribution is characteristic of carbonium-ion cracking. The conversion of c<-rrethylnaphthalene with Ti-CLB is lower than that with RE-Y catalyst, but cooparable to that with Al-cLB. On a surface area basis, the activity differences are not obvious. Naphthalene and the components in the more than Cll range are the main products in the c<-rrethylnaphthalene conversion over Ti-cLB and there is nearly no appearance of products in the C1 to C9 range (see Fig. 6B), whereas 33.6% of the products over the RE-Y catalyst are the cooponents in the C1 to C9 range. The selectivity for naphthalene product of Ti-CLB is higher than those of Al-CLB and RE-Y catalyst, which shows that the reaction of c<-rrethylnaphthalene over Ti-cLB is possibly a disproportionation reaction. Like Al-cLB, the Ti-CLB showed rapid deactivation in the cracking of hydrocarbons with contact tirre, which was related to the easy formation of coke deposits on the Ti-CLB surface. TABLE 3 Cracking conversions for n-nexanetcg ) , n-hexadecane(C16) and c<-rrethylnaphthalene (Cll) over Ti-cLB, Al-cLB and RE-Y cat. at 400 oC. Sarrple
Acidity (uV)
Conv. for C6(%)
Ti-cLB 507 Ti-cLB 711 Ti-cLB 912 Al-CLB 425 Re-Y cat. 1385
0.23 0.1 6.6
Cll cracking
C16 cracking
Conv. C3+C4 C5-C9 Conv. (%) (%)* (%)* (%) 24.8 46.7 52.4 59.8 100.0
73.8 60.4 64.7 68.0 75.2
26.2 38.5 33.8 32.0 22.8
27.9
C** 10 >Cll <.ClO (%)* (%)* (%)* 89.7
10.3
32.6 76.1 68.3 56.2
23.9 10.2
0 0 33.6
* (wt ) % of total converted product. **C1O - Naphthalene. It has been reported that the hydrotreating catalysts supported by the
carriers cooprising titanium dioxide showed higher catalytic activites than the conventional alumina based catalysts (ref. 4). Because the Ti-CLB has more considerable acidity, higher surface area and larger pore-size than the
657
conventional Ti02,
the effect of Ti-CLB used as hydrotreating catalyst
carrier on catalyst performances will probably be greater than that of the Ti02' In the present study, the pyridine hydrodenitrogenation activities of NiO-W03 catalysts supported by Ti-CLB carrier were measured. The experimental results are listed in Table 4. TABLE 4 Hydrodenitrogenation conversions for pyridine over NiO-WO§ catalysts supported by Ti-CLB carrier. Catalyst
Conversion for pyridine (%)
NiO-W03/Ti-CLB with binder
34.1
NiO-W°3/ Ti-CLB
25.8
Industrial hydrotreating cat.
22.7
The data in Table 4 indicate that the NiO-W03 catalysts supported by Ti-cLB
carrier have higher hydrodenitrogenation activities than the industrial catalyst,
and NiO-wD3 /Ti-cLB catalyst in which a binder was added shows much
higher hydrodenitrogenation conversion than that without the binder. This shows that it is of significance to further investigate the effects of Ti-cLB used as catalyst carrier on catalyst performances. CONCLUSIONS
Direct exchange reaction of Na-bentonite with Ti(S04)2 solution under some strictly given conditions gave a cross-linked product with a basal spacing 16.2 ,., and surface area 204 M2/g. '!he nitrogen adso:rption isothenn of the Ti-cLB is of Langmuir type, indicating that the high surface area is attributed to the microporous structure fomed by the hydroxy-Ti pillars. Like Al-cLB, the Ti-cLB is a weak acidic material and has both Lewis and Bronsted acidity at lower terrperatures, whereas mainly Lewis acidity at higher temperatures. Its acidity decreases with increasing terrperature. Paraffins cracking reactions over the Ti-cLB proceed via carbonium-ion reaction mechanisms. Its cracking activity is generally in line with its acidity neasured by ADT. Methylnaphthalene on Ti-cLB undergoes probably a disproportionation reaction and the selectivity for naphthalene product is higher than that of the RE-Y zeolite catalyst. The pyridine hydrodenitrogenation conversion over NiO-wD3 catalyst supported
by Ti-cLB carrier is higher than that over the industrial hydrotreating cata-
358
lyst. A binder added to the NiO-W03/Ti-CLB
catalyst can obviously irrprove
the pyridine hydrodenitrogenation activity of the catalyst.
REFEREIICES 1 2 3
4 S 6 7 8 9 LO L1
D.E.W. Vaughan, R.J. Lussier and J.S. Magee, Stabilized pillared interlayered clays, U.S. Patent 4,248,739 (1981). J. Shabtai, Class of cracking catalysts acidic forms of cross-linked smectites, U.S. Patent 4,238,364 (1980). Liu Zhonghui and Sun Guida, Factors affecting acidity and basal spacing of cross-linked smectites, in: B. Dr:£aj, S. Hocevar and S. Pejovnik(Eds.), Studies in Surface Science and Catalysis 24 "Zeolites", Elsevier, Amsterdam-oxford-New York-Tokyo, 1985, pp. 493-500. R.J. Mikovsky and A.J. Silvestri, Use of catalyst comprising titania and zirconia in hydrotreating, U.S. Patent 4,186,080 (1980). Tsuji, Jiro, Organic synthesis with the titanocomplex, Kagaku(Kyoto), 36 (2) (1981) 153-156 (Japanese). J.F. Duncan and R.G. Richards, Hydrolysis of titanium (IV) sulphate solution. I. Precipitation of hydrous titanium dioxide, New Zealand Journal of Sci., Vol. 19 (1976) 171-178. Research Group 108 of Research Institute of Petroleum Processing, Ammonia adsorption differential thermal method for determination of surface acidity of catalysts, Petroleum Refining, No.1 (1979) 6-12 (Chinese). N. Lahav and U. shani, Cross-linked smectites. 1. Synthesis and properties of hydroxy-aluminum-montIooril1onite, Clays and Clay Minerals, Vol. 26, No.2 (1978) 107-115. G.P. Barrett, L.G. Joyner and P.H. Halenda, The detennination of pore volume and area distribution in porous substances. I. Camputation from nitrogen isotherms, J. AIrer. Chern. soc., 73 (1950) 373-380. S. Yamanaka and G.W. Brindley, High surface area solids obtained by reaction of montIoorillonite with zirconyl chloride, Clays and Clay Minerals, Vol. 27, No. 2 (1979) 119-124. H. Kodama and S. Shan Singh, Hydroxy aluminum sulfate-montIoorillonite complex, Can. J. soil Sci., 52 (1972) 209-218.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
659
EFFECT OF HYDROXY-ALUMINUM POLYMERIC CATIONS ON ACIDITY OF CROSS-LINKED HYDROXY-ALUMINUM SMECTITES
SUN
GUIDA 1,
YAN
FUSHAN1,
SUN
DEHAI1 and LIU ZHONGHUI2
1Departrrent of Petro-Chemical Engineering, Fushun Petroleum Institute, Fushun, Liaoning Province, P.R. China 2Research Institute of Petroleum Processing, Beijing, P.R. China
ABSTRACT Three cross-linked hydroxy-aluminum bentonites were synthesized by reactions of Na-bentonite respectively with three hydroxy-AI polymers with different polymeric degrees. Acidity of cross-linked hydroxy-AI bentonites with same amounts of bentonite and aluminum entering the clay interlayers increases with increase of polymeric degree of hydroxy-AI polymer cations. Dehydrationdehydroxylation of interlayering hydroxy-AI pillars between 300o-600OC causes the polymeric structure of the pillars progressively to decompose and tf'.8 acidity decreases with increasing terrperature. This further assumes that the acidity is mainly from the polymeric structure of interlayering pillars affected by the clay layers. In stoichiorretric cross-linking reactions the higher the polymeric degree of hydroxy-A! cations, the IOOre is the amount of aluminum exchanging into bentonite. INTRODUcrION Cross-linked srrectites are a new class of IOOlecular sieve-like materials with a larger pore-size, in which the cross-linked hydroxy-A! srrectites can catalyse cracking reactions of hydrocarbons and their advantage over conventional zeolites becares IOOre marked with gradual increase in the kinetic diameter of hydrocarbons, which bring to attention of sane research workers. This indicates the potential of cross-linked srrectites for cracking of heavy feedstocks (ref. I, 2, 3, 4). Therefore, it is of significance to systematically investigate their acidities. The present study is part of a program to study acidities of cross-linked srrectites. In our previous study, it has been explored that the acidity of cross-linked srrectites is usually weak and directly
related to the kind and
amount of their interlayering pillars, and that the cross-linked hydroxyA! srrectites treated at over 3000C have Lewis type acidity mainly (ref. 5). The present study is concerned with the relationship between acidity of crosslinked hydroxy-A! bentonite (A!-cLB) and hydroxy-AI polymeric cations.
660
EXPERIMENTAL Clay materi.cil the clay material used was natural calcium form Heishan bentonite produced in Liaoning Province, China and possessed a cation exchange capacity of 0.73 meg/g. The water slurry of homoionic fonn bentonite (Na-B) with particle size ~
2 pm used for the preparation of Al-CLB was obtained from the clay slurry
by ion exchange with sodium fonn resin and gravity sedimentation. Hydroxy-Al polymers Two hydroxy-Al solutions respectively with OH/Al molar ratios of 1.1
and
1. 85, Al( 1.1) and AI( 1. 85). were prepared by dropwise addition of a NaOH solution to an AlCl3 solution. The resulting solutions were aged at room terrperature for 7 days. Another hydroxy-Al solution, Al{NH3), was prepared by dropwise addition of NH3 water to the hydroxy-Al solution of Al{1.85) and the pH of the resulting system was controlled at 6.0. Preparation of cross-linked bentonite The hydroxy-Al solution was dropwise added with vigorous stirring into
the Na-B slurry at a certain Al/Bentonite ratio (Al/B). After stirring at 750C for 2 hours, the exchange was essentially conplete. The resulting system
was filtered, washed and oven-dried at 120OC. Characterization techniques Acidity of cross-linked products was detennined by an anmonia adsorption differential theJ:mal method (ADT) (ref. 6) and expressed relatively in tenperature
difference potential (IlV), which corresponded to exothenn of the anmonia
adsorption differential thennal analysis curve. The pv value of total anmonia adsorption can be used to stand for relative total acid amount of the product and that left after desorption with nitrogen for 30 minutes for relative stronger acid amount. Infrared spectra were obtained using a Perkin-elmer 580 B IR spectrometer. The basal spacings (d001) of the products were determined by X-ray diffraction analysis, which was carried out with a Geigerflex D-9C X-ray diffractometer. BET surface areas were obtained using
aCarId
Erha sorptanatic instrument of series 1800. Differential thermal analysis
was carried out with a Model DA-2A instrument. An atomic adsorption spectrometer was used for the detennination of Na20 and CaD contents. RESULTS AND DISCUSSION
In equilibrium system of AlClrNaOH solution, molecular weight of hydroxy-
Al polyrreric cations increases with increase of OH/Al molar ratio or pH of
661
the solution generally (ref. 7, 8). In order to further verify this point, we have determined freezing point depression values of the hydroxy-Al polyrrer solutions, then calculated average polymeric degrees of the hydroxy-Al polyrrers (average number of aluminum atoms in a hydroxy-Al polyrreric cation) with the freezing point depression constant Kf' which was obtained by calculating from the freezing point depression value of a
NaCl solution with the similar
ionic strength to the hydroxy-Al solutions. Both hydroxy-Al solutions of each group in Table 1 were adjusted to have same Al +3 concentration and same nonpolymeric ion concentration. Table 1 shows that the polymeric degrees of hydroxy-Al polymers became sequentially larger in the order of Al(1.l), Al(1.85) and Al(NH3)' This can also be shown from the dOOI values in Table 2, because in general, the larger the hydroxy-Al polymeric cations, the larger is the basal spacing of their cross-linked product. TABLE I Freezing point depression values of Al(1.l), Al(1.85) and Al(NH3) solutions and average polymeric degrees of their polymeric cations. Group
Hydroxy-Al polyrrers
pH
I
Al(1.I) Al(1.8S)
4.0 4.2
0.665 0.605
6
Al( 1.85) Al(NH )
4.2 6.0
0.700 0.690
6 >13
2
Freezing point depression (OC)
Average polymeric degree
2
In the previous study, it has been indicated that acidity of cross-linked smectites is directly related to the kind and amount of their interlayering pillars and consistent with their cross-linking degree (ref. S). Therefore, the samples which were used for the investigation of effect of hydroxy-Al polymeric cations on acidity of Al--cLBs should contain same amount of bentonite and same amount of aluminum in the interlayering pillars. In order to obtain the cross-linked products with same amount of bentonite and same amount of aluminum in the interlayering hydroxy-Al pillars, a lower Al/B ratio (0.9 rrM/ g) was selected for preparation of Al( 1. 85 )-CLB and Al(NH3)-CLB. 'Ihe aluminum contents in the pillars of Al(1.I)-CLB with Al/B=1.5 and Al--cLBs with Al/B=0.9 in Table 2 indicate that they contain same amount of bentonite and same amount of aluminum in the pillars basically. Therefore, the difference among their acidities is caused by the different hydroxy-Al polymers. Their acidities measured by ADT method in Table 2 and the infrared
662 spectra in Figure 1 show that the larger the polyrreric degree of interlayering hydroxy-Al polyrrers, the higher is the acidity of their cross-linked products. This indicates that the acidity of Al-eLB is also related to the polyrreric degree of its interlayering hydroxy-AI pillars. TABLE 2 Properties of Al-cLBs treated at 500OC. Saq>le
Al/B d001<1,,) (11M/g) (dried)
Acidity(pV)
Na20
CaO Al content
(wt%)(wt%) in pillars
Total Stronger Al( 1.85 )-cLB Al( NH3)-CLB Al(1.1)-CLB Al( 1.85 )-cLB Al(NH3) -CLB Na-B
0.9 0.9 1.5 1.5 1.5
18.20 18.35 15.65 18.47 18.64 12.63
302 354 285 425 502 86
75 96 73 108 150 5
(mg/1gB)
0.09 0.08 0.06 2.10
0.006 0.006 0.006 0.04
22.1 24.0 23.7 33.5 40.3
Moreover, it is noteworthy that the stoichiometric cross-linked bentonites syntlY:lsized with different hydroxy-Al solutions have not sarre anPunt of aluminum in the interlayering pillars in bentonite of 1 g (see Table 2). The Na20 and CaO contents of the Al-cLBs with Al/B=1.5 in Table 2 are negligible, so they can be considered to be stoichiometric cross-linked products. The emountis of aluminum in tlY:lir pillars show that in the stoichi~ic
Al-cLBs,
the higher the polyrreric degree of interlayering hydroxy-AI cations, the larger is the enourrt of aluminum in the pillars of their cross-linked product. The possible reason is that average positive charge of an Al atom in hydroxy-
Al polyrreric cations varies with their polyrreric degree, Le. the greater the polyrreric degree, the lower is the average positive charge of an Al atom in hydroxy-AI polymeric cations. Therefore, the amount of aluminum in the hydroxy-Al cations with higher polyrreric degree which exchange into bentonite is larger in stoichiometric cross-linking reaction. The experiment results show that an Al atom in the hydroxy-Al cations with lower polyrreric degree has larger average positive charge, and vice versa. This is in line with the
fonnulas of some hydroxy-Al polyrreric cations given by Occelli et al (ref. 7) and Vaughan et al (ref. 8). In order to investigate the relationship between acidity of Al-eLB and polyrreric structure of its interlayering hydroxy-Al pillars, some properties
663
c
L
t~456 1600
1500
1400
Wavenumber (cm- 1)
1600
1500
1400
Wavenumber(cm- 1)
"454 1600
1500
1400
wavenumber(cm- 1)
Fig. 1. Infrared spectra of pyridine adsorbed at 250 0 (1) and 450 0 C (2) on Al(1.l)-eLB (Al, Al(1.85)-eLB (B) and Al(NH3)-eLB (C) after treatment at 3500C and lXlO-5 torr for 4 hours.
of Al-eLBs treated at different temperatures were measured. Table 3 gives the results of infrared analysis in the OH region of Al--cLBs and indicates that the errount of water and hydroxyl lost in Na-B between 250 0 and 5000C was smaller than that in Al--cLB obviously. This shows that for the Al-eLB, the water and hydroxyl were mainly lost fran its interlayering pillars. Fran Figure 2 it can be seen that the differential thermal analysis curves of AlCLBs are different fran that of Na-B in the tE!llIJerature range of 300 0-600OC, and that for the Al-CLBs an endothermic reaction corresponding to water and hydroxyl loss fran the interlayering pillars took place between 300 0 and 600OC. Meantime, it can be seen that the endothermic peaks of Al-eLB at 6710C corresponding to structural hydroxyl loss fran the clay surfaces were smaller than that of Na-B at 690OC, and the hydroxyl content of Na-B treated at 4500C was larger than that of Al-eLB (see Table 3). This suggests that the hydroxyl of the clay surface in Al-cLB was easier lost than that in Na-B, or that
near-covalent bonding between the hydroxyl and hydroxy-Al polyrrers was formed. 'Ibis is probably corresponding to the guess of Lahav et al (ref. 91, who indicated that heat treatment of Al-eLB at terrperatures of 2500C results in gradual transition fran ionic to near-covalent bonding between the polyrreric species and the clay layers. Because the high surface areas of cross-linked smectites are due to the microporous structure of the interlayered space formed by the stable pillars (ref. 10), variation of surface area of Al--cLB should
664 be caused by change of its hydroxy-Al pillars. Table 4 indicates that surface
area of Al-cLB is dependent on treatment terrperature, which is caused by decomposition of polymeric structure of its hydroxy-Al pillars. The above discussion is consistent with the results given by Yamanaka et al (ref. 10) and KOdama et al (ref. 11). It is noteworthy that Vaughan et al (ref. 8) and Occelli et al (ref. 7) indicated that the dehydration-dehydroxylation reaction of the
interlayering polymeric cations is formulated as follows:
and at high telTQ;l8I'ature the prcxluced protons are capable of leaching Al+-~
from
the hydroxy-Al pillars. This can explain the fact that at high temperature the Al-<:LBs contain lewis acid sites mainly and have no or trace amount of Bronsted type acidity. TABLE 3
Hydroxyl contents of Al(1.85)-<:LB and Na-B treated at different terrperatures. Sample
Relative hydroxyl content (peak area rrm2/10 mg)
*A--3740cm-1. **e--3645cm-1 for Al-<:LB and 3620am- 1 for Na-B. TABLE 4
Surface areas of Al-cLBs and Na-B treated at different tefll>8ratures Sample
Surface area (M2/g ) 3000C 5000C
Al(1.1)-CLB Al( 1.85 )-CLB Al (NH )-<:LB 3
Na-B
88 135 155 56
62 110 131 50
Temperature (Oe) Fig. 2. IJI'A curves of Al-cLBs and Na-B.
665 The acidities of Al-cLBs treated at different temperatures in Table 5 and
the infrared spectra of pyridine adsorbed on them
in
Figure
1
show
that in the temperature range of 300 0 -600OC acidity of Al-CLB decreases with increase of treatrrent temperature. From the above discussion, it has been known that in this temperature range dehydration-dehydroxylation reaction
causes the polymeric structure of the pillars to decorrpose , In addition, the hydroxy-Al solid obtained by evaporating the hydroxy-AI polymer solution (Al(1.85»
possessed lower acidity (c ,a, 100 pV) than the Al-cLB with the
same amount of aluminum as the solid, and the previous study indicated that acidity of cross-linked srrectites is mainly due to their interlayering pillars. This further shows that acidity of Al-CLB is directly related not only to polymeric structure of its interlayering pillars, but also to effect of the bentonite surface upon the interlayering pillars. Therefore, it is reasonable to assume that the acidity of Al-CLB is mainly from the polymeric structure of its interlayering pillars affected by the bentonite layers. In view of the above, we propose a IlPdel of the Lewis acid sites in hydroxyAl polymeric structure in the clay interlayers: OH
OH
OH
OH 1.
I I I -O-M-O-W\-O-~.::. -0-1:\-0, (H;2C»" '"
heat
00000
..
'si'" 'si"si'" 'si"St'
/\/\/\/\./\ (Clay surface)
This Lewis acid site can be transfonred into bronsted acid site by adsorbing water at lower terrperature, which can explain the fact that at below 3000C the Al-cLB has marked Br8nsted type acidity. TABLE 5
Acidities of Al-cLBs treated at different temperatures. Acidity- (PV)
sanple 3000C
4000C
5000C
6000C
Tot.* Str.** Tot.* Str.** Tot.* Str.** Tot.* Str.** Al(l.l)-cLB Al(1.85)-cLB Al( NH3)-cLB
530 746
220 350
492 631 640
* Tot. - Total acid aIOOUIlt. **Str. - Stronger acid amount.
111 130 160
285 425 502
73 108 150
179 303 351
48 67 73
666
',Tbeachievable acidity of treated Al-CLB is variable, but can be controlled and maximized. The primary variables used to obtain this optimization are polyrreric degree and amount of interlayering hydroxy-AI polymers, cation exchange capacity and structure of clay. Hydroxy-Al polymers with high polymeric degree, which are able to increase acidity of cross-linked product and improve stability of the product, and suitable clay, which can stabilize polyrreric structure of interlayering pillars, should be used in AI-cLB preparation, so as to obtain cross-linked SIrectites with high acidity and stability. CONCLUSION Acidity of Al-cLBs is affected not only by amount of interlayering hydroxyAl polyrrers but also by their polyrreric degree, Le. acidity of Al-CLB increases with increase of polymeric degree and amount of the polymers. When Al-CLB is treated in the temperature range of 300 o-600oC, its acidity decreases with increasing temperature. This acidity decrease is due to deconposition of polyrreric structure of the interlayering pillars caused by their dehydration-dehydroxylation. Thus, it is assumed that acidity of the cross-linked hydroxy~aluminumbentonites
is mainly from the polyrreric structure of inter-
layering hydroxy-e.hmirum pillars,which are affected by the smecti.tes interlayers. For the stoichiorretric Al-cLBs, the amount of aluminum entering the bentonite interlayers increases with increase of polyrreric degree of the hydroxy-Al polymeric cations. The hydroxy-AI polymers with higher polymeric degree are able to improve the acidity and stability of the cross-linked SIretites.
REFERENCES
1 2 3
4 5
6 7
D.E.W. Vaughan, R.J. Lussier and J.S. Magee, pillared interlayered clay materials useful as catalysts and sorbents, U.S. Patent 4,176,090 (1980). J. Shabtai, Class of cracking catalysts acidic forms of cross-linked SIrectites, U.S. Patent 4,238,364 (1980). J. Shabtai, R. Lazar and A.G. Oblad, Acidic forms of cross-linked smect.Lt.es , A novel type of cracking catalysts, in: T. Seiyama and K. Tanabe (Eds.), Proc. 7th Inter. Congr. Catal., Tokyo, 1980, KodanshaElsevier, Tokyo, 1981, pp. 828-837. J. Shabtai and R. Lazar, Catalytic cracking properties of novel crosslinked SIrectites, Preprints, Vol. 24, No. 2 (1979) 622-623. Liu Zhonghui and Sun Guida, Factors affecting acidity and basal spacing of cross-linked snectLtes , in: B. Dr!l!aj, S. HoCevar and S. Pejovnik (Eds.), Studies in Surface Science and Catalysis 24 "Zeolites", Elsevier, Amsterdam-Dxford-New York-Tokyo, 1985, pp. 493-500. Research Group 108 of Research Institute of Petroleum Processing, 1lrmPnia adsorption differential thennal rrethod for deteI:Illination of surface acidity of catalysts, Petroleum Refining, No.1 (1979) 6-12 (Chinese). M.L. OCcelli and R.M. Tindwa, Physicochemical properties of IlOntmoril-
667
8 9 10 11
lonite interlayered with cationic oxyaluminum pillars, Clays and Clay Minerals, Vol. 30, No.1 (1983) 22-28. D.E.W. Vaughan, R.J. Lussier and J.S. Magee, Stabilized pillared interlayered clays, U.S. Patent 4,248,739 (1981). N. Labav and U. Shani, Cross-linked smectites.1. Synthesis and properties of hydroxy-aluminum-montmori1lonite, Clays and Clay Minerals, Vol. 26, No.2 (1978) 107-115. S. Yamanaka and G.W. Brindley, High surface area solids obtained by reaction of montmorillonite with zircony1 chloride, Clays and Clay Minerals, Vol. 27, No.2 (1979) 119-124. H. Kodama. and S. Shan Singh, Hydroxy aluminum sulfate-montmorillonite complex, Can. J. Soil Sci., 52 (1972), 209-218.
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DISCUSSION J.W. JOHNSON: Can you comment on the thermal and steam stability at elevated temperature of your titanium bentonite samples? SUN GUIDA: In our study, we determined some properties of all cross-linked samples treated at 500°C. The experimental results show that cross-linked bentonites treated at 500°C have considerable surface acidity, cracking activity, surface area and basal spacing. We measured the steam-thermal stability of Al-CLBs merely. After the steam-thermal treatment of Al-CLB at 600°C for l8h, its acidity and cracking activity decreased obviously, and after the steam-thermal treatment at 700°C for l8h, its acidity and cracking activity were approximately those of Na-bentonite. Therefore, we can predict that the thermal and steam stability of Ti-CLB at elevated temperature (at more than 600°C) is also low. This shows that the thermal and steam stability at elevated temperature of cross-linked smectites should be improved. According to our experience with the preparation of cross-linked smectites, we think that there are mainly two possibilities for improvement of the thermal and steam stability of cross-linked smectites. The first possibility is to increase the polymeric degree of the cross-linking agent (hydroxy-metal polymers). But it is rather difficult to increase the polymeric degree of hydroxy-Ti polymers. The second possibility if to use other clay materials with a special structure, which is able to stabilize the polymeric structure of the hydroxy-metal pillars. We shall make a further study in this respect.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
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CONTROLLED PREPARATION OF MONODISPERSE BIMETALLIC Pd-Au COLLOIDS WITH THREE DIFFERENT MICROSTRUCTURES AND THEIR USE IN PREPARING SUPPORTED BIMETALLIC CATALYSTS J. B. Michel- and J. T. Schwartz-
Central Research and Development Department, E. I. du Pont de Nemours and Company, Experimental Stat ion, wilmington, 19898, U.S.A.
DE
SUMMARY Bimetallic monodisperse Pd-Au particles with three different microstructures have been reproducibly prepared, in pure form, as colloidal sols. The three microstructures constitute a complete set for the bimetallic Pd-Au system: 1. alloyed Pd-Au particles 2. layered, Au-coated Pd particles 3. layered, Pd-coated Au particles This is the first report of microstructures 1 and 3. The three microstructures have been unambiguously characterized using high resolution analytical electron microscopy. Supported Pd-Au/C catalysts have been prepared by adsorption of the colloidal Pd-Au particles onto carbon supports. Characterization of the catalysts with AEM indicates the Pd-Au particles are deposited without change of microstructure onto the support. Thus, supported bimetallic catalysts have been prepared with a unique degree of control over metals microstructure. INTRODUCTION Supported bimetallic catalysts are finding an increasing number of applications in industrial catalytic processes. Important examples are alumina-supported Pt-Re and Pt-Ir catalysts used in the reforming of naphtha to gasoline (ref. 1). Because of their importance and unique catalytic properties, there is tremendous interest in understanding, and ultimately optimizing, the relationship between microstructure and catalytic performance (ref. 1,2). Correlations between catalyst microstructure and performance ideally require precise control of metals-microstructure, i.e. the size, composition, and elemental distribution of each supported bimetallic particle. This degree of control is generally not feasible using conventional catalyst preparation -Authors to whom correspondence should be addressed.
670
techniques such as coimpregnation or stepwise deposition of metal salts onto a support (ref. 3). The chemistry employed does not guarantee the formation of a homogeneous microstructure with all particles bimetallic. Furthermore, the support can play an active role in the formation of the metal particles. Pt/Ti0 2 is a convincing example of support participation in the microstructure formation of supported metal particles (ref. 4). Several different strategies have been developed in order to achieve more precise control of microstructure during catalyst preparation. A surface science approach has been used to epitaxially deposit a thin layer of metal onto one face of a single crystal of the second metal, producing a model bimetallic catalyst with a microstructure as precise as a single crystal (ref. 5). The physical and chemical properties of these model catalysts are being studied using powerful surface science techniques and under simulated catalytic conditions. A strategy adopted for practical bimetallic catalysts is to deposit bimetallic organometallic clusters on a support (ref 6). This is an attractive way of preparing sites which are indeed bimetallic; however, the characterization of these small metal clusters is quite difficult. The purpose of this paper is to describe a novel technique for the preparation of supported bimetallic catalysts, which permits a unique degree of microstructure control. Bimetallic monodisperse Pd-Au particles have been prepared as colloidal sols in the absence of a support. The size, composition, and elemental distribution of each bimetallic particle can be controlled. After formation of these well-defined microstructures, the bimetallic particles are deposited from the sol, unchanged, onto a support. We also report new chemistry of colloidal bimetallic PdAu. A complete range of microstructures has been prepared and characterized: 1) "alloyed" Pd-Au particles 2) layered particles, with an Au coating over a Pd core 3) layered particles, with a pd coating over an Au core Prior to this work, microstructures 1 and 3 had not been reported. All have been unambiguously characterized using high resolution analytical electron microscopy.
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EXPERIMENTAL Preparation of colloidal sols and supported catalysts For all of the preparations described here, the water used was distilled and then filtered through a water purification system, equipped with mixed bed ion exchangers and an organic/colloid removal column. The glassware used was washed with aqua regia and then rinsed with the purified water. All solutions were filtered through a 0.22~ Millipore filter prior to use. Colloidal Pd and colloidal Au sols were prepared by reduction of the metal salt with sodium citrate following the procedures of Turkevich (ref 7,8). Alloyed PdxAul-x sols with three different compositions, x = 0.2, 0.5, and 0.8 were prepared using a two-step procedure. First, a solution of Pd(II), HAuC1 4 and H2 0 was brought to reflux. Then a 1 % sodium citrate solution was added, and the solution refluxed for four hours. The quantities required are shown in Table 1. TABLE 1.
Reagent quantities for preparation of PdxAul-x sols Reagent
Volume Weight Volume Volume
Pd(II) solution 1 (cc) HAuC1 4·H2 0 (mg) H2 0 (cc) 1% sodium citrate (cc)
Sol
75 85 885 120
115 38
485 400
170 15 430 400
lpd(II) solution: 0.165 g Pd(II)CI 2 + 20 cc IN HCl diluted to 1 liter TwO different layered Pd-Au microstructures were prepared. Au-coated Pd particles were prepared using the procedure of Turkevich and Kim (ref. 7). In this procedure, Au(III) is reduced by hydroxylamine and deposited onto colloidal Pd "seeds". The preparation to make Pd-coated Au particles was carried out in the following way: 250 ml of 9.3 x 10-4 M PdCI 2 , 500 ml of 3.4 x 10- 2 M sodium citrate and 500 ml of AU colloid were added to a round bottom flask. The Au colloid had been filtered by a Millipore ultrafiltration system to lower the ionic strength. This mixture was allowed to reflux for six hours. The coating of Pd onto Au was followed with optical spectroscopy.
672
Supported catalysts were prepared by deposition of colloidal metal particles onto various carbon supports. Fine carbon powder was added to the colloidal metal sol, the slurry stirred, and the removal of colloidal metal from the sol by adsorption onto the support was monitored with optical spectroscopy. Supports included Calgon8 PCB 12 x 30 coconut carbon and American Cyanamid wide pore carbon (ref. 9). Granular carbon supports were ground and sieved to 325 mesh before use. Physical characterization 1. Optical absorption spectra were recorded from 200 to 800 nm using a HP8450A diode array spectrophotometer. 2. Analytical electron microscopy (AEM): Sol and supported catalyst samples were studied using a Vacuum Generators HB501 scanning transmission electron microscope (STEM). This 100 KV STEM is specially designed for high resolution imaging coupled with high resolution energy dispersive x-ray microanalysis. X-rays are collected by a Kevex detector, interfaced with a Tracor Northern 2000 system for data collection, storage, and analysis. In addition, the computer can be used to control the 20K electron beam during analysis, in order to generate digital x-ray images (x-ray dot maps). This analytical tool allows the imaging and microanalysis of particles 20! in size (ref. 10). RESULTS Stable colloidal suspensions of alloyed PdxAul-x particles have been synthesized by reduction of an aqueous solution of Pd(II} and Au(III) with citrate ion. The ratio of Pd/Au in these sols can be varied over a wide range, provided higher levels of citrate are used when Pd/Au~l (Table 2). TABLE 2.
Concentrations in PdxAul-x Sols
Sol
Pdo . 2 AuO . 8 Pdo. s AuO . 6 Pdo. 8 Au O . 2 Pd Au
Concentrations (roM) [Pd]
[AU]
0.065 0.187 0.158 0.190
0.220 0.106 0.042 0.254
[Pd]+[Au]
[citrate]
[citrate] [Pd]+[Au]
0.285 0.293 0.200 0.190 0.254
3.78 13.60 13.60 13.60 1. 70
13.3 46.4 68.0 71.6 6.7
673
1.5 1.4 1.2
~
1.0
~ 0.80 ~ 0.60
0.40 0.20 0.0 230
300
400
5
WAVELENGTH lnml
Fig. 1. Optical absorption spectra recorded during formation of PdO.2Auo.e colloidal sol. Numbers indicate hours at reflux.
1.50 1.40
1.20
...... ......
----L
......
.....
-'--~'"';--=--....-
~----~
0.20
-'--......
• 300
700
800
Fig. 2. Comparison of optical spectra of alloyed PdAu sols (broken lines) to monometallic sols' spectra (solid lines). Legend: a - AU~ b - Pdo.2Auo.e; c - Pdo.sAuo.s; d - Pdo.eAuo.2; e - Pd.
674
The formation of the PdO.2AuO.8 sol was monitored during synthesis with optical spectra, as shown in Fig. 1. The spectral data suggest that the reaction is complete after three hours of reflux. The optical spectra of three sols with Pd/Au = 0.25, 1.0, and 4.0 are superposed with the spectra of pure Pd and pure Au in Fig. 2. The alloyed nature of the Pd-Au sols is apparent, since their spectra are not combinations of the spectra of the pd and Au sols. Analytical electron microscopy was used to characterize the microstructure of the PdxAul-X particles. The uniformity of particle size of a PdO.5AUO.5 sol is shown in a STEM electron micrograph in Fig. 3a. The bimetallic nature of individual particles was determined by x-ray microanalysis in the STEM. A typical result for PdO.2AuO.8 is shown in Fig. 3b.
Ar•• I
Fig. 3. a) STEM micrograph of PdO.5AuO.5 particles. b) STEM image and x-ray emission microanalysis of PdO.2AuO.8 colloidal sol. Layered Pd-Au colloids were also prepared. A layer of Au was deposited on colloidal Pd particles using the procedure of
675
Turkevich and Kim (ref 7). The layered microstructure was confirmed with analytical electron microscopy, as shown in Fig. 4.
Ar •• 2 Au
Au
Fig. 4. Annular dark field STEM image and x-ray emission microanalysis of layered, Au-coated Pd colloidal particles. The formation of the Pd-coated Au particles was accom- , plished by seeding a colloidal Pd preparation with 2001 Au colloid particles. The colloidal AU particles may be seen in Fig. 5 which is a micrograph obtained with a transmission electron microscope. The formation of the novel Pd-coated Au microstructure was followed by optical absorption spectroscopy. The spectra are shown in Fig. 6. As refluxing began, the characteristic peak of colloidal Au was apparent at 518 nm. As the reaction progressed, the decrease of th~peak at 518 nm was observed with th~
676
Fig. 5. TEM image of colloidal Au particles.
0.40 0.20 Pd---
O.O=-_~:--_---:-::300 230
400
_ _-;-J:"-_ _-=-==:--_--=~ 500 600 WAVELENGTH lnml
700
__-;;;\
800
Fig. 6. Optical absorption spectra of colloidal Pd (solid line) compared to spectra obtained during formation of Pd-coated Au sold (broken lines: hours of reflux indicated). appearance of a new peak at 254 nm which is characteristic of colloidal Pd. This microstructure was also unequivocally determined by high resolut ion AEM as shown in Fig. 7. Fig. 8 is a STEM image obtained at a magnification of 500,000, and it shows several Pd-coated Au particles. As can be seen from this micrograph, all of these particles are monodisperse. Analyses carried out on several of these particles confirmed the thin coating of
677
Are.
Are. 3
Are. 2
1
Au
! I.
ii
:.
Pd
J
~ulf'
)~ IPd
Ii! )\~L_
Pd
Aul
r
, i\ ll .A, ....
I"
Fig. 7. STEM image and x-ray emission microanalysis of layered, Pd-coated Au particle.
Fig. 8. STEM image of layered, Pd-coated AU particles.
678
Pd on Au. Additional evidence for the coating may be seen in the bright field image by the variation in contrast due to differences in electron densities. Supported metal catalysts were prepared by adsorption of colloidal metal particles onto finely-divided carbon suspended in the sol. Metal loadings of up to 3wt% were achieved, depending on the sol and the adsorption characteristics of the carbon support. Analysis of the supported catalysts in the STEM indicated that the colloidal metal particles were deposited intact, with little agglomeration in most cases, and that the bimetallic particle microstructure formed in the colloidal sols was retained in the supported catalysts. Examples of carbon-supported catalysts are shown in Fig.'s 9 and 10.
Fig. 9. STEM image of colloidal Pd on carbon
DISCUSSION Colloidal metal sols have been studied since the time of Faraday, who reported the preparation of colloidal gold in 1857 (ref. 11). Research has led to a variety of applications for metal colloids, including electroless plating (ref. 12), labelling of charged sites in cells (ref. 13) and others. Current research has been stimulated by the discovery of surface enhanced
679
Fig. 10. STEM image of layered, Pd-coated Au particles on carbon.
Raman scattering (SERS) of molecules bound to colloidal Cu, Ag, or Au particles (ref. 14). This tradition of research has established that the chemistry of particle formation in colloidal sols can be manipulated with great precision. particularly striking demonstrations of this precision are provided by the work of Matijevic (ref. 15) and Turkevich (ref. 3d, 16). In this paper we report the first preparation of a complete range of microstructures of colloidal bimetallic particles. Alloys and both bilayered combinations of colloidal Pd-Au have been reproducibly prepared at temperatures of ca. lOODC. The low temperatures employed permit isolation of the layered microstructures, thermodynamically unstable at high temperatures. These colloidal Pd-Au particles have been carefully characterized, and deposited without change of microstructure onto supports. The bimetallic catalysts are produced wi th a unique degree of control over their metals microstructure. Preparation of Pd-Au Colloidal Sols Reproducible syntheses of Pd and Au colloids developed by Turkevich and coworkers (ref. 7,8) provide a strong foundation for developing the chemistry of the bimetallic colloids. The monometallic colloidal sols are prepared by refluxing an aqueous
680
solution of the metal chloride with an excess of sodium citrate, Nag[(02C) (OH)C(CH2C02)2] (concentrations shown in Table 2). One of the roles of the citrate ion is to act as a surfactant by adsorbing on the surface of the metal particles. This stabilizes the dispersion against aggregation. A higher ratio of citrate/metal is required by Pd (Table 2), reflecting its lower colloid stability. In this work, "alloyed" PdxAul-x sols (x = 0.2, 0.5, 0.8) have been prepared by coreducing Pd (II) and Au(II 1) wi th citrate ion. Consistent with the relative colloidal stabilities of Pd and AU, a higher ratio of citrate to metal is required for the two sols with high Pd content, Pdo.5AuO.5 and Pdo.aAuO.2. These sols have been characterized with optical absorption spectroscopy (200 - 800 nm) and analytical electron microscopy. Optical spectra recorded during synthesis indicate sol formation is complete after 3 - 4 hours of reflux. This reaction time is consistent with the reflux times required for the formation of the monometallic sols, 0.5 hr. for AU, 4 - 6 hr. for Pd. The "alloyed" microstructure of the sols can be deduced by a comparison of their optical spectra to those of colloidal Pd and AU, shown in Fig. 2. The alloys' spectra can not be obtained by combination of the spectra of colloidal Pd and colloidal Au. Instead, the alloy sols' spectra exhibit weaker absorbance in the vicinity of the absorbance maxima of the pure metals (Pd - 250 nm; Au - 520 nm) , but stronger absorbance at all other wavelengths than expected from a weighted combination of the Pd and Au spectra. This stronger absorbance is particularly obvious for the Pdo. 2AuO. a sol. Turkevich has observed a similar "spreading of absorbance" over a wide range of wavelengths in the spectra of alloyed Pt-Au colloids (ref. 16). However, in studies of alloyed Ag-Au colloids, Papavass i.Li ou has observed a single peak which shifts linearly with composition (ref. 17). The microstructure of the novel alloyed Pd-Au colloidal particles was directly observed and characterized using analytical electron microscopy. The uniformity of particle size of a PdO.5AuO.5 sol is shown in Fig. 3a. Some aggregates are observed, but most particles are relatively monodisperse, with an average particle size of ca. 1601. In addition to particle imaging, high resolution microanalysis of individual particles as small as 10-201 is feasible in this analytical electron microscope (ref. 10). Composition profiles across individual
681
colloidal particles can be measured using the microscope's finely-focussed 20X electron beam. Profiles are measured by collecting x-ray emission spectra at several points along a particle's diameter. Microanalys is of a part icle f rom a Pdo. 2AuO. 8 sol is shown in Fig. 3b. Analyses at two points along the part icle' s diameter are equivalent, indicating there is no segregation of Pd and Au within the particle. In addition, the particle's analysis is consistent with the overall composition of the sol. Twenty or more particles from each sol were analyzed. Typically, only small variations in particle compositions were observed. To the best of our knowledge, this is the first time that the elemental composi tion and distribution of individual colloidal bimetallic particles has been reported. Bimetallic particles with layered microstructures complete the range of Pd-Au microstructures. Both possible bilayer microstructures have been prepared, reproducibly, in pure form: a Pd core with an Au coating, and an Au core, with a Pd coating. Colloidal particles with the first microstructure, an Au coating on a Pd core, were prepared using a procedure developed by Turkevich (ref. 7). Reduction of Au(rrr} by hydroxylamine in the presence of colloidal Pd causes deposition of an Au layer onto the Pd seed particles. The layered microstructure of the Pd-Au particles is established by characterization in the analytical electron microscope. The segregation wi thin many of the particles is clearly visible in the annular dark field image,. shown in Fig. 4. Proof of the Pd-Au segregation is provided by x-ray microanalysis. The x-ray emission spectra are included in: Fig. 4. Focussing the 20X electron beam near the particle edge (area l), penetrates only the Au coating - no Pd signal is detected in the x-ray emission spectrum. Shifting the beam to the particle'S center (area 2), the beam now penetrates both the Au coating and the Pd core, so both Pd and Au x-rays are detected. prior to this work, the Pd-coated Au colloid microstructure had not been known. This layered structure then completes a full series for the pd-Au bimetallic colloids. The preparation to form the novel Pd-coated Au colloid particles involves the reduction of a Pd salt in the presence of 200X Au colloid particles. These monodisperse Au colloid particles may be prepared following the method of Turkevich. The formation of the pd-coated Au microstructure was followed by opt!cal absorption spectros-
682
copy. Fig. 6 shows the spectra as a function of time of reflux. Layer formation was indicated by the gradual decrease of the peak at 5l8nm, characteristic of colloidal Au, and the increase of a peak at 254nrn which is characteristic of colloidal Pd. The spectrum at six hours of reflux still displays some slight features of the colloidal Au structure. Since the Pd coating is ca. 25X thick (see below), the combination of Pd and Au features in the spectrum is consistent with the observations made by Morriss and Collins (ref. 18) on the optical spectra of Ag-coated Au particles. The optical spectra of Au particles coated with a thin layer of Ag (-20X) still displayed some features of the Au colloid. However when thicker coatings were deposited (-70X), the optical spectra began to look more like pure Ag. The coated microstructure was also unambiguoisly characterized using high resolution AEM. Fig. 7 is a bright field image of the 250X coated particle. X-ray emission spectra are shown for three different areas on the particle. Spectra one and three are an analysis of the periphery of the particle on opposite sides, and spectrum two represents an analysis through the center of the particle. In the center of the particle the Au peak is much larger than the Pd peak. This is what one would expect since the 20X beam is going through the 25X Pd coating on each side of the particle while most of the signal arises from the 200X Au core. On the periphery the Pd/Au ratios change with the Pd signal being the dominant one. Due to mUltiple scattering events analysis on the periphery of the particle also detects some Au. Preparation of Supported Catalysts Supported Pd-Au catalysts were prepared by adsorption of the colloidal bimetallic particles onto two carbon supports. In order to determine the adsorption characteristics of these carbon supports, the deposition of monometallic 75X colloid Pd particles was examined first. Fig. 9 is an electron micrograph of a Pd/C catalyst prepared by adsorption of the colloidal Pd onto an American Cyanamid "controlled pore" carbon, which has a narrow pore size distribution (ref. 9). A uniform distribution of Pd on the support was obtained, and the monodisperse microstructure of the colloidal metal sol was retained in the catalyst. Fig. 10 shows Pd-coated Au particles deposited on the same carbon support • a uniform distribution of metal particles was obtained. .~gain,
683
The layered microstructure of the metal particles is still apparent in the supported catalyst, and may be seen as contrast differences due to differences in electron densities. The layered particle microstructure of this catalyst persists, even after treatment in H2 at 200°C. "Alloyed" Pd-Au/C catalysts were prepared similarly. Although some agglomeration of the metal particles was observed in one of the alloy Pd-Au/C catalysts, x-ray microanalysis in the analytical electron microscope indicated that the bimetallic particle microstructure of the "alloyed" colloids was retained in the supported catalysts. Thus, we conclude that the colloidal bimetallic particles are deposited by adsorption onto the support without change in microstructure. SUMMARY AND CONCLUSIONS Bimetallic monodisperse Pd-Au particles with three different microstructures have been reproducibly prepared, in pure form, as colloidal sols. The three microstructures constitute a complete set for the bimetallic Pd-Au system: 1. alloyed Pd-Au particles 2. layered, Au-coated Pd particles 3. layered, Pd-coated Au particles This is the first report of microstructures 1 and 3. The three microstructures have been unambiguously characterized using high resolution analytical electron microscopy. Supported Pd-Au/C catalysts have been prepared by adsorption of the colloidal Pd-Au particles onto carbon supports. Characterizat ion of the catalysts with AEM indicates the Pd-Au particles are deposited without change of microstructure onto the support. Thus, supported bimetallic catalysts have been prepared with a unique degree of control over metals microstructure. ACKNOWLEDGMENTS The authors wi th to acknowledge the expert technical assistance of E. T. Jones, J. E. Moore, D. L. Smith and I. R. Hartmann. Helpful discussions with U. Chowdhry, A. W. Sleight and C. E. Lyman are also recognized. We thank N. Budynkiewicz for accurately typing this manuscript. REFERENCES 1
J.H. Sinfelt, and. Applic~tions,
Bimetallic Catalysts: Discoveries, Concepts, John Wiley and Sons, New York, 1983, p.130.
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2
3
4 5
6 7 8 9 10 11 12 13
14
15 16 17 18
a) V. Ponec, Adv. Catal.,32 (1983) 149-214. b) H. Charcosset, Int. Chem. Engr.,23(2) (1983) 187-212. c) W.M.H. Sachtler, Catal. Rev.-Sci.Eng.,14(2) (1976) 193-210. a) G. Cocco, G. Carturan, S. Enzo and L. Schiffini, J. Catal.,85 (1984) 405-414. b) K. Aika, L.L. Ban, I. Okura, S. Namba and J. Turkevich, J. Res. Inst. Catalysis Hokkaido Univ.,24(1) (1976) 54-64. c) K.P. deJong, B.E. Bongenaar-Schlenter, G.R. Meima, R.C. Verkerk, M.J.J. Lammers and J.W. Geus, J. Catal.,81 (1983) 67-76. d) J. Turkevich, Actas Del 8° Simp Iberoamericano de Catalsis (1982) 27-60. R.T.K. Baker, E. B. Prestridge and R.L. Garten, J. Catal.,59 (1979) 293-302. a) J.E. Houston, C.H.F. Peden, D.S. Blair and D.W. Goodman, Surf. Sci., in press. b) C.H.F. Peden and D.W. Goodman, J. Catal., in press. c) P.F. Carcia and A. Suna, J. Appl. Phys.,54(4) (1983) 2000-2005. M. Kaminsky, K.J. Yoon, G.L. Geoffroy and M.A. vannice, J. Catal.,91 (1985) 338-351. J. Turkevich, G. Kim, Science,169 (1970) 873-879. J. Turkevich, P.C. Stevenson, J. Hillier, Disc. Faraday Soc.,ll (1951) 55-74. Kindly supplied by Dr. J.L. Schmitt; Chemical Research Division, American Cyanamid Co., Stamford, CT; U.S. Patent 3,978,000. C.E. Lyman, Am. Chem. Soc. Symp. Ser.,248 (1984) 311-333. M. Faraday, Phil. Trans.,147 (1857) 145. a) C.R. Shipley, Jr., U.S. Patent 3,011,920 (1961). b) N. Feldstein, U.S. Patent 4,323,594 (1982). a) H. Gerber, M. Horisberger and H. Bauer, Infec. Immun.,7(3) (1973) 487-492. b) H. Bauer, D.R. Farr, M. Horisberger, Arch. Microbiol.,97 (1974) 17-26. c) W.P. Faulk and G.M. Taylor, Immunochem,8 (1971) 1081-1083. a) o. Silman, L.A. Bumm, R. Callaghan, C.G. Blatchford and M. Kerker, J. Phys. Chem.,87 (1983) 1014-1023. b) J.A. Creighton, M.S. Alvarez, D.A. Weitz, S. Garoff and M.W. Kim, J. Phys. Chem.,87 (1983) 4793-4799. c) M. Mabuchi, T. Takenaka, Y. Fujiyoshi and N. Uyeda, Surf. Sci.,119 (1982) 150-158. E. Matijevic, Ann. Rev. Mater. Sci.,15 (1985) 483-516. R.S. Miner, Jr., S. Namba and J. Turkevich, in: T. Seiyama and K. Tanabe (Eds.) Proceedings of the VII International Congress on Catalysis, Elsevier, New York, 1981, pp. 160-169. G.C. Papavassiliou, J. Phys. F:Metal Phys.,6(4) (1976) LI03-105. R.H. Morriss and L.F. Collins, J. Chem , Phys.,41(11) (1964) 3357-3363.
685
DISCUSSION M.J. LEDOUX: a) What is the lowest size that you can obtain for the particles obtained with this method? b) Which are the parameters that limit the low size of the particles? J.T. SCHWARTZ: a) We have not determined the minimum particle size which can be prepared. However, we speGulate that the minimum size for the layered Pdcoated Au particles is 30-50 A, sinc~ Turkevitch has reported that the nucleus of colloidal gold particles is 30-40 A in size (see ref. 8). We estimate a similar size for the Au-coated Pd particles. The minimum size of the alloy particles would have to be determined experimentally. b) The principal factors controlling particle size are the metal involved and the ratio of citrate ions to metal ions in the reaction mixture. By changing the relative amounts of reactants, changes in the relative rates of the two independent processes of nucleation and growth of the metal particles are effected. G. JANNES : Your layered bimetallic catalysts are well characterized and seem to be stable enough to deserve some catalytic testing. Do you have some experiments able to shed some light on the geometric or electronic interpretation on the alloy effect in bimetallic catalysis?
J.T. SCHWARTZ : These catalysts have been tested in several different catalytic applications. However, the results are proprietary, so we are unable to discuss them at this time. J. KIWI: How do you get rid of the citrate rest when you prepare your gold or layered sols? How much do you loose of your colloids by your cleaning procedure? Do your sols vary in size with the time of citrate reduction? Have you done chemisorption on your particles mainly on Pd-Au alloys, to verify you are really dealing with an alloy? J.T. SCHWARTZ: The only case reported here where the sol was cleaned is the layered, Pd-coated Au sol. The gold sol used to make the layered particles was cleaned using a millipore ultrafiltration system to lower the ionic strength prior to layer formation. There is no significant loss of Au colloid during filtration as determined by UV-VIS spectroscopy. We have characterized the alloy microstructure with high resolution STEM. Chemisorption was not used to characterize the bimetallic Pd-Au particles. Colloidal Pd/C catalysts were studied using 02-02 titration to verify that the colloid's particle size was retained in the supported catalyst. J. OTTERSTEOT : Have you tried to deposit noble metals on other colloidal particles such as silica or alumina? In our laboratory we have deposited Pt on colloidal silica and alumina and found the resulting material to be efficient catalysts for the oxidation of CO. J.T. SCHWARTZ: Turkevitch reported the deposition of colloidal Pd onto colloidal alumina. We have not deposited these metal particles onto colloidal oxides. J.M. HERRMANN: 1/ What are the different oxides that you use as supports? 2/ Your colloidal bimetallic particles can be deposited by mere adsorption on a support. Do they provide stable enough supported bimetallic catalysts: for example, can they resist sonication and drying without leaving the support? J.T. SCHWARTZ: 1/ We have deposited colloidal metal particles onto MgO, Si02, and A1203. 2/ The adsorbed colloidal metal particles adhere strongly to the support. We have not observed any evidence of separation of the particles from the support
6~
after drying, grinding, treatment at 300°C in H2' or liquid phase catalytic testing. A.F.H. Wie1ers : I believe you have presented a very elegant method for the production of small alloy particles. As to the possible application of this material as supported catalysts, I have two questions: 1/ As you are depositing reduced metal particles onto a support, it appears that metal-support interactions are rather weak. Have you information on the thermostabi1ity of this catalyst? 2/ Can you give some information about the use of different oxidic support materials (Si02, A1203). Are metal-support interactions increased? 3/ Finally, I wonder whether the relatively large (10 nm) alloy particles can be homogeneously distributed over a porous support material? J.T. SCHWARTZ: 1/ These catalysts have been treated in HZ at temperatures of 200-300°C. We have not observed significant changes in mlcrostructure using electron microscopy. 2/ The colloidal bimetallic particles are adsorbed by supports with positively charged sites. We have adsorbed colloidal metal particles onto MgO, functiona1ized Si02 and A1203' We have not measured the relative strengths of metal-support interactions in this work. 3/ The homogeneity of particle distribution depends on the nature of the support. We have observed some particle agglomeration in a few cases. However, a relatively uniform distribution is generally obtained. M.A. MARTIN-LUENGO :When you speak of precise controlled microstructures, do you mean that your particles have all the same colloid size or also that they have well defined crystallographic planes? J.T. SCHWARTZ: The components of microstructure which have been controlled using the colloid technique are - particle size - elemental distribution within each particle. We have not characterized the crystallography of these bimetallic particles. L. GUCZI : Have you studied the conditions (temperature, atmosphere, etc .•. ) under which segregation of one metal to the particle surface starts? J. T. SCHWARTZ: No, we have not attempted to alter the elemental distribution within the bimetallic Pd-Au particles. C.J.G. van der Grift: 1/ Can this preparation method also be applied to the production of other bimetallic catalysts, especially containing less noble metals? 2/ Can you comment on the remarkable stability of the sphere distribution over the support? Since probably the contact area between bimetallic sphere and support is small and furthermore the interaction between support and alloy particles doesn't seem extensive because otherwise the support would be wetted by the metallic phase with a change in microstructure of the spheres. J.T. SCHWARTZ: 1/ Yes. We have used the citrate reduction method to prepare other bimetallic particles, including Pd-Ag. We are not aware of non-noble metal particles prepared by this technique. 2/ We believe the attraction between the particle and the support can be quite strong, suggested by the rapid adsorption of the colloidal Pd-Au particles by some supports. We believe this attraction anchors the colloidal particles to the adsorption site on the support. The anchoring can be quite firm, serving as the basis for numerous applications of colloidal metal particles for labelling charged sites of surfaces, biological cells, etc. We have not studied the particle-support interface. However, we believe the particle-support
~7
contact area is considerably larger than expected from a flat plane-sphere model, since the supports used are high surface area, with rough porous surfaces.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV
689
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
CONTROLLED SURFACE REACTIONS FOR THE PREPARATION OF DIFFERENT TYPES OF ALUMINA SUPPORTED Sn-pt CATALYSTS E.KERN-TALAS, M.HEGEoOS, S.GOBOLOS, P.SZEDLACSEK and J.MARGITFALVI Central Research Institute for Chemistry of the Hungarian Academy of Sciences, H-1525 Budapest POB 17, Hungary ABSTRACT Different types of SnPt/A1 203 catalysts were prepared in wich the oxidation state (ionic or metallic) and the environment (platinum or support) of the corresponding forms of tin were controlled by using different types of Controlled Surface Reactions (CSRsl. These catalysts were studied in n-hexane conversion. By using these catalysts additional new evidences were obtained with respect to (i) the mode and ways of control the selectivity of Pt/A1203 catalysts, (iil the alterations of the reaction routes in hydrocarbon conversions and (iiil the role of tin in reforming catalysts. INTRODUCTION Supported Sn,Pt catalysts have been widely used in naphta reforming or aromatization (refs. 2-9) and dehydrogenation of paraffins (refs. 10, 11)). There are different approaches for the preparation of these catalysts. The coimpregnation (ref.2) and two-step impregnation (refs. 6, 12) with H and SnC1 are the 2PtC16 2 most general methods. The use of different tin and platinum containing ionic bimetallic clusters (refs. 13-15) and the application of anchoring techniques (refs. 16, 17) were also studied. The simpliest way for tin introduction is the impregnation of a Pt/A1 203 catalysts with SnC1 (ref. 9). 4 Despite of the tremendous amount of works done in the field discussed there are many questions to be answered. The general questions are: (i) in what forms (metallic or ionic) does tin exist in a working catalyst, (ii) what is the role of the different forms of tin in these catalysts? These questions can be answered if all surface interactions taking place during the preparation of these catalysts could be controlled. It should be mentioned that with the exclusion of the anchoring technique (refs. 16, 161 mentioned earlier none of the methods reviewed can guarantee the exclusive formation of a given tin containing surface species. We believe that the lack of control of surface interactions during the preparation of these catalysts is responsible for all of the disagreements and disputes in connection with the form, state and role of tin in these catalysts. In this contribution, as a continuation of our previous work (ref. 1),we are presenting additional new experimental data obtained in preparation of different
690
types of alumina supported tin platinum catalysts. Our studies are based on a new approach using different types of Controlled Surface Reactions (CSRs) (ref. 1). In this approach the chemical nature of a Pt/A1 catalyst itself is used 203 to control the exclusive formation of a given tin containing surface species. SURFACE CHEMISTRY Preparation of catalysts with direct tin-platinum interaction. Hydrogen chemisorbed on the platinum sites was used as a driving froce to guarantee the control for the formation of bimetallic surface species with direct Sn-pt interaction. These catalysts were prepared by using a new surface reaction between different tin alkyls and the hydrogen chemisorbed on the platinum (reaction (1)). In a subsequent reaction step the primary formed surface complex (I) was decomposed in a hydrogen atmosphere. For Sn(C2HS)4 the surface chemistry can be written as follows (ref. 1): (1) ptHa + Sn(C2HS)4 ----. Pt-Sn(C 2H S)4_x (I) + x C2H6 (I)
(2)
Pt-Sn
Preparation of catalysts with tin alumina interaction. These catalysts were prepared by anchoring SnC1 2 or SnC1 4 into a lithium modified alumina. The primary formed surface complexes were decomposed in a hydrogen atmosphere at high temperatures. Upon using SnC1 4 as tin precursor the following surface reaction were suggested: ~-oH + C4H g l i - §I-oli (II) + C4H10 (3) n -oLi
+ (III)
SnC1 4 -
( -0-)nSnC14_n
-H
2
( -0-) n Sn
(III)
(IV)
+
n LiCl
+ (4-n) HCl
(4) (S)
EXPERIMENTAL PART In this work two types of base Pt/A1
catalysts (Pt1 and Pt2) with O.S w% 203 Pt content prepared by impregnation of Ketjen CK-300 type alumina with H2PtC16 in HCl solution were used. Characteristic properties of these catalysts are given in Table 1. The platinum black used to model surface reaction (1) was a low surface area (0.5 m2 /g) catalyst. Reaction (1) and (3) were monitored by volumetry and parallel GC analysis of the reaction products. Reaction (2) was studied by Temperature Programmed Reaction (TPR) technique (heating rate SOC/min). In this study three different types of alumina supported tin-platinum catalysts were prepared: Type A: catalysts with direct tin platinum interaction, Type B: catalysts with tin-alumina interaction,
691
Table 1 Designations and properties of catalysts used Precursor Modifier Sn Li Catalysts or catalysts w% precursor w% b Pt1 b 0 0 Pt2 0 0 Ai Pt2 Sn(C 2H s)4 0.035 0 A2 Pt2 Sn(C2H sl 4 0.15 0 A3 Ptz Sn(C2H sl 4 0.21 0 Li1 PtZC C4HgLi 0 1.15 Pt2d Li2 C4H gLi 0 81 Li1 SnC1 4 0.12 1.29 8Z Li1 SnC1 2 O.1Z 1.04 83 LiZ SnC1 4 0.10 1.05 84 Li2 SnC1 2 0.13 1.10 85 0.16 Sn/A12 D3 0.96 l(CH~3Pt8r]4 A2c LiA2 C4H gLi 0.15 C1 LiAZ 0.Z5 SnC1 4 1.15
Cl
w%
H/Pt treatment a Dz-H Z HZ
1.33 1.06 1.04 1.10 1. 05 1. DB
0.30 0.41 0.29 0.35 0.12
0.30 0.42 0.45 0.41 0.19
1.53 1.44
0.13 0.15
0.21 0.15
0.38 0.15 0.11
0.39 0.42 0.24
0.05
0.35
-
a) oxygen treatment at 400°C. hydrogen treatment at SSOO; b) the initial H/Pt ratio of the bas.e Pt 1 and Pt2 catalvsts are, 0.10 and 0.S7 , respectively; cl dehydroxylation at SooOC; d) dehydroxylation at 1SooC. Type C: catalysts with tin platinum and tin-alumina interactions. Catalysts type C were obtained by combination of the methods used to prepare catalysts type A and B. A catalyst type A was prepared first followed by procedures used for the preparation of catalysts type 8. The characteristic features and designation of all catalysts prepared are given in Table 1. Catalysts prepared were tested in n-hexane conversion at SZOOC under atmospheric pressure using the modified slug pulse technique described elsewere (ref. 18) • RESULTS Surface reactions involving hydrogen adsorbed on metals Preparation of catalysts type A (i) Study of the reactivity of hydrogen chemisorbed on platinum sites. The reactivity of hydrogen adsorbed on platinum black was studied at SOOC using SnLC ZHS)4' The formation of ethane and ethylene is shown in Figure 1a. There are two different parts on the kinetic curves. At the very beginning of the reaction the fonmation of ethane is very ~ast and it is· the only reaction product. As the reaction proceeds the amount of adsorbed hydrogen for reaction (1) is consumed and the catalytic decomposition of Sn(CzH S)4 prevails with parallel formation of ethane and ethylene. However, the rate of this side reaction is much smaller than that of the reaction (1l. In our previous work using Pt1 type base catalyst (ref. 1) reaction (1) was
692
(I) 10-6 mol
.,x-
.--
.-
'C_
x-
e:
C2H e
.(
8
, I
"
e "I
//A
4
,
"I
C2H
I
4
2
50
100
150
200 t,mln
300 T/OC
Fig.1. Reactivity of hydrogen adsorbed on Pt black. A. Study of reaction (1); t=500C, Pt=2.4 g,lSn(C zHs)4]O = 2.1 x 10- z mol.dm- 3, solvent: benzene; B. Decomposition of surface complex (I) in Hz by TPR TechniqueJ (1) amount of ethane in arbitrary units. 10-e mol
15 o
o
eo
t,mln
Fig.2. The influence of the initial concentration of Sn(CZHS)4 on the rate and selectivity of reaction (i); t ~ 50°C, 0, A, 0 - C?H 6, ., " • - CZH 4. Designation and concentrations, in mol.dm-3,1o- 4: 6,. - 2.62 (Ai), OJ. 12.9 (A2J, 0,. - 38.8 (A3). not studied in such detail The fonnation of ethane and ethylene was monitored and the appearence of ethylene was used to terminate reaction (1). Upon using Pt2 type base catalyst special care was taken to follow both steps (reactions (1) and (2)) used in the preparat~on. In these experiments the concentration of the Sn(C2HS)4 has been varied. The kinetic curves of ethane and ethylene fonnation are given Figure 2. Upon increasing the concentration of Sn(CZHS)4 the rate of reaction (1) and the total amount of ethane formed has
693
been increased. In sample Ai the reaction is terminated duo to the total consumtion of Sn(C2HS)4 introduced. These kinetic curves strongly resemble that of obtained upon studying platinum black (see Figure 1a). Parallel to the formation of ethane ethylene was also detected in all samples, however, pronounced ethylene formation was observed only at high concentration of Sn(C2HS)4 (sample A3Pt). The formation of ethylene is the indication for (i) the total consumption of ador (iii) the sorbed hydrogen, (ii) the catalytic decomposition of Sn(C 2HS)4 loss of control in surface reaction (1) due to an unknown surface reaction. A. ( i i ) Decomposition of the primary formed surface complex (Il, (TPR studies) • a) Unsupported platinum. The decompostion of the primary formed surface complex (I) was studied between SO and 4000C in a hydrogen atmosphere. In the presence of hydrogen only ethane formation was observed. The TPR curve of ethane formation is shown in Figure 1b. The small peak around 2300C was attributed to the decomposition of the solvent adsorbed or chemisorbed on the platinum, while the large complex peak at lower temperatures was attributed to the decomposition of surface complex CIl. b) Pt1 type base catalyst and pure alumina. Typical TPR results, i.e. the formation of ethane and ethylene, observed upon preparing catalysts type A with different Sn/Pt ratios are given in Figure 3. For comparison results obtained on an AIZ03 sample containing adsorbed Sn (C2HS)4 are also presented. No tin was added into sample 1 but all procedures used for tin introduction were simulated. The intensitiy of the high temperature peak around 2700C was independent of the tin content provided the Sn/Pt atomis ratio was bellow O. S (ref. 1). The integral intensity of the low temperature complex peak between 7o-ZoooC strongly depended on the Sn/Pt ratio. From this result the high temperature peak in sam; pIes 1 and Z was attributed to the hydrogenolysis of the adsorbed solvent used in reaction (1). The low temperature peak was assignated to the decomposition of surface complex (I). In sample 4 parallel to the formation of ethane ethylene was also formed. This fact and the high Sn/pt atomic ratio of this sample as well as the high intensitiy of the ethane and ethylene peaks around 2800C may indicate on the tin-alumina interact ion. cJ PtZ type base catalyst. Decomposition of surface species formed in reaction (1) are given in Figure 4. In this series using sample A2 (see Figure 2.) the influence of the reaction time on the amount and the types ·of the primary surface complexes formed under condition of reaction (1) was studied. Upon increasing the reaction time the integral intensities of both ethane peaks, i.e. the low temperature one corresponding to the Sn-pt interaction and the high temr perature one characteristic for the tin-alumina interaction, increased. The appearance of the high temperature peak is characteristic for other highly dis-
694
150
Fig. 3. The type of TP~ curves obtained on Pt1 type catalysts, 1 - blank experiment (no tin added),2 - sample with Sn-pt interaction (Sn/pt = 0.31), 3 - Sn(C2H s ) 4 adsorbed on pure A1 203, 4 - sample with Sn-A1203 interaction (Sn/Pt = 1.5). Fig. 4. The influence of the time of reaction (1) on the formation and decomposi tion of primary surface complexes (Sample A2) (see Figure 2). Results obtained by TPR technique. persed platinum catalysts (ref. 19). Further studies including labelling technique will be required to elucidate this phenomenon. Surface modification involving reactions of the support surface hydroxy Is Preparation of catalysts type B (i) Preparation from Pt2 type base catalyst. It is known that the rate of reaction (3) and the concentration of surface species (II) are strongly influenced by the temperature of dehydroxylation of the support. The amount of C4H 10 fonned in reaction (3) while preparing samples Li2 was twice as much than that of formed for sample Li1. Thus, the concentration of surface species (II) could be changed creating favourable or unfavourable conditions to anchore tin chlorides. In reaction (4) the reactivity of tin precursors increased in the following order: Sn(C2HS)2CI2 « SnCl4 < SnCI2" It was also observed that the HCI formed in reaction (5) partly adsorbed into the alumina with the formation of new types
695
of acidic sites. Catalysts 81 and 82 were prepared from sample Li1 using SnC1 and SnC1 4 2, respectively. Sample Li2 was used to prepare catalysts 83 and 84 by anchoring SnC1 4 or SnC1 2. (iiJ Preparation using tin modified alumina (catalyst 85). To prepare this catalyst the pure alumina was dehydroxylated at 1500C followed by lithiation and anchoring of SnC1 4. This tin modified alumina was impregnated with [(CH3)3Pt8r]4 as described elsewhere (ref. 20J. Preparation of catalysts Type C Catalysts type C were prepared from catalyst A2 after its dehydroxylation at 0C
500
and subsequent lithiation. Experimental conditions were identical to those
used for the preparation of catalyst Li1. The lithiated catalyst (LiA2) was reacted with SnC1 4 to prepare catalyst C1 containing both ionic and metallic forms of tin, in 0.10 and 0.15 w%., respectively. Catalytic reactions In our previous study (ref. 1) it was shown that upon tin introduction the initial rate of n-hexane conversion decreased on all types of catalysts, and different forms of tin resulted in different alterations in the selectivity of n-hexane conversion. Catalysts prepared in this work have essentially the same properties. As a general observation the formation of 2-methylpentane, 3-methylpentane, methylcyclopentane, and toluene were almost entirely suppressed on catalysts Li1, Li2, LiA2, 81-85 and C1, i.e, on catalysts containing surface species (11)-(IV). The changes in the surface acidity of these catalysts can be found elsewhere (ref. 19). In this study more attention is paid to analyze the selectivity conversion dependences. In this respect the selectivity of the formation of 1-hexane, benzene, and light hydrocarbon (C will be discussed. 1-C5) The conversion~selectivity dependence for benzene (8), light hydrocarbons
6)
(Cr) and 1-hexene (C are given in Figure 5. For the clearity results obtained on one A type and three 8 type catalysts (A2 and 81, 82, 83) are presented and discussed. All of the A type catalysts strongly resembled the base Pt2. The only difference is a slight increase in the selectivity of light hydrocarbons observed at high conversions. On catalyst A2 constant Scr values were obtained up to 20 % conversion. It is an indication that light hydrocarbons are primary products on this catalyst. On all catalysts type A the formation of isohexanes, methylcyclopentane and toluene were only slightly altered.
696
_-A"
40
40
20
20
2
4
8
__+-
8
4
2
8
8
sc= . 6
.....
'- ~
20
20
... n-a.. ~
+...... +
.... ,. + °b"-........ - - __ . . . -... -.. -~ -_H-_
"'::ta~
..........
2
4
6
- *- :::-a--EI- __._
....... 121-. -0 __
2
8
4
6
- +-
8
0 _ - - 0- -
.... .... -
........ ...... ---
20
---+-
.... ,-10
10
'-
-...02$~~--
121-
......
"'--*-_••-.- - - - _.- ::= *.=>*-.-=:2
4
6
8
2
4
6
8
Convere. on 'lb
Fig.5. Selectivity-conversion dependences. Reaction of n-hexane at 5200 catalysts: * - pt2, t:, - Li1, A- Li2, 0 - A2, • - LiA2, 0 - 81,. - 82,111- 83, + -C1. Catalyst Li1, Li2, 81, 82 and 83 have an increased selectivity for benzene formation. The higher the selectivity of benzene formation the lower is the selectivity of the primary formed 1-hexene. All these facts and the strong suppression of the formation of isohexanes and methylcyclopentane can be attributed to the changes in the route of benzene formation. On these catalysts the
697
consecutive dehydrogenation mechanism via formation of hexenes and hexadienes prevails. The low benzene selectivity of the 82 catalyst observed at high extent of conversion and the high Scrvalues observed on this catalyst can be attributed to the hydrocracking of the precursor of benzene on acidic sites formed from surface species (III), This high hydrocracking activity could be strongly reduced anchoring SnC1 4 to sample Li2 containing higher amount of -O-Li surface species (see catalyst 83). The modification of the A2 catalyst by lithium and subsequent tin introduction (catalysts LiA2 and C1) strongly reduced the selectivity of benzene formation with parallel increase of the selectivity of 1-hexene and other unsaturated hydrocarbons (e.g. hexadienes). On these catalysts the formation of isohexanes anti methylcyclopentane is completely blocked by surface species (II) or (IV) and the readsorption of the primary and secondary formed unsaturated hydrocarbons is hindered due to the tin-platinum interaction. The selectivity~conversion dependences of light hydrocarbons on catalysts Li1. Li2. B1-B3 and C1 have different character than that obtained on the Pt2 base and A2 catalysts. The increase of Scr values with conversion observed on these catalysts may indicate that different intermediates formed from n-hexane are involved in further hydrogenolysis or hydrocracking reactions. CONCLUSIONS Our experimental data have shown that by using different types of surface reactions the fonn of tin (ionic or metallic) and its environment (platinum or support) can be controlled in Sn-Pt/A1 203 catalysts. Based on results obtained in n-hexane conversion it seems that the presence of ionic type surface species with tin-alumina interaction is crutial in the perfonnance of these bimetallic catalysts. In these catalysts the formation of benzene takes place via unsaturated hydrocarbons via consecutive dehydrogenation. Tin in metallic state, i.e. the fonnation of surface species with tin-platinum interaction (catalysts type A), have only secondary importance as only minor selectivity changes were observed on these catalysts. It is suggested that on bifunctional catalysts (e.g. the base Pt and catalysts type A) the consecutive dehydrogenation have less contribution in the fonnation of benzene than the alternative routes via fonnation of isohexanes or methylcyclopentane. ACKNOWLEDGEMENTS The authors thank or.L.Koltai for Sn and Li detennination by AAS and Mrs. H.MilliAn and I.Turi for technical assistance.
698
REFERENCES
2 3 4 5 6
7 8 9 10 11 12 13 14 15 16 17 18 19 20
J.Margitfalvi, M.Heged~s, S.G~bolos, E.Kern-Talas, P.Szedlacsek, S.Szab6 and F.Nagy, in Proc. 8th Int.Congress on Catalysis, Berlin (West), July 2-6, 1984, Vol. IV. Verlag Chemie, Weinheim, 1984,pp. 903-914. French Patent 2,031,894 (1968), to CFR U.S. Patent 3,577,473 (1971), to Chevron Ger.offen 2,113,520 and 2,222,250 (1972), to UOP V.N.Selesnev, Y.V.Fornichev and M.Ye.Levintver, Neftekhimiya, 14 (1974) 205-208. R.Bacaud, P.Bussiere, F.Figueras and J.P.Mathieu, in B.oelman, P.A.Jacobs and G.Poncelet (Eds.), Proc.lnt.Symp. on Scientific Bases for the Preparation of Heterogeneous Catalysts, Brussels, October 14-17, 1975, Elsevier, Amsterdam, 1976, pp. 509-523. B.Davies, G.Westfall, J.Watkins and J.Pezzanite, J.Catal., 42 (1976) 247-256. F.M.Dautzenberg, J.N.Helle, P.Biloen and W.M.H.Sachtler, J.Catal., 35 (1980] 119-128 R.Burch and L.C.Garla, J. Catal., 71 (1981) 360-372. Z.A.Sadyhova, N.V.Nekrasov, V.R.Gurevich and S.L.Kipenman,Kinet.i Katal., 22 (1981) 396-402]. S.B.Kogan, N.M.Podklyetnova, O.M.Oranskaya, I.V.Semenskaya and N.R.Bursian, Kinet. i Katal., 22 (1981] 663-667. G.T.Baronetti, S.R.Miguel, O.A.Scelza, M.A.Fritzler and A.A.Castro, Appl. Catal., 19 (1985) 77-85. V.Keirn, H.Leuchs, B.Engler, Forschungsbereichte den Landes Nordhein Westfalen, BRD 2838 (1979) 1-39. V.I.Zaikovskii, V.I.Kovalchuk, Yu.A.Ryndin, React.Kinet.Catal. Lett., 14 (1980] 99,,103. B.N.Kuznetsov, V.K.Duplyakin, V.I.Kovalchuk, Yu.I.Ryndin and A.S.Bely, Kinet. i Katal., 22 (1981) 1484-1489. B.N.Kuznetsov, in Yu.I.Yenmakov (Ed.), Supported Metallic Catalysts for Hydrocarbon Reactions (New Approaches to their Preparation and study (in Russian] Institute of Catalysis, Novosibirsk, 1978, Vol.2. pp. 43-66. V.K.Duplyakin in Yu.I.Yenmakov (Ed.), Supported Metallic Catalysts for Hydrocarbon Reactions (New Approaches to their Preparation and Study (in Russsian], Institute of Catalysis, Novosibirsk, 1978, Vol. 1, pp. 74-89. J.Margitfalvi, P.Szedlacsek, M.HegedOs and F.Nagy, Appl. Catal., 15 (1985] 69-78. J.Margitfalvi, E.Kern-Talas and G.Resofszki, in preparation J.Margitfalvi, S.G8bOlos, M.HegedUs and F.Nagy, React.Kinet.Catal. Lett. 21 (1982) 541-548.
699
DISCUSSION G.J. ANTOS: Have you investigated the stability of your Sn-Pt sites to oxidation? Do you believe Sn migrates off the Pt? How far? J.L. MARGITFALVI: 1/ The Sn-pt sites formed in reactions (1) and (2) can be completely altered by applying oxygen treatment. The role of oxygen treatment has been discussed in our previous work (ref. 1). Our primary XPS data and the selectivity changes obtained in n-hexane conversion strongly support the idea of B. Davis (J. Catal. 89, 371 (1984)), i.e. after oxygen treatment and final reduction, a tin-aluminate-like phase can be formed without tin-metal-platinum interaction. 2/ We may demonstrate that Sn can migrate on the platinum at high temperatures (above 500-520°C) in hydrogen atmosphere. This process is more pronounced on small Pt particles than on larger Pt crystallites. However, basic additives or the presence of ionic forms of tin strongly influences this process. With respect to the question how far, unfortunately I cannot answer, we have only speculation. J.P. BOURNONVILLE : 1/ Could you be more explicit about the size effect on the nature of the chemical reaction? Have you worked with particles less than 1 nm in size? 2/ Have you performed any successive impregnation of tin in order to reach a saturation value, for instance? J.L. MARGITFALVI : 1/ We have observed selectivity changes both in reaction (1) and reaction (2). Reaction ~1) On PtlA12 3 with high dispersion, i.e. with particle size around 1.2-1.5 nm (catalyst Pt2) the formation of ethylene was more pronounced than that of catalyst with low dispersion (Ptl). On catalyst Ptl methane formation was not observed while methane formation was detected on catalyst Pt2. Reaction (2) The total ethane selectivity on catalyst Pt2 was lower than that on catalyst Ptl due to the formation of methane and ethylene. In addition the new TPR peak around 250°C observed on catalysts Pt2 (see Fig. 4) is a further indication for the alteration in the surface chemistry. We have additional experimental evidences that on Pt/A1203 catalysts, with high dispersion parallel to reaction (2) two side reaction can take place: (i) carbon-carbon bond scission and (ii) hydrogen abstraction from "-CH2-" or "-CH3;:" moieties. 2/ The driving force in our system is the surface reaction between adsorbed hydrogen and tin alkyls (reaction (1)). The introduction of tin into the platinum strongly reduces the amount of hydrogen adsorbed on Pt, and consequently it strongly hinders reaction (1) too. As a result, upon successive tin introduction only impregnation of the support with tin alkyls will take place resulting in an increase of the total tin content, but tnis form of tin will be introduced into the support and not onto the metal. L. GUCZI : During the TPD measurement you observed a second peak. result of the reaction between tetra-alkyl and A1203 OH groups?
Is it not a
J.L. MARGITFALVI : The second peak around 250°C or above in the TPR experiment can be attributed to different surface reactions as follows: (i) Hydrogenolysis or hydrocracking of the trace amount of adsorbed solvents (benzene, hexane) used in reaction (1) (see curve 1, Fig. 3). (ii) Decomposition of SnR4 adsorbed on pure A1203 (see curve 3, Fig. 3). (iii) Decomposition of SnR4 adsorbed on Pt/A1203 at high Sn/Pt ratios (see curve 4, Fig. 3). (iv) Rehydrogenation and hydrogenolysis of "dehydrogenated" C2Hx groups in surface complex (I) suggested on Pt/A1203 catalyst (Pt2) with high dispersion (see Fig. 4).
700
By calculation of the total balance of tin. ethane and ethylene in reactions (1) and (2), the extent or the contribution of the reaction between tin alkyl and A1203 can be evaluated. Further information could be obtained by comparing TPR data obtained in hydrogen and nitrogen atmospheres.
B. Delrnon, P. Grange. P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV 1987 Elsevier Science Publishers B.V.• Amsterdam - Printed in The Netherlands
701
PREPARATION AND PROPERTIES OF ANCHORED Pt-Mo!SiO Z BIMETALLIC CATALYST YANG YASHU, GUO XIEXIAN, LI HUIMIN, DENG MAICUN and LIN ZHIYIN Dalian Institute of Chemical Physics, Chinese Academy of Sciences, Palian, P.P.China
SUl>:MARY bimetallic catalysts of 3wt%Pt-1.1wt%~o prepared by ana impregnation technioues are examined and compared. ~'PP, ESP and TEM measurements' and n-pentane conversion are used to characterize the surface properties and the c at a.Iy td.cc be-e hevior of different pre~aration respectively. The results clearly domonstrate the significant effect of preparation method. With anchoring technique, uniform surf8ce snecies interpreted as Ft",!'"oZ+C-Si==)?, is for-med whi ch is e venl y cli.stributed in the surf1)ce layer and responsible for the high hydrogenalysis actiVity of the resulted catalyst. For the impregnated preparation, the formation of such kind of stoichiometric surface complexes is almost completely excluded.'ilJae separate part.i.c Les of platinum, molybdenum oxidies and Pt-Mo alloys are the sur face species which indic ate a strong synergy effect in n-pentane hydrogenolysis. Pt-Mo!SiO~
anchorin~
INTRODUCTION The study of Pt-Mo bimetallic catalysts becomes increasingly interesting to many authors in the last decade. Yermakov et al. (ref. 1) have successfully prepared the anchored Pt-Mo!Si0 2 catalysts and pointed out that the formation of chemical bonds between Pt atoms and Mo Z+ ions is presumably the reason of the great increase in catalytic activity of ethane hydrogenolysis. Gandhi et al.(ref. 2) indicated that the catalytic properties of pt-Mo/r~1203 prepared via co-impregnation method were very close to that of Rh!r-A1 20 or Ir/r-AI 20 catalysts in the reduction of NO by 3 3 F 2• For Pt-Mo!Y-zeolite and Pt-Mo!Si0 2 catalysts, as reported respectively by Tri et al.(ref. 3) and Leclercq et al.(ref. 4), synergetic effect between platinum and molybdenum occurs which facilitates the actiVity in butane hydrogcnolysis. In this paper it is attempted to study the difference in sur~ facfl. property and catalytic behavior between Pt-Mo!Si0 2 catalysts prepared by anchoring and conventional impregnation methods with the aid of physical measurements such as TPR, ESP, TEM and test reaction of pentane conversion for catalytic ch ar-ac t.eml.zsatd on.•
702
EXPERIMENTAL AND RESULTS Prenaration of catalysts Z.96wt%Pt-l.08wt%Mo/SiOZ(PM-1) and o.68wt%MO/SiOz(M-1) anchored catalysts were the samples obtained after the decomposition of ~-allyl Pt-Mo and single ~-allyl Mo-surface complexes anchored on the surface of SiO Z as described in (ref. 5). 3wt%Pt-l.06wt%Mo/ Si0 2(PM-2) and 1.06wt%Mo/Si0 2(M-2) catalysts were prepared by stepwise or single impregnation of SiO Z with aqueous solution of (NH4)6M070Z4' drying at lZOoC, calcining at 550 oC(stop here for M-2), the resulting material was SUbsequently impregnated with aqueous solution of HzPtC1 6, dried and finally reduced in flowing HZ' TPR }1easurements The experiments were performed in a~flow system using a 10% HZ in Ar gas mixture of flow rate of 25 ml/min and sample heating rate of 8 °C/min for reduction. The profiles observed on oxidized samples M-1(Fig. 1, a) and M-Z(Fig. 1, b) reveal obvious differen-
b
500 700 °c Fig. 1. fPR of Mo/SiO Z reoxidized at (a) 300 0e for M-1i (b) 500°C for M-Z. ces in reduction property between anchored and impregnated samples. Fig. 1a consists of two symmetric peaks of approximately equal areas at 590 and 700 oe. Fig. 1b exhibites three distorted peaks a at 557, 715 and 800 0e respectively. These results are in close agreement with that of Yermakov (ref. 1) and Le~lercq (ref. Z) with the implication that for oxidized M-1, Mo ions are uniformly anchored on Si0 2 with two doubly bonded oxygen atoms attached to each Mo 6+ ion, while for OXidized M-Z, the molybdenum oxide is not evenly deposited on the surface. Fig. 2 is the TPR spectrum of oxidized Pt/Si0 2 catalyst which displays two reduction peaks at 130 and 445°C. The TPR curves for oxidized PM-1 and PM-2 are
703
shown in Fig. 3a and Fig. 3b. Comparing to Fig. 2, Fig. 3a has a low temperature peak of much greater area which reflects the increasing rate of reduction of platinum oxide in the presence of anchored molybdenum oxide. It is interesting arid easy to reason that the
~I~ 200
P::
600 "c
400
Fig. 2. TPR of Pt6Si02 reoxidized at 300 C.
a
!1z;;J 200
400
600
800 "c
200 400 600 Fig O 3. TPR of Pt-MO/SiO Z reoxidized at (a) 300 0C for PM-1i (b) 500 G for PM"'!2. reduction of Mo 6+ (to MO Z+ ) is greatly aided by platinum due to its combination with Mo-oxide surface moieties and the reduction spectrum for anchored molybdenum oxide becomes broadened with the two reduction peaks of Fig. 1a collapsed into a single band shifted to lower temperature region. It may be suggested that three main kinds of metal-oxygen bond, such as Pt-O-Mo, Mo=O and Mo-OPt-O-Mo, are present in the complex of oxidized PM-l. For the oXidized sample PM-2 (Fig. 3b) there remains a high temperature branch for platinum similar to Fig. 2, but this branch is over~ lapped with the complicated reduction spectrum of molybdenum oxide (see Fig. 3b). It seems obvious that individual but intermingled Pt and Mo-oxide particles exist on the surface of PM-2 such that the TPR pattern is mainly determined by the irregular sizes of individual particles of various composition rather than any characteristic surface species with definite reduction peaks.
ESR spectra The sample was pretreated in a quartz vassel at different conditions before ESR measurement. After oxidation at 450 oC, the sample Was outgassed at 10-5 torr, 450°C, and cooled to room temperature, then a few injections of anthracene in heptane were introduced to the entire wetness of the catalyst sample. ESR spec-
704
tra were recorded at room temperature with a JES-FEZXG spectrometer. The variations of the time-dependent in signal intensity of ~05+ peak (at g=1.93) Wfre shown in Fig. 4a~d. It is seen
c
~ 1.936
er... 1. 903
0C preoxidized catalysts Fig. 4. E.SR spectra of M0 5+ on 450 wetness of anthracene ( a) on JVI-1, _,0 5+ signal (1) after being staid for 5 min, height of 11-' (2) 1 day, = i. " 2 days, 0) If " (b) on M-2, (1) after being staid for 20 min, Height of 1-',05+ signal (2) 1 day, If " 2 l1ays, 0) (c) on PM-1, " " Height of Mo 5+ signal ~1)2) after being staid for 517 min, hrs, If " 0) 31 hrs, If " (4) 55 hrs, If " (d) on PM-2, (1) after being staid for 20 min, Height of M0 5+ signal (2) 1 day, If If 2 days, 0) If If
after nil
19;§ , 9.3 11 41.5; 5.5 40 72 98.5; 10.5 27.5 30.5.
that the observed rate of signal intensity increasing with time is faster on M-2 (Fig. 4b) comparing to ]v!-1 (Fig. 4a). However, the trend of Mo 5+ signal changes was adversely affected in the pre-
705
sence of platinum, the intensity on the anchored catalyst PM-1 (:Cig. 4chather than the impregnated sample PI'!-Z (Fig. 4d) was increasing far more rapidly with time. It is likely that the rapid increment of M0 5+ observed on PM-1 is closely related to the facilitation of the slow process of reduction of the ~06+ ions by anthracene due to the connection between platinum and molybdenum ions mentioned above which is also in conformity with the above TPR measurements. Electron micrograph and Diffraction pattern A high resolution electron microscope of Model HITACHI H-600 with point resolution limit of 2.4 ~ waS used. The reduced sample machanically grinded in air and treated with 30% HF solution, and then the extracted powder was suspended in ethenol by ultrasonic dispersion technique. The suspension was used to prepare the standard carbon replica mounted on a copper screen. The electron morphology of sample PM-l was shown in Fig. 5a which displays a high uniformity of particles with a size distribution of fairly narrow range of 20~0 ~ (Fig. 5b). The electron diffraction pat-
% 50 30 10 . 30 Fig. 5a. Trasmission electron micrograph of PM-1. (lX150jOOO)
50
Fig. 5b. Particle size distribu tion on PH-l.
terns obtained by selected area method were depicted for var.ious catalysts in Fig. 6. The Fig. 6a (Pt/Si0 2) and Fig. 6b (H00 3/SiO Z) samples exhibit respectively the normal diffraction of crystal lattice of platinum and mo Lybderiaa The diffraction pattern of Fig. 6c for the impregnated catalyst PM-2 is mainly composed of 1 the diffraction rings of Fig. 6a and 6b. This indicates that the major part of the supported phase in PM-2 catalyst is crystallographically a mixture of intermingled individual particles of platinum and various molybdenum oxides. Moreover, as depicted in Fig. 6c
2,
the diffraction pattern observed in the same surface of
706
a
b
d
Fig. G. Electron diffraction0Cpatterns, (a) Pt/Si0 2; (b) ~o03~§iOZ; (c 1-c Z) PM-Z,reduced at 350 in HZi (c PM-Z,reauced at 4~) C 3) in HZ; (d) PM-1, reduced at GOOoe in H2• PM-Z but of selected ar'e a different from those of Fig. GC 1 indicates that the presence of local region in the surface where only Mo0 or M00 could be detected without coexistence of any 3 3/SiO Z appreciable amount of platinum. When the impregnated catalyst PM-Z is reduced at high temperature such as at 450°C, it should be mentioned that obvious interaction between platinum and molybdena Can also be affected. This is illustrated by the diffraction diagram of Fig. GC which exhibites a set of additional rings beside 3 the basic pattern of Pt and ~oC3' So far we are not able to find out an adequate assignment for this new pattern in accordance vdth ASTM standards. It is very plausible that a small portion ~106+ ions is probably c ap ab I.e be Lng reduced to zero valent and some kinds of clustering or alloying species formed by the intimate neighboring platinum end molybdenum atoms might be able to add some new rines to the diffraction pattern as depicted in Fig. 6c The diffraction rings of Fig. 6d of the extracted pow3• der from anchored catalyst PM-1 are completely different from Fig. Gc. Here the diffraction rings of neither platinum nor molybdena crystallites are detected. A new set of diffraction rings
707
with clear and intense edges is observed instead. A semiquantitative analysis with reference to ASTM Standard Data substantiates that a uniform surface complex containing Pt group (see 3M0 2 table 1) is dominating over the supported phase of PM-1. The XPS
TABLE I Data for electron diffraction pattern of Pt nos. 2
3 7 8 9 10
exper. 2.46 2.11 1.06 0.94 0.86 0.82
3Mo 2
H
ASTI-l 2.413 2.124 1.061 0.938 0.859 0.84
40 100 16 25 2
40
surface analysis of the anchored Pt-MO/Si0 2 catalyst (ref. 1) has inferred the preponderance of Mo 2 + as the exclusive state ofmolybdenum in the surface of reduced sample. It is therefore suggested that for the anchored preparation of PM-1, a uniform and stoichiometric surface complex of Pt/Mo=3/2 is formed with the divalent Mo 2 + ions anchored through oxygen bonds on to 5i0 2 surface. Reaction test The conversion of n-pentane was determined for catalyst evaluation. 50 mg of catalyst sample Was used for test run. The reaCtion was carried out in a flow microreactor at temperature of 340 0C and hydrogen pressure of 1.5 kg/cm 2• Dosages of 2JlI each were successively injected into hydrogen with a flow rate of 80 ml/min. The results are shown in table 2. I t is interesting that while the molybdenum catalyst of M¥l and M-2 are of nil activity for pentane conversion, the hydrogeno lysis as well as the total conversion of n-pentane is remarkably higher on Pt-Mo bimetallic catalysts than the monometallic platinum catalyst and the PM-1 is the most active in particular. Otherwise, it can be seen that a higher ratio C1/C of 1.5 on 4 PM-2 than those on PM-l, and Pt/Si0 2 is observed in product distribution. DISCUSSION The strict differences in surface property between PM-1 and PM-2 evidently emphasizes the importance of the effects of prepa-
708
TABLE 2 Conversion of n-pentane on Pt/Si0 2 and Pt-Mo/Si0 2 catalysts a product distribution of hydrogenolysis (mol%)
convercata- H/Ptb sion lyst (mol%) P-2 C d 0.45 PM-3 <0.1 PM-l 0.51
H
4.39 7.01 11.75
. -
11.6 38.0 38.0 12.40.94 22.1 32.1 30.8 15.0 1.4fl 17.8 33.7 34.0 14.5 1.22
I
68.8 20.8 10.4 87.0 9.1 3.9 86.6 7.6 5.8
2
0
a. react. temp. 340 C, p~ 1.5kg/cm , vel. H2 80 ml/minj b. hydrogen chemisorptio~at room terr.p. in a pulse flow technique (ref. 6); c. 3wt%Pt/SiO prepared by impregnation method, reduced at 450 oC; d. 3wt%Pt-3.3~MO/§i02 prepared by separate impregnation method, reduced at 450 Cj e. H hydrogenolysis, I isomerization, Cy cyclization. ration methods. For the anchored Pt-Mo catalyst PM-1, the following essential points are summarized about the nature of the surface. First, a uniform surface species with stoichiometric ratio of Pt/Mo (atom ratio) =3/2 is formed with the M0 2+ ions bonded to the oxygen atoms of Si0 2 writen as Pt3MO~+{O-Si=)2 groups which may have the same diffraction of crystal lattic as that of Pt 3M0 2 alloy detected by electron diffraction technique shown in table 1 when these surface complexes are being extracted off the surface of Si0 2• Second, the behavior of red~ction on reoxidized catalyst PM-1 exhibited from ESR and TPR measurements may be implicated that due to the mutual interaction with platinum, the rate of partial reduction of M0 6+ to M05+ with anthracene is significantly higher on PM-1 (Fig. 4c) than on molybdena alone (Fig. 4a), the reduction profile at higher temperatures (~3000C) in TPR curve (Fig. 3a) becomes broadened and shifted to lower temperature region comparing to Fig la of anchored molybdenum oxide. It is speculated that a group of Pt;Mo~+Ox~-S~)2 (I) x+7, may eXist in the oxidized state of Pl':-l, and therefore the phenomena described above can be accounted for as follow: ESR
0
0
'O/~o'0/Pt'o,R/PV
~
+
0/'0
;;j/~;Wi!I Il!. (I)
Anthracene
-
709
(Mo 5+ ) + Anthracene +
TPR
(1)
,
<, ~;
° Pt M pt) CJ 9, IiI "O("..,) 0, I LQ'"(. ,_/
Pt" ".
lJ' _.
-' T1a
0"0
I
II
0 -. l
+ HZ
----4.....
Pt 3Mo2+ Z
(Z)
_
It is noticed that the anchored Pt-Mo catalyst PM-1 is the most active one in total conversion as well as hydrogenolysis for reaction of n-pentane (see table 2). According to our previous work (ref. 7), the enhanced hydrogenolysis on bimetallic catalyst could be interpreted as the consquence of increased severity of dehydrogenation of the reaction intermediates. In the other words, it is anticipated that platinum is affected electronically by molybdena and becomes more favourable for the formation of mUltiple dehydrogenated intermediates. The increase of total conversion and hydrogenolysis selectiVity on PM-1 is thus understandable in this aspect since the bond cleavage process on PM-1 ratains the feature approach to that on Pt/SiO Z' as revealed by the lower value of C1/C ratio observed in the hydrogenolysis products. 4 For the impregnated Pt-Mo catalyst PM-Z, the TPR, ESR and TEM measurements show that due to the presence of the intimated mixed individual particles of platinum and various molybdenum oxides, the reduction temperatures of molybdenum oxides are sharply shifted toward to lower temperatures (Fig. 3b) comparing to molybdenum oxide (Fig. 1b), and the lower rate of increment of M0 5+ ESR signal in PM-Z (Fig. 4d) comparing to PM-1 may be caused by the coreduction of individual particles of ptO and M0 6+ to Pto y (x>y) x and M0 5+ . When the catalyst PM-Z is reduced at higher temperature such as 4500C, some kinds of Pt-Mo alloy~ may be formed as depicted in Fig. 7c3. Therefore, the intermingled individual particles of platinum, molybdenum oxides and Pt-Mo alloys are responsible for the higher activity and hydrogenolysis selectiVity of n-C
5
conversion, and C1/C4 (1.5) ratio (table Z) in hydrogenolysis pro-
710
ducts which may show the presence of multiple splitting of npentane. CONCLUSIONS Pt-Mo bimetallic catalyst prepared through the step of anchoring of ~-allyl complexes can be adequately applied to realize a high dispersion of bimetallic ingredients by the formation of uniform surface complexes with stoichiometric composition of Pt3MO~+~0-Si~)2 groups which are responsib~ for the high activity and hydrogenolysis selectivity of n-pentane conversion. Pt-Mo bimetallic catalyst obtained by conventional technique have non-uniform surface composition. The major part of the supported phase in PM-2 is crystallographic ally mixture of intermingled individual particles of p'l at.i.num , various HoO x and Pt-Mo alloys, the synergy effect is closely related to the higher activity and hydrogenolysis selectivity of alkane conversion than that on platinum catalyst. ACKNOWLEDGEMENT The authors wish to thank Mr. SONG, YONGZHE for assistance with the ESR measurements and Mr. CAl, HAlLIN for the electron microscope studies.
REFERENCES
2
3 4
5 6 7
Yu. I. Yermakov, B. V. Kuznetsov and V. ·A. Zakharov, lICat.alysis by Supported Complexes" in "Studies in Surface Science and Catalysis 8", Elsevier, Amsterdam-Oxford-New York, 1981, pp. 345-350. H. S. Gandhi, H. C. Yao and H. K. Stepien, in "ACS Symposium Series, No. 178, Catalysis under Transient Conditions", A. T. Bell and L. Louis Hegedus (Eds.), American Chemical Society, 1982, p. 143-162. T. M. tri, J. P. Candy, P. Gallezot, J. Massardier, M. Primet, J. C. Vedrine and B. Imelik, J. Catal., 79 (1983) 396. G. Leclercq, T. Romero, S. Pietrzyk, J. Grimblot and L. Leclercq, J. Mol. Catal., 25 (1984) 67. Guo Xiexian, Yang Yashu, Deng Maicun, Li Humin and Lin Zhiyin, J. Catal., 98 (1986). Zhang Yihua, Lin Zhiyin and Yang Yashu, J. Catal.,(China) 1 (1980) 299. Guo Xiexian, Yang Yashu, Shi Yingzhen, Yuan;Zhenkui, Li Huimin and Deng Maicun, in "Catalysis on the Energy Scene", S. Kaliaguine and A. Mabey (Eds.), "Studies in Surface science and Catalysis 19", Elsevier, Amsterdam-Oxford-New York-Tokyo, 1984, pp. 389.
711
DISCUSSION G.C. BOND: Have you integrated your TPR curves to establish the oxidation states of the Mo and of the Pt after your oxidation treatments? YANG Y. : When Pt- and Mo-oxide are Co-simultaneously present, it is difficult to determine the valence of Pt and Mo by integrating the area under TPR curve. We used BRD and oxygen chemisorption to establish the oxidation states of the Mo and of the Pt before and after our oxidation treatments. L. GUCZI : Have you observed segregation of the bimetallic particles during or after the catalytic reaction? YANG Y. : We have not studied the segregation of the bimetallic catalysts during or after the catalytic reaction. J.P. CANDLIN : Although you state that the hydrogenolysis of pentane is more efficient with the anchored Mo/Pt catalysts, is there any difference in the hydrogenolysis products depending on whether an anchored method or aqueous impregnation method is used? (e.g.
C
~
, C,' C
5 2
C + C 3 2)
YANG Y. : The product distributions of hydrogenolysis of n-pentane on Pt-Mo-Si02 catalysts prepared by anchoring and impregnation methods are different form each other; the Cl/C4 atomic ratio is higher with the latter than that with the former catalyst. It indicates that the terminal C-C scission is favoured on Pt-Mo-Si02 catalyst prepared by impregnation method.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV 1987 Elsevier Science Publishers B.V" Amsterdam - Printed in The Netherlands
713
CONTROLLED PREPARATION OF BIMETALLIC HYDROGENATION CATALYSTS
O.A. FERRETTI
1,
L.C. BETTEGA de PAULI 4 J.P. BOURNONVILLE
2,
J.P. CANDy
3,
G. MABILLON
4
and
1CINDECA (UNLP-CONICET), U7 nO 257, 1900 LA PLATA, ARGENTINE 2pETROQUISA, av Buenos Aires, 40, 20070 RIO de JANEIRO, BRAZIL 3INSTI'IUT DE RECHERCHES SUR LA CATALYSE, 2, av Albert Einstein, 69626 VILLEURBANNE Cedex, FRANCE 4INSTITUT FRANCAIS DU PETROLE, B.P. 311, 92506 RUEIL MALMAISON, Cedex, FRANCE
ABSTRACT Tetrabu tyltin reacts selectively wi th oxidised or hydrogen covered transition metals such as Rh, Pd, Ni at mild temperatures with evolution of butane am butenes. This reaction allows to master the macroscopic homogeneity of the bimetallic phase and probably the microscopic one. Activation at high temperature in hydrogen leads to the evolution of the residual butyl radicals with diffusion of tin through the metallic particles. The bimetallic phase appears to be very selective in hydrogenation and dehydrogenation reactions owing to the suppression of hydrogenolysis.
INTRODUCTION One of the most important goals of catalysis is the improvement of the selectivi ty of catalysts in order to lower the energetic and economic cost of industrial processes. For the hydrogenation or dehydrogenation of hydrocarbons this is generally achieved by using bimetallic catalysts (1). The simplest way to prepare such catalysts is the coimpregnation of salts of the metals. Though it is very simple this method does not necessarily give good results because the obtained catalysts can be heterogeneous at the atomic scale. As the selectivity of ca talysts seems to be dependent for a number of reactions on the degree of association of the metals, other methods of impregnation have already been studied. For example YERMAKOV (2) has found that the impregnation of bimetallic complexes such as (pt(SnC12)2Cl)2- on alumina gives better dehydrogenation catalysts than those obtained by two step impregnation of tin chloride on alumina followed by adsorption of platinum chloride. For the hydrogenation of esters to alcohols TRAVERS (3) has shown that the modification of rhodium on alumina or silica with tetrabutyltin leads to highly active and selective catalysts. Tin and rhodium appeared to be closely associated as evidenced by STEM analysis. Nevertheless the nature of the reaction between supported rhodium and tetrabutyltin was not studied. Startim from the previous results this work aimed to study and develop this organometallic route of preparation of bimetallic catalysts am particularly to precise the effect of the nature and the oxidation
714
state of the transition metal on its reactivity towards tetrabutyltin. EXPERIMENTAL Preparation of the monometallic catalysts Nickel, palladium and rhodium were deposited on silicas by cationic exchange of ammine complexes of these metals at pH 10-11. Silicas were purchased from Shell. The resulting catalysts were activated by calcination at 573 K, and/or reduction between 500 and 773 K. Preparation of the bimetallic catalysts The activated catalysts were contacted under argon or hydrogen with tetrabutyltin (Merck) dissolved in heptane. The temperature was progressively raised to heptane reflux. After around 6 hours the solid was filtered, washed in heptane and dried under vacuum. Characterisation of the catalysts Metal contents were determined by X ray fluorescence and intraparticle distribution by microprobe analysis. S.T.E.M. analyses were performed on a VG HBS. For Ni containing catalysts magnetisation was determined by the Weiss extraction method. Catalytic tests Hydrogenation of ethyl acetate to ethanol was performed in a differential reactor working at 5 MPa pressure and temperature in the range 521-573 K. Dehydrogenation of isobutane was performed in a similar device working at atmospheric pressure and elevated temperature. RESULTS Characteristics of the monometallic catalysts The main characteristics of the three monometallic catalysts are listed in table 1. Similar rhodium catalysts have been prepared with different rhodium contents and dispersions to test the effect of these parameters or the reactivity of SnBu
4•
Effect of the oxidation state of the transition metal The nature of the interaction between SnBu and supported metals has been 4 studied at different stages of their preparation : after impregnation of the ammine complexes and drying, calcination, reduction and chemisorption of hydrogen 'or oxygen. The resulting catalysts were thoroughly washed with heptane to eliminate physisorbed species. In these conditions a blank experiment showed that the silica support fixed only a very limited amount of tetrabutyl-
715
TABLE 1 Characteristics of the monometallic catalysts
Ni
Rh
2.4
2
40
50
Pd
Carrier
Metal content
0.55
Wt. %
Dispersion %
55
Dried Rh catalyst is almost unreactive to SnBu (table 2). But after calci4 nation this same catalyst can react and fix 2,3 % tin. After reduction and subsequent flushing in air at room temperature the reactivity remains steady. A comparable result is obtained if after reduction the catalyst is cooled and reacted with SnBU in hydrogen. But if after reduction the catalyst is flushed 4 in argon at 573K to desorb hydrogen and cooled in argon its reactivity towards SnBU
4
decreases sharply.
TABLE 2 Effect of the chemical state of the transition metals on the fixation of SnBu Pretreatment
Sn introduced wt%
Sn fixed wt%
Rh
a b c d e
2.3 2.3 2.3 2.3 2.3
0.20 2.3 2.4 0.6 2.3
Ni
f g
5 5
0.15 3.5
Pretreatment
4
a, dried - b, dried and calcined - c, dried, calcined, reduced and reoxydised - d, dried, calcined, reduced and desorbed e, dried, calcined and reduced - f, dried, reduced and passivated g, dried and reduced.
716
Reduced Ni catalyst passivated in diluted oxygen reacts to a very limited extent with SnBU
But if after reduction the catalyst is maintained in hydrogen 4. its reactivity becomes high. Calcined or reduced and put in air Pd catalyst reacts easily with.SnBU
4• All these reactions lead to an evolution of a mixture of butane and butenes
when performed in argon, and only butane when performed in hydrogen. Stren;1th of the interaction between .transition metal and SnBU 4 The isotherm of fixation of SnBu on oxidized rhodium is presented in figu4 re 1. It shows an initial steep increase in the amount of fixed to introduced tin, and a plateau for a tin content of 2.3 % equivalent to an atomic ratio
= 2.3 (Rhs s behaViour. sn/Rh
= surface
Rh atom). Palladium catalyst exhibits a similar
For hydrogen covered rhodium or nickel there is no such well defined plateau in the range of concentrations studied.
•
3.00
2.00
• RH-H
• fb-{)
• NI-H
1,00
• RHil3
0.00 +----,,--,.--..---,.--.--r---r-,---,,---1
0.00
2,00
1.00
Fig. 1. Isotherms of fixation of SnBu
3.00 4
on different supported transition metals
The selectivity of the reaction between snBu shown in figure 2 for Rh/Si0 and dispersions (0.50
2
and Rh/A1
and the transition metals is 4 catalysts of different metal contents
203 The direct proportion ali ty
between tin content
and surface rhodium content indicates that there is no spill-over effect of SnBU on the carrier induced 4
by the presence of the transition metal.
717
SN
r----------------..,
3.0
2.0
a: tga==2.45
0.5 Fig. 2. Effect of rhodium content and dispersion on the amount of fixed tin
The strengbn of the interactions. their specificity and their irreversible behaviour (evolution of gases) allow to speak of true chemical reactions between SnBu
and metallic oxides or hydrides, These reactions are limited to the 4 surface of the metallic particles in their oXide or hydride form and cause
the evolution of about two to three of the four butyl radicals of snBu The 4, remaining ones are decomposed when the catalysts are activated in hydrogen. This leads to the formation of metallic tin at a temperature of about 500 K where tin can rapidly diffuse into the transition metal particles. After the activation the catalyst can be represented by a global formula Met Sn /Si0 ; x y 2 = Rh. Ni or Pd and y/x depends on the dispersion of the transition metal.
Met
Characterisation of the bimetallic catalysts (a) Microprobe analysis Either the transition metal (Ni or Rh) or tin is uniformly distributed in the whole volume of the silica balls (fig. 3). Moreover in a case where the concentration profile of nickel shows a certain heterogeneity that of tin follows the same variations. similar results have been obtained for Pd-Sn/Si0
2•
718
Fig. 3. Repartition profiles of Rh, Ni and Sn in the bimetallic catalysts
(b) STEM analysis The observed homogeneous repartition of the two metals at the scale of microprobe (resolution
limitrvl~m)
does not necessarily imply an homogeneity
of the composition of each metallic particle. STEM analysis of Rh-Sn or Ni-Sn catalysts (fig. 4 a, b) indicates that tin is never observed alone on the carrier, and for each catalyst the signals corresponding to the two metals is simultaneously detected when the metallic particles are analysed. Unfortunately this technique did not allow us to undoubtly measure variations of composition from one bimetallic particle to another. Thus we do not know if the composition is homogeneous at the particle level or if there are variations associated for example with the size of the particles. (c) Magnetic properties The study of the magnetic properties of nickel is a powerful tool of characterisation of this metal : percentage of reduced atoms, dispersion, interaction with adsorbates and promoters can be easily studied (4-5). Thus we have studied the evolution of the saturation magnetisation of nickel at different stages of the preparation of Ni-Sn. After reduction at 773 K and desorption the saturation magnetisation is 93 % that expected for the total reduction of ionic species to metal atoms (fig. 5). The adsorption of hydrogen on this catalyst decreases the saturation magnetisation of 25 % against 30 % expected from the dispersion determination by
719
4A -4a- Rh-Sn
4. -4b- Ni-Sn ~ig.
4. STEM analysis of Rh-Sn and Ni-Sn particles
720
electron
r\~
microscopy.
NI
90.00 80.00
NI-H
70.00 60.00 50.00
NI-H
40.00
+
SNBu4
30.00
~3711O
20.00
(NI-SN)-H
on 10
10.00 0.00
Fig. 5. Decrease of the magnetization saturation at the different stages of the prepara tion of Ni-Sn catalyst
The reaction of this hydrogen covered nickel with tetrabutyltin at 371 K leads to a strong decrease of the saturation magnetisation. As at this stage tin still holds one butyl group; it seems very unlikely that it diffuses through nickel. The magnetisation decrease is more probably associated with a long distance effect of tin on the magnetic properties of the nickel atoms of the second and third layers of the particle. This is consistent with the very low content of tin (less than 10 %) necessary to suppress the magnetism of nickeltin alloy. Activation of the bimetallic catalyst in hydrogen at 773 K fUlly suppresses the saturation magnetisation of nickel. Vacuum or argon treatment at 573 K does not restore the magnetic properties of the metallic phase. As was previouly said metallic tin is generated during the activation treatment. Thus the suppression of the magnetic properties of nickel can be associated with the diffusion of tin through the nickel particles. Moreover this result is confirmed by the detection of X ray diffraction patterns characteristic of intermetallic compounds such as Ni
3sn,
Ni
3Sn 2,
NiSn or Ni
3Sn4•
(d) Catalytic properties Two reactions have been used to check the effect of the specific interaction between tin and group VIII metals on the catalytic properties of supported
721 metallic phases
hydrogenation of ethyl acetate and dehydrogenation of iso-
butane. As we can see in tables 3 and 4, the addition of tin notably enhances the selectivity in hydrogenation (table 3) or dehydrogenation products (table 4), whatever the group VIII metal may be. Moreover, in the case of hydrogenation of ethyl acetate, tin not only decreases hydrogenolysis but also increases the hydrogenating activity of the metallic phase (6). This special behaviour has been connected with the particular interaction between tin and group VIII metal in the metallic particles after activation in flowing hydrogen at high tempera-
ture (7). Finally the comparison between nickel and rhodium based bimetallic catalysts emphasizes the superiority of the rhodium based one. TABLE 3 Ethyl acetate hydrogenation P : 5 MPa - H : 9 - Molar conversion 2!ACOEt Catalyst
10 %
Molar Selectivity (% mole)
T (K)
Hydrogenation (a) Products
Hydrogenoiysis(b) Products
% Rh
543
46.8
53.2
% Rh % Sri
573 529
93.7 97.9
6.3 2.1
2.9% Ni 3 % Sn
523
93.8
6.2
1 2
(a) hydrogenation products : almost only ethanol (b) hydrogenolysis products : methane and ethane TABLE 4 Isobutane dehydrogenation
Catalyst
(K)
T
WHSV (h- 1)
Conversion (mole %)
Selectivity (mole %, IC = C + C + C 4 1 2 3
1
% Rh
473
2.7
3.6
1 2
% Rh % Sn
773
2.8
11.6
95
5
2.9 % Ni % Sn 3
773
8.3
7.5
66
34
99.8
722
CONCLUSION The interaction between SnBU and oxidised or hydrogen-covered transition 4 metals is a true chemical reaction leading to the evolution of butane and butenes and to the formation of supported bimetallic complex still holding butyl radicals. The nature of this liquid-solid reaction has not been fUlly understood but appears to be strongly dependent on the chemical state of the transition metal. Nickel oxyde reacts with SnBU to a very limited extent whereas 4 rhodium and palladium oxides react to a larger extent; the reaction rate between SnBU and PdO being higher than with Rh As SnBU is known to be a 4 4 203, reducing agent and as the reaction seems strongly dependent on the reducibility of the supported metallic oxyde it may be an oxido-reduction. The composition of the mixture of gaseous
products has been interpreted in a similar way (8)
though it remains difficult to write a global reaction scheme. With hydrogen covered metals the only product evolving during the fixation of SnBU
is butane.
4
For one leaving butyl group the reaction equation is : Met H +
SnBu4----~
MetSnBU + BuH 3
It looks like a dinuclear reductive elimination (9). Activation in hydrogen leads to the evolution of the remaining butyl groups and to the diffusion of tin in the transition metal particles with, at least in the case of Ni-Sn formation of intermetallic compounds.
This method of modification of transition metals by tin looks very interesting because it allows to modify selectively the metals whereas conventionnal inorganic methods such as impregnation of SnCl in HCl strongly modify the pro2 perties of the support. The resulting catalysts are homogeneous at the macroscopic scale. On the other hand, the metallic particle size cUstribution strongly influences the:microscopic homogeneity. Because the ratio fixed Sn/Met
s is
nearly steady, the amount of fixed tin depends on the metallic dispersion: the smallest particles contain more tin than the biggest. Therefore a narrow particle size.distribution of the transition metal should lead to a very homogeneous composition of the bimetallic phase. Moreover, tin addition causes a drastic decrease of hydrogenolysis reactions in hydrogenation of ehhyl-acetate and dehydrogenation ofisobutane and also an increase of the rate of hydrogenation of ethyl-acetate into ethanol. REFERENCES J.P. Boitiaux, J. Cosyns, M. Derrien and G. Leger, Hydrocarbon Processing (March 1985) p. 51. 2 Y.P. Yermakov, B.N. Kuznetsov and V.A. Zakharov, catalysis by supported complexes, Studies in surface science and catalysis, Elsevier, Amsterdam, 1983, pp 351-363. th 3 C. Travers, J.P. Bournonville and G. Martino, Proc. of the 8 Int, Congress on Catalysis, Berlin (West), West Germany, July 2-6, 1984, Verlag Chemie Ed,
723 (IV) pp. 891-902. 4 J. Dalmon, J. Catal., 60 (1979) p 25. 5 J. Richardson, J. Appl. Phys., 49 (1978) p. 1786. 6 J.P. Candy, O.A. Ferretti, J.P. Bournonville and G. Mabilon, J. Chern. ·Soc., Chern. Com. (1985) pp. 1197-1198. 7 O.A. Ferretti, Thesis, Paris (1986). 8 A. El Mansour, J.P. Candy and J.M. Basset, in preparation. 9 J. Nortcn, Acc. Chern. Res., 12 (1979) p. 139.
724
DISCUSSION A.F.M. WIELERS : I have a question with regard to your magnetic measurements. You have shown that upon alloying with SnMn saturation magnetization of nickel slowly decreases. Have you calculated how many Ni atoms are magnetically affected in the introduction of one Sn atom? As this alloy formation between Ni and Sn is a rather exothermic process, it can be expected that this might be rather drastic. Can you comment on this? J.P. BOURNONVILLE : We have not calculated the number of nickel atoms affected by one tin atom. Nevertheless, our results are in very good agreement with the literature data concerning the tin addition effect on the magnetic properties of nickel-tin alloy: one tin atom can suppress the magnetism of the surrounding fifteen nickel atoms. We have not taken into account the thermodynamic aspect of the alloy formation, which is probably favoured by the diffusion of the metal inside the particle at high temperature (773 K) under reducing atmosphere (H 2).
B. Delmon. P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV '987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
725
©
CHARACTERIZATION AND CATALYTIC PROPERTIES OF COBALT FISCHERTROPSCH CATALYSTS PREPARED BY CHEMICAL REDUCTION AND USED IN A LIQUID PHASE. C. BECHADERGUE-LABICHE', S. MAILLE', P. CANESSON 1, M. BLANCHARD 1 and D. VANHOVE 2. luniversite de Poitiers - U.A. C.N.R.S. 350, Catalyse en Chimie Organique, 40 Avenue du Recteur Pineau, 86022 Poitiers (France). 2Universite Claude Bernard, U.A. C.N.R.S. 231, E.S.C.I.L., 43, Boulevard du 11 Novembre 1918, 69622 Villeurbanne (France).
SUMMARY Chemical reduction of an organic salt of cobalt leads to active and selective catalysts for hydrocondensation of carbon monoxide into light olefins. Magnetic and X.P.B. measurements have shown that the active part of the catalyst is constituted of thin metallic particles inbedded in an alumina gangue with an average diameter less than 100 ~. It is possible to obtain a catalyst with roughly identical properties by another preparation method if it leads to these thin cobalt particles. INTRODUCTION Catalytic synthesis of low molecular weight olefins from carbon monoxide and hydrogen is of both practical and academic interest in C1 chemistry, and a number of recent studies are devoted to this subjet (refs. 1-2). In order to improve the selectivity into light olefins, various modifications for the catalyst have been proposed (refs. 3-5). Metallic cobalt deposited on a support is a well known catalyst for the Fischer-Tropsch synthesis. It usually leads to the formation of a large cut of linear paraffins, and it is possible to reduce this spectrum by a carreful choice of the support (ref. 6). Recently, it has. been shown that chemical reduction of an organic salt of cobalt, namely cobalt acetylacetonate, leads to the formation of very thin metallic particles these particles used in a liquid phase as a slurry are active and selective for the transformation of syngas into light olefins (ref. 7). This reaction is also observed
726
when
the
CO +
is
that
can
be
H mixture 2 considered as
liquid
replaced
by
syngas
methanol
for
(ref.
8)
this
particular
catalytic
properties
reaction. Since
the
same
metal
has
different
when it is deposited on a support or used without this support in
a
has
liquid been
with
phase,
the
investigated,
some
exact
and
nature
results
physico-chemical
of
this
reported
last
here
characterizations
catalyst
are
that
dealing
have
been
tentatively correlated with catalytic properties. METHODS Catalyst preparation a)
Catalysts
reduction. at
6.5
of
the
g of Co
CR
type
{acac)2
cobalt
(1.5 g of Co)
acetylacetonate previously dried
50°C under vaccum were dissolved into 235 ml of dry benzene
under
an
{C 2H5)3 added
Ar
atmosphere.
(AI/Co =
in a
occurs
of
of
g
benzene
After
diluted
was
immediately
butadiene
CO +
added.
was
H 2
reactor under
50
at ml
The
the
g
syngas
the
was
3) then
distillation
slurry
temperature
until
2
of o.terphenyl by
obtained
the
after
solution
eliminated
and
and
(butadiene/Co
room temperature and 20 Benzene
cooling into
GOC,
4.32
g
of
Al
of dry benzene were
few minutes with the aid of an air-tight styringe.
Reduction 4.1
1.5)
of
was the
reaction
then
minutes, into
a
15
allowed
to
(solvent) under was
temperature
of
a
at
added.
stream into
rised
(190°C)
dry
warm
was
transfered
catalyst
solution
ml
of the
slowly
has
been
obtained, the first heating being achieved on two days. b) tion.
Catalysts of
the CD type
cobalt carbonyl decomposi-
8.25 g of fresly distilled aluminum ethanolate was dissol-
ved into 300 ml
of
refluxing dry benzene
after the dissolu-
tion has been complete, the solution was cooled at room temperature and then 5.8 g of Co the
solution
was
partly
was
introduced
into
The
temperature
of
in
the
same
the the
conditions
were added. After dissolution, 2(CO)8 evaporated and the liquid mixture reactor
reactor than
with was
those
20
then
for
g
of
rised
o.terphenyl. under
syngas
the CR type catalysts.
Catalytic reactions catalytic
reactions
have
been
performed
made in pyrex glass represented in figure 1.
in
an
apparatus
727
EXIT
A
Fig. 1. Schematic diagram of the experimental vessel. A-B-C sursaturator condenser system for methanol introduction D gas flow meter ; E : gas intrances ; F : reactor ; G exit to gas chromatographs. All the experiments have been carried out under atmospheric pressure in dynamic conditions. Quantitative analysis were performed by on line gas chromatography. For CO, H2, CH 4, CO 2 a catharometer Fischer Gas Partitionner equipped with two columns in series (the first one, 3 m, 1/4", filled with 30 % HMPA on chromosorb P and the second, 2 m, 1/4", filled with 13 X molecular sieves) was used. Hydrocarbons were analysed using a GIRDEL 3000 gas chromatograph equipped with a 2 m, 1/8", 30 % squalane on spherosil column. From time to time, an analysis with a 30% SE 30 on chromosorb P column has been made in order to check if heavy hydrocarbons are present in the effluent at the reactor exit. Physico-chemical measurements a) Magnetic measurements. Magnetization variations versus magnetic field have been studied in the 0-21 kOe range by axial extraction of the sample (ref. 9). 0.3 ml of the catalyst slurry were introduced under argon into a sealed Pyrex tube and then placed in the magnetic field. A rapid extraction fo the sample far from the magnetic field induced a current
728
in a coil. The integrated signal was converted using a magnetic picture factor and plotted versus magnetic field on one side and the inverse of this field on the other side, giving measurements for : i) the mass of magnetic species (metallic cobalt) ii) the value of the size of metallic particles from low fields magnetization (excess value) iii) the value of the size of metallic particles from high fields magnetization (default value) iiii) the remanent magnetization measurement that is related to the amount of metallic particles with a size greater than a limit fixed by the experimental temperature measurement ° at room temperature in the case of cobalt). 120 A b) XPS measurements. Superficial concentrations of cobalt and aluminum were determined from XPS analysis, using a HP 5950 A apparatus equipped with an Al anode. After washing under argon with numerous portions of benzene in order to remove all the solvent, the catalyst sample was deposited on a gOld grid, the 4f7/2 of the support being used as a reference for the calibration of the binding energy scale. The superficial AI/Co ratio has been determined after integration of the Al2s and Co 2P 3/2 peaks. RESULTS AND DISCUSSION Catalysts of the CR type During the catalyst preparation, the first heating has been realized under syngas (1 CO 2 H2). Magnetic measurements during this thermal treatment allow to know both the content into metallic cobalt and particles size results obtained are summarized in figure 2. Metallic cobalt begins to appear at 100°C, and the size ° of these metallic particles is quasi homogeneous around 20 A, demonstrating that after chemical reduction by AI(C 2H5)3 cobalt is dispersed at a quasi atomic level by complexation with butadiene. Decomposition of this complex starts around 100°C, leading to the formation of very thin particles. The size of these particles does not change until 150°C and the average diameter increases slightly from 20 to 35 by heating at
A
A
200°C. Besides
these
particles,
magnetic
measuremens
show
also
729
DimenSion(%)
40
20
10
e---e----e1-
1.-_-
--_------'0
+-
nt:)
200
150
100
Fig. 2. Effect of the temperature treatment during the first heating on the particles size. X : higher value for particles size (from low fields magnetization). + lower value for particles size (from high fields magnetization). extent of cobalt into large metallic particles ( > 120 ~). that a part of the metal is agglomerated into large metallic ° The particles with an average diameter greater than 120 A. percentage of metal into these particles is rather low (less than 10 %) and does not change by increasing the temperature until 200°C. Analysis of the gas composition at the reactor exit during this
preparation
starts
around
shows
140°C.
is very high (> 80
that
At
% into
carbon monoxide
this
temperature,
the C + cut)
hydrocondensation
olefins
selectivity
and the methane
2
ratio
is rather low, as evidenced on figure 3. When
the
catalyst
temperature
increases,
the olefins
selecti-
decreases and the methane ratio increases. At 200°C, the olefins selectivity and the methane percentage are roughly
vity
equal at 50 %. After this
preparation,
the
stable for long periods of time. As evidenced on figure 4,
obtained
is
very
period, dimension increases together with the particles. After one month
during
and dispersion of particles size percentage of cobalt into large
catalyst
this
730
ACTIVITY
SELECTIVITY
mol/hxg
0/0 80
...
.....
...... .
_~.....
X
>c''
103
.!
./
... ...
1.5
/
......... ...... ~ -i
60
/
..............
:/
1.0
Ji/ ...
40
0.5 20 T (oC)
160
180
200
o
Fig. 3. Evolution of the reaction products during the catalyst preparation. : olefins selectivity. : methane selectivity. x : catalyst activity. on stream, cobalt into these large particles represents about 40 % of the total metallic content. Nevertheless, it must be noted that this evolution of the catalyst does not affect the olefins selectivity and only the methane ratio increases. This catalyst, unlike classical cobalt deposited on a support, leads to the transformation of syngas into light olefins with a good selectivity, and a typical hydrocarbons repartition at 190°C is represented on figure 5. The main hydrocarbon is propene and the chain length is limited at CG. As it has precedently claimed (ref. 8), this catalyst is also effective for the transformation of methanol into light hydrocarbons. In this reaction, methanol can be considered to react as liquid syngas, and on figure 5 is also represented hydrocarbons repartition starting from
731
0;' Dimenslon(
d> 120
U
X
)(
40
60
'/ -/1
50
V
+
+
40
20
Ie
30
+ ~
----
,..-Jt
20
TIME ON STREAM (Days)
0
0
30
20
10
Fig. 4. Influence of time on stream on particles size. x higher value for particles size. + : lower value for particles size. extent of cobalt into large metallic particles. CH Since this kind of catalysts is not acidic, the reaction 30H. mechanism is different from methanol transformation into gasoline
(Mobil
is
a
is
also
(ref.
process)
classical observed
10)
that
acidic with
takes
catalyst. cobalt
place on HZSM5 This
deposited
methanol on
zeolite which transformation
alumina
catalysts
and a proposed mechanism for this reaction has been
given elsewhere (ref. 11). As
it
has
.
been observed during the catalyst preparation,
the olefins selectivity is better if the reaction is carried out at low temperature (figure 6). There is a sharp decrease in olefins productions when the reaction temperature is rised from 160°C to 170°C. For higher temperatures, the olefins selectivity in the C + cut and methane ratio are rather stable
z
around 55 %. This observed selectivity into light olefins can be originated either from the absence of any support for the metallic species or from the use of a liquid phase that acts as a dispersing agent for the active phase. Experiments run with supported cobalt on alumina used in a liquid phase have shown that
in these conditions syngas is transformed into a mixture
732
%
CO+ 2 H2
r,
I
40
1
PARAFFINS
\ I
I I I I I I I I I
OLE FINS
,
~
CH30H
,.., I
•
M
H
~1tql !ll
30
>'I ;~)
:.,
•1 I
~~I
I I I 1
;tt f'
I I I I I 1
' ~.I
l"'1 I
r" 20
1 I I I
\ I
r-
1J
l.~~ ~~
,.~4I..
~
",
10
~
~:t ~~
I'II,"
~~1 1
Fig. 5. type.
2
Hydrocarbons
',1
;"1
",~I
:/\
':1 ~I
r-
~ ~'1
~j
»1 ~;1
~
.,
., f.
\\1
I
~!
:'1 ~~
;':1
"I
i'~l .;. ~.
3
...1
~~
l~ ,II ~!111 I
r'I
I
I
I ..
4
nC
5
repartition
with
catalysts
of
the
CR
of light paraffins (ref. 12). Frorr: the comparison of these two sets of experiments, the liquid phase seems rather to intervene as a dispersing agent limiting the chain growth, leading to a norrow cut of light hydrocarbons, and not as a main factor for olefins selectivity. X.P.S. analysis of the surface after reaction shows that the surface of catalyst particles is greatly enriched into A13 + ions. Since during the preparation, aluminum is introduced with the reducing agent, a theoretical ratio AI/Co of 1 .3 is expected. Bulk chemical analysis of the catalyst has confirmed this ratio but from X.P.S. measurements, it can be concluded that the superficial AI/Co ratio is about 12. So, the surface of the real catalyst seems to be constituted
733 SELECTIVITY
ACTIVITY
%.
!
100
mol/hag x 103
!
80
/
1.5
60
1.0
40
0.5 20
160
180
200
Fig. 6. Influence of the reaction temperature on activity and selectivit¥ of CR type catalysts. : olefins selectivity inside the C2 cut. methane selectivity. x catalyst activity. of metallic cobalt species well dispersed into an alumina gangue. Moreover, chemical and X.P.S. analysis show that a very important amount of carbonaceous species (50 % by wt.) is included in the catalyst particles. From microscopic studies, it can be seen that particles are agglomerated into relatively large conglomerates the size of which increasing notably with working time (10-50 ~m after 27 days). Catalysts of the CD type Since the real catalyst is a dispersion of metallic cobalt into an alumina gangue, starting from coo and A1 3+ ions, it may be possible to prepare active and selective catalysts
734
for
the
transformation
of
syngas
(or
methanol)
into
light
olefins. Cobalt CO + 2 walls
of
If CO
carbonyl
HZ
leads the
(CO)a
Z
to
decomposition the
reactor.
formation
The
only
alone of
a
under
a
stream
of
cobalt mirror on the
obtained
product
is
methane.
is mixed to aluminum ethanolate before decomposi-
tion, only a part of metallic cobalt is deposited as a mirror, the
other
the
solvent.
part
is
as
Results
thin
metallic
obtained
for
particles
dispersed
hydrocarbons
into
selectivites
at 190°C with this catalyst are summarized in figure 7.
% 40
CHa OH
r,
,.,
PARAFFINS
I I
OLEFINS
I
I I I
I I
I
30
I I
I
, I
20 ~.,
I
I
I I I I I I I I
10
nC 1 Fig. 7. type. As
Hydrocarbons
repartition
with
catalysts
of
the
CD
for the other experiments the chain growth is limited
at C7, and the selectivi ties, except for methane, are similar whatever the starting products are syngas or methanol. Here also,
ole fins
selectivity is
rather good,
but lower than that
735
observed with catalysts of the CR type. Moreover, catalytic activities for the two types of catalysts are roughly the same. Various studies of the preparation parameters have shown that the most important factor for the obtention of an active an selective catalyst is that the formation of a cobalt mirror must be minimized, and the more important is this mirror formation, the higher is the methane selectivity. The best conditions are obtained if the AI/Co ratio is at least of 1.5 and the temperature rising during the cobalt carbonyl decomposition very slow. Some experiments run with a Raney cobalt catalyst have shown that syngas is transformed mainly in methane, the olefins selectivity being rather low (20 % of the C2 + cut). Moreover, in this case catalyst activity is about 15 times lower than that observed with CR and CD types catalysts. These results are in good agreement with the others since the Raney cobalt catalyst is mainly constituted of rather large metallic particles, leading to methane formation. From this study, it can be concluded that the use of a liquid phase brings about a limitation in the chain growth of hydrocarbons. This is probably a consequence of the dilution of the metallic active species into an inert phase. The active species for obtention of light olefins are thin metallic cobalt particles which are embedded into an alumina and carbonaceous gangue that prevents metallic particles to collapse, large particles being responsible for methane formation. ACKNOWLEDGEMENTS One of us (C.B.) thanks C.d.F. Chimie for a grant. The authors thank the C.N.R.S. (GRECO "Call) for financial support. The aid of Dr. MERIAUDEAU (I.R.C. Villeurbanne, France) for X.P.S. analysis is highly appreciated.
REFERENCES
2 3
B. Bussemeier, C.D. Frohning and B. Cornils, Hydrocarbon Process., 55 (1976) 105 and references cited therein. C.D. Frohning in J. Fabbe (Ed), New Syntheses with Carbon Monoxide, springer-Verlag, Berlin, 1980, p. 356. J. Barrault, C. Forquy and V. Perrichon, Applied Catalysis, 5 (1983)
119.
736
Y. Do i , H. Miyake, A. Yokota and K. Soga, J. Catal., 95 (1985) 293. 5 R. Hemmerich, W. Keirn and M. Roper, J. Chern. Soc. Chern. Commun., (1983) 428. 6 D. Vanhove, Zhang Zhuyong, L. Makambo and M. Blanchard, Applied Catalysis, 9 (1984) 327. 7 M. Blanchard, D. Vanhove, F. Petit and A. Mortreux, J. Chern. Soc. Chern. Commun., (1980) 908. 8 M. Blanchard, D. Vanhove, R.M. Laine, F. Petit and A. Mortreux, J. Chern. Soc. Chern. Commun., (1982) 570. 9 J.A. Dalmon, Thesis, Lyon, 1971. 10 M. Blanchard, L. Makambo and D. Vanhove, Nouveau J. Chimie, 6 (1982) 459. 11 L. Makambo, P. CanessonthD. Vanhove, M. Blanchard, A. Mortreux and F. petit, Proc. 8 Into Congr. Catal., Berlin, 1986, Dechema, 1986, pp. 11-207-219. 12M. Blanchard and D. Vanhove, Metal-Support and Metal-Additive Effects in catalysis, B. Imelik et al., Elsevier, Amsterdam, 1982, p. 219. 4
737
DISCUSSION K. NOACK: Can your method of preparation be used for other metals? P. CANESSON : Yes, this method of preparation can be applied to other metals, and we have prepared by chemical reduction iron, ruthenium, rhodium and chromium catalysts. Moreover, it can also be applied to bimetallic catalysts starting from a mixture of organic salts of the two desired metals. G. JANNES : Having in mind the dependence of activity and selectivity of your catalyst on the size of the small metal particles, could you comment on your choice not to support the catalyst? P. CANESSON : We have tried to conduct the same preparation method by impregnation of an alumina previously treated by Al(CZH)3 with cobalt acety1acetonate. In this case, the selectivity into olefins is worse than that observed with the slurry catalyst; this result may be a consequence of a higher hydrogenating power for supported catalysts. J.G. van OMMEN: From your research, it appeared that the actual small size of Co particles is important for the selectivity. Will it be possible to use the "COO" complex because there you have the smallest particle possible, or otherwise to stabilize it by putting the organometallic Co complex on a support? P. CANESSON : The use of a support for tentatively stabilize the Coo complex has not been successful yet. The support increases the hydrogenation rate leading to a decrease into olefins selectivity. A. VANNICE: First, a comment: we have supported Co(CO)a clusters on carbon and they behave very similarly to your carbonyl precursors. My questions are: have you studied chemisorption of CO and HZ? What are the turnover frequencies of your catalysts compared to literature values? P. CANESSON : We have not tried to study carbon monoxide or hydrogen chemisorption with our slurry catalyst. Since we have not determined the number of supposed active sites, it is not possible to give a value of the turnover frequencies. Nevertheless, the specific activity of the slurry catalyst is of the same order of magnitude than that of supported cobalt (or other metals) catalysts, namely at ZOO°C under atmospheric pressure,about 10-3 mole of carbon monoxide transformed per hour and per gram of metal. XU XIAODING : 1/ Did you see water gas shift reaction operating? Z/ In the presence of HZO, oxidic Co can form. Did you see oxidic cobalt species and what is the1r function besides the effect on Co particle size? P. CANESSON : 1/ At ZOO°C, 0.1 MPa, we have not observed the water gas shift reaction since we have detected only few percentages of carbon dioxide in the reaction products. 2/ With the real working catalyst, we have no experimental evidence on the real oxidation state of cobalt species: either metallic or oxidized. For the XPS measurements, the solvent must be eliminated and during this treatment and transfer into the high vacuum chamber some modifications of the surface state of cobalt can occur and I prefer not to draw any conclusions from the XPS analysis on this subject. D.M. HERCULES: I agree that use of ESCA to study oxidized Co species is very difficult. I would suggest the use of Raman spectroscopy which can be used in situ; some cobalt species have high Raman scattering cross-sections. P. CANESSON : Thank you very much for your suggestion.
This page intentionally left blank
739
B. Delmon.P. Grange. P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V.• Amsterdam - Printed inThe Netherlands
COMPARISON OF THE QUANTITATIVE STUDIES BY STEM OF HYDRATED HYDROXYCARBONATES AND RELATED MIXED OXIDES CATALYSTS FOR CO HYDROGENATION TO ALCOHOLS R. SZYMANSKI, Ch. TRAVERS, P. CHAUMETTE, Ph. COURTY, D. DURAND Institut Frangais du Petrole, I & 4 Avenue de Bois Preau - B.P. 311 92506 RUEIL-MALMAISON CEDEX (FRANCE) ABSTRACT Through the example of investigation of Zn A1 4 and Cu Cobzn AId mixed 1 20 oxide catalyst precursors, a method of characterizat~on is depi~tea, ~omple mentary to the overall methods (XRD, XPS, •.. ). This method is based on the analysis by STEM of thin foils of the solids, obtained by ul tramicrotomy. It enables to fully characterize all the different phases formed in such complex systems by their respective morphology, chemical composition, structure and proportion, and to visualize at a quasi-macroscopic scale their distribution. The new available information constitutes a guide to correlate catalytic properties with related characterization and to improve knowledge on catalyst preparation procedures. INTRODUCTION In parallel to its R and D program on methanol and C alcohols processes 1-C6 (ref. 1) IFP has developped new characterization techniques to study the corresponding
catalysts.
Among
these
techniques,
quantitative
high
resolution
Scanning Transmission Electron Microscopy (STEM) appeared to be a very powerful tool to follow the preparation procedure and correlate the different steps of catalyst preparation with catalyst properties for alcohol synthesis. Since a part of these correlations have already been disclosed (ref. 2,
3), the
present paper emphasises the utility of the analysis by STEM of thin foils of the solid as a complement to other techniques (XRD, XPS, ATD, ATG) to determine the morphology,
structure,
chemical composition, distribution and pro-
portion of the different phases detected in mixed oxide catalysts and related precursors. This quantitative evaluation of the phase compositions is completely illustrated in the case of a dried calcined zinc-aluminate mixed oxide. The same technique is also applied to a CuaCobZncAl catalyst to obtain information d on the evolution of the catalyst precursor during drying or calcination before and after a1ka1inisation. This information is particularly useful in view of catalysts optimization for methanol and C C alcohols synthesis. 1-' 6
740 EXPERIMENTAL Materials The ZnA1
and CuaCobzncAl mixed oxides have been prepared by copred 204 cipitation of the metal nitrates (CU, Co (II), Al, Zn) in 0.5-1.5 M aqueous solutions with 0.5-1. 5 M alkali metal carbonate solution under controlled and constant pH conditions at temperatures between 293 and 363 K. The precipitates were subsequently washed, dried and calcined (ref. 4). Among other procedures (ref. 12, 13), alkalinisation of the calcined precursor is carried out via a conventional aqueous wetting followed by ovendrying and calcination. Methods - XRD powder patterns were collected using a PHILIPS goniometer with the Co-K radiation
(A=
0.179 nm)
- the characterization by STEM was performed on a high resolution VG HB5 dedicated STEM equipped with a KEVEX energy dispersive X-ray spectrometer, a GATAN 607 electron energy loss spectrometer and a high sensitivity TV camera for the collection of electron microdiffraction patterns (ref. 5). Thin folls of the compounds (t
information
was
obtained from
the X-ray emission spectra
using a quantitative method developped in our laboratory (ref. 6) which is derived from the Cliff-Lorrimer approach for the use of the so called k-factors to correlate the peak intensities with concentrations in the thin film approximation.
The use of ul tramicrotomy foils made the thin film criterion valid
in most analyses. The
TV-camera
facility
enabled the collection of electron diffraction
(ED) patterns on the dried precursors in spite of instability observed for small electron probe sizes. RESULTS AND DISCUSSION 1 - Zn Q support lA1 2 4 characterization
an example of STEM capabilities for mixed oxides
• XRD : the X-ray diffraction (XRD) patterns of the dried and calcined precursor are shown in figure 1 (ref. 7). - The XRD pattern of the dried precursor (fig. la) matches that of an hydrotalcite-like phase crystallized in a hexagonal structure with the lattice
741
Figure 1 XRD powder patterns of the dried (a) and calcined (b) precursor of the ZnAl support 204
parameters
~
= 0.305 nm and
when preparing S).
c
= 2.24 nm. This type of phase is often observed
(Cu)(Co)Zn Al mixed oxides by coprecipitation methods (ref.4,
The stoechiometric form of hydrotalci te-like phases,
isostructural with
the minerals hydrotalcite, pyroaurite and stichtite is M~IM;II(OH)16C03-4 H20 and is generally described as an alternation of positively charged brucite layers IM~I M;II (OH)16 12+ , in which the cations are distributed among the octahedral position, and negatively charged interlayers !C03.4HfoI2-, (ref. 9). According to our nominal stoechiometry in metals MII(Zn), M II(AI) (Zn/AI=0.5) the observed hydrotalcite-like phase is likely to contain a number of vacancies, at least if we suppose the sample to be homogeneous. The lower lattice parameters values than those measured on the Zn C0 phase (a = 6AI2(OH)16 3.4H20 0.3075 nm, ~ = 2.275 nm Iref.SI) support this statement, the presence of vacancies being consistent with a compression of the structure. The background underlying the hydrotalci te pattern is likely to be due to the presence of an amorphous phase in the sample, which cannot be identified accurately by only the XRD method. - The XRD pattern of the calcined precursor (figure Ib) reveals the disappearance of the hydrotalci te-like structure.
The new pattern matches that
of an AB
spinel (A = Zn, B = AI) in agreement with the nominal composition, 204 the broad lines observed being indicative of the amorphous and/or divided
state of the phase resulting from the the decomposition of the dried precursor phases. It must be pointed out that this line broadening makes it difficult to obtain accurate values for the lattice parameters or to detect the presence of other phases which could result from the decomposition of the dried precursor.
742
STEM Typical bright field and dark field images obtained at low and high magnification on the dried and calcined precursor are shown in figure 2. Thanks to the use of microtomy cuts of the solids, it was possible to visualize unambiguously the different morphologies present at each step and moreover their respective distribution on a quasi-macroscopic scale inside the grains (figure 2c). Accordingly, two distinct morphologies were observed in the precursor at each step of the preparation (drying Ifigure 2al and calcination these
I figure
2d I ). The STEM capabilities enabled us to fully characterize
morphologies
(structure,
composition,
proportions).
They
correspond
to two distinct phases formed after the drying step which decompose independently during the calcination step into two other specific phases. Dried precursor The dried precursor is composed of large platelets, showing a
hexagonal
shape in some cases and whose size varies from 0.1 to 1 micrometer, randomly dispersed among a granular phase (which looks like an agglomerate of very small cristallites). The large platelets appeared well crystallized, showing well defined Electron Diffraction (ED) spots, which were indexed with the hexagonal symetry (figure 2b). The second phase was found amorphous (no ED). In agreement with the XRD patterns, it is concluded that the large platelets correspond to the hydrotalcite-like phase. • The Zn/ Al atomic ratio of each phase was investigated by X-ray microanalysis. The platelets and the granular phase showed specific compositions and were found to be very homogeneous over the sample. The mean Zn/AI ratios measured on each phase are 0.9 for the platelets (hydrotalcite-like phase) and 0.35 for the granular phase (amorphous), to be compared to the 0.5 value of the nominal composition of the sample and the value of 3 for the stoechiometric hydrotalcite structure • • Due to the homogeneity of each phase over the sample, their respective proportions could be evaluated from the set of equations obtained by writing the respective repartition of Zn and AI, given the nominal composition. The calculated proportions,
nearly 30 at. % of hydrotalci te-like phase versus
70 at. % amorphous phase, are in agreement (in a rough approximation) wi th the statistical distribution of the platelets observed in the imaging modes of the STEM over the sample. It should be noted that a simply quali tative interpretation of the peak intensities on the XRD pattern would lead to overestimating the proportion of hydrotalci te-like phase and the overall homogeneity of the dried precursor.
743
Figure 2 : STEM characterization of' the ZnAl 0 support a) Bright field image of the dr i.ed precursor ; ~r ED pattern of the platelets in the dried precursor; c), d) Dark field image at low magnification and bright field image of the calcined precursor j e), f) XES spectra and ED pattern of the granular phase in the calcined precursor ; g), h) XES spectra and ED pattern of the holey platelets.
744 Calcined precursor . The biphasic character of the dried precursor is kept after calcination, as illustrated by figure 2d. Large platelets are still present but are now holey ; they are distributed among a granular phase similar to that observed in the dried sample, which seems composed of larger micrograins (Ill X-ray
microanalysis
revealed
#4
that
nm) . their
compositions
are
directly
related to those measured on the parent phases in the dried precursor. The X-ray spectra in figures 2e and 2g illustrate the spectra obtained respectively on the granular phase and on the holey platelets over the calcined precursor. A great homogeneity in the composition of each phase was observed, the mean value of the respective ZnlAI atomic ratios being the same as after the drying step, 0.35 for the granular phase and 0.9 for the platelets. Consequently, the
calculated phase proportions are unchanged (70 at. % granular,
30 at.%
platelets). • It is clear that the two phases in the calcined sample result from the decomposi tion of the respective phases in the dried precursor. Accordingly, the structural evolution of each phase after its decomposition under the calcination conditions was investigated by electron microdiffraction. The ringed ED patterns collected on the ganular phase
(figure 2f) cor-
respond to the spinel type structure detected by XRD on the calcined precursor (broad lines)
: the amorphous phase of the dried precursor decomposed into
well crystallized micrograins of the expected spinel structure in spite of
= 0.35 or ZnO• 7 0 O. 3Al204) • collected on the holey platelets (figure 2h) are more
its unstoechiometry (ZnlAL The
ED patterns
troublesome: the zone axes indexed (nearly 10011 in the hexagonal symetry) are quite similar to those observed on the dried precursor which correspond to the hydrotalci te-like phase. Accordingly one would expect to observe the associated XRD pattern as in the dried sample. In fact, the structure of the holey platelets is not yet clear. However, as
it
has
been
shown
that
the
decomposition of hydrotalcite-like
phases
proceeds by a compression along the C-axis which involves the disappearance of the associated lines (10061,
10031, ••• ) in the XRD pattern (ref. 7, 8),
it is assumed that the observed ED patterns and crystal shape (holey platelets) result from an incomplete decomposition of the hydrotalcite-like phase under the
calcination
conditions.
The associated XRD pattern would be convoluted
with that of the spinel granular phase
745
Figure 3 XRD powder patterns of the dried and calcined precursor of Cu CObZncAld mixed ~xiaes before (a, b) and after (c, d) alkalinisation
ZlII
10.110
11.110
21.110
30.110
".110
10.110
".110
11.110
2 - cuaCobZncAl
mixed oxides catalyst precursors : investigation of the d effect of alkalinisation in aqueous solution on the precursor homogeneity
• XRO : the XRD powder patterns of the catalyst precursor are given in figure 3
(ref. 7). The XRD pattern of the dried precursor (figure 3a) matches that of a well
crystallized hydrotalci te-like phase of the type previously depicted
(with lattice parameters a~o.305 nm,C=2.25nm) which is tranformed in an AB 204 spinel type phase after calcination (figure 3b). Line broadening makes difficult the accurate determination of the spinel lattice parameters. The alkalinisation in aqueous solution of the calcined precursor results after drying in the recristallization of a hydrotalci te-like phase besides a spinel phase and CuO (figure 3c). This new hydrocalcite-like phase differs from the former
(a:
0.305 nm
,c
2.28 nm
), the differences between the
spinel phases being more difficult to appreciate.
746
After calcination,
the XRD pattern of one (or more ?) spinel type phase
is detected besides well crystallized CuO (figure 3d). XRD shows
that
the
alkalinisation
in
aqueous
solution of the calcined
precursor has involved a segregation of the CuaCObZncAl by
spinel phase, caused d the hydrotalci te-like-spinel phase transformation, which seems to present
a reversible character at least for a part of the calcined precursor spinel phase.
Such a phenomenon was
previously described for alumina (ref.
10)
as
well as for Zn Al and Cu Zn Al mixed oxides (ref. 4). However, the slightly different lattice parameter values and the presence of CuO indicate that the term "reversible" should be improper, the composition of the new hydrotalcitelike phase being probably different from the former phase • • STEM : thin foils of the precursor (obtained by ultramicrotomy) were investigated by STEM at each step of the preparation in order to precise the segregation mechanism shown up by XRD and especially the active metals repartition, which is of great importance to explain the final catalytic properties. Typical bright field images of the precursor before alkalinisation are shown 4b).
in One
figure type
4,
after
drying (figure 4a)
of morphology
was
and after calcination
observed at each step,
(figure
respectively large
platelets (0.1 to 1 micrometer) showing sometimes a hexagonal shape (figure 4a) in the dried precursor and large holey platelets in the calcined precursor. The observed morphologies corresponding to the hydrotalcite-like phase (figure 4b),
are quite similar to these observed for the hydrotalci te-like phase and
associated spinel in the Zn A1
support (figures 2b 20 4 are in agreement with the XRD powder patterns. The
was
apparent homogeneity
confirmed,
shown by
and 2d). The results
the presence of one-type morphology,
at each step, by X-ray microanalysis. Figure 4c illustrates
the X-ray emission spectra obtained allover the sample after drying or calcination, and corresponding to the nominal composition. • Figure 4d shows the morphologies observed after alkalinisation in aqueous solution the
and
drying
morphology
of
the
calcined precursor.
corresponding to
the
One can easily distinguish
recry.stallized hydrotalci te-like
phase
detected by XRD (figure 3c), whose type (platelets, hexagonal shape) is exactly that already observed for hydrotalcite-like phases (figure 4a). Its structure was confirmed by electron microdiffraction. Besides the
this
phase,
a granular type phase is observed corresponding to
spinel phase of the XRD pattern (figure 3c). These two phases form an
intimate mixture over the sample (CuO cristallites were sometimes observed). X-Ray microanalysis revealed a great homogeneity of each phase over the sample.
The
respective
X-ray emission spectra show that the recrystallized
hydrotalcite-like phase (figure 4f) nearly lost all of its Co content (compared
747
Figure 4 : STEM characterization of CUaCobZncAl mixed oxides d a), b) Bright field images of the dried and calcined precursor; c) associated XES spectra, before alkalinisation ; d), e) Bright field images of the dried and calcined precursor f), g) XES spectra of the platelets and granular phase at each step, after wetting.
748 to the nominal composition, figure 4c) and that the granular phase, which lost a part of its Cu content, is Co rich (figure 4g) • • After calcination, the biphasic character (plus the CuO crystallites) is kept (t'ig\.lr.e
~.e).
The hydrotalci te-like phase decomposed into a spinel
t.ype phase of the same morphology as in the calcined precursor before
a]l~a
linisation (figure 4b), besides a granular spinel phase. The calcination did not change the Cu Co Zn Al content of each specific phase. By using the same approach as that described for the Zn Al
0 support, 2 4 it was moreover possible to evaluate the respective proportions of each phase in the sample. • This study, which has confirmed the complementarity and the agreement between XRD and STEM information, illustrates the type of data one might expect by applying such an approach to the investigation of mixed oxides catalyst precursors.
In the depicted example, it has been shown that a further alka-
linisation leads to the formation of large amounts of a Cu-rich spinel phase and
CuO cr.stalli tes.
Under the reacting medium,
these phases undergo the
formation of very large metallic copper crystallites on one hand and,
to a
lesser extent, large metallic cobalt-based ones on the other, both of them being responsible for significant increases of hydrocarbons formation at the expense of alcohol selectivity in CO hydrogenation (ref. 2, 3). It has been claimed and confirmed by experiments that homogeneity of such complex phases is the key to gain the best alcohol selectivity (ref. 13). Therefore, such available information is very useful for catalysts preparation and catalytic behaviour understanding. This type of study is also in progress in our laboratory for the characterization of Cu-Zn-Al methanol synthesis catalysts. CONCLUSION The main purpose of the paper was to present a powerful method of investigation of mixed oxides catalyst precursors, complementary to the overall characterization tools such as XRD and XPS (X-ray Photo Electron Spectroscopy). This method, based on the characterization by STEM of thin foils of the solids, can provide unique information, very useful to improve the preparation procedures of catalyst precursors or to further explain the catalytic behaviours. Through the example of the Zn Al 0 support, it is clearly shown that 2 4 only a XRD characterization would have led to incomplete and/or erroneous conclusions about the sample composition. The new information available by the method enables us to control the preparation of such supports in order to obtain them with the textural and acidic properties required for specific catalytic processes.
749
The description of a CU a COb Znc AId catalyst precursor at each step of its preparation illustrates the type of information made available on such very complex systems. One can follow the genesis of the metallic active phase from the precursor preparation initial step to the final reduction step. Moreover the same approach can be used to describe the evolution of the active phase under the reacting conditions. Accordingly, we recommend this method which is currently applied inside our Rand D programs to characterize mixed oxide based catalysts and which allowed us
to prepare catalysts of homogeneous composition, highly active,
selective and stable, for the synthesis of methanol (ref. 11) or C
I-C 6
alcohols
(ref. 12, 13) from synthesis gas. Acknowledgments : the authors thank their colleagues from IFP for the collaboration in XRD (especially Mrs C. DURAND) and Transmission Electron Microscopy. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12 13
J.P. Arlie, J.P. Cariou, Ph. Courty, A. Forestiere, Ph. Travers, VI Symposium on alcohol fuels technology, Ottawa (1984). Ph. Courty, D. Durand, E. Freund, A. Sugier, J. Mol. Cat. 17 (1982), p , 241P. Grandvallet, Ph. Courty, E. Freund, Proceedings 8th International Congress on catalysis (Berlin), vol. II, p. 81. Ph. Courty and Ch. Marcilly, "A scientific approach to the preparation of bulk mixed oxide catalysts", Preparation of catalysts II, Poncelet, Amsterdam (1983) J. Lynch, J. Microsc. Spectrosc. Electron., 8 (1983), p. 481. R. Szymanski and J. Lynch, "Quantitative X-ray microanalysis of divided solids in the STEM", in preparation for the 11th Int. Congo X-ray Opt. and Microanal., London (Canada), (Aug. 1986). C. Durand et al. unpublished results. C. Busetto, G. del Piero, G. Manara, F. TRIFIRO AND A. Vaccari, J. Catal. 85 (1984), p. 260. R. Alman, Acta Crystal1ogr., B24, (1968), p. 972. a) R. Tertian, D. Papee, J. Chim. Phys., (1958), p. 341. b) D. Papee, J. Charrier, R. Tertian, R. Houssemaine, Congres de l'Aluminium, Paris, Juin 1954. US 4,552,861 to IFP US 4,122,110 and 4,291,126 to IFP GB 2,118,061 and 2,158,730 to IFP
750
DISCUSSION A. VACCARI : 1/ How do you justify the unusual Zn/Al ratio (- 0.9) in your hydrotalcite-like phase, taking into account that you have large excess of positive charges ? 2/ It is clear that a water alkalinization step gives rise to heterogeneity of the system. Have you tested different methods and what are your ideas on the role of water and drying and/or calcination conditions? P. CHAUMETTE : 1/ The Zn/Al of -0.9 is an experimental result and crystallographic investigation as well as TGA-TDA analyses are still in progress to explain the measured Zn/Al ratios. 2/ There is not only one answer to that question. Depending on the nature and composition of the dried or calcined precursor, different methods and thermal activation procedures are to be used to avoid heterogeneity of the system. STEM analysis allows us to follow the homogeneity of the solid during all these steps, and to determine the best procedure to be used. ZHAO J.S. : 1/ You mentioned in ZnA1204 that the holey platelets result from an incomplete decomposition of the hyarotalcite-like phase under the calcination conditions. Have you any evidence? What is the calcination temperature? On ED patterns with hexagonal symmetry, could it be that ZnO crystallites exist because ZnO is an hexagonal transparent thin platelet? 2/ On the XRD pattern of Cu-Co-Zn-Al mixed oxides, have you found rosasite or malacnite-like phases? P. CHAUMETTE : 1/ Complementary studies have been performed on different zincaluminate samples. TGA analysis up to BOO°C, and X-ray diffraction and microscopy clearly show the evolution of the hydrotalcite-like phase to holey platelets during thermal activation. Details will be given in a future publication. The holey platelets presented in this paper were obtained after calcination in the range 350-450°C. The ED patterns observed could effectively be confused with that of ZnO, but all the ED patterns were associated with XPS spectra showing the presence of Zn and Al in a constant Zn/Al ratio on the different morphologies (granulae, platelet, holey platelet) and this, on different samples of the same type at each step of the preparation. 2/ This phase has a particular morphology and contains no aluminum. It was not detected by STEM on the Cu-Co-Zn-Al mixed oxide presented in this paper. K. FOGER : Have you studied the decomposition of the hydrotalcite phase as a function of temperature, heating rate? When does spinel form and how does the thermal treatment affect the morphology of the spinel phase ? P. CHAUMETTE : As mentioned to Dr. Zhao we followed by TGA-DX-STEM the decomposition of the hydrotalcite phase. Its morphology evoluates from platelets to holey platelets and finaly small crystallites after a treatment at 750°C. This evolution is closely related to the decomposition of carbonates introduced at the coprecipitation step. E.B.M. DOESBURG : We investigated the Ni/alumina system prepared by coprecipitation by TEM and high resolution electron microscopy. In all stages of preparation (dried, calcined, reduced material), we see the same morphology and the pictures ressemble very much the pictures you made. That is rather surprising. Did you also look at the reduced catalysts and do they also show the same morphology as the calcined and the dried catalysts? P. CHAUMETTE : It is true that the same morphology is observed after drying, calcination and also reduction of our catalysts with only slight changes after each treatment as you found in your own materials. We are now investigating the evolution of different catalysts after syngas treatment with the same techniques.
751
J.W. GEUS : In the electron microscope, the decomposition of the hydrotalcite structure has to be envisaged. Combination of XRD and electron diffraction can indicate whether the decomposition proceeds to a significant extent in the vacuum and the electron beam of the electron microscope. We have got indications that embedding specimens liable to decomposition can prevent or can diminish decomposition. My questions are : 1/ Was there a significant decomposition in the electron microscope? Z/ If there was decomposition, did you compare the extent of decomposition for embedded and sectioned and not embedded specimens ? P. CHAUMETTE : A significant decomposition of the hydrotalcite-like phases occurred in the electron microscope working at high magnification with high current probes. A high sensitivity TV-camera coupled with the TEM enabled imaging with low current probes to minimize decomposition. In the STEM, the ED patterns could be obtained especially thanks to our specific recording device using a TV-camera (see ref. 5). A. VANNICE : Have you found any chemisorption technique whose adsorption value correlates with catalytic activity? P. CHAUMETTE : It is very difficult to interpret chemisorption measurements on bimetallic copper-cobalt catalysts. Work is under progress in the field of HZ and CO chemisorption in our laboratory to find a correlation with catalytic activity but no such correlation has been obtained at the time being.
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B. Delman, P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
753
PREPARATION OF MULTI COMPONENT CATALYSTS FOR THE HYDROGENATION OF CARBON MONOXIDE VIA HYDROTALCITE-LIKE PRECURSORS. S. GUSI 1, F. PIZZOLI
1,
F. TRIFIRO,l, A. VACCARI
l
and G. DEL PIER0
2
lIstituto di Tecnologie Chimiche Speciali, Viale del Risorgimento 4, 40136 BOLOGNA (Italy). 2ENIRICERCHE, Via F.Maritano 26, 20097 SAN DONATO MILANESE (Italy).
ABSTRACT Precursors with hydrotalcite-like structures were successfully prepared by precipitation as a function of the stereochemistry and the ionic radius of the different elements. It was shown that the addition of a second bivalent element (e.g. Mg, Co, and Zn) forced the copper to enter into ternary precursors. The nature of the elements plays an important role in the stability of the precursors, their decomposition temperature and the compounds obtained by calcination. The presence of another element may modify the catalytic activity of the Cu/Cr system. In particular, zinc promoted methanol synthesis, while cobalt showed different effects depending on its amount. No effect was detected for magnesium, despite its forming a HY precursor with copper. Such varying performance patterns may be attributed to the formation of mixed phases, which are -stable under catalytic conditions thus preventing the formation of CuCr0 2• INTRODUCTION The two most important processes in the industrial hydrogenation of carbon monoxide are the synthesis of hydrocarbons (1-3) and the synthesis of methanol, used as both raw material and fuel (4-6). The last few years have seen an intensification of research efforts on catalysts in the low-temperature synthesis of methanol (7-9), with special emphasis on formulations for the simultaneous synthesis of higher alcohols (10,11). Multicomponent, copper-containing catalysts are employed and structural non-homogeneity andlor chemical segregation are detrimental to proper activity and selectivity. Previous works have shown that hydrotalcite-like phases (HY, having general formula (CuXZn6_x)A12(OH)16C03'4H20) may be useful precursors of catalysts for the synthesis of methanol at low temperature (12,13). In these precursors all cations are randomly distributed in positively-charged brucite-like layers tiered 2with negatively charged (C0 interlayers (14). 3·4H20)
754
A study of this technique as a general method in preparing multicomponent catalysts is reported herein. To this end HY precursors containing different bivalent and trivalent elements were prepared. EXPERIMENTAL The precipitates obtained by coprecipitation, i.e. placing a solution of the nitrates of the elements into a solution of NaHC0 at 333K, were washed to a mi3 nimum sodium content, dried at 363K and calcined at 623K for 24h (12). XRD powder patterns were collected using a Philips goniometer equipped with a stepping motor and automated by means of a General Automation 16/240 Computer. The radiation was nickel-filtered CUK a (A =0.15418 nm) with silicon as internal standard. The phase compositions and crystal sizes were determined by a profile fitting method, comparing the observed profiles with the computed ones, calculated according to Allegra and Ronca (15). A Carlo Erba Sorptomatic mod. 1826 with N adsorption was used to measure the 2 surface area, volume and radius distribution of pores, the latter being calculated by means of the Pierce's method (16). The DSC analyses were done on a PerkinElmer 2C calorimeter, automated by a Perkin-Elmer Data System 10 Computer. TG analyses were performed by a Perkin-Elmer thermoba1ance TS6-2, in a N flow with 2 a heating rate of 10K min-l. The infrared spectra were recorded with the KBr disk technique by a Jasco A-202 spectrophotometer. The calcined catalysts were reduced in the reactor using a H mixture, with 2/N2 the hydrogen content and temperature being progressively increased. The catalytic tests were performed in a copper piston flow reactor, 493-573K range, with a GHSW of 15,000-18,000h-l, using a H v/v gas mixture. The reaction pro2:CO:C02=65:32:3 ducts were analyzed in-line by a Perkin-Elmer F30 gas chromatograph, equipped with FID and two (1/8in. x 1.5m) columns fitted with 80-120 Poropak QS. After condensation of the other compounds at 273K, the gases were analyzed with a Carlo Erba Fractovap GT gas chromatograph equipped with TCD and two (1/4in. x 2.0m) columns fitted with 100-120 Carbosieve S. The catalytic data were collected and analyzed using a Perkin-Elmer Chromatography Sigma 15 Data Station. RESULTS AND DISCUSSION Nature of the precipitates Table 1 shows that the copper precipitates preferentially as malachite, with
755
TABLE I Compositions, compounds identified by XRD and surface area values for the precipitates dried at 363K. Sample
Composition
Atomic ratios
Phases identified
(%)
Cat Cat Cat Cat Cat Cat Cat Cat Cat Cat Cat
1 2 3 4 5 6 7 8 9
10 11
Cu/Cr Zn/Cr Co/Cr Cu/Zn/Cr Cu/Mg/Cr Cu/Co/Cr Cu/Mn/Cr Cu/Zn/A1 Cu/Zn/Fe Cu/Zn/A1/Cr Cu/Co/ln/Cr
76.0:24.0 76.0:24.0 76.0:24.0 38.0:38.0:24.0 38.0:38.0:24.0 38.0:38.0:24.0 38.0:38.0:24.0 38.0:38.0:24.0 38.0:38.0:24.0 38.0:38.0:12.0:12.0 36.0:2.0:38.0:24.0
Surface area m2 g-l
M»HY HY HY HY HY HY MnC03; HY HY Au HY HY
84 120 10 106 55 58 68 9
64 2 56
M= malachite-like phase; HY= hydrota1cite-1ike phase; Au= auricha1cite-1ike phase only a small amount of HY phase. This can be explained by its stereochemistry (17), taking into account that in the HY phases the cations are randomly distributed among the octahedral positions of the brucite-like layers (14). Magnesium, cobalt, and zinc, which can form regular octahedral structures (17), not only give rise to pure HY compounds but also force the copper to enter into ternary HY phases, that may be obtained without side phases for Cu/Me(II) ratios
~
1.0 (12,
14). It is worth noting that the chromium forms, even if in a small amount, a HY phase with the copper, unlike what has previously been observed with aluminum (14). However with the iron, copper and zinc precipitate essentially as auricha1cite, even if a small amount of a HY phase cannot be excluded (Figure 1). Chromium compounds have lower crystal size and greater cell volume than those with aluminum (Table 2). Cell volume is in agreement with the dimensions of the trivalent element (18) but practically independent of the nature of the bivalent ions. Nevertheless, the
~/~
ratios were always smaller than those reported for
both natural hydrota1cite and stichtite. This was due to smaller values for the parameter
~
and larger ones for
~,
indicating a closer packing of the layers and
a greater disorder in the cationic brucite-like layers. It may be deduced that such compounds have an elastic structure which is compatible with slight alterations linked to the different compositions.
756
TABLE 2 Crystal size
and crystallographic parameters of some hydrotalcite-like phases. 0
Crystal size (nm)
Cell volume ($.3)
c (A)
a (A)
cia
Cat 2 Cat 4 Cat 5 Cat 8 Cat 10 Hydrotal cite a Stichtite b
3.2 4.5 4.5 30.0 n.d.
192.7(7) 192.7(7) 192.6(8) 183.5(2) 186.7(2) 188.7 192.7
22.70(3) 22.71(3) 23.01(4) 22.45(1 ) 22.54(2) 23.120 23.235
3.131(5) 3.130(5) 3.109(5) 3.071 (1) 3.093(2) 3.070 3.095
7.25 7.26 7.40 7.31 7.29 7.53 7.51
a) NBS 14-191
b) NBS 14-330
Sample
By contrast, owing to the higher value of Mn(II) ionic radius (18), only a small amount of HY phase was obtained together with MnC0 for Cat 7. This last 3 contained some copper as suggested by the lower cell volume compared to that reported in the NBS 7-268 (307.6 vs. 309.7 A3, respectively). In all cases the HY phases were already present in the first precipitates and no aging procedure was necessary for their formation. It is worth noting that the nature of the Me(II) also affects the stability of the HY phases. The Cat 5, which was washed at length to minimum sodium content, gave rise to amorphous hydroxycarbonates on account of the strong tendency to trap CO of the compounds 2 containing magnesium (19,20). Therefore, a sample containing more sodium was employed for Cat 5, so as to start in this case too from a HY precursor. Calcination processes The DSC-curves of some precursors are shown in Figure 2. The thermograms showed the presence of different endothermic peaks as a function of the nature of the precursors and of the composition. In particular, for the HY precipitates, the first three peaks may be attributed, as previously reported (12,21), to losses of water and the decomposition of the hydroxycarbonates to oxides. These transitions were accompained by a weight loss of about 25% for all samples, with changes only in rate of weight loss. The temperatures and peak intensities changed for each sample, probably on account of different electrostatic interactions between basic layers and interlayers (22) and the different bond strength in the basic layers associated with the different compositions. The endothermic peaks
757
Cat 2
o
'C
c
W
Cat 10
...
'C
C"
'C
Cat8
410
550
690' 830
970
Temperature (K)
Fig. 1. XRD powder patterns of the samples containing different trivalent elements after drying at 363K.
Fig. 2. Differential scanning calorimetric curves of some precursors.
at temperatures higher than 800K were due to the reaction among oxides in forming spinel phases. In this case, too, the intensity and the temperature correlated to the nature of the elements present. Cat 7 shows a broad thermogram where the peaks of the HY-phase decomposition are still detectable. These transitions were associated with a weight loss of about 10%. The main endothermic peak at about 700K, with a weight loss of about 18%, may be attributed to both the decomposition to oxides (mainly of MnC0
and the forming of chromite phases, as confir3) med by the infrared analysis of the discharged samples. It is worth noting that, in all cases, the HY phases decomposed to oxides at a lower temperature than did hydroxycarbonates (23).
758
TABLE 3 Surface area, cumulative pore volume and most frequently occurring pore radius for the catalysts calcined at 623K for 24h. Sample
Cat Cat Cat Cat Cat Cat Cat Cat Cat Cat Cat
1 2 3 4 5 6 7 8 9 10 11
Surface area (m2 g-l)
Pore volume (cm 3 g-1 )
62 124 96 119 94 116 51 65 84 138 88
0.390 0.631 0.337 0.635 0.383 0.406 0.253 0.329 0.887 0.765 0.912
Pore radius (nm) 3.4 8.3 2.5 4.7 3.9 2.9 9.4 4.7 3.1 5.5 10.1
Calcined samples characterization A strong increase of the surface area was observed for the catalysts obtained from HY precursors by calcination at 623K for 24h. The cumulative pore volumes seemed to correlate more to the nature of the elements, evidencing a promoting effect of zinc and cobalt (Table 3). On the other hand, the pore-size distribution curves for all the catalysts obtained from HY precursors showed a narrow peak centred around the most frequently occurring pore radius (Figure 3). In Figure 4, the XRD powder patterns of some calcined samples are shown. Also in this case the nature of the different elements plays an important role. For instance, the cobalt, even if present in a small amount, strongly reduces the compounds, segregation of the oxide phases and favours the formation of spinel~like 3 probably on account of a partial oxidation to C0 + (compare Cat 4 with Cat 6 and Cat 11). Oxidation of the Mn(II) to Mn(III) occurred with formation of a mixed Cu/Mn/Cr phase for Cat 7 during calcination. A further calcination to 673K for 10h completed this reaction, giving rise to a compound with a well defined XRD pattern very similar to that reported for CuCrMn0
(NBS 24-355). 4 The different performance of chromium and aluminum (cf. Cat 4 and Cat 8, ha-
ving the same composition) may be explained by taking into account that, for all the chromium-containing catalysts, the infrared spectra of the calcined samples revealed amounts of amorphous chromates (24). Therefore, it is possible to hypo-
759
_.... 60
IE
-
,..-----,------.-----r-...,
c , I
'01
Fig. 3. Pore size distributions for catalysts obtained from different precursors (full line = Cat 1; dashed line = Cat 10).
M
E E
, I
40
I
,
..
I
I I
~20
I
>
I
\
I
I I , I
,
\
'0
\
O'--_--I..:.-
--'
,
'- , ... , ..;>...,;;;;;::;O&.--'
10
100
F t nm)
Cat 1
•
•
CatS
•
to
20
30
40
50
60
70
28 -
Fig. 4. XRD powder patterns of some catalysts calcined at 623K for 24h.
10
20
30
40
50
60
70
28 -
Fig. 5. XRD powder patterns of the samples heated in vacuum at 900K for 8h (e CuCr0 2; ... spinel-like phase).
760
thesize that during calcination the oxidation of chromium to chromates takes place followed by reduction to chromite, as previously reported for Cat 2 (21). These redox steps reduce the specificity of the nature of the precursors and may favour the segregation of the elements. Catalytic activity The catalytic activities and selectivities in methanol synthesis for some catalysts are reported in Table 4. The rate of methanol formation is referred to the unit weight of both catalyst and copper in order to show the influence of the other bivalent elements on the activity of copper. In the same Table are the compounds identified byXRD analysis after the catalytic tests. An high degree of synteri ng was generally observed for all catalysts, when di spersi on of the samples before the tests was taken into account. The presence of copper oxides may be attributed to the reoxidation by air of a reactive fraction of copper (13). Under these experimental conditions, Cat 2 and Cat 3 showed very low activity, while Cat 9 was active, giving rise only to hydrocarbons. The results showed that, in determining catalytic behaviour, the relative stability of the mixed Cu/Me(II) phase in the working conditions is more important than the nature of the precursor. In fact Cat 5 showed the same catalytic behaviour as Cat 1, despite its being obtained from a HY precursor. These two catalysts were the least active and the only ones for which a CUCr0 phase was iden2 tified after reaction. Similar results were obtained by heating the precipitates at 900K for 8h in vacuum. The XRD powder patterns (Figure 5) for all the samples show a high sintering degree with segregation of different phases, but the main feature is the presence of a CUCr0
phase (marked with a dot) for Cat 1 and Cat 5 2 only, as previously reported for the catalysts after reaction, while the presen-
ce of spinel-like phases was observed for the other samples. It should be noted that our catalytic data for Cat 1 are in good agreement with those reported by Apai et a1. (25) for a CulCr =2.0 catalyst. By assuming a linear correlation of catalytic activity with the pressure, it was possible to calculate from the data of Apai et al. a rate of methanol formation of 0.074 kg 1 h- k9~~t for similar experimental conditions. The absence of correlation between the catalytic activity and the presence of a CUCr0 phase, does not justify its 2 role as active phase proposed by Apai et a1. (25,26). Cat 7 was, like Cat 1, selective in the methanol synthesis and only a little
761
TABLE 4 Catalytic data for the synthesis of methanol and compounds indentified by XRD analysis after reaction (T=533K, P=1.2MPa, GHSV=16500h- 1). Sample
Selectivity (%)
Cat Cat Cat Cat Cat Cat Cat
1 4 5 6 7 8 11
99.3 99.7 99.6 2.7 99.5 98.8 90.0
Rate of methanol formation -1
(kg h k9 0.076 0.166 0.043 0.007 0.051 0.160 0.005
-1
Cat)
(kg h k9
Cu)
0.099 0.423 0.082 0.017 0.121 0.372 0.013
Compounds identified by XRD analysis after reaction CuO, CU20, CuCr02 CuO, ZnO, spi ne1 -1i ke phase CuO, CU20, MgO, CUCr02 CuO, CU20, spinel-like phase Cu, quasi-amorphous phases CuO, Cu, ZnO CuO, ZnO, spinel-like phase
more active. Before reaction, this catalyst showed the presence of a Cu/Mn/Cr phase due to manganese oxidation, but it was not stable in the catalytic test conditions. The absence after reaction of a CUCr0
phase may indicate that the 2 chromium was still interacting with the manganese, forming an amorphous spinel. The presence of zinc strongly increases the activity in the CO coversion in
comparison with that of Cat 1, without modifying the methanol selectivity. Under these experimental conditions, the effect of chromium and aluminum (Cats 4 and 8) was negligible. Cobalt (Cat 6) gave rise to an increase of activity, but with a high selectivity to hydrocarbons. When it was present in a small amount, however, a strong deactivation was observed (cf. Cat 4 and Cat ll)similar to that reported by Lin and Pennella for impregnated catalysts (27). These patterns may be attributed to specific Cu/Zn or Cu/Co interactions connected with the formation of mixed phases which are stable in the working conditions. It should be noted that the presence of mixed phases with a homogeneous distribution of the elements was recently reported by Clausen et al.(28) for Cu/Zn/A1 catalysts on the basis of surface analysis. In this connection, it is worth noting that the role of a Cu/Zn interaction in the methanol synthesis is in agreement with the mechanism proposed by Henrici-01ive and Olive (29). CONCLUSIONS Chromium was better at promoting the formation of HY phases. The addition of a second bivalent element may force the copper to enter into these structures, which were obtained without side phases for Cu/Me(II) ratios
~
1.0. The stabili-
762
ty of the HY phases was related to the nature of the elements present. The HY phases decomposed to oxides at a lower temperature than did hydroxycarbonates; but for chromium-containing compounds, it must be observed that during calcination an oxidation to chromate followed by a reduction to chromite took place, thus reducing the specificity of the nature of the precursors. The Cu/Cr system was active and selective in methanol synthesis and no modifications were observed by the addition of magnesium. These were the only catalysts for which the presence of a CUCr0 2 phase was detected after reaction. By contrast, the presence of other elements (such as Co and Zn) modified strongly the activity andlor selectivity of the Cu/Cr catalyst. We attributed this effect to the formation of mixed phases which are stable in the catalytic conditions. In fact, a CUCr0 phase was not observed either after the catalytic tests or 2 after heating in vacuum at high temperature. ACKNOWLEDGMENT The financial support from SNAMPROGETTI (Milan, Italy) is gratefully acknowledged. REFERENCES
2 3 4 5 6 7 8 9 10 11 12 13
M. Dry, in J.R. Anderson and M. Boudart (Eds.), Catalysis Science and Technology, Vol 1, Springer-Verlag, Berlin, 1981, pp. 159-256. J. Haggin, Chern. Eng. News, 59 (1981) 22-32. M.E. Dry, J. Molec. Catal., 17 (1982) 133-144. I. Wender, Catal. Rev. Sci. Eng., 26 (1984) 303-321. W.H. Calkins, Catal. Rev. Sci. Eng., 26 (1984) 347-358. J.H.C. van Hooff, in R. Prins and G.C.A. Schuit (Eds.), Chemistry and Chemical Engineering of Catalytic Processes, Nato Advanced Study Institute Series, E-39, Sijthoff & Noordhoff, Netherlands, 1980, pp. 599-619. G. Natta, in P.H. Emmett (Ed.), Catalysis, Vol III, Reinhold, New York, 1953, pp. 349-411. H.H. Kung, Cata1. Rev. Sci. Eng., 22 (1980) 235-259. K. Klier, in D.D. E1ey, H. Pines and P.B. Weisz (Eds.), Advances in Catalysis, Vol 31, Academic Press, New York, 1982, pp. 243-312. P. Courty, D. Durand, E. Freund and A. Sugier, J. Mo1ec. Cata1., 17 (1982) 241-254. P. Courty, D. Durand, A. Sugier and E. Freund, British Pat. (1983) 2,118,061A, 28 pp. P. Gherardi, O. Ruggeri, F. Trifiro, A. Vaccari, G. Del Piero, G. Manara and B. Notari, in G. Ponce1et, P. Grange and P.A. Jacobs (Eds.). Preparation of Catalysts III, Elsevier. Amsterdam, 1983, pp. 723-731. S. Gusi, F. Trifiro, A. Vaccari and G. Del Piero, J. Catal., 94 (1985) 120-127.
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14 C. Busetto, G. Del Piero, G. Manara, F. Trifiro and A. Vaccari, J. Cata1., 85 (1984) 260-266. 15 G. Allegra and G. Ronca, Acta Crysta11ogr., A34 (1978) 1006-1013. 16 C. Pierce, J. Phys. Chern., 57 (1953) 149-152. 17 A.F. Wells, Structural Inorganic Chemistry, Clarendon, Oxford, 1975, 1095 pp. 18 R.D. Shannon and C.T. Prewitt, Acta Crystal1ogr., B25 (1969) 925-946. 19 G.J. Ross and H. Kodama, American Mineralogist, 52 (1967) 1036-1047. 20 G.W. Brindley and S. Kikkawa, American Mineralogist, 64 (1979) 836-843. 21 G. Del Piero, M. Di Conca, F. Trifiro and A. Vaccari, in P. Barret and L.C. Dufour (Eds.), Reactivity of Solids, Elsevier, Amsterdam, 1985, pp. 1029-1034. 22 O. Marino and G. Mascolo, in D. Do11imore (Ed.), Proceedings Second European Symposium Thermal Analysis, Heyden, London, 1981, pp. 391-394. 23 C.W. Beck, American Mineralogist, 35 (1950) 985-1013. 24 J.A. Campbell, Spectrochim. Acta, 21 (1965) 1333-1343. 25 J.R. Monnier, M.J. Hanrahan and G.R. Apai, J. Cata1., 92 (1985) 119-126. 26 G.R. Apai, J.R. Monnier and M.J. Hanrahan, J. Chern. Soc. Chern. Commun. (1984) 212-213. 27 F.N. Lin and F. Pennella, in R.G. Herman (Ed.), Catalytic Conversion of Synthesis Gas and Alcohols to Chemicals, Plenum, New York, 1984, pp. 53-63. 28 B.S. Clausen, B. Lengeler and B.S. Rasmussen, J. Phys. Chern., 89 (1985) 2319-2324. 29 G. Henrici-Olive and S. Olive, Catalyzed Hydrogenation of Carbon Monoxide, Springer-Verlag, Berlin, 1984, pp. 131-141.
764
DISCUSSION XU Xiaoding : What is the selectivity of Cat. 6? Are there C1-C2 alcohols formed? Could you comment on the oxidation states of Co and Cr of Cat. 6 after reduction? A. VACCARI : Catalyst 6 showed a selectivity of 97.3% in hydrocarbons (paraffins and olefins) and of 2.7% in methanol. In the reaction conditions reported and without doping no higher alcohols were observed. After calcination the XRD powder patterns of Cat. 6 showed the presence of only a sDinel-like phase, in which cobalt was present partially as CollI and with a CrV1 content comparable with the values reported for the Zn/Cr systems (Del Piero et al., in P. Barrett and L.C. Dufour (eds.), Reactivity of Solids, Elsevier, Amsterdam, 1985, 1029-1034). Such Cr VI species were, at least in part, located on the surface. After reduction, no spectroscopic evidence of CrVI was found. The spinel-type phase showed an evolution towards a rock-salt type structure, that may be explained assuming the reduction of CollI to CoIl and taking into account the higher ionic radius of the latter. No free metallic cobalt and/or cobalt oxides were observed (Fornasari et al., I.E.C. Submitted). P. CHAUMETTE : What is your own opinion on the effect of the cobalt content on CO hydrogenation activity of Cu-Co catalysts? A. VACCARI : The presence of small amounts of cobalt in the catalysts for lowtemperature methanol synthesis has a drastic deactivating effect, without change of selectivity. Increasing the cobalt content, an increase of activity was observed, with a synergetic effect between cobalt and copper correlable to the presence of a non-stoichiometric spinel-type phase (Fornasari et al., I.E.C. submitted). On the other hand, according to what has been reported by other authors (for instance, Courty et al., J. Mol. Catal. 17, 1982,241-254), the preparation procedure, the alkalinization, the activation and conditioning steps havea determining effect and may modify strongly the final activity and selectivity. ZHAO J. : What is the relationship between activity, composition and precursor structure? The content of aluminum is extremely high in your experiments. Can you get higher activity if lowering the Al content? A. VACCARI : In a previous paper, the correlation between catalyst composition and catalytic activity in the low-temperature methanol synthesis was reported in detail (Gusi et al., J. Catal. 94, 1985, 120-127). The choise of a particular Cu/Zn ratio depends mainly on tne gas mixture composition, but in all cases aluminum in concentration higher than 10% inhibits the activity. With the stoichiometric H2/CO ratio, the maximum of activity, expressed both per kilogram or liter of catalyst, was observed for a Cu/Zn ratio = 2, i.e. for samples obtained from mixed precursors. On the other hand, according to other data presented in this Symposium (Doesburg et al.), the hydrotalcite-like phases are catalyst stabilizers, favouring the formation of small and stable oxide crystallites (Busetto et al., J. Catal. 85, 1984,260-266). E.B.M. DOESBURG : The effect of Al on the methanol synthesis is quite dramatic. Do you see the same effect if you plot the copper surface area as a function of the aluminum content? A. VACCARI : We did not observe a correlation between the copper surface area and the activity in the methanol synthesis, while a good fit was observed both with the sum and the product of the amount of CuO and undetected copper, determined by XRD, in the spent catalysts (Gusi et al., J. Catal. 94, 1985, 120-127). of the Furthermore the effect of aluminum cannot be explained on the~asis copper surface area, but may be better correlated to the amount of copper interacting in a 1:1 atomic ratio with aluminum and presenting a low reactivity
765
towards the reduction (Gusi et al., Reactivity of Solids, in press). This interaction is responsible for the small and stable copper crystal size and for the absence of surface acidity also for the catalyst with 31% aluminum content.
J. BARRAULT : The methanol or alcohols synthesis needs one catalyst and Z or 3 reagents (CO, C02 and HZ)' You tried to establish some relation between catalytic activity and catalyst composition, but do you have some information about the variations of reagent activation, especially CO and HZ, when you change the composition or the structure of your catalysts? A. VACCARI : We have not investigated specifically this point. However, the catalytic behaviour in the methanol synthesis was determined using two different HZ/CO ratios: the stoichiometric ratio and a higher one employed industrially in the recycl ing loop. In all cases, only cobalt (with a change of selectivity) and zinc modify markedly the activity of the copper; furthermore, an excess of zinc with respect to the copper gives rise to less active catalysts, in agreement to that reported in this Symposium by Rasmussen et al .. The differences of activity in the methanol synthesis were more evident with the stoichiometric HZ/CO ratio than with the higher one, showing an increase of activity attributable to the zinc and a maximum of activity with a Cu/Zn ratio = Z (Gusi et al., J. Catal. 94, 1985, 120-lZ7). On the basis of the data of Agny and Takoudis (I.E.C., Prod. Res. Dev., 24, 1985, 50-55), the hypothesis reported by HenriciOlive on the role of the ilii"c ("Catalyzed Hydrogenation of Carbon Monoxide", Springer-Verlag, Berlin, 1984, Ch. 8) and that on the role of copper particle size (Doesburg et al ., this Symposium) may account for the differences in the catalytic activity.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
767
PREPARATION AND CHARACTERIZATION OF COPPER/ZINC OXIDE/ALUMINA CATALYSTS FOR METHANOL SYNTHESIS E.B.M. DOESBURG, R.H. HOPPENERa, B. de KONING, XU XIAODING and J.J.F. SCHOLTEN b Laboratory of Inorganic and Physical Chemistry, Delft University of Technology, P.O. Box 5045, 2600 GA Delft, The Netherlands apresent address: DSM, Research and Patents, P.O. Box 18, 6160 MD Geleen, The Netherlands. bAl so affiliated to DSM, Research and Patents, P.O. Box 18, 6160 MD Geleen, The Netherlands.
ABSTRACT A series of Cu/ZnO/A1203 catalysts for the low pressure methanol synthesis has been prepared by coprecipitation with a sodium carbonate solution from solutions of a mixture of the corresponding metal nitrates, followed by ageing in the mother liquor, drying, calcination and reduction. The catalysts and their precursors were analyzed by techniques like X-ray diffraction (XRD), X-ray line broadening (XLB), differential scanning calorimetry (DSC), chemical analysis, adsorptive decomposition of N20 and B.E.T.-measurements. The catalytic activity for the methanol synthesis was determined in a flow reactor under industrial conditions. Depending on the ratio of the metal ions in the initial metal nitrate solutions different compounds were formed during coprecipitation, like rosasite, (Cu, Zn)2(OH)bCOO' malachite, CU2(OH)2C03' Cu, Zn-hydrotalcite, which we called rOd Its (Cu, Zn)6A12C 3( H)16.4H20 and a ternary compoun~ structure is uoknown and it contains, be Zn +, up to 28 at. % Cu + and up to 17 at. %A1J+. Addition of 7 at. %Mg 2ides + stabilizes the Cu, Zn-hydrotalcite structure but leads to a drastic decrease in catalytic activity. The rate of methanol production depends on the phase composition of the precursors. Rosasitecontaining precursors give the highest activity; hYdrotalcite proves to be an excellent catalyst stabilizer which evokes the formation of small Cu and ZnO particles. MgZ+ inhibits methanol production.
2rite.
INTRODUCTION For many years Cu/ZnO/A1 203 catalysts have been used in industry for the low pressure methanol synthesis, operating at 70 bar (= 7 MPa) and 2200C-2400C (ref. 1). The life time of these catalysts is limited due to sintering of both copper and zinc oxide; the decrease of the free-copper and total surface area after three years' use in the reactor is of the order of a factor of two. To maintain the methanol production at a high level, loss of activity with time is finally compensated by gradually increasing the pressure and temperature. In doing so industrial catalyst life time is ca. 3-4 years on an average. Besides the demand for a stable activity the fuel-methanol catalysts have to satisfy also the requirement for stable selectivity towards higher alcohols
768
(refs. 2-4). The most promising preparation method for such multicomponent catalysts to satisfy these requirements is the coprecipitation method by which a precursor is made where all the elements are in the same structure. Before starting the preparation of modified methanol synthesis catalysts for the production of fuel-methanol, we first investigated the preparation of the classical Cu/ZnO/A1 203 catalysts. The industrial catalysts are made by coprecipitation techniques (ref. 1). Busetto et al. (ref. 5) and Gusi et al. (ref. 6) investigated the preparation of the precursors for the Cu/ZnO/A1 203 catalysts by an inverse coprecipitation technique. They found that a (Cu,Zn).phase .1S formed for 2.2 ~ (Cu 2+ + Zn 2+ )/Al 3+ ~ 3.2 and Cu 2+ /Zn2+ hydrotalc1te 1.0 or 0.5. Outside this range of metal ion ratios, other compounds like rosasite, (Cu,Zn)2(OH)2C03' and malachite, Cu2(OH)2C03' are formed. In the present investigation we prepared the precursors of the Cu/ZnO/A1 203 catalysts by coprecipitation at constant pH and temperature. We examined the effect of metal ion ratios. the precipitation temperature, the ageing temperature and Mg 2+ addition, on the composition of the precursor phase. Chemical analysis by Atomic Absorption Spectrometry (AAS). X-ray diffraction (XRD). high temperature X-ray diffraction. X-ray line broadening measurement (XLB). differential scanning calorimetry (DSC), chemosorptive N20 decomposition and N2-physisorption were used to characterize the products formed after coprecipitation. calcination and reduction. The catalytic activity for methanol synthesis was determined and related with the composition of the precursor phase and the catalyst preparation conditions. EXPERIMENTAL Catalyst preparation The catalyst precursors were prepared by coprecipitation at constant pH and temperature. A 1.0 1 aqueous solution of copper,zinc and aluminium nitrate(0.5M) and an aqueous solution of sodium carbonate (1.0 M) were added simultaneously with constant stirring to 0.5 1 of demineralized water. The rate of addition was 0.5 ! 0.005 l/hr. The temperature was maintained at 600C. 700C. 800C or 900C. respectively. with an accuracy· of! 0.4 0C. The pH was kept constant during coprecipitation of 7.0 ! 0.005 by adjusting the rate of adding the sodium carbonate solution. After precipitation the coprecipitate was aged in the mother liquor for one hour at a temperature of 600C, 70 0C or 800C. respectively, with an accuracy of ! 0.4 0C. After ageing the coprecipitate was filtered on a G-4 glass filter and thoroughly washed with 2.5 1 of demineralized water with the same temperature as that during ageing. The precipitate was dried at 75 0C. All chemicals used were analytical grade. SamplesM-5, K-9, K-ll. K-7, 0-1 and 0-4 are prepared by a slightly different coprecipitation method described elsewhere (ref. 7 ). During preparation of sample M-5 the metal nitrate solution
769
contained also magnesium nitrate. Sample K-12 was precipitated at 60°C ~ 50C and subsequently kept overnight in the mother liquor at room temperature before filtration and washing. Dried coprecipitates were calcined overnight in air at 350°C. Reduction was performed in situ in 5 vol. % H2 in N2 according to the following reduction programme: - heating up to 2400C at a heating rate of 0.5 0C/min.; - reduction at 2400C for 8 hours; - gasflow 1.2 l/hr. Catalyst characterization X-ray diffraction (Enraf-Nonius, Guinier-De Wolff camera mark II) was used to determine qualitatively the phase composition of all the samples. The decompositon of the coprecipitated precursors in N2 at a heating rate of l oC/min. was studied in a Nonius Guinier-Lenne camera. In situ reductions were also performed in the same camera. Copper crystallite sizes were calculated from the line broadening of the CU(lll) diffraction line using the Scherrer equation (ref. 8). Diffractograms of calcined precursors were recorded with a Siemens D500-B goniometer between 29 _. < 28 < 42 degrees. CuK a radiation was used in all cases. Differential scanning calorimetry (DSC) was used to investigate the decomposition of the precursors in air at a heating rate of 100C/min. The free copper surface area was determined applying the N20-method (ref. 9) in a closed volumetric system (ref. 7). After ending the N20 chemosorptive decomposition the then passivated sample was transferred and used for XRD and XLB. The B.E.T.surface areas of the reduced catalysts were calculated from nitrogen adsorption isotherms at -196oC, taking 0.162 nm 2 for the cross-sectional area of an adsorbed nitrogen molecule. The catalytic activity for the methanol synthesis was measured in a microreactor at 70 bar (= 7 MPa), at 2400C with a space velocity of 35.000 h- 1. The composition of the feed gas was CO/C0 2/H2/inert: 4/3.5/82/10.5 (v/v). The catalyst volume was 3.7 ml, the catalyst particle diameter between 1.2 and 3.5 mm and the measurement was carried out outside the region of diffusional retardation. RESULTS AND DISCUSSION Coprecipitates In table 1 the X-ray diffraction results of the fresh coprecipitates after one hour of ageing are given. The ZM 2+/A1 3+ ratios based on the weights of chemicals used and also those calculated from the results of chemical analysis (A.A.S.) are given in two separate columns. The difference between these two values can be explained by the fact that not all the A1 3+-ions are precipitated during the coprecipitation as was shown by analysis of the clear filtrate after filtration of the precipitate. The samples of which 2.1 2 ZM 2+/A1 3+ ~ 3.1 give
770
TABLE 1 Analysis results of the coprecipitates. Sampl e Code
Composition T p6ec. C Cu/Zn/Mg/Al in at. %
B-1 B-3 B-4 B-5 B-8 B-9 B-lO B-13 M-5 K-9 K-ll B-14 B-15 B-18 B-25 B-26 B-29 K-lO D-9 D-8 K-12 K-7 D-1 D-4
30/32/0/38 II
34/34/0/32 II
35/35/0/30 II
32/37/7/29 36/36/0/29 48/24/0/29 37/37/0/26 II II
49/24/0/26 II
38/38/0/25 50/25/0/25 25/51/0/24 58/24/0/18 28/55/0/17 56/27/0/17 67/33/0/00
60 80 60 60 80 60 60 80 80 80 80 60 60 80 60 60 80 80 80 80 80 80 80 80
TAg5ing C
60 80 60 80 80 60 80 80 80 80 80 60 80 80 60 80 80 80 80 80 80 80 80 80
EM 2+/M3+ Theor.
From A.A.S.
1.6 1.6 2.1
2.1 2.7
II
II
2.3
2.9
II
II
2.6 2.5 2.5 2.8
3.1 3.0
II
2.8 II
3.0 3.0 3.2 4.6 4.9 4.9
3.7 3.7 3.1 3.8
X-ray analysis Precursor Phases* H H H H H H H H H H H + (Ros. ) H+(Ros.) H + Ros. H + Rod. H + Ros. H + Ros. H + Ros. H+ (Ros.) H+ (Ros.) Rod + (H) Ros. + H Rod. Ros. + (H) Ros.
~H = hydrotalcitej Res. = rosasite. Rod.= roderite, minor amounts indicated by
brackets.
only the (Cu,Zn)-hydrotalcite compound as is expected (ref. 10) and confirms the results of Gusi et al. (ref. 6). Depending on the cation ratios other compounds are fonned as well (ref. 7). For Cu 2+ /Zn2+ .:: 1.0 and (Cu 2+ +Zn 2+)/Al 3+ > 3.1 rosasite, (Cu,Zn)2(OH)2C03' together with (Cu,Zn)-hydrotalcite are formed. At the Zn 2+-rich side and for (Cu~+ + Zn 2+)/A1 3+ > 3.1, a ternary compound which we called roderite, with an unknown structure but with a characteristic X-ray diffraction pattern, is formed together with (Cu,Zn)-hYdrotalcite. For Cu 2+/Zn2+/A1 3+ = 28/55/17 at. %, for Zn 2+/A1 3+ = 90/10 at. %and for Zn 2+/A1 3+ = 87.5/12.5 at. % (ref. 7) the pure roderite phase is fonned. This compound was detected earlier by Ketchik et al. (ref. 11) and the X-ray diffraction pattern is similar to the diffraction pattern of hydrozincite. A pure hydrotalcite compound is also formed for (Cu 2+ +Zn 2+ +Mg 2+)/A1 3+ =3.1 and with 7 at. %Mg 2+ added. We prepared a series of coprecipitates with Mg 2+ added. Although the addition of Mg 2+ does not strongly influence the region of
771
existence of the pure hydrotalcite compound, it has a strong influence on the ratio hydrotalcite/rosasite in favour of hydrotalcite. As can be seen in table 1, in most cases there is hardly any effect of the precipitation temperature or ageing temperature on the composition of the coprecipitate phase as revealed by X-ray diffraction. An exception was observed in the series of Cu 2+/Zn 2+/A1 3+ = 37/37/26 at. %, the coprecipitates 60/60 (precipitated at 600C and aged at 60oC) and 60/80 (Cu,Zn)-hydrotalcite together with rosasite is formed but the coprecipitates 70/70, 70/80 and 80/80 of the same series contain roderite instead of rosasite together with (Cu,Zn)-hydrotalcite. The XRD results are confirmed by the DSC results. In figure 1 the DSC patterns of the pure compounds CU,Zn-hydrotalcite (B-1), rosasite, (Cu,Zn)2(OH)2C03 with Cu 2+/Zn2+ = 4 and malachite Cu2(OH)2C03 are presented.
B-1
236
137
E ~
Cu/Zn:4
190 267
~
~
w
377
1
365
Malachite
314
0
500 T(·C)~
Fig. 1. DSC-results of the pure compounds. 2+ 2+ a. (Cu,Zn)-hydrotalcite (B-1). b. rosasite (Cu,Zn)2(OH)2C03 (Cu /Zn = 4) and c. malachite Cu2(OH)2C03' The numbers written in the DSC pattern indicate the positions of the peaks or shoulders in degrees centigrade. The DSC pattern of the Cu,Zn-hydrotalcite is in accordance with the DTA results of Myata (ref. 12) for Mg-hydrotalcite, but peaks are shifted towards lower temperatures.
772
382
8-14
249 138
8-4 252 8-18
I
.~
145
E
:
0x
W
8-5 257
~
w
138
147 8-25 230
362
B-8
138 145 T('C)-
500
Fig. 2. DSC results of samples B-14, B-18 and B-25.
0
T('C)_
500
Fig. 3. DSC results of samples B-4, B-5 and B-8.
In fig. 2 the DSC patterns of B-14, B-18 and B-25 are presented. B-14 and B-25 show the hydrotalcite pattern together with a small peak between 3400C and 365 0C which belongs to rosasite. Sample B-25 appears to contain more rosasite than B-14 as is also confirmed by X-ray diffraction results. The DSC pattern of sample B-18 is quite different from that of sample B-14 and the large peak at about 3000C can be ascribed to the decomposition of roderite. This is confirmed by the results of high temperature X-ray diffraction which shows a change in the diffraction pattern from roderite to AMOM (X-ray Amorphous Metal Oxide Mixture) at the same temperature. The effect of the temperature during coprecipitation and ageing on the composition of the coprecipitate is shown in fig. 3 which presents the DSC patterns of samples B-4, B-5 and B-8. All three coprecipitates have the hydrotal cite structure as is shown by X-ray diffraction. B-4 and B-5 are precipitated at 600C but B-4 is aged at 600C and B-5 at 800C. We see no difference in the DSC patterns so it appears that there is no effect of the ageing temperature on the composition of the coprecipitate. B-8 is precipitated at 800C and aged at 80°C. The DSC pattern of B-8 is different from the pattern of B-5 so there is an effect of the precipitation temperature on the composition of the coprecipitate although this is not confirmed by X-ray diffraction as can be seen in table 1. However as XRD only shows the identity of crystalline materials B-8 may contain an amorphous phase together with hydrotalcite. By comparing the DSC pattern of
773
B-8 with the DSC pattern of B-18 in fig. 2 we conclude that this amorphous phase is most probably amorphous roderite. Calcined precursors The compounds detectable in the calcined precursors depend strongly on their phase compositions before calcination. Roderite and hydrotalcite gave after calcination an AMOM as can be seen in fig. 4 for sample K-9.
2800.0 .....------,--......,.--....,...---,....---,---..,...---,....-----,,...-----.-----,
K-11 2440.0
.'
2080.0
CPS
t
....
1720.0
1360.0
·····..··········-:······:·..···CUO '" ASTM
"'ZnO
30.000
-
Fig. 4. Diffractograms between 27 overnight in air at 3500C.
38.000
<
28
<
39.000
42.000
42 degrees of K-9 and K-l1 calcined
This AMOM has, for the Zn-rich samples, an XRD pattern similar to ZnO with broad diffraction lines which are shifted. A precursor containing rosasite and hydrotal cite (K-ll) gave after calcination a mixture of AMOM and CuO (see fig. 4). This AMOM is stable up to about 500 0C above which temperature ZnO and CuO are formed as was confirmed by high temperature X-ray diffraction. The X-ray diffraction pattern of sample K-ll (Cu/Zn/Al = 47/23/29 at. %), of which the dried precursor phase consists of a mixture of hydrotalcite and rosasite, is that of AMOM and of superimposed CuD originating from rosasite (fig. 4). The diffraction lines of CuD are slightly shifted with regard to the ASTM-values of CuD.
774
Reduced catalysts As can be seen in (ref. 9) of the NZO decomposition of NZO equation: ZCuS + NZO ~ (CuZO)S
fig. 5 the reduced catalysts do not show a sharp endpoint chemosorptive decomposition due to a slow continuous after a fast one. During the fast process according to the +
NZ
a strongly held monolayer of CuZO is formed. But in the presence of ZnO or A1 Z03 a slow oxidation of bulk copper will take place as was shown by Evans et al. (ref. 13).
-_.--- -------"...:--=--------......--~K-12 K-l0
..
• K-ll
N 20-Method 90t
so
-
Torr H
20
Tlme(hr)
°
Fig. 5. Results of NzO-chemosorptive decomposition on reduced catalysts as a function of time. Straight lines give the extrapolation to t = hr to estimate the free copper surface areas. Assuming that the formation of the oxygen monolayer corresponds with the fast part of the NzO decomposition, we can roughly estimate the free Cu-surface area by extrapolation to t = hr as demonstrated in fig. 5. By taking a croSS sectional area of 0.058 nm Z for each CUzO we calculated the free Cu-surface
°
775
areas which are listed in table 3. After the Cu-surface area determination the then passivated sample was transferred and was used for XRD and XLB. The Cuparticle sizes of the catalysts calculated by using the Scherrer equation are listed as dCu in table 3. By using a half-sphere model the average copper particle sizes were also calculated from the copper surface areas, and they are listed in the fifth column of table 3. There is a reasonable agreement for samples M-5, K-10, D-9, D-8 and K-12. The discrepancies for other samples can be explained either by the fact that the XLB method is restricted to particles bigger than 3.0 nm or/and that there may be different particle size distributions which are very dependent on the method of preparation. TABLE 2 Characterization results of some reduced catalysts. Sample code
Composition Cu/Zn/Mg/Al
Tpoec. C
~. gelng
B-4 B-5 B-8
34/34/0/32 34/34/0/32 34/34/0/32
60 60 80
60 80 80
Precursor phases
S~ET
S~u
m /gcat m /g Cu
°C
H H H
85.9 43.3 37.7
13.0 13.6 16.5
Another reason, which is also stated by Gusi et al. (ref. 6), can be the large interaction between support and the Cu particle for catalysts originating from the ternary precursors hydrotalcite or roderite. The development of Cuo-exposed surface is different per unit weight of copper (between 97.5 m2/g cu for M-5 and 23.2 m2/g cu for K-11), which is very much dependent on the method of preparation, and the type of precursor and the pretreatment of the catalysts . For example a different temperature of coprecipitation or/and ageing could indeed influence the total B.E.T. surface area as well as the Cu surface area. This is demonstrated in table 2 for samples B-4, B-5 and B-8. There is a large effect of the ageing temperature on the B.E.T. surface which ismostprobably caused by recrystallization of the fresh precipitated species during ageing. Activity measurements The results of the microreactor experiments are presented in table 3, where the activity is expressed per gram of catalyst. From a technical point of view this is an important quantity and sample K-12 appears to be the best catalyst in this respect. Both samples M-5 and K-9 stem from a monophasic hydrotalcite precursor and their atomic compositions are the same, except for the replacement of 4 at. %of Cu and of Zn by 7 at. % Mg in M-5. Due to this difference sample M-5 exhibits an
..., ...,
TABLE 3 Some properties of the catalysts.
Q')
d
Sample code
Composition Cu/Zn/Mg/Al
SBET 2 m /gcat
S 2 Cu m /gcat
CY from N20
M-5 K-9 K-ll K-10 0-9 0-8 K-12 K-7 0-1 0-4
32/32/7/29 36/36/0/29 48/24/0/29 38/38/0/25 50/25/0/25 25/51/0/24 58/24/0/18 28/55/0/17 56/27/0/17 67/33/0/0
50.8 60.7 65.0 96.0
31.0 13.1 10.9 20.1 18.3 14.4 21.0 14.8 24.2 16.6
3.5 8.8 14.5 5.9 9.0 5.3 9.0 5.7 7.5 12.4
------59.0 317.0
-------
H = hydrotalcite; Ros. = rosasite; Rod. = roderite. Minor amounts indicated by brackets.
(nm) from XRO 3.3 4.0 9.9 6.6 11.5 4.8 8.0 3.8 10.8 20.5
Activity gMeOH/gcat· h gMeOH/gcu· h 0.24 0.97 1.1 1.1 1.2 0.90 1.9 0.79 1.45 0.98
0.75 2.84 2.34 3.09 2.47 3.95 3.38 3.16 2.70 1.60
Precursor phases
H H H + (Ros.) H + (Ros.) H + (Ros.) Rod. + (H) Ros. + H Rod. Ros. + (H) Ros.
777
activity which is five times lower than found for K-9. This proves the strong suppressing effect of Mg 2+ on activity and it points to a direct involvement of the oxidic part of the catalyst surface in the synthesis of methanol, the magnesium ions being dissolved in the lnO. Another possibility is that, due to the presence of Mg-ions in the oxidic support, the copper crystallites are less well fixed by which a more rapid sintering of the copper crystallites under testing conditions occurs. Gusi et al. (ref. 6) pointed out that the rate of methanol formation per unit catalyst weight and at a constant A1 3+ content of 24 at. %, gradually increases with increasing Culln ratio (see fig. 2 in their paper, for a H2/CO/C02 ratio of 86:8:6). Our results show the same trend: 0-8 ,Culln ratio 0.5, activity 0.90 K-10 , 1.0, 1.10 0-9 , 2.0, 1.21 Furthermore Gusi et al. (ref. 6) found that both an inhibiting effect and an activating effect of A1 3+ are possible, depending on the Cu/ln ratio, in accordance with the results of Shimomura et al. (ref. 4). Our results are in accordance with this: at a Culln ratio of 0.5 (compare samples K-7 and 0-8) A1 3+ has an activating effect, whereas at a Culln ratio of 1 (compare K-10 and K-9) an increase of the A1 3+ content works inhibitive. At a Culln ratio of 2.0 (compare 0-4, 0-9 and K-11), at 25 at. %A1 3+ an activating effect is found, whereas a further increase to 29 at. % Al3+ inhibits the reaction rate. In fig. 6 the activities of the catalysts, expressed per m2 free copper surface area, are plotted as a function of the average copper particle size, dw' It is seen from this figure that seven out of ninp catalysts show a nearly (excluding the Mg 2+-containing constant activity of about 0.06 g CH30H.mc~.h-1 sample M-5). This value is twice as high as the value of 0.03 g CH30H.mc~.h-1 which we calculated from fig. 4 of the paper of Chinchen et al. (ref. 15 ). Samples K-11 and K-12 are the exception, their activities being 1.5 and 1.7 times higher respectively. No direct explanation for this is at hand. We speculate that these samples, which both have a bi-phasic precursor of rosasite and of hydrotalcite, are stabilized by the presence of zinc aluminate, originating from the hydrotalcite, and highly active due to the presence of an A1 3+-free component, originating from the rosasite precursor. Another plausible explanation can be given by looking at fig. 7, which gives the initial methanol production rate per gram of copper (also listed in table 3) as a function of the average copper particle size dCu(111) ' Now it appears that there is a maximum in the activity at a copper particle size of about 7 nm. This can be explained by the results of Van Hardeveld et al. (ref. 16) and Pritchard (ref. 17) by assuming that the dissociative chemisorption of H2 can be activated on rough planes (containing high index faces) of the copper
778
•
0.10
K-11
• K-12
K-9. .0-8
t
.0-9
• •
0.05
K-7
•
0-1
0-4
K-10
M-5
Ol-------L....-------I~---
o
10.0 -
d"CU(111)
20.0
[nm]
Fig. 6. Initial methanol production rat~ per_m 2 of free copper surface as a function of the average particle size (dw = dCu(lll))' particles. Van Hardeveld proved that Ni (having a f.c.c. structure like Cu) particles between 7.0 and 8.0 nm have an optimal exposition of high indexed faces. By taking into account the inaccuracy of the XLB method the optimal value of 7 nm is quite reasonable. The low activity of the Mg 2+ containing catalyst can be explained by a too small copper particle size, or it can be a sintering effect as stated before. CONCLUSIONS 1. The ternary precursors hydrotalcite and roderite give catalysts with very small copper particles. 2. In most cases all the copper surfaces seem to be equally active for methanol synthesis (0.06 g CH30H/m~u.h ). Two catalysts, however, with a bi-phasic rosasite/hydrotalcite precursor show a relatively high activity per unit free copper surface area. Most probably the r.osasite precursor supplies the more active part of the catalyst. 3. The catalytic activity expressed per gram of copper shows an optimal activity at a copper particle size of about 7 nm which might be explained by the
779
4
, I
,
~,
I
\
I •
. hr
g
I
K,!.7
~-9
Cu
, I
,.
,
K-12
\
•
\
\ K-1<\
.
". " 0-1
, I
,
• ',0-9 K-1'1", .. , ....
2
........
'- -
"-4
•
M-5
O~-------~------_.....I-_-
o
10
ii
20 (nm)
Cu(11\)
Fig. 7. Initial methanol pro~uction rate per gram of copper as a function of the average particle size (aw = dCu(lll))' higher activity of the high index faces of the copper particles for methanol synthesis. 4. The addition of Mg 2+-ions dramatically suppresses the catalytic activity. 5. Further studies are necessary to discriminate between both hypothesis whether there is an effect of the precursor composition on the catalytic activity or there is a copper particle size effect. ACKNOWLEDGEMENTS The authors wish to thank Drs. W. Glasz and Dr. E.C. Kruissink from DSM for measuring the activities for methanol synthesis and for valuable discussions. Mr. J. Teunisse for the surface area measurements. Mr. B. Sonneville for preparing some catalyst precursors and Mr. N. van der Pers and Mr. J. van Lent for XRD-measurements. REFERENCES 1 J.T. Gallagher, J.M. Kidd, Brit. Pat. 1159035 (1965). D. Cornthwaite, Brit. Pat. 1296212 (1969). B.M. Collins, Brit. Pat. 1405012 (1972). 2 M.l. Greene, A.I.Ch.E., C.E.P., Aug. (1982) 46.
780
3 Ph. Courty, J.P. Arbe, A. Convers, P. Mikikenko and A. Sugier, Hydr. Proc., Nov. (1984) 105. 4 Ph. Courty, D. Durand, E. Freund and A. Sugier, J. Mol. Catal., 17 (1982) 24l. 5 C. Bussetto, G. Del Piero, G. Manara, F. Trifiro and A. Vaccari, J. Catal., 85 (1984) 260. 6 S. Gusi, F. Trifiro, A. Vaccari, G. Del Piero, ibid., 94 (1985) 120. 7 R.H. Hoppener, E.B.M. Doesburg and J.J.F. Scholten, Appl. Catal., (1986) to be published. 8 B.E. Warren, "X-ray diffraction", 1966, 253. 9 ,J.J.F. Scholten and J.A. Konvalinka, Trans. Faraday Soc., 65 (1969) 2465. 10 R. Allmann, Chimica, 24 (1970) 99. 11 S.V. Ketchik, L.M. Plyasova, T.M. Yurjeva, L.I. Kuznetsova, T.P. Minuyukova, Izv. Sib. Otd. Akad. Nauk. SSSR, Ser. Khim. Nauk., 6 (1983) 36. 12 S. Myata, Clays and Clay Miner., 28 (1980) 50. 13 J.W. Evans, M.S. Wainwright, A.J. Bridgewater, O.J. Yound, Appl. Catal., 7 (1983) 75. 14 K. Shimomura, K. Ogawa, M. Oba and Y. Kotera, J. Catal., 52 (1978) 191. 15 G.C. Chinchen, K.C. Waugh and D.A. Whan, Appl. Catal. (1986) to be published. 16 R. van Hardeveld and F. Hartog, Adv. Catalysis, 22 (1972) 75. 17 J. Pritchard, T. Catterick and R.K. Gupta, Surf. Sci., 53 (1975) 1.
781
DISCUSSION Y. OGINO : You showed the MeOH forming activity of your catalysts as a function of crystallite size and of surface area of metallic copper. 00 you consider that the metallic copper is responsible for the methanol synthesis activity? E.B.M. OOESBURG : We indeed think, in accordance with research workers from the ICI "New Science Group" in England (see our ref. 15) that methanol synthesis runs exclusively over the surface of the copper crystallites. On that copper surface both Cu(O) and CU(I) sites are necessary to catalyse the reaction. The main function of lnO seems to be the strong fixation of the copper crystallites. E.K. POELS : At Unilever Research we also found that activity of rrethanol synthesis catalysts is proportional to the copper metal dispersion as measured by CO chemisorption. However, from literature it is known (e.g. Trifiro et a1., Whan et a1.) that both metal and copper oxide are necessary for methanol synthesis and as it is not unlikely that the amount of reoxidisab1e copper is proportional to dispersion; could you please elaborate on what you mean exactly by stating that activity is related to the metallic copper? E.B.M. OOESBURG : The linear relationship between the catalytic activity and the initial copper surface area is an indication that all copper atoms at the surface have the same activity for the methanol synthesis reaction; copper surface only is involved in the reaction. We agree that part of the initial copper surface is oxidized under reaction conditions as proven by Chinchen et a1. (see our ref. 15). The degree of oxidation of the copper surface depends on the feed gas composition. In all activity tests this feed gas composition was the same and hence this factot will not influence the linear relationship between activity and copper surface area. However, there will be an influence on the slope of the line, as indicated in our paper. Note added to the discussion by one of the authors : J.J.F. SCHOLTEN: I think that research workers have to be encouraged to plot m~thanol activity : a) per gram of catalyst, b) per gram of copper and c) per m free-copper surface area. Activity per gram of catalyst is, of course, an important technical quantity. The activity per gram of copper learns s~mething about the "copper economy" of the process. Finally. the activity per m free-copper surface area is of high scientific significance; the ICI "New Science Group" found strong evidence that the activity is strictly proportional to the extent of the free-copper surface area, independent of the type of support. Furthermore, more attention is needed, for the influence the presursor composition might have on catalyst stability. For this purpose we need an accelerated stability test, as the industrial life-time is nearly three years. B.S. RASMUSSEN: The new extra phase which you call roderite, you describe to be a compound which contains Cu, In and A1 and to have a X-ray pattern which is similar to the pattern of hydrozincite. Could it not be possible that these diffraction lines really originate from a hydrozincite-like structure and the Al is present in an amorphous form not detectable by X-ray diffraction? Therefore I ask you if you have support for the presence of Al in this roderite phase - for instance from selected area analysis in the electron microscope. E.B.M. DOESBURG : The atomic composition of the ternary compound which we called roderite was determined by chemical analysis. The X-ray diffraction pattern of this compound ressembles the pattern of hydrozincite but it contains some extra strong lines which cannot be ascribed to hydrozincite. The same pattern is found for compounds with the atomic compositions In/A1 = 90/10 at %, and In/Al = 87,5/11,5 at %. We have no direct evidence that the Al-part is not in an amorphous form. However, compounds with about the same amount of Al present but with other Cu/ln ratios (K-12 and 0-1) give hydrotalcite together with rosasite. The roderite compound is prepared in the same conditions so we do
782
not expect a separate amorphous alumina phase. H. SCHAPER: It is well known that the properties of coprecipitated compounds depend on precipitation parameters such as pH and temperature. Have you investigated the effect of such variations ? E.B.M. DOESBURG : We did not investigate the effect of the pH on the nature of the compounds formed during precipitation contrary to the effect of temperature during coprecipitation and ageing. The temperature during coprecipitation and ageing was varied between 60°C and ao°c. There was hardly any effect of the temperatur on ~he phase composition of the coprecipitate except in the series of CuZ+/Zn2+/Al + = 37/37/26 at % (see table I of our paper). In this series coprecipitation at 70°C and ao°c gives roderite instead of rosasite together with hydrotalcite. The effect of the temperature on the BET-surface area and Cu-surface area is more pronounced as can be seen in table 2 of our paper. We believe that the large effect of the ageing temperature on the BET-surface area is most probably caused by recrystallization of the fresh precipitated species during ageing. B. NOTARI : You find that the relation between the surface area of copper and the catalytic activity exists for all catalysts but the MgO containing one: the hypothesis you advance to explain this irregularity is that a fast sintering of the copper occurs on the catalyst during the first hours on stream. Since in my experience this fast sintering occurs only under special conditions and away from conditions under which these materials are normally used, I would like to contribute to the solution of the problem by proposing a second hypothesis, namely that the MgO hinders the accessibility of the reactants to the surface of the copper particles. This could arise as a consequence of the different behaviour of MgO with respect to ZnO during the reduction step : ZnD would move away easily and copper would form small aggregates, while MgD would find it difficult to move, thus covering the copper particles while they are forming and constituting a barrier to the reagents. When in such a system the measurement of the surface area of copper is carried out, the diffusion of oxYgen ions through the thin MgO layer and the long measurement times (as compared to reaction times) give the indication of a high copper surface as if the MgO layer did not exist. E.B.M. DOESBURG : I thank you for this interesting suggestion. thesis deserves further experimental verification.
This new hypo-
B. NOTARI : I would like to contribute to clarify the complex situation recalling that there is general agreement on the fact that in order to operate as a catalyst, copper must change its oxidation state: the result will be that of a dynamic equilibrium and the average oxidation state will depend on many factors like composition of the gas phase, temperature, pressure and so on. The point in discussion has been whether all this occurs on the CU(I) species dissolved in the ZnO crystals as proposed by Klier or, alternatively, on the surface atoms of the copper particles which are formed during the reduction of the mixed Cu-Zn oxides as suggested by Dr. Kochloefl and by myself (7th International Congress on Catalysis, New Horizons in Catalysis, Vol.7, p. 486-487, Elsevier 1980) and also by Dr. Andrew (Post Congress Symposium of the 7th International Congress on Catalysis, Osaka, Japan, July 1980). All evidence which has accumulated and the new one presented here are in favour of this second hypothesis. E.B.M. DOES BURG : At present we are in favour of the ICI view that the process runs over the surface of the copper crystallites. Part of this metal surface is covered by oxygen under reaction conditions (see our ref. 15). Hence we conclude that the function of Cu(O) is the (activated) dissociative chemisorption of hydrogen and the function of O(ads) on copper is the chemisorption of CO 2 to form a C0 3 (ads) species.
783
Reaction between H(ads) and C03 (ads) on copper finally leads to the formation of CH30H. Furthermore we refer to Klier's results on the influence of C02 or H20 in the feed (G.A. Vedage, R. Pitchai, R.G. Herman and K. Klier, Proc. 8th Int. Congr. on Cat., Berlin, 1984, 11-43). Increasing the concentration of C02 (or H20) initially leads to an enhancement of the reaction rate. After having passea a maximum activity, more C02 or H20 leads to a decline of activity. Obviously, a delicate balance between the number of Cu(O) and CU(I) sites or the surface is necessary to arrive at optimum activity.
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B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV
785
© 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
THE ACTIVITY OF COPRECIPITATED Cu. Zn CATALYSTS FOR METHANOL SYNTHESIS Birgitte S. Rasmussen, P.E. H0jlund Nielsen, J0rgen Villadsen, John B. Hansen Haldor Tops0e Research Laboratories, Nym011evej 55, DK-2800 Lyngby, Denmark
SUMMARY The aim of 'this work has been to elucidate the influence of Cu and Zn on the activity for methanol synthesis. A series of catalysts with Cu/Zn contents from 100/0 to 0/100 has been prepared. The precursors have been characterized by chemical analysis, X-ray powder diffraction and analytical electron microscopy. The activity for methanol synthesi's of the calcined precursors suggests that the most active catalyst is obtained when the mole percent composition of Cu/Zn is close to 75/25. INTRODUCTION The industrial low-pressure catalysts for methanol synthesis are predominantly based on the compositions: Cu/ZnO/Cr 203 or Cu/ZnO/A1 203 (ref. 1). Depending on the preparation method and the content of Cu, Zn and Al, the coprecipitated Cu-Zn-Al catalysts may consist of one or several different structure types (refs. 1-5). It is, however, very difficult to get an tdea from the 1iterature of the optimum composition of the catalysts and of the role of each of the present components. Frolich et al. (ref. 6) were the first to report catalyst activity for methanol synthesis of different Cu and Zn compositi'ons, where the catalyst was prepared by precipitation with ammonia. They found that the most active catalyst (350 °C, 204 atm.} contained approximately 35 mole percent CuO and 65 mole percent ZnO. Almost the same result was obtained (250 °C, 75 atm.) by Herman et al. (ref. 1), who investigated methanol catalyst made by coprecipitation of Cu and Zn nitrate by sodium carbonate. Shimomura et al. (ref. 7) found the most active methanol catalyst (300 °C, 50 kg/cm 2) to have the Cu/Zn/Al atom percent composition of 60/35/5, both for a catalyst series prepared by a coprecipttation and a kneading method, and an optimum composition with respect to Al, when 10 atom percent is present. The same tendency was found by Gusi et al. (ref. 8), who from Cu, Zn and Al containing catalysts, prepared by the coprecipitation method, found maximum
786
activity for catalysts containing 10 percent Al or less and a Cu/Zn ratio greater than or equal to one depending on the synthesis gas composition (523 K, 1.2 MPa). The aim of this work has been to elucidate the influence of Cu and Zn on the activity for methanol synthesis. In order to do this, a series of catalysts with various Cu, Zn contents has been prepared. Alumina has been added in order to stabilize the catalysts during testing. EXPERIMENTAL SECTION Catalyst Preparation The catalysts were prepared by coprecipitation of Cu, Zn nitrate, following procedures described in the literature (refs. 4, 9-12). The filter cake was dried at 100 °c for about 16 hours. Calcination took place in air at 350 °c for about 2 hours. Chemical analyses for Cu, Zn and Al were performed upon the dried filter cake. X-Ray Diffractometry X-ray powder patterns were recorded, using Cu Ka radiation generated at 40 kV and 40 rnA on a Philips vert teal X-ray diffractometer, which was equipped with an automatic divergence slit, diffracted beam graphite monochromator, proportional counter and pulse height analyzer. A scan speed of 2° 26/min. was used. Electron Microscopy The electron microscope used was a JEOL 100 CX STEM/SCAN system with a KEVEX energy dispersive X-ray spectrometer for the AEM analysis. The samples for electron microscopy studies were made by placing a drop of ultrasonically dispersed powder in toluene on the sample grid (formvar/carbon). The microscope was operating at a direct magnification of 300,OOOx. Activity Measurements In order to measure the methanol activity, the calcined filter cake was crushed down, and 2.0 9 of the 7-25 mesh fraction (0.7-2.8 mm) was loaded into a tubular reactor system. Feed gases were taken from cylinders, mixed, compressed and stored in high pressure storage tanks. From the high pressure storage tanks, the pressure was reduced to the desired reaction pressure (68 kg/cm 2g). The gas was then passed to the reactor made from a copper-lined stainless stee1 tube with an internal di ameter of 6.3 mm. The reactor was heated with 3 heating elements, each controlled by a separate temperature controller. Ap-
787
proximately isothermal conditions were obtained for all runs. Normal reaction temperature was 222 °c with a maximum difference of 2 °c over the reactor. After the reactor, the pressure was released, and the gas was passed to a cooler at -55 °c and after that to a separator. Liquid products were collected for analysis. The volume of dry product gas was measured with a wet gas meter. A constant bleed of feed gas was taken. Feed gas and dry exit gas were analysed by on-line gas chromatography (Hewlett-Packard 5840A with Porapak Q column). Liquid product was analysed, using a Hewlett-Packard 5840A gas chromatograph equipped with a Carbopack C column. The mean reaction rates were calculated from gas flow and analyses. and checked by the amount of liquid product formed. The feed gas used had a composition of 10 vol.% CO, 1.25 vol.% CO 2, 3.0 vol.% Ar, and was balanced by H2. The space velocity was adjusted to give an exit dry gas containing approx. 6.0 vol.% CO. Under these test conditions. the amounts of methanol formed (g MeOH/h/g catalyst) were a direct measure for the catalyst activity. as both the methanol synthesis and shift reaction are far from chemical equilibrium. The choice of this particular synthesis gas ensures very high reaction rates (CO/C0 2) ~ 8 as shown by Klier (ref. 13). RESULTS The chemical analyses for Cu, Zn and Al in the dried precipitates are given in Table 1. The residual amount of alkali metal does not exceed 1000 ppm in any of the catalysts. The X-ray powder patterns of the precipitates show - besides a boehmite phase (A100H) - one or two of the following phases: a malachite-like phase, where some zinc has substituted copper in the structure (Cu. ZnL2(OH)2C03 (JCPDS card No. 10-399'. an aurichalcite-Hke phase (Cu. Zn}5(OH)6(C0 3}2 (JCPDS card No. 17-743) and a hydrozincite-like phase (Cu. Zn)5(OH)6(C0 3)2 (JCPDS card No. 19-1458) (ref. 14). Although aurichalcite and hydrozincite have the same chemical composition. one has to distinguish between the two minerals, since aurichalcite is orthorhombic and hydrozincite monoclinic. but with related structures. This can be seen from Figure 1. where C-O represents the hydrozincite. and C-30 the aurichalei te phase. The malachite-like phase is present in the more copper-rich samples (C-100 and C-90 in Figure 1). With increasing zinc content. the crystallites of the rna lachite-l ike phase becomes smaller simultaneously with increasi ng changes of the cell parameters (C-81 and C-72). At a certain limit it is not possible
788
for the malachite to accomodate more zinc, and the more zinc-rich phase, aurichalcite, is also formed (C-61 and C-51 in Figure 1). No phase containing Al together with Cu and/or Zn has been detected in the X-ray powder patterns. The phases detected in the catalysts after calcination are boehmite together with CuO and/or ZnO. From the width of the (100)-diffraction line of ZnO and (200)-diffraction line of CuO, the ZnO and CuO crystallite diameters have been estimated for the various catalysts (Table 1). It is seen that the crystallites of CuD and ZnO in all the catalysts are of approx. the same size, although a slight increase in size is seen for the pure Cu-Al and Zn-A1 catalyst. Analytical electron microscopy (AEM) measurements were carried out on at least 10 different areas (~ 100Al of the catalyst C-81, C-72, C-51 and C-30. The resulting average value and the standard deviation of the Cu, Zn and A1 contents by molar ratios are given in Table 2. The scattering in values from area to area is quite large for all the catalysts. This is mainly due to large TABLE 1 Chemical analysis of the precipitated, dried catalysts and the phases detected by X-ray powder diffraction of the catalysts before and after calcination.
Catalyst Relative Atom % Chemical Analysts Cu : Zn : Al Afte& Drying at 100 C %Cu %Zn %Al 0 10 19 28 39
41 42 45 45 45
43.8 39.3 35.4 32.0 27.2
4.3 8.7 12.7 18.1
7.6 7.8 8.4 8.5 8.4
M, M, M, M, M,
C-51
51: 49
49
22.4
21.8
9.2
M, A, B
C-30
30: 70
56
12.8
30.8
10.2
A, B
c-o
0:100
40
47.9
7.7
H, B
C-lOO C-90 C-81 C-72 C-61
100: 90: 81 : 72: 61 :
Phases after Phases after 0 Dryiag at Calc. at 350 C 100 C
M= Malachite-like, (Cu, Zn)2{OH1 2C03 A = Aurichalcite-like, (Cu, Zn)5(OH1u(C0 3)2 H = Hydrozincite-like, (Cu, Zn)5(OH)6(C03)2 B = Boehmite, A10OH.
B B B B A, B
CuO(104A), B CuO{81A), B CuO(74A), B CUO(74A), B CuO(73A), ZnO(87A), B CuO(73A), ZnO{82A), B Cu0(76A) , ZnO(80A), B ZnO(139A), B
789
C-100 C-90 C-81 C-72
C-61 C-51 C-30 C-O 50
40
30
20
10
29 (degrees)
Fig. 1.
X-ray powder diagrams of the precipitated, dried catalysts.
790
variations in Al present. The variations in Cu and Zn - when Al is excluded (Table 2) - are very small for the C-8l and C-30 catalyst and a little higher for the C-72 catalyst. The C-5l catalyst shows a much higher variation in Cu and Zn, suggesting that several phases containing Cu and Zn are present. The average molar ratio of Cu and Zn found by the AEM measurements is very close to the ratio found by chemical analysis. The average amount of Al found, as compared to Cu plus Zn present, is smaller than expected and may be due to inhomogeneous distribution of Al. The activity for methanol synthesis (g MeOH/h/g catalyst) after approx. one week on stream is shown in Figure 2. These activities suggest that the most active of these catalysts is obtained when the atomic ratio between Cu and Zn is close to 3:1. Analysis by gas chromatography of the liquid reaction products revealed that - besides small amounts of water - they contain minute quantities of byproducts such as ethanol, l-propanol, 2-methyl-1-propanol, 1- and 2-butano1, methyl formate and dimethyl ether. Analysis of the dry exit gas showed only the presence of dimethyl ether; no methane or' any other hydrocarbons were detected. The amount of by-products formed, calculated relatively to the converted CO plus CO 2, was maximum 0.11%. Very small variations were seen, and no correlation between the activity and by-products formed was seen. TABLE 2 Summary of AEM data from different areas of catalyst C-81 , C-72, C-51 and C-30.
Atomic Ratio Zn A1
Atomic Ratio Cu Zn
Catalyst
Cu
C-81, average (std. dev.)
80.8 (9.5)
19.2 (3.0)
20.9 (28.7)
80.9 (1. 3)
19.1 (1. 3)
C-72. average (std. dev.)
71.4 (7.4)
28.6 (4.1 )
32.2 (20.3)
71.4 (3.4)
28.6 (3.4)
C-51, average (std. dev.)
48.8 (15.8)
51.2 (15.6)
37.7 (61. 1)
48.2 (8.9)
51.8 (8.9)
C-30. average (std. dev.)
31.4 (6.4)
68.6 (12.8)
36.2 (45.6)
31.4 (1. 3)
68.6 (1. 3)
791
gMeOH/h/g 3.0 @
2.0
e ®
® ®
1.0
@
0
20
40
60
80
100
Zn % Zn+Cu Fig. 2. Catalyst activity for methanol synthesis (g MeOH/h/g catalyst) as a function of the mole percent of In relative to In + Cu present in the catalysts. DISCUSSION The results obtained by X-ray powder diffraction (XRD) seem to be confirmed by the AEM analysis. This means that all the catalysts are estimated to be single-phase precursors with respect to Cu and In present, except for C-51 and C-6l. C-72 may not be purely single-phased as it may consist of small amounts of other phases not detectable by XRD, since the AEM analysis shows a greater variation in Cu and In content from area to area (Table 2). Catalyst precursors containing only Cu and In have, for most of the variable Cu/ln ratios, been reported to consist of more than one phase (refs. 1, 4 and 5}. All three authors have reported tfte catalysts either to be pure hydrozincite or to be mixed with a malachite-li'ke phase in the range where Cu/Zn is approximately equal to 30/70. However, in a later report by Himelfarb et al. (ref. 15), it is recognized that the most active catalyst (Cu/ln = 30/70) in the work of Herman et al. (ref. l) is a synthetic single-phase auri-
792
chalcite precursor, which we have also calculated for our catalyst with the same composition. We have not detected any phase similar to the hydrota1citetype phase, (Cu, Zn)6A12(OH)16C03.4H20, described by Susetto et al. (ref. 2) and Gherardi et a1. (ref. 3), where Al is precipitated together with Cu and/or Zn. XRD of the calcined catalysts shows that an addition of ZnO into CuO induces a reduction of the CuO crystallite s rze and vice versa. The same tendency - but much more pronounced - has been observed by Okamoto et al. (ref. 5). This smaller variation in the crystallite size of CuO and ZnO in our catalysts may be due to the A1 present. The actual size of the oxides cannot be compared, since no corrections have been made for instrumental broadening in the work by Okamoto et a1. (ref. 5).. The composition with respect to Cu and Zn, which gives the maximum activity for methanol synthesis, was here found to be close to 75:25. This is in conf1i ct with the results obtained both by Froltch et a1. (ref. 6) and Herman et a1. (ref. 1) in studies of alumina-free catalysts. The agreement is much better with the studies performed on alumina-containing catalysts studied by Shimomura et a1. (ref. 7) and Gusi et a1. (ref. 8). The latter authors have, however, not found the optimum Cu/Zn ratio to be as high as 3, but to be closer to 2. For catalysts with approximately the same CuLZn ratio, large variations in methanol yield have been reported in the literature. The testing conditions used by the various authors vary considerably, but not enough to explain the differences in methanol yield. One can thus conclude that the optimum chemical composition alone is not enough to ensure high catalytic activity. A number of other parameters greatly influence the catalyst activity, notably the method of preparation and activation. Furthermore, our results show that a reproducible preparation as such could be a very difficult thing, since the precipitated catalyst may consist of several phases. A notable result is the fact that both of the precipitated, mixed phases (malachite or auricha1acite) are precursors to the active phase. With respect to activity, our results may be expressed in a simplified manner: Cu rich: activity
~
Zn rich: activity
~
XZn XCu
This indicates the mutual promoting action, which Cu and ZnO show, but does not indicate where the activity is located. However, from the shape of the curve one may infer that ZnO is about 5 to 10 times as good a promoter for Cu
793
than Cu is for ZnO. This seems to be in contradiction with Klier's hypothesis (ref. 16) to the effect that very small amounts of Cu dissolved in ZnO constitute the active phase. ACKNOWLEDGEMENTS The authors wish to thank O. S0rensen and J. Refslund Andersen for making the electron microscopy analyses, E. Dankvard S0rensen for making all the activity measurements, H. Tops0e, N. Tops0e, B.S. Clausen, J.K. N0rskov and P. Stoltze for valuable discussions, and finally L. Kj01bye for typing the manuscri pt. REFERENCES
2 3 4 5 6 7 8 9 10 11 12 13 14 15 16
R.G. Herman, K. Klier, G.W. Simmons, B.P. Finn, J.B. Bulko and T.P. Kobylinski, J. Catal., 56(1979)407-429. C. Busetto, G. Del Piero, G. Manara, F. Trifira and A. Vaccari, J. Catal., 85(1984)260-266. P. Gherardi, O. Ruggeri, F. Trifira, A. Vaccari, G. Del Piero, G. Manara and B. Notari, in: G. Poncelet, P. Grange and ~.A. Jacobs (Eds.), Preparation of Catalysts III, Elsevier, Amsterdam, 1983, p, 723. G. Petrini, F. Montino, A. Bossi and F. Garbassi, in: G. Poncelet, P. Grange and P.A. Jacobs (Eds.), Preparation of Catalysts III, Elsevier, Amsterdam, 1983, p. 735. Y. Okamoto, K. Fukino, 1. Imanaka and S. Terantshl, J. Phys. Chern., 87(1983)3740-3747. P.K. Frolich, M.R. Fenske, P.S. Taylor and C.A. Southwich, Jr., Ind. Eng. Chern., 20(1928)1327-1330. K. Shimomura, K. Ogawa, M. Dba and Y. Kotera, J. Catal., 52(1978)191-205. S. Gusi, F. Trifiro, A. Vaccari and G. Del Piero, J. Catal., 94(1985)120127. E. Ramaroson, R. Kieffer and A. Kiennemann, Appl. Cata1., 4(1982)281-286. Y. Amenomiya and T. Tagawa, Proc. 8th Int. Congr. Catalysis, Berlin, July 2-5,1984, Vol. 2, p. 557. D. Duprez, J. Barbier, Z. Ferhat.Hamtda and M. Bettahar, Appl. Catal., 12 (1984)219-225. T. Tagawa, G. Pleizter and Y. Amenomiya, Appl. Catal., 18(1985)285-293. K. Klier, Adv. Catal., 31(1982)243-313. Powder Diffraction File (1984). JCPDS International Centre for Diffraction Data, 1601 Park Lane, Swarthmore, PA 19081, U.S.A. P.B. Himelfarb, G.W. Stnnnons, K. Klier and R.G. Herman, J. Catal., 93 (1985)442-450. K. Klier, Appl. Surf. Sct., 19(19841267-297.
794
DISCUSSION S. TAMHANKAR : Do you have any mechanistic explanation for the trend in activity seen as a function of the Cu/ln ratio ? B.S. RASMUSSEN: Both the theory about SMSI between Cu and lnO and the theory about proportionality between activity and surface area of the copper metal phase could be possible explanations of the trends seen in activity for methanol synthesis as a function of Cu/Zn ratio. XY Xiaoding : Since you measured your Cu surface area by X-ray line broadening method, were you not surprised that you did not find a linear relationship between the activity and the copper surface area as you mentioned in your answer. B.S. RASMUSSEN : I said that from the X-ray line broadening method used to estimate the Cu surface area of the used catalysts there could be a linear relationship between the activity and the calculated copper surface area; however, the method is very doubtful. C.S. BROOKS : The normal use for methanol synthesis catalysts is for steady state use in a reduced state. Can you predict the physical stability of your catalyst pellets exposed to repeated alternating cycles of H2 reduction and air oxidation? B.S. RASMUSSEN: No, we have not examined the alternating reduction/oxidation of these catalysts. E.8.M. DOESBURG : You did not mention the rosasite phase which you should expect if more zinc is introduced in your malachite structure, because malachite, zinc containing malachite and rosasite do not form a continuous series of isomorphous structures. Is rosasite not formed in your experiments ? B.S. RASMUSSEN : We have never experienced that rosasite was present in any of the precursor structures. +This is based ~n the fact that the diagnostic diffraction line at d = 2.62 - 0.01 A(34.2 - 0.2° 26 using CuKa radiation) has never been observed. Furthermore, we have never in the literature seen any proof of rosasite being prepared synthetically. See e.g. : A.C. Roberts, Powder Diffraction 1 (1986), 56-57. Y. OGINO : Usually the activity of methanol synthesis catalysts is very sensitive to the conditions of reduction. Could you show the reduction conditions for your catalysts ? B.S. RASMUSSEN: They have been reduced by heating to 220°C in nitrogen containing approx. 3 vol. % synthesis gas. B. NOTARI : You state that no hydrotalcite phase has been found during the preparation of your catalyst precursors, while most research works report that hydrotalcite is formed when solutions containing appropriate amounts of Cu, Zn and Al salts are precipitated at controlled pH with carbonated or bicarbonates. It seems to me that what you find is the logical consequence of using pseudoboehmite as the source of Al : being insoluble under precipitation conditions it is not surprising that Al has not entered into the structure of the new precipitate which was forming from Cu and In in solution and that no hydrotalcite has been found. In order to obtain the formation of hydrotalcite a soluble Al salt must be used in the proper range of concentrations. B.S. RASMUSSEN: Yes, that is correct.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors). Preparation of Catalysts IV 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
THE
PREPARATION
AND
CHARACTERISATION
OF
SEQUENTIALLY
PRECIPITATED
795
AND
COPRECIPITATED NICKEL-ALUMINA CATALYSTS AND A COMPARISON OF THEIR PROPERTIES H.G.J.LANSINK ROTGERINK, J.G.VAN OMMEN AND J.R.H.ROSS Department of Chemical Technology, Twente University of Technology, P.O. Box 217, 7500 AE Enschede, The Netherlands.
SUMMARY A series of nickel-alumina catalysts with a range of NilAl ratios have been made by first precipitating the alumina component and then the nickel component. The effect on the properties of these samples of various pretreatments have been examined and the results have been compared with those for equivalent coprecipitated materials. At higher NilAl ratios, the sequentially prepared materials contain two phases and the resultant catalyst contains large Ni crystallites. This contrasts with coprecipitated samples where a single phase is obtained. A model for the catalysts is proposed. INTRODUCTION Earlier work [1-4,11] on nickel-alumina coprecipitates has shown that these materials give rise to very active and stable catalysts after calcination and reduction as long as the NilAl ratio is between 2 and 3. Originally developed for steam-reforming of naphtha, this type of catalyst was later studied in relation to its possible use in methanation of gas streams with a high content due to the high reaction enthalpy, temperatures up 7000C or more 2; can be reached in the catalyst bed.
of CO and CO
Precipitation from a solution of nickel and aluminium nitrates gives a nickel-aluminium-hydroxy-carbonate with the hydrotalcite (Feitknecht [7]) structure in which the nickel and aluminium ions are randomly distributed between two sheets of hydroxyl-groups in a brucite-like layer and the carbonate ions and the water of crystallisation are located in an interlayer. The structure thus consists of two types of layer which alternate in the c-direction. Several explanations have been given in the literature for the high stability of coprecipitated nickel-alumina catalysts. For example, Alzamora et al. [3] have suggested that small nickel crystallites (diameters in the order of 5 nm) of the reduced catalyst are surrounded by even smaller alumina crystallites in such a way that there is no possibility of direct physical contact between two adjacent nickel crystallites and therefore sintering is hindered. The model has since been modified by Doesburg et al. [4]
in the light of the suggestion of
Puxley et al. [5] that the nickel crystallites are paracrystalle due to stress in the crystals induced by the presence of aluminium ions. Other examples of
796
paracrystallinity are to be found in the catalysis literature, for instance in Fe/Al
materials, used for ammonia synthesis [6], and in Cu/ZnO/Al materZ03 Z03 ials used in the production of methanol from synthesis gas. One major disadvantage of the coprecipitated nickel-alumina catalysts (Ni/Al
ratios between Z and 3) is that the temperature required for reduction is high (500
0
-
600 °C). For low temperature hydrogenation processes, for example
in HZ + N streams to be used for NH Z Z 3 production, there is a need for catalysts which can be reduced at temperatures
methanation of traces of CO and CO
of the order of those used in the reactor itself since the reduction is carried out in-situ. The catalysts do not have to be as stable against sintering as those used for high temperature methanation. We have shown [8] that it is possible to lower the reduction temperature for coprecipitated samples by increasing the Ni/Al ratios to values higher than 3, and that these coprecipitated materials have high activities in the methanation of CO, as long as the catalyst or its precursor has not been used at temperatures exceeding 400 0C. For Ni/Al ratios higher than 3 or 4, Puxley et al [5] and Kruissink et al [1] report two separate phases in the precursor, one being the Feitknecht compound (Ni6AlZ(OH)16C03.4HZO) and the other nickel hydroxide. One would therefore expect, for example, to find two reduction peaks in temperature programmed reduction of the calcined phase resulting from such a precipitate. Since we did not observe such a reduction behaviour but instead single-step reduction, we concluded that it was most likely that the structure of the oxidic phase is one in which there is a single nickel-oxide phase which containing some alumina and that there is no formation of a separate (pure) nickel-oxide phase. The amount of the alumina dissolved or trapped in the nickel-oxide phase depends upon the overall Ni/Al ratio. It has been claimed in the patent literature by Futami [9] that seguential precipitation of nickel-alumina materials is superior to the normal coprecipitation route. The way in which the precipitates are made would result in the covering of the alumina component of the precipitate by the nickel component. According to Futami [9], the intimacy of the contact between the aluminium and the nickel components will be lower and the activity will be higher for sequentially prepared precipitates than for coprecipitates. The aim of the work presented here was to prepare and characterise sequentially precipitated catalysts with high nickel/aluminium ratios (in the range 3 to 20) and to compare these with equivalent coprecipitated samples. Particular attention is given to the influence of temperature of calcination and reduction on the phase composition and properties (Ni crystallite size and methanation activity) of the resultant materials. A model is given for the sequentially precipitated catalysts.
797
EXPERIMENTAL Catalyst Preparation The preparation of catalysts-precursors made by coprecipitation has been described previously [8]. The sequential precipitates were made in an analogous way except that the alumina was now precipitated first and this was followed by the nickel. An aqueous solution (0.9 M) of Al(N03)3.9H20 was added dropwise, with constant stirring, to a Pyrex vessel containing ISO ml deionised water at 80 0C. At the same time, a Na solution (1M) was also added at such a rate as 2C03 to keep the pH at a constant value of about 7 during the precipitation. The precipitation was complete after 30 min. when the precipitate was further aged 0C at 80 (pH 7) for about 30 min. The Ni 2+ solution was then added dropwise to the aged precipitate under the same conditions as described above but the pH was now maintined at 10. After the last drop of solution had been added, the precipitate was aged for a further 30 minutes in the mother-liquor (also at BOoC and pH 10). The precipitate was then filtered and washed thoroughly with 1 dm3 oC. deionised water (80°C) and dried in an oven overnight at 130 It was then washed and dried once again in order to remove the last traces of sodium ions, as these have a detriminal effect on the properties of the catalyst [2]. Calcination of the precipitates was carried out in a flow of nitrogen in a Stanton Redcroft tube furnace maintained at a temperature of 350°C (21 h) or 600°C (3 h), after an initial temperature increase of 2oC/min (Stanton Redcroft Temperature Programmer). Each sample was then reduced in a 1:1 H stream (the 2/N2 0C final temperature, after heating at the same rate, being 400 (5 h) or 600°C (3 h», and passivated in a 1% O2 in N stream at room temperature overnight. 2 Characterisation The sodium contents of the precipitates were analysed with atomic absorption spectroscopy and the carbonate content was measured by dissolving the precipitate in phosphoric acid and titrating the CO 2 evolved. Thermogravimetric studies of the calcination and reduction were carried out with a DuPont system (990 Control unit and 951 TG system). The heating rate for calcination and reduction was 200C/min, the sample weight being about 5 mg. The gases used were identical to those mentioned above. The degrees of reduction of the samples pre-reduced in the tube furnace were also determined with TGA; a sample of the passivated material was re-reduced to 400°C in the H mixture 2-N2 and then re-oxidised to constant weight at 400°C. The degree of reduction could then be calculated from the known composition of the catalyst sample and the weight increase during the re-oxidation step. Total surface areas were determined using a flow method for the adsorption of argon at 78 K [13].
798
Phase Analysis and Particle Size Measurements Phase analysis measurements on the precipitates and on the calcined and reduced samples were carried out by X-ray diffraction (XRD) using a Philips PW 1370 diffractometer, using Ni-filtered CU radiation. X-ray line broadening Ka calculations of the particle sizes of the oxidic and reduced phases were carried out using the Scherrer equation applied to the following diffraction lines: for the calcined material, the (220) line of NiO (ASTM: d 220 = 1.477 A) and the (200) line of NiO (ASTM: d = 2.088 A); for the reduced material, the (200) 200 line of Ni (ASTM: d200 = 1.762 A). Activity measurements Activity measurements were carried out with a differential scanning calorimeter (DuPont 910 DSC) using exactly the same procedure described previously [8]. Repeat measurements carried out with duplicate samples were generally within circa 10% of the average values for that sample composition. RESULTS AND DISCUSSION The Precipitates Phase analysis using XRD (Fig. 1) showed that the sequentially precipitated
3/1 sample has the hydrotalcite (i.e. Feitknecht) structure. When the NilAI ratio is increased, the main diffraction peak of the hydrotalcite structure
20
9
6
4
3
60
50
40
30
20 20 10
Fig.I. X-ray diffractograms of sequentially precipitated samples with different NilAI ratios.
799
(i.e. the (003) peak at an angle of about 12
0
)
decreases markedly in intensity.
The Feitknecht compound is the only visible phase present up to a Ni/Al ratio of 9. However, for the two samples with the highest Ni/Al ratio (20 and a new set of three peaks (at 17.6 of 5.03, 2.65 and 1.57
0
,
33.8
Arespectively;
0
and 58.9
0
)
00
)
we see
corresponding to d-values
the sample without aluminium contains
13.6 wt% C0 = and probably consists of some sort of nickel-hydroxy-carbonate. 3 The sodium content of the precipitates was highest for the sample without aluminium, it being 0.03 wt%. Calcination Fig. 2 gives DTG results for the decomposition of the precipitates. At the lowest Ni/Al ratio, we see three distinct peaks. The first is caused by the removal of adsorbed water, the second (at ca. 180oC) by removal of water of
dW
- dt Ni/Al 00
3 100.
260
300
Fig.2. DTG curves for the calcination of sequentially precipitated samples. crystallisation [1]. The third and main peak is due to the irreversible decomposition of the hydroxy-carbonate into the oxidic phase. On increasing the Ni/Al ratio we observe two phenomena: firstly, the amount of water of crystallisation is lowered and, secondly, the main peak shifts to a lower temperature. The Calcined Material After calcination at a temperature, Tc , of 350 0C, the XRD-patterns indicate that the structure is thar nf poorly crystalline NiO. For the samples with Ni/Al-ratios of 3 and 4, the observed lattice spacings (d 200 = 2.080 A) are lower than for pure NiO (according to ASTM-standards: d 200 = 2.088 A), indic3 ating that the NiO-phase contains a certain amount of the smaller A1 + cation
800
[3]. The diffraction peaks are very broad. Assuming that the broadening (after correction for instrumental line broadening) is caused entirely by the effect of small crystallites, we calculate with the Scherrer formula the particle sizes given in Table 1. It can be seen that the particle sizes calculated from the TABLE 1 NiO particle sizes (in
A)
for sequentially precipitated samples calculated from
the (ZOO) and (ZZO) reflections for different samples. Temperature of oC. calcination was 350 NilAl ratio Crystallite Size I
A
re f l , ZOO ref l , D ZZO D
ZO
(pure NiO)
3
4
6
9
47
6Z
67
76
lZ8
154
81
88
114
116
91
79
0:>
(ZOO) reflection for low Ni/Al ratios are lower than than those calculated from the (ZZO) reflection. Both series of values increase with increasing Ni/Al ratio. This observation is consistent with the results obtained for the coprecipitated samples [8]; however, it should be noted that the particle sizes for
A).
the latter samples are somewhat smaller (D = 35-45 ZZO The BET surface areas of the samples were also measured and they were found to decrease with increasing Ni/Al ratio and range from 174 mZ/g for Ni/Al III mZ/g for the sample without aluminium (T = 350oC). c
= 3 to
Reduction Reduction experiments carried out in the TGA system with two series of samples, one calcined at 350 0C and the other at 600 oC, showed similar trends (DTG curves in Figs. 3 and 4). The first peak in the DTG diagrams at 80
0C
is
caused by the removal of adsorbed water, the other peaks by removal of oxygen from the samples. On increasing the Ni/Al ratio, the temperature at which reduction occurs at a maximum rate decreases. In both series, we observe a splitting of the reduction peak for samples with Ni/Al
=6
and 9. The additional peak now
found is at a temperature which is nearly the same as that found for a pure NiO sample prepared and pretreated under the same conditions. This peak-splitting behaviour is in contrast to our results for coprecipitated samples for which only one reduction peak was observed for the whole composition range. Reduced Phase Table Z gives the degrees of reduction of the prereduced samples, these being determined thermogravimetrically by reoxidation of prereduced samples, as described in the Experimental section. With some of the samples which had high
801
dW
-It
Ni/Al ----20
Ni/Al
~:::=:=== '/
~-~-::::::~---/
00
20 9
~6
~;
r-----
300
100
Fig.3. DTG curves for the reduction of sequentially precipitated samples, oC. with T = 350
500
300
100
500
Fig.4. As for Fig.3 but with T c
c
TABLE 2 Percentage of nickel oxide reduced to nickel for sequentially precipitated samples, as a function of the Ni/Al ratio and the temperatures of calcination (T and reduction (T c) r). Ni/Al Ratio T IOc c
T jOe r
350 350 600 600
400 600 400 600
3
59 90 16 83
4
6
9
20
Degree of Reduction
I %
75 94 49 95
97 94 97 101
88 86 76 94
91 83 83 93
00
(pure NiO)
88 96 102 92
Ni/Al ratios, the re-oxidation was slow and a constant weight was reached only after some time. The degree of reduction calculated for these samples may therefore be slightly too low. It is clear that reduction at 4000C can give quite high degrees of reduction as long as the Ni/Al ratio is high enough, especially for T = 350oC. Reduction at 6000C gives almost complete reduction. The parc
ticle sizes, calculated from X-ray line broadening of the (200) peak of Ni, are given in Table 3 for the samples calcined at 350 0C and reduced at 600°C. The results are compared with those for comparable coprecipitated samples [8]. For the sequentially prepared materials, as for the coprecipitated ones, increasing the Ni/Al-ratio leads to larger crystals. However the sequentially precipitated samples have much larger particle sizes than do the coprecipitated ones.
802
TABLE 3 Ni particle sizes (A) calculated from the (200) diffraction peak for sequentially and coprecipitated samples of different NilAl ratios. The samples were calcined at 350°C and reduced at 600°C. The figures given for the coprecipitates are obtained from reference [8]. 3
4
6
160 50
162
330
NilAl ratio Crystallite Size I A
Sequential Coprecipd.
9
20
90
700 310
w(pure NiO) 930
Activity Measurements Since the activation energy in CO-methanation was found to be practically the same (94 ± 5 kJ mol-I) for all the precipitates, we shall compare the behaviour of the different samples in terms of their activities at 300°C. The results of the activity measurements are shown graphically in Fig. 5. With a reduction temperature of 600
0C
the optimum activity of the sequen-
tially precipitated samples was found for a NilAl ratio of about 4; however, with a reduction temperature of 400°C, the maximum activity was found at a
1.0
0.8 0.6
0.4
0.2
Fig.5. Activity in CO-methanation at 300
(in mol CH 4 (g.catalyst h)-I) as a + Al); • : TC = 350, T = 600; R
= Ni/(Ni R = 400°C.
function of the nickel mol fraction (XNi o : TC = 350, TR = 400; x : T = 600, T
C
0C
~3
Ni/Al ratio of 9. This result is similar to the results obtained for coprecipitated samples [8]. Since after reduction at 600
0C.
the degree of reduction of
all the sequentially precipitated samples was about 85 - 100 % and the nickel particle sizes increased with increasing Ni/Al ratio (see Tables 2 and 3). the sequentially precipitated samples with the lower Ni/Al ratios can be expected to have the largest nickel surface area and therefore the highest activity. as is indeed observed. When the reduction temperature is 400
0C.
the situation is somewhat differ-
ent: the degree of reduction for samples with low Ni/Al ratios is lower than with higher ratios (Table 2). We do not have particle size measurements for samples reduced at 400 0C; however, by analogy with the coprecipitated materials [8J. we expect the particle sizes to be lower but otherwise to behave in the same way as those after 600
0C
reduction. Hence. we would expect the active metal
area. and thus the activity. to go through a maximum, as indeed is shown in Fig. 5 to be the case. Comparison of Coprecipitated and Sequentially Precipitated Samples Although the activities of the coprecipitated and the sequentially precipitated samples are very much the same, there are differences between the two types of catalysts. The first difference is the occurence of two reduction peaks for the sequential precipitates. whereas for the coprecipitates there is only one. For the coprecipitates. this led us to the conclusion (see Introduction) that the most likely structure for coprecipitated nickel-alumina catalysts of higher Ni/Al ratios is one in which A1 3+ ions are included in the NiO phase and that no separate "pure" NiO phase exists. It is clear from the results described above that this model can not be applied to equivalent sequentially precipitated nickel-alumina samples. We shall now try to explain why there is this difference between the two preparation methods. We know from the literature [1,3,5] that the hydrotalcite structure is stable between Ni/Al=2 and 3 or 4. At ratios of 6 or higher, the excess amount of nickel is expected to form a separate phase of Ni(OH)2' as we did indeed observe for both coprecipitated and sequentially precipitated samples, but only after applying a hydrothermal treatment. In the case of the coprecipitates. a phase separation could only be observed after applying a hydrothermal treatment but, for the sequential precipitates. two phases were observed without this extra treatment, some sort of nickelhydroxy-carbonate being formed in addttion to the hydrotalcite structure. In the coprecipitation method, metal tons precipitate continuously in a non-varying environment; however. this is not the case in the sequential precipitation method in which the first drops of Ni 2+ solution see a very different precip-
804
itation environment than do the last drops. In the early stages of the precip2+ itation of Ni in the latter circumstances, there is much exposed aluminium hydroxide with which to react to give the stable hydrotalcite structure. However, as soon as all of the available alumina has reacted to give hydrotalcite, there is nothing with which the excess nickel can react further to give a nickel-aluminium compound and, since the pH is high enough, a nickel hydroxide or nickel-hydroxy-carbonate is precipitated. The whole precipitation process is a delicate balance of crystal formation, dissolution and crystal-growth. It is clear that the coprecipitates having Ni/Al ratios above 4 represent meta-stable structures in which a non-ideal nickelalumina phase exists; during sequential precipitation, no such phase can be formed. Upon calcination, the coprecipitates are transformed into a more or less homogeneous nickel-oxide/aluminium-oxide system in which the nickel oxide phase contains some alumina. From the reduction experiments performed on the calcined sequentially-precipitated samples, it is clear that there are now two nickel oxide containing phases: one in which reduction occurs at a low temperature (comparable to pure NiO) and one reduceable at a higher temperature (comparable to the coprecipitated samples). After reduction, the coprecipitated samples will contain only paracrystalline nickel (as well as alumina), whereas the sequentially precipitated samples will contain both non-paracrystalline (derived from the "pure" NiO phase) and paracrystalline nickel. Another interesting difference between the two types of precipitate is that although the sequential precipitation route gives larger crystals after calcination and reduction than does the coprecipitation route, the activities for the methanation of CO are nearly equal. Since the methanation reaction is regarded to be structure-insensitive except perhaps for very small crystallites, it must be concluded that the sequentially precipitated nickel-alumina catalysts have almost the same active nickel surface area as comparable coprecipitated catalysts, despite the difference in nickel particle size for the two types of sample. One possible explanation for this is that a sizeble proportion of the smaller crystallites of the coprecipitated materials lies within pores which are not, perhaps due to pore closure, accessible to the rection mixture. However, a more likely explanation is based on differences in nickel particle size distributions. We have concluded above that the sequentially precipitated samples contain both paracrystalline (thus small) and non-paracrystalline (thus larger) nickel crystallites and they will therefore have a bimodal crystallite-size distribution. This is in contrast to the coprecipitated samples for which a uniform particle size (unimodal distribution) is expected. Particle sizes measured with XRD are volume-averaged. Since part of the nickel in the sequentially precipitated samples is found in the larger crystals, we find a
805
rather large mean particle size for this type of precipitate. However, the smaller crystallites have much larger surface to volume ratios and therefore contribute more to the surface area. One would therefore expect a higher surface area for the sequential precipitates than that calculated on the basis of X-ray line broadening and in consequence also a higher activity. Confirmation of whether or not this explanation of the behaviour is correct must await the results of planned hydrogen chemisorption experiments. It should be noted that we have here used the X-ray data only for a qualitative comparison of the different samples. It is clear from the results presented that a hydrotalcite structure is formed in addition to a nickel hydroxy-carbonate structure under the precipitation conditions used here. We therefore conclude that the model suggested by Futami [9] in which the precipitate particles consist of an aluminium-rich core covered by a nickel-rich layer is confirmed by our measurements. After calcination and reduction, the activities of sequentially precipitated samples are nearly the same as those of coprecipitated samples. However, in analogy with experiments with lanthanum-promoted nickel-alumina catalysts which have shown that the thermo-stability of La-containing co precipitated samples is higher than for comparable sequential precipitated samples [12], we conclude that the sequentially precipitated Ni-Al materials are also likely to be relatively unstable. Such instability will arise from the presence of large nickel crystallites which are not stabilised by the aluminium present in the formulation. ACKNOWLEDGEMENTS We should like to thank Sally Salemink for performing part of the experimental work, Jaap Boeijsma for his help with the XRD measurments, and Hans Weber and Wim Lengton for doing the chemical analysis of the samples. REFERENCES 1 E.C. Kruissink, L.L. van Reijen and J.R.H. Ross, J. Chem. Soc., Faraday Trans I, 77 (1981), 649. 2 E.C. Kruissink, H.L. Pelt, J.R.H. Ross and L.L. van Reijen, Appl. Catal. 1 (1981), 23. 3 L.E. Alzamora, J.R.H. Ross, E.C. Kruissink and L.L. van Reijen, J. Chem. Soc., Faraday Trans. 1, 77 (1981), 665. 4 E.B.M. Doesburg, P.H.M. de Korte, H. Schaper and L.L. van Reijen; Appl. Catal. 11 (1984), 155. 5 D.C. Puxley, J.J. Kitchener, C. Komodromos and N.D. Parkyns, Preparation of Catalysts III, Eds. G. Poncelet, P. Grange and P.A. Jacobs, Elsevier, Amsterdam" (1983), 237. 6 J.M. Schultz, J. Catal., 27 (1972), 64. 7 W. Feitknecht and M. Gerber, Helv. Chim. Acta, 25 (1942), 131. 8 H.G.J. Lansink Rotgerink, H. Bosch, J.G. van Ommen and J.R.H. Ross, Appl. Catal., in press. 9 H. Futami, U.S.Patent 4,215,998 (1980). 10 ASTM-index. 11 J.R.H. Ross, Catalysis, Vol. 7; Royal Soc. Chern., Specialist Periodical Report (1985), 1. 12 M.R. Gelsthorpe, B.C. Lippens, J.R.H. Ross and R.M. Sambrook, Proc. IXth Iberoamerican Symp. Catalysis; Ed. M.F. Portela (1984), 1082. 13 H. Bosch and A. Peppelenbos; J. Phys. E, Sci. Instr., 10 (1977), 605.
006
DISCUSSION H. SCHAPER: How do you visualise the reaction between precipitating nickel hydroxide and aluminum hydroxide to form a hydrotalcite-like compound, and, if this involves a (partial) dissolution of the aluminum hydroxide, why is it not possible to obtain Ni/Al ratios higher than three? H.G.J. LANSINK ROTGERINK : The aluminum hydroxide which is precipitated first is very amorphous and will therefore be subject of dissolution-precipitation cycles. One can thus expect that free aluminum is present in solution during the first stages of the addition of the nickel solution. The precipitation process is a delicate balance of the dissolution (of aluminum hydroxide), crystal formation and growth. If, on the one hand, the dissolution rate is fast, the hydrotalcite (Ni/Al = 3) will be formed almost exclusively during the initial part of the precipitation of the nickel-containing solution. If, on the other hand, the rate of dissolution is relatively slow as compared to the rate of addition of the nickel solution, then formation of a nickelaluminum-hydroxy-carbonate phase (Ni/Al > 3) can be expected; however, such a compound is itself also subject to dissolution-precipitation cycles, aiding the process of formation of pure hydrotalcite (Ni/Al = 3). The experimental conditions which we used probably correspond to the first case, leading to the formati'on of a hydrotalcite phase (Ni/Al = 3-4) and aluminum-free nickelhydroxy-carbonate. K. JOHANSEN: Have you recorded X-ray diffraction diagrams of samples calcined at 600°C? If it is so that a nickel-aluminate phase is identified, your curves could be explained. H.G.J. LANSINK ROTGERINK : Yes, we did make X-ray diffraction diagrams of samples calcined at 600°C. The calcined precipitates, however, do not contain nickel-aluminate. From work done by other investigators, we know that the formation of NiA1204 takes place at higher temperatures than 600°C. Moreover, nickel-aluminate cannot be reduced below 600°C. This is therefore not a likely explanation for the observed two-step reduction process. A. VACCARI : 11 Have you tested other techniques to evidence the possible presence of side amorphous phases for instance for the sample with Ni/Al = 9.0, that has a very unusual ratio for pure hydrotalcite phase? 21 What is the reason to calcine your compounds to 600°C, taking into account that, according to the literature, you obtain an inreversible transformation of the hydrotalcite-like phase, therefore reducing the specificity of the precursors ? H.G.J. LANSINK ROTGERINK : So far, the only evidence we have for the existence of separate amorphous phases comes from the DTG-curves of the reduction process. We are well aware that this is only indirect evidence and we will try by using STEM or HREM to prove that the precipitates with Ni/Al ratios higher than 3 or 4 do indeed contain separate (amorphous) aluminum-free phases. J.R.H. ROSS: This type of system was described in detail by Puxley et al. in the last meeting of this series. The hydrotalcite structure does not exist in the calcined material. The question we attempt to answer is how the structure of the final catalyst relates to that of the original precipitates. D.M. HERCULES: When you prepare sequentially precipitated catalysts, do you have the Ni present in solution when you precipitate aluminum at pH = 7 ? H.G.J. LANSINK ROTGERINK: No, the sequence is as described in the paper: aluminum is precipitated at pH 7; the pH is then increased to 10 as the nickel is added.
807
E.B.M. DOESBURG : Do your precipitates consist of pure nickel-alumina hydrotalcite with amorphous material present or just Ni-Al-hydrotalcite with a different Ni/Al ratio? If you perform a hydrothermal treatment, you will see different phases, all of which were also present in your starting material. That will give you an idea of the phase-composition of your material. H.G.J. LANSINK ROTGERINK : We think that our sequentially precipitated samples consist of a nickel-alumina hydrotalcite with amorphous material also present. The bases for our model are the reduction profiles of the samples (DTG-curves). From these, one can see that upon increasing the overall Ni/Al ratio, formation of a non-stabilized nickel oxide phase takes place together with an aluminastabilized phase. If the second possible model which you suggest (Ni-Alhydrotalcite with a different Ni/Al ratio) is correct, one would not expect a two-step reduction process, but only a shift in the DTG curves to lower temperatures on increasing the Ni/Al ratio. As to your suggestion of hydrothermal treatment, we think that this will not work out the way you suggest. From a parallel investigation on coprecipitated samples with high Ni/Al ratios, we know that these samples consist of a meta-stable nickel-aluminum-hydroxy-carbonate phase which is more or less homogeneous. Hydrothermal treatment, however, causes a phase separation into pure hydrotalcite and nickel hydroxide. Hence, we do not obtain information about the structure of the precipitate phase. One technique which might be fruitful is STEM of HREM. We will try to perform such measurements in the near future. G.M. PAJONK : I suppose you have measured the total surface area of your samples. What is the order of magnitude of the surface area you obtained? What is dependence between the surface areas and the Ni/Al ratio? H.G.J. LANSINK ROTGERINK : The samples calcined at 350°C have total surface areas in the range 174 m2/g (for Ni/Al = 3) to 111 m2/g (for the samples without Al), decreasing smootly with increasing Ni/Al ratio.
This page intentionally left blank
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
809
PREPARATION AND CHARACTERISATION OF THICK LAYERS OF SEMICONDUCTORS OXIDES FOR GAS CHEMISORPTION AND DETECTION C.LUCAT. F.
MENIL.
M. DESTRIAU, J. 5ALARoENNE and J. PORTIER
University of Bordeaux I
(FRANCE)
ABSTRACT The sensitivity to methane of layers of semiconducting oxides 5r screen-printed with mineral insulating binders has Ca Fe0 3_ x 1_ beenYsh~wn to increase smoothly up to volumic concentrations of binder of about 30%, then very sharply between 30 and 40% before reaching a maximum above this concentration. A possible explanation is based on the formation of a thin gas-sensitive shell ar the grains of semi-conducting oxide after the liquid binder has first filled the empty spaces between these grains during sintering.
j
INTRODUCTION During the chemisorption of a gas on a solid semiconductor, the electron transfer between gas and solid modifies the resistance R of the solid phase. With semiconductor oxides containing transition elements, this modification is generally associated with the presence of two oxidation states for these elements. In the case here considered of a reducing gas (methane) chemisorbed on a p-type semiconductor (MFeD
3_ x
or similar), the resistance R of the solid
phase increases. This phenomenon is used for gas detection in chemical sensors. = ~R/Ro' where ~R = R (resistance in 1% of CH 4 - R (resistance in air), at 470°C, is very small if there o are too many carriers (e.g. in 5rFeO The addition of a glassy Z. 8)' composition, based on Ca, B, Al and 5i oxides, which "melts" and
The sensitivity s in air)
may be reacts, at the firing temperature, also necessary to improve the mechanical properties, may modify the sensitivity by modifying the number of carriers in the detection zone. We have tried to check the whole phenomenon : - by various chemical substitutions (e.g. of Sr by Ca and Fe by Al in the semiconducting oxide) - and by modifying interactions between oxide and glassy composition.
810
EXPERIMENTAL We have prepared oxide powders with the general formula Sr 1_ x Ca Fe Al 03 sintered 24 hours at 1200°C. cooled to room tempex 1-y y -x rature at a rate equal to 10 0 C / h o u r . all ground in the same conditions
(grinding giving grains of submicron size).
An "ink" was pre-
pared by mixing controlled quantities of the oxide with an organic compound and a glass. Then this "ink" was screen-printed on an alumina substrate and refired. A first set of investigated oxides is listed below : SrFe0 2. B1 Sr5/Sca1/SFe03_x Sr2/3Ca1/3Fe03_x
Sr1/3Ca2/3Fe03_X Sr1/SCa5/SFe03_x CaFe0 2. 5
Sr1/2ca1/2Fe03_x X- Ray analysis of these oxides shows evolution from a distorted perovskite to the brownmillerite structure when replacing Sr by Ca, 4 Meanwhile Mossbauer analysis indicates that the amount of Fe + tends to zero. X- Ray analysis of the samples after mixing with a calcium containing glass shows a new phase, the quantity of the oxide phase decreasing, and becoming undetectable when volumic percentage of the calcium containing glass is equal to about 40%. but without any displacement of the rays. In the same conditions the resistance of the samples increases. slOWly from zero to about 30% of this volumic percentage,
then
strongly above this value (see Fig. 1). The sensitivity, s, layers to 1% CH 4 in air. ~R/Ro' ~ 35% of glass. However. the
of the
follows the same trend as R up to o ~R/Ro values exhibit a maximum at a-
bout 40% of glass and decrease for larger percentages (see Fig. 2), whereas the Ro values continue to increase sharply. On the other hand. the ratio ~R/Ro is much more sensitive to the volumic percentage of the calcium containing glass than to the substitution of strontium by calcium in the semiconducting oxide. In our previous work (ref. 1. ref. 2) the increase of sensitivity of the shells, compared to the cores, was attributed to a decrease of the density of carriers in the shells. associated with the decreasing content of tetravalent iron. itself associated with the substitution of strontium by calcium in the shells. This substitution mechanism is obViously inaccurate to explain the improved sensing properties with increasing glass content.
811
A second set of oxides with substitution of Fe
by Al shows ef-
fects similar to those seen above for substitution of Sr by Ca.
log Ro 6
G --------I(
~~x
4
2
0.2
,
10 9 Ro
lC t(
/~
6 4
2
0.2 Fig. 1 : Variation of log Roversus the volumic percentage, the glassy binder. (a) For Sr1/SCaS/SFe03_x' (b) For SrFe0 3_ x,as example~ The firing time as well as the grinding time 60 hours)
C
b,
of
(going from 1 to
do not produce any meaningful modification.
The interface SrFe0
was simulated by depositing and 2. S/g1ass firing a layer of glass on a sintered pellet of srFe0 The ana2• S' lysis of the interface by Castaing microprobe showed that the elements of the semiconducting oxide
(i.e.
Sr and Fe)
are alsopresent
812
on the glass side whereas those of the glass (i. e. Ca, B, AI, Si) are not present on the semiconducting oxide side.
SY.
10
l(
0.4
0.2
0 5 ~
Cb )('
.1\ I
20
)(
G
,It )f
I
x
x
It
K
Ie
\
,
It
l('
10
~/ )(
X
/>f II
====;:::: 0.2
K
I
0.4
Cb
Fig. 2 : Variation of the sensitivity. s, versus the volumic percentage, C of the glassy binder. b, (a) For Sr1/ScaS/SFe03_x' (b) For SrFe0 ,as e x arnp Lar, 3_ x DISCUSSION
We make the hypothesis that during the treatment, the glass "melts", first fills the volumes between the grains of oxide, then surrounds the grains of oxide forming a grain core-grain shell structure (see Fig. 3). The core of the grain contains mainly the starting semiconducting oxide whereas in the shells a new phase
813
resulting from the chemical reaction between the semiconductive oxide and the glass precipitates. This phenomenon doesn't modify greatly either the resistance or the sensitivity for a volume of glass less than about 30% of the whole volume, that is to say before filling by the glass between the grains of oxide (independently of the grain size). Beyond this value the resistance and the sensitivity are mainly controlled by the shells. The grain shells become more resistive than the cores, the sensible phase seeming to be the new phase coming from the interaction of glass and oxide.
Fig. 3 : Schematic model showing the growth of the new phase as C b increases (the spaces between the grains of semiconducting oxide' are filled when C > 0,26). b REFERENCES 1.
2
F. Menil, J. Portier, C. t u c e t , C. Miquel and J. Salardenne, Proceedings of the International Meeting on Chemical Sensors of Fukuoka, Kodansha and Elsevier, Amsterdam. 1983, pp. 193-197. C. Lu c e t , F. MeniL C. MiqueL M. Destriau. J. Salardenne and J. Portier, Rev. Chim. mt n e r-e Ls , 21 (1984) 194-201.
814
DISCUSSION G.M. PAJONK : As your sensor is based upon a redox system, do you know if CH4 and O2 react on its surface and what are the products? Have you measured its surface area? M. DESTRIAU : The gas composition has not been analyzed during the detection experiment. However the resistance of the sensor has been shown to stabilize quickly for any concentration of CH 4 smaller than 3% in air. Moreover we have shown that such a stabilization cannot be obtained when the carrier gas does not contain any oxygen. For CH4 concentration smaller than 3%, it may be thought that CH4 and 02 somehow react at the surface of the sensor and lead to a stationary state. The value of the specific area of the powder of the semiconducting oxide, measured with a B.E.T. apparatus, was around 8m 2.g- l.
B. Delman. P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
815
STANDARDIZATION OF CATALYST TEST METHODS BY THE COMM ITTEE ON REFERENCE CATALYST OF THE CATALYSIS SOCIETY OF JAPAN T. HATTORI l+, H. MATSUMOT0 2+ and Y. MURAKAMI 1* IDepartment of synthetic Chemistry, Faculty of Engineering, Nagoya University, Chikusa, Nagoya 464, JAPAN. 2Research and Development Division, JGC Corporation, Bessho 1-14-1, Minami-ku, Yokohama 232, JAPAN *Chiarman of the Committee. +Secretary of the Committee. SUMMARY This paper reports the manual of CO-pulse method for metal surface area of supported metal catalysts standardized by the committee and an interim report of standardization of NH 3-TPD method for zeolite acidity. In the CO-pulse method, much stress has been 1aid on the rapi d measurement; it takes about three hours to measure the amount of adsorbed CO including pretreatment. The results obtained by the method agreed well with those obtained by the other methods. In the NH3-TPD method, it has been clarified that the contact time of carrier gas at a sample cell has a large effect on peak temperature. The Committee on Reference Catalyst of the Catalysis Society of Japan was started at 1978with the following aims: (1) the distribution of reference catalysts, (2) the collection and report of the data on reference catalysts, (3) the standardization of catalyst test methods (ref. 1). Eleven oxides (A1 203, Si0 2, Si0 2-A1 203, Ti0 2 and MgO) and ten zeol ites (Mordenite and Y with different Si/Al ratio) were selected as the reference catalysts, and they have been distributed to more than two hundreds users including those in foreign countries. Any members of the Catalysis Society of Japan can obtain the reference catalysts through the committee members without any obligations and can take part in the activity of the committee. Another sample of MgO and six samples of ZSM-5 supplied by Mobil are going to be distributed. Thirty one supported catalysts of Pt, Pd, Ru, Rh and Ni were prepared for the project of meta 1 surface area. For the second aim, the committee have organized ten symposia on the subject shown below, and a special session on the reference catalystwas organized in the annual meeting of the Catalysis Society of Japan in October 1984 (ref. 2). The results have been publ ished as the preprints of the symposia as well as the scientific papers submitted by each user. (1) Characterization of Alumina; Oct. 1979. (2) Metal Surface Area; Oct. 1980. (3) Metal Surface Area; June 1981. (4) Support Effect; Oct. 1981.
816
(5) Metal Surface Area; Dec. 1982. Alumina; Aug. 1983.
(7)
of Zeolites; Oct. 1984.
(6) Metal Surface Area and Properties of
Metal Surface Area; May 1984.
(8) Characterization
(9) Characterization of Zeolites; July 1985.
(10)
Characterization and Catalysis of Magnesia; Oct. 1985. Two catalyst test methods have been standardized; the determination of BET surface area by NZ adsorption and the rapid measurement of metal surface area of supported metal catalysts by CO-pulse method. Further the standardization is in progress for the measurement of acid property of zeol ites by temperature-programmed desorption (TPD) of NH 3. The manual s for the BET surface area by the volumetric method and by the flow method were publ ished in "Shokubai (Catalyst)", a bull etin of the Catalysis Society of Japan (ref. 3).
In the manual s , more stress was laid on the "rapidity" of determination
than the manuals published by SCI/IUPAC/NPL (ref. 4) and ASTM (ref. 5).
It
has been shown that the accuracy of the single point-flow method is similar to that of the mul ti point-vol umetric method, if enough attention is paid to the cal ibration of the detector and to the pretreatment of catalysts. In the present symposium, we report the process of standardi zation, the standardized manual of CO-pulse method for metal surface area and an interim report of standardization of TPD method for zeol ite acidity. STANDARDIZATION PROCESS General process of standardization is shown in Fig. 1.
The committee does
not decide in advance the subject of standardization, and the users can use the di stributed reference catalysts as they 1 ike.
Therefore, the results
reported in the symposium include various subjects measured by various methods and procedures.
In the symposium, the discussions were done about
the correlation between the results of different subjects and about the comparison of the results obtained by identical method.
The former leads to
cooperative work, and the 1 atter leads to the standardization of catalyst test methods.
The test method to be standardized is selected through the
discussion and every participant can take part in the decision.
Then, the
detailed procedures are discussed to decide tentative procedures. In the next symposium, the results obtained by the tentative procedure are reported.
If the results agree well with each other and with the results
obtained by the other methods, and if the consensus is reached on the procedure, then the tentative procedure becomes standardized one.
If not,
improved procedures are proposed, and the same process is repeated.
In the
case of the standardization of CO-pulse method, four symposia have been devoted to the decision of final procedure.
817
I Distribution
of Reference Ca ta ly st ]
J,
Symposium
Results on Various Subjects by Various Methods ~ Selection of Subject of Standardization
J,
Propose Tentative Procedure
.L ICata1yst Test by Proposed Procedure .L Symposium
R'P~R'::'t'
CONSE~
Improved Procedure -
.~es ISTANDARDIZED PROCEDUREI Fig. 1
Schematic Diagram of Standardization Process
CO-PULSE METHOD FOR RAPID DETERMINATION OF METAL SURFACE AREA As reported previousl y (ref. 1), the tentatively standardized procedure But the standardi zation of detai led gave good agreemen t of the resu 1ts. procedure and the comparison with the other methods were remained for further examination. New series of catalysts shown in Table 1 were prepared and four symposia were held for the development of the standardization. The results and the procedure were summarized and edited by N. Nojiri, and it was pub l i sned as the manual for supported Pt catalyst (ref. 6). Modified manuals for supported Pd, Rh and Ru catalysts were proposed, and the reproducibility of the results is going to be examined at the special session in the annual meeting of Catalysis Society to be held in coming October. It was pointed out that the samples No. 10 - 20 contain significant amount of chloride, and a series of Cl-free supported metal catalysts (No. 29 - 31) were prepared and added to the list. The application of the manual to the Cl-free catalysts a1so wi 11 be exami ned at the speci a1 sess i on. Another project was started for supported Ni catalyst (ref. 7). In the present report, only the results on the supported precious metal catalysts were described. Brief Description of the Standardized Manual The method covers the rapid determination of the amount of adsorbed CO at room temperature on supported Pt catalysts by a conventional pu1 se-adsorption apparatus with He carrier gas. It takes only about three hours, even if the pretreatment and the measurement are done in series.
818
TABLE 1 List of New Series of Supported Precious Metal Catalysts No.
Mark
Metal
Metal Content
10 11 12 13 14 15 16 17 17-2 18 19 20
JRC-M -0. 5Pt(1.0) a JRC-A4-0.5Pt(0.5)a JRC-M -0. 5Pt( O.l)a JRC-A4-5.0Pt(I.0)a JRC-A4-5.0Pt(0.gJ a JRC-M-0.5Rh[2] JRC-S3 -0. 5Rh JRC-A4-0.5Ru JRC-M -0. 5Ru[2]b JRC-S3-0.5Ru JRC-A4-0.5Pd[2]b JRC-S3-0.5Pd
Pt Pt Pt Pt Pt Rh Rh Ru Ru Ru Pd Pd
0.5 wt% 0.5 wt% 0.5 wt% 5.0 wt% 5.0 wt% 0.50 wt% 0.49 wt% 0.38 wt% 0.49 wt% 0.51 wt% 0.48 wU 0.52 wt%
29 30 31
JRC-A4-0.5Pt[3]c JRC-A4-0.5Pd[3]c JRC-A4-0.5Rh[3]c
Pt Pd Rh
0.50 wU 0.50 wt% 0.49wU
Support
Mated a1
a expected dispersion, b second batch of preparation, c third batch of preparation. No. 10 - 14 were prepared by Uchijima & Kunimori (Univ. of Tsukuba), No. 15 - 20 were prepared by Nippon Engel hard, and No. 29 - 31 were prepared by Toyota Central R&D Labs. Apparatus Typical example of apparatus is a conventional flow apparatus for BET surface area equipped with a thermal conductivity detector, an injector val ve or a serum cap for CO pul se injection, and a sampl e tube. Although the purification of carrier gas (He) had only a 1ittle effect, it is recommended to purify the carrier gas by a U-tube trap containing molecular sieve at liquid nitrogen temperature or by a column of oxygen scavenger such as manganese oxide, chromium oxide or reduced copper. Pretreatment (1) Charge 100 - 200 mg of catalyst to a sample tube which has been weighed. (2) Raise sample temperature to 673K in flowing air, and hold for 15 minutes. (3) Purge air with flowing He for 15 minutes. (4) Reduce the catalyst in flowing H2 for 15 minutes. (5) Let flow He again for 15 minutes, and remove furnace to cool the sample tube to room temperature in flowing He. In these procedures, a flow rate between 20 and 40 cm 3/min is preferred for a11 the gases. Measurement of CO uptake It is recommended to start measurement as soon as possible after cooling. (1) Inject CO pul ses with an interval between 2 and 3 minutes until the amount of exit CO pulse reaches a steady value, and then inject two more pulses to confirm the steady value. Pulse size should be determined so
819 that the steady value can be obtained within 3 pulses.
Pulse size between
50 and ZOO mm 3 is preferred for the cata 1ys t of meta 1 1oadi ng of ca. 0.5 wt%. (Z) Detach the sampl e tube from the apparatu s, wei gh, and ca 1cul ate sampl e weight. (4)
Calculate the amount of adsorbed CO as follows: [CO]ads
= L([CO]steadY - [CO]exit)'
where [CO]steady is an average of the steady value of exit CO pulse, and [CO]exit is the amount of exit CO pulse before the steady value is attained. (4) Report CO uptake (cm 3[STP]/g-catalyst) or percent exposed (dispersion) by assuming the stoichi ometry (CO/Pt = 1). Results Ql the Standardized Method Figure Z shows a frequency distribution of the resul ts obtai ned by the standardized method and those obtained by omitting Oz pretreatment. of Oz pretreatment gave smaller CO uptake and larger scatter.
The omit
The samples
may be contaminated by organic molecules, which may result in the smaller CO uptake and larger scatter without Oz pretreatment.
The results on the
sampl es No. 10 and 11 agreed wi th each other withi n a reasonabl e error, but the sample No. 12 gave relatively large scatter.
The small ratio of
adsorbed CO to injected CO would lead to large error.
In Fi g. 3, the average of CO uptake by the s tandardi zed method was compared with the results by the other methods including uptake of H2, 02 and HZS by pul se method, H2' CO and Oz uptake by static method, and HZ-C ZH 4 titration.
The resul ts
of CO-pul se method agreed well with those of HZ
adsorption by static and pulse methods and CO adsorption by static method. Furthermore, a good correlation was observed between the CO-pulse method and the Oz adsorption by pulse and static methods.
These results indicate that
the standardized method can be used for the rapid measurement of CO uptake. App 1i cat i on to the Other Metal The manual was appl ied for supported Pd, Rh and Ru catalysts, and the results were shown in Fig. 4.
It appears that the purification of carrier
gas is necessary, and that Oz pretreatment does not always give good results. Further examination has been done, and the following modification has been proposed (ref. 8). (1)
The purification of carrier gas is necessary.
(Z) Oz treatment should be omitted.
Thus, raise sample temperature in
flowing HZ' and reduce catalyst for 15 minutes.
Then let flow He for 15
820
minutes, and then remove furnace to cool sample tube to room temperature in flowing He. (3) The following temperatures are recommended for the reduction; 573-673K for Pd, 673-723K for Rh and 723K for Ru. The modified manuals are going to be examined at the special session in the annual meeting of Catalysis Society.
No. 10 No. 15 %
0.8 COl Rh
0.68
NO~ a 0.2 0.4
20 % 0.26
No. 12
Q
'
CO / Ru
I
80 %
-80
No.13
0.60
n
-40
rT.1
n, ~Y:-' -20
a
. JQ!8CJ l. o. waI: '~.l . j'
Ig,
20
,..
I~
N~ 0.2
n
0.6
0.4
0.6 CO / Pd
0.8
Fig. 4 Frequency distribution of CO uptake on Rh, Ru and Pd catalysts measured by standard-80 -60 -40 -20 40 60 % ized method for Pt catalyst. X'd square, with 02 treatment; Fig. 2 Frequency distribution of CO upteke relative to averaged values measured by stand- bold square, carrier gas was ardized method. X'd square, with 02 pretreat- purified. ment. No. 14
0.10
+-'
c... 1.0 <,
g
" 0.8
+-'
c...
<, 0
"
0.6
+-' c, <, 0
0.4
~
0.2
u
0.. <, ::I:
a
0/ Pt o pulse o static S / Pt ~ pulse
H/ Pt ~ o pulse 'B st at i c titration ~ COl Pt 0~
CI)
~':/Ou ~o-rJ a
0.2
0.4
0.6
0.8
CO / Pt
1.0 0
0.2
0.4
0.6
0.8
1.0
ov pulse (average)
Fig. 3 Comparison of the standardized method with the other methods.
821
NH 3-TPD METHOD FOR DETERMINATION OF ACID PROPERTIES OF ZEOLITES Ten samples shown in Table 2 were selected as the reference catalyst, and 34 reports were presented in two symposia held in 1984 and 1985. They include adsorption, ion exchange property, catalysis, external surface area, acid property, and structure by TEM, SEM, XPS, NMR and IR. Among these, the external surface area and the acid property by temperature programmed desorption (TPD) of NH 3 attracted attention from the viewpoint of standardization. As for the external surface area, four groups reported the results calculated from the amount of adsorbed N2 at 77K or from the adsorption kinetics of N2 onto zeo1 ites whose pore had been fi lled with pore fi 11 ing agent (ref. 9). As shown in Table 2, some disagreement was observed. Further examination will be necessary to reach final consensus on the method. TABLE 2 Reference Zeolites and Thier External Surface Areas
Mark
JRC-Z-Y4.8 JRC-Z- Y5. 6 JRC-Z-MI0 JRC-Z-MI5 JRC-Z-M20 JRC-Z-HY4.8 JRC-Z-HY5.6 JRC-Z-HMI0 JRC-Z-HM15 JRC-Z-HM20
SifA1
Type a
4.8 5.6 9.8 15.0 20.1 5.2 5.6 9.9 14.9 19.9
NaY NaY NaM NaM NaM HY (99) HY (72) HM (98) HM (99) HM (99)
External Surface Area (m 2fg)b c n-C 4 H20 c H2 0c C6 H6c H2 0c S-Mp d S-Mpd F-Spe F-Spe S-Mp d
4.0 8.6 (8.8) (6.9) 4.3 8.5
6.7 5.2 4.8 6.2 9.0 9.8 10.0 5.3 7.6 11.7
6.9 7.2 8.9 11.0 13.0 17.0 10.6 9.8 13.5 15.2
6.5 6.0 9.0 14.7 19.4 12.2 8.1 10.7 13.4 16.9
11. 7 5.6 10.5 13.2 16.8
and numerals in parentheses denote a Y and Mdenote zeolite Y and morde~ite, the extent of proton exchange (%). measured by fi 11 ed pore method except those in parentheses which were measured by adsorePtion kinetic method. c pore fi 11 ing agent: n-C4 denotes n-butene. mul ti point BET method by static apparatus. e single point BET method by flow apparatus. TPD spectra of NH3 reported in the first symposium were measured by various methods and procedures. Desorbed NH3 was detected by thermal conductivity detector, mass spectrometry, vacuum gauge, or gravimetry. Zeolite bed was pumped to evacuate the desorbed NH 3, or it was under the reduced or atmospheric pressu reo The pretre atment conditi on and the heati ng rate also were different. These differences may lead to the significant scatters in the desorbed amount as a measure of acidity and the peak temperature as a measure of aci d strength. However, the trend agreed we 11 wi th each
822
other, which encouraged us to develop the standardization, and a tentatively standardized procedure was proposed. Tentati ve Manua 1 and Results on TPD of NH3 Following experimental conditions only were unified as a tentative manual. (1) pretreatment, evacuation at 773 K for 1 hr. (2) adsorption and evacuation of NH 3 at 373 K. (3) increase of bed temperature, at a rate of 10 K/min from 373 K, until all NH 3 is desorbed. In the next symposium, 12 groups reported the resul ts of TPD of NH 3 measured according to the tentative manual. The outline of TPD spectra and the Sequence of the acid strength agreed well with each other. TPD spectra on H type zeolites had two desorption peaks (1- and h-peaks at low and high temperatures, respectively) on not only mordenite but also Yzeal i teo On the other hand, Na-mordeni te possessed on 1y the 1-peak, whose amount was much greater than that on the decationized one. The peak area of h-peak agreed pretty well with each other and with NH 3 uptake measured by microca Ior imetry, But the 1arge scatter was observed in the peak area of 1peak. These results led to the view that the 1-peak of the mordenite does not originate from molecules adsorbed on the weak-acid site but from physically or very weakly chemisorbed molecules, and that only the h-peak reflects the aci di c property of zeo 1i tes. The comparison of peak temperatures of the h-peak consistently afforded the fo11 owing sequence for the aci di ty strength, and thi s sequence was in good agreement with that obtained microca10rimetrica11y. HM-15
> HM-20 > HM-10 » HY-4.8 > HY-5.6
However, the peak temperatures did not agree at all. For example, the peak temperature on HM15 was distributed from 673 K to 8B3 K. It was suggested that the peak temperature seemed to depend on the contact time. As for the manual, the following v iews were offered. (1) Zeol ite samp1 e should be evacuated at least to 10- 4 Torr to remove residual NH 4+ ions. (2) Since heating rate has not so large effect, the heating rate in the tentative manual is preferred for practical usage. Detailed investigation was conducted by a project group, and a brief summary of the resu 1ts, which are going to be pub 1i shed (ref. 10), is shown below. The detailed investigation of TPD spectra measured by 12 groups indicated that the contact time of carrier gas W/F has a significant effect on the peak
823
temperature and the peak area of l-peak. As shown in Fig. 5, the peak temperature gradually increased with the contact time W/F, where Wand F denoted the zeol He weight and the flow rate of carrier gas at a pressure in the TPD sample cell, respectively. When the readsorption of a desorbed mol ecule occurs freely, the peak temperature (Tm) is theoretically correlated with the contact time (ref. 11): 210gTm - log(dT/dt) = t.H/2.303RTm + logB(W/F) B = (1/d)(1-e)2(t.H/R)exp(-t.S/R) where dT/dt and e show the rate of temperature increase and the coverage at peak maximum, t.H and t.S are heat and entropy of adsorption, and d is the density of sol id, respectively. When t.S is assumed to be 100 J K- 1 mc l " ", the simu1 ated curve fo 11ows the relation between the peak tempera ture and the contact time, suggesting that NH 3 is freely readsorbed on these zeolites in the TPD experiment. The following two items wi 11 be confirmed in the next symposium schedul ed for October, 1986. (1) Only the h-peak has to be accounted for to measure the acidity of zeolites. (2) The peak maximum temperature has to be compared with those measured by the other apparatus on referring to the equation shown above .
.......
::>::: 900
L.
::J .j..J
a 800
L.
cu
Cl.
E
.:::.c. 700
a
0..
10- 5
10- 4
10- 3
10- 2
WI F (g,min/cm3 ) Fig. 5 Correlation between the peak maximum temperature of the h-peak and the contact time on JRC-HM-15. The theoretical curve was simul ated on the parameters: dT/dt = 10 K/min, t.H = 150 kJ/mo1, d = 0.5, and e = 0.35.
824
REFERENCES
2 3 4 5 6 7 8 9 10 11
Y. Murakami, in: G. Poncelet, P. Grange and P.A. Jacobs (Eds), Preparation of Catalysts III, Elsevier, Amsterdam, 1983, p.775. Shokubai (Catalyst), 26 (1984) 280. Shokubai (Catalyst), 26 (1984) 495. D.H. Everett, G.D. Parfitt, K.S.W. Sing and R. Wi lson, J. Appl. Chern. Biotechnol., 24 (1974) 199. ASTM 0 3663-78, ASTM Standard, 25 (1980) 964. Shokubai (Catalyst), 28 (1986) 41. M. Yamada, in ref. 2. E. Yasui, N. Tumaki, 1. Imai, in ref. 2. 1. Suzuki, S. Namba, T. Yashima, J. Catal.,81 (1983) 485, M. Inomata, M. Yamada, S. Okada, M. Niwa and Y. Murakami, J. Catal., in press. M. Niwa, M. Iwamoto, K. Segawa, to be submitted. R.J. Cvetanovic, Y. Amenomiya, Advan. Catal., 17 (1967) 103.
Committee Members Y. Murakami (Chi arman, Nagoya Univ.), H. Matumoto (Secretary, JGC Co.), T. Hattori (Secretary, Nagoya Univ.), H. Hattori (Hokkaido Univ.), S. Okazaki (Ibaragi Univ.), J. Take (Tokyo Univ.), H. Niiyama (Tokyo Inst. Tech.}, M. Niwa (Nagoya Univ.), S. Yoshida (Kyoto Univ.), T. Imanaka (Osaka Univ.), H. Arai (Kyushu Univ.), and 1. Furuoya (takeda Chern. Ind., Ltd.). Participants in CO-Pulse Method for Metal Surface Area of Precious Metals K. Kunimori, Y. Ikeda, E. Yamaguchi, 1. Hashimoto and 1. Uchijima (Univ. of Tsukuba), T. Nakata (Nippon Engel hard), T. Imai and K. Iida (Mitsubishi Heavy Ind.), N. Nojiri, M. Kurashige and K. Yaguchi (Mitsubi shi Petrochem.), S. Sago, 1. Deguchi and S. Nakamura (Sumitomo Chern.), T. Inui, A. Sakamoto, T. Otowa and Y. Takegami (Kyoto Univ.), T. Suzuki (Cosmo Oil), T. Mori and H. Masuda (Gov. Ind. Res. lns t., Nagoya), S. Abe (Ni kki Chern.), N. Tsumaki and K. Takeda (Wakayama Coll ege of Technol.), E. Yasui and F. Haga (Nippon Oi 1), M. Ipponmatsu, O. Okada and S. Takami (Osaka Gas), J.O. Hernandez, E. Choren, G. Arteaga and L. El Chaar (Univ. Del Zulia, Venezuela) Participants in External surface Area of Zeolite 1. Suzuki, H. Ezure and K. Saito (Utsunomiya Univ.), 1. Takahasi and M. Ozaki (Kagoshima Uni v.), M. Inomata, M. Yamada, S. Okada and H. Yamamoto (JGC Corp.), M. Niwa (Nagoya Univ.).
Participants
~
NH 3-TPD Method for Zeolite Acidity
K. Segawa and Y. Kurusu (Sophi a Uni v.), Y. Nai -Ju (Qinghua Univ.), M. Nakamura, Y. Inoue, Y. Mitarai and Y. Takami (Sumitomo Metal Mining), T. Imai and K. Iida (Mitsubishi Heavy Ind.), M. Iwamoto, M. Tajima and S. Kagawa (Nagasaki Univ.), S. Nakamura, M. Ishii and K. Matsuoka (Shindaikyowa Petrochem. Ind.), H. Sato, K. Hirose and H. Ichihashi (Sumitomo Chem.), M. Ushio and T. Ishi i (Nippon Oi 1), T. Kawamura, S. Baba and T. Shimizu (Cosmo Oi 1), Y. Takasaki and 1. Shibuya (Utsunomiya Univ.), K. Mahos, Shin Chang Sub and H. Niiyama (Tokyo Inst. Technol.), T. Takahashi and M. Ozaki (Kagoshima Univ.), S. Nakata, M. Sohma, K. Nagaoka and S. Asaoka (Chiyoda Chem. Eng. Construction), and M. Sawa and M. Niwa (Nagoya Univ.).
825
DISCUSSION N. PERNICONE : Can you quantify, on a relative basis, the participation of industrial and academic research institutions to the activities of the Japanese group ? T. HATTORI : The relative participation of industrial and academic institutions depends on the subject. As for the BET surface area, the number of industrial institutions is almost half of the academic institutions. The standardization of static BET method was carried out mainly by the academic institutions, while the industrial institutions participated in the standardization of dynamic method. As for the metal surface area, many academic institutions participated in the initial stage, but the CO-pulse method was standardized mainly with the help of industrial institutions. Almost equal numbers of industrial and academic institutions participate in the zeolite project. A. FRENNET : In the comparison of the standardized method with the other methods in the characterization of Pt supported catalysts by chemisorption, I am surprised to see (fig. 3) that the ratios H/Pt, CO/Pt, O/Pt, S/Pt are the same. This is in contradiction with the observations made on very clean model samples like single crystal faces (see the work of J. Oudar) for example. I would like to hear your comments on this point. T. HATTORI: It is not surpris,ing that H/Pt, CO/Pt and O/Pt are the same. It has been frequently reported that CO/Pt agrees well with H/Pt on supported catalysts (for example, A. Renouprez, C. Hoang-Van, P.A. Compagnon, J. Catal., 34 (1974) 411). O/Pt was a little smaller than CO/Pt and H/Pt, which agrees well with the published results cited in the comment of Dr. Herrmann. The adsorption of sulfur is not frequently used for the determination of metal surface area. S/Pt in Fig. 3 was calculated from the sulfur content of the catalyst which had been treated in an atmosphere of 10 ppm H2S/H2 at 573 K for 10 hours. The sulfur content was measured by X-ray fluorescent spectroscopy. In the symposium on Reference Catalyst, a question was presented about the possibility of the sulfurization of bulk Pt as well as surface Pt, but it is hard to sulfurize bulk Pt in such a low concentration of H2S as 10 ppm. J.M. HERRMANN: On the right side of Fig. 3, representing a plot of (O/Pt) vs (CO/Pt), you find, from the slope of the dashed line, a ratio (O/CO) equal to 0.76. A close value, substantially lower than unity was also found for (O/H) ratio on various Pt/Si02 catalysts with the same pulse technique (T. Uchijima, J.M. Herrmann, Y. Inoue, R.L. Burwell, J.B. Butt and J. Cohen, J. Catal.50, (1977) 464), as well as on Pt/Ti02 (J.M. Herrmann, M. Gravelle and P.C. Gravelle, J. Catal. (accepted)). Do you have a collective interpretation for the origin of this chemisorption stOichiometric ratio? If not, could this question be a problem to be studied in your Committee? T. HATTORI: Wilson and Hall (J. Catal., 17 (1970) 190,24 (1972) 306) showed that the stoichiometric ratio of O/Pt coula be lower thanIH/Pt depending on the dispersion: the O/H ratio could be 0.5 for totally dispersed Pt catalysts, while the ratio could be 1.0 for Pt catalysts of relatively low dispersion. The ratio shown in Fig. 3 is just intermediate between both extremes in spite of relatively low dispersion. More extensive work will be necessary to give a collective interpretation. At present, the task group has not any concrete plan to examine 'it, although it is possible that some members will personally do it. L. GUCZI :' Among the metals for which CO chemisorption was applied to determine the metallic surface area, I have a little doubt on the reliability of Ru. Especially, under oxidizing condition ruthenium tends to be sintered (Prep. Cat. II, p. 391, Elsevier, Amsterdam, 1979). Thus pretreatment at 300°C seems to me a little drastic handling. Moreover, depending on the environment some
826
of the ruthenium may stay in oxide form which cannot be counted by CO chemisorption. T. HATTORI : In the symposium on Reference Catalyst. it was pointed out that Ru tends to be sintered in the presence of oxygen. And this is one of the reasons why oxygen treatment has been omitted in the tentatively standardized procedure. Further examination on Ru has been carried out, and one of the members found that the preoxidation at lower temperature greatly increased CO uptake. Further discussion will be done in the coming special session at the annual meeting of the Catalysis Society of Japan. XU Xiaoding : 1/ When different metal surface areas are measured by the CO method, which is the CO/metal ratio used to calculate the metal surface area? 2/ When the problem concerns the strength and the distribution of acid sites with different acid strengths, how would you use the NH3-TPD results since different base probes used lead to acid sites (amount) of different acid strengths. T. HATTORI: 1/ As described in the text. it is recommended to report the result of CO-pulse method in terms of CO uptake in cm 3(STP)/g-catalyst or dispersion in the ratio of adsorbed CO molecules to total metal atoms. The metal surface area may be reasonably calculated on the assumption of CO/surface metal being unity. as shown in various papers such as in the paper presented by Prof. Bond at this symposium. 2/ It is not necessary to use various base probes for the determination of acidic properties by TPD method, because TPD method gives both the strength and the amount of acid sites; the peak area stands for the amount of acid sites and the peak temperature stands for the acid strength. It is recommended to use a probe base of small molecular size. If a large probe molecule is used. the diffusion resistance may have significant effect on the TPD spectrum. Actually, TPD of pyridine on mordenite gave desorption peaks at much higher temperatures than those in TPD of NH3. Therefore, we recommend to use only NH3 as the probe base. ZHAO Jiusheng: Sites with different acid"strengths are suitable for different reactions. Different probes, such as NH3' pyridine and other basic compounds will determine different sites of the catalyst. This will give incorrect information when you use NH3 as the only standard probe molecule. What do you think of this? T. HATTORI: See answer to question (2) of Dr. XU Xiaoding. H. WETSTEIN : Shouldn't the zeolite acidity be measured in addition to ammonia also by some other base and those values reported? T. HATTORI : See answer to question (2) of Dr. XU Xiaoding. A. FRENNET : In the titration of the acidity of zeolites. TPD measurements of NH3 you presented were conducted up to 800°C. I am wondering if there is not ammonia decomposition at such high temperatures. Have you some information, some experimental evidence that such a decomposition is negligible? T. HATTORI : One of the members confirmed that nitrogen was not formed in the TPD experiment in which desorption was completed below 600°C. In one of the TPD experiments conducted up to 80eoC, mass spectrometer was used to detect desorbed NH3' Thus. the TPD spectrum consisted only of NH3' but it did not show decomposition products. The TPD spectra measured by mass spectrometry agreed well with those measured by using TCD, indicating that the decomposition of NH3 is n~gligible in the TPD experiment. It should be noted that this remark can be applied to zeolites which do not contain any metals and metal oxides.
B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet (Editors), Preparation of Catalysts IV © 1987 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
EUROPT-l:
827
THE FIRST PLATINUM ON SILICA REFERENCE CATALYST
Geoffrey C. BONO l and Peter B. WELLS 2 loepartment of Chemistry, Brunel University, Uxbridge, UB8 3PH, United Kingdom 20epartment of Chemistry, University of Hull, Hull, HU6 7RX, United Kingdom
ABSTRACT A 6% platinum on silica catalyst has been prepared and thoroughly characterised by workers in a number of European laboratories. In the 'asreceived' state the platinum is present as a disordered oxide which is however easily reduced to metal without change in dispersion. The arithmetic mean metal particle size, estimated by transmission electron microscopy, is 1.8 nm. Hydrogen chemisorption on this catalyst is a complex phenomenon, temperatureprogrammed desorption revealing at least three adsorbed states. Some of the hydrogen consumed in volumetric experiments may be used to break Pt-O-Si bonds. Derivation of a monolayer volume by extrapolation to zero pressure is shown to have no theoretical basis. Chemisorption of carbon monoxide or of oxygen may, with certain limitations, be used to obtain estimates of metal dispersion. INTRODUCTION Progress in the science of heterogeneous catalysis and in the development of improved catalysts for industrial use depends upon obtaining a deeper understanding of the correlations between chemical composition, physical structure and catalytic performance. To achieve this we need accurate and sensitive techniques for measuring and describing those aspects of structure upon which catalytic activity and selectivity depend. Each method employed must inevitably provide information which is unique and peculiar to that method; each involves highly specific pretreatments, conditions of measurement and assumptions in deriving conclusions from the observations; each contributes something to the store of knowledge of the material under study. As techniques for physical characterisation were developed through the 1960s and early 1970s, each laboratory had necessarily to devise the apparatus and procedures it believed to be the most appropriate. Inevitably these differed to some extent, and nowhere was this more true than with the method of selective gas chemisorption, for which no commercial equipment was available. Differences were smaller where this was to hand (e.g. with electron microscopy, X-ray diffraction, and various spectroscopies), although even here the methods of sample preparation and pretreatment could vary considerably. Evidence was therefore lacking that results obtained in various laboratories were strictly comparable.
828
These considerations led Professor E.G. Derouane in 1975 to bring together a group of European scientists interested principally in catalysis by metals. This group subsequently formed the Research Group on Catalysis as one of several constituting the Study Group on Surface Chemistry, reporting to the Council of Europe's Committee on Science and Technology. At its meeting in London in 1976 during the Sixth International Congress on Catalysis the Group decided to commission the preparation of a large batch (~6 kg) of a platinum/silica catalyst and invited its members to participate in a programme for its thorough physical characterisation. This programme had two principal objectives: (1) to compare and contrast results obtained by the techniques used in the various laboratories of the Group members; and (2) to provide a very well characterised material of value to the Group members, and subsequently to the wider scientific community. No attempt was made to impose an arbitrary specification of the methods to be used: each laboratory employed its normal facilities and procedures, and performed those measurements for which it was best suited. This work was carried forward enthusiastically during the following five years, after which the results were subjected to extensive critical analysis. They have now been published in 'Applied Catalysis' [1-5] and this paper is a distillation of the prodigious amount of work which went into the project. Many laboratories participated in it and it is estimated that about one hundred persons contributed their time and skills, in addition to their normal duties. In keeping with the policy of the Group, not all of these persons are named as authors of this paper. The authors of this paper are the spokesmen for those who did the work, although one of us (P.B.W.) acted as project coordinator and was the principal author of the detailed publications [1-5]. The contributions of the various laboratories are denoted by codes and the appendix identifies them by their addresses and the names of the Group members currently representing them. THE CATALYST The catalyst selected for this study was platinum on silica; a metal content of about 6% was decided upon, in order to facilitate estimation of metal particle size by electron microscopy. As it turned out, this was extremely useful in determining adsorption stoichiometries. The catalyst was prepared by Johnson Matthey Chemicals Ltd under the supervision of Dr D.E. Webster. 6 kg of silica (Sorbosil AQ U30, Crossfield Chemicals) was stirred with 60 1 of 0.01 M [Pt(NH 3)4]C1 2 solution, and the pH was maintained at 8.9 by addition of a solution containing [Pt(NH 3)4]C1 2 (0.01 M) and [Pt(NH 3)4](OH)2 (0.1 M). The material was filtered, washed, dried and reduced in hydrogen at 673 K. It is convenient for characterisation work: it is
829
an easily-handled granular material, showing only a very small weight loss on heating to 473 K (0.1%) [2J. The platinum content was determined by several different methods (spectrophotometry, atomic absorption spectroscopy, PIXEl and may be expressed as 6.3 ±O.l% by weight. One feature of this catalyst which is less than wholly satisfactory is that the metal content appears to vary slightly with grain size: the fraction larger than 500 ~m contains 5.7 ±O.l% metal by weight, so the smaller grains must have relatively higher concentrations. The ideal reference catalyst would comprise only grains within a narrow size range: at the very least, the finest fraction should be removed by elutriation at the start of the preparation [2]. Highly consistent results were obtained for the total surface area. Nitrogen physisorption, gas chromatography [6] and mercury porosimetry all afforded results expressible as 185 ±10 m2g-1. Only single-point BET measurements gave values outside this range. The pore volume is 0.77 cm 3g-1 and mean pore diameter is 14 nm. A point of some interest is that the surface area of the catalyst is only about one-half that of the support: the loss of surface area is tentatively attributed to disappearance of mesoporosity in consequence of hydrothermal sintering during the drying stage. This is a facet of catalyst preparation not often considered, but may be important if it results in the encapsulation of some of the metal precursor [2]. METAL PARTICLE SIZE DISTRIBUTION It was somewhat surprising when an EXAFS study [7J revealed that the platinum in the as-received catalyst was principally in the form of a disordered oxide (Pt-O bond length, 0.20 nm). Oxidation must have occurred following the reduction step, but there is evidence to show that re-reduction, which is easily effected, does not lead to any measurable change in particle size [3]. The catalyst was subjected to extensive characterisation by transmission electron microscopy: all reports agree that visible particles are between 0.9 and 3.5 nm in size, and that 75% are less than 2 nm in size. The reported size distributions (Figure 1) lead consistently to an arithmetic mean size of about 1.8 nm, from which dispersions of 56-59% for a cubic model or 63-66% for a hemispherical model may be derived [3J. The catalyst is thermally stable under reducing conditions; thus for example heating under hydrogen at 873 Kfor 5 h is without measurable effect on particle size distribution, although sintering evidently occurs at and above 1073 K and more specially in the presence of oxygen [3J.
830
60 [a]
[b]
[c]
[e]
[I]
30
60 Ed]
30
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Q)
.D
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:l
1
2
3
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0
3
0
2
3
4
5
4
[h]
[g]
z
20 0
20 [i] 1
4
2
3
5 6 7 8 Particle size/nm
9
2
3
2
3' - 4
5
0
0 20 [j]
0
0
10 .11
12
Fig. 1. Size distributions of the platinum-containing particles determined by high resolution transmission electron microscopy. Histograms refer to the following samples: (a) to (e) as-received Europt-l as measured in five laboratories; (f) EuroPt-l re-reduced in hydrogen at 623 K for 2 h (zero sintering condition); (g) to (j) sintering observed when EuroPt-l was re-reduced in hydrogen at, respectively, 873 Kfor 5 h, 1073 Kfor 6 h, 1273 Kfor 4 h, and 1273 K for 15 h as measured in four laboratories.
831
CHEMISORPTION OF HYDROGEN Volumetric methods are commonly employed to measure hydrogen chemisorption isotherms, with a view to obtaining useful information on the dispersion, area and particle size of the metallic component. A recent very detailed article on "The Measurement of Catalyst Dispersion" [8] contains the following unambiguous statement. "Chemisorption of hydrogen on platinum is instantaneous at ambient temperature and it readily reaches a complete monolayer on the exposed platinum surface." This statement is then qualified in three ways: (i) the H/Pt stoichiometry has to be assumed; (ii) there must be no physical adsorption of hydrogen on either the metal or the support; and (iii) there must be no other complications such as dissolution or spillover. The programme here described addresses the validity of the above statement and its qualifications. Conventional volumetric procedures were employed by nine laboratories, using the conditions in vogue at the time. After being evacuated the sample was reduced at a temperature between 428 and 673 K and then evacuated again at some temperature between 573 and 773 K. To convey a sense of the degree of reproducibility attained, eight values for the amount of hydrogen chemisorbed at 40 Torr equilibrium pressure lay between 163 and 190 ~mol g-l (mean 178 ~mol g-l) [4J. Some of the isotherms are shown in Figure 2; they cover various pressure ranges, but none of them shows a completely flat plateau region. Since the amount of hydrogen chemisorbed increases continuously, at least up to 300 Torr equilibrium pressure, each of the isotherms shown is a portion of a general
.....0 e0
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e
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o
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300
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20
30
0
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0
0.2 0.4 -0Equilibrium pressure of "2 / Torr
Fig. 2. Isotherms for the adsorption of hydrogen on Europt-1 at room temperature as measured in three ranges of pressure in four laboratories.
832
.... I
0> N
200
:I:
---
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o
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10910{Equilibrium pressure of H2 / Torr) Fig. 3. Isotherms for the desorption of hydrogen from EuroPt-l at three temperatures over a very wide range of equilibrium pressure as measured in one laboratory.
isotherm, and any semblance of a linear region which may be safely extrapolated to zero pressure is therefore accidental. Clearly if such a procedure is used the intercept obtained will increase with the maximum equilibrium pressure attained. A particularly careful and detailed study of hydrogen chemisorption on EuroPt-l was conducted in laboratory 83 [4,9]. The quantity adsorbed is shown as a function of log(equilibrium pressure/Torr) at various temperatures in Figure 3, from which it is clear that (i) the Temkin isotherm is obeyed at the lowest temperatures and (ii) there is no saturation limit below atmospheric pressure. From the results shown in Figures 2 and 3, it is evident that the H:Ptt ot ratio at and above 40 Torr pressure is significantly greater than unity (by about 10% at 40 Torr). Extrapolation to zero pressure as commonly practicised would give an uptake of 170 ±10 ~mol g-l which would correspond to an H:Ptt ot ratio of approximately unity. These results thus indicate a dispersion of 100% and are therefore clearly at variance with the estimate of particle size by electron microscopy which gave a dispersion in the region of 60%. This discrepancy may perhaps be resolved by reference to measurements of the thermal desorption of adsorbed hydrogen; these reveal at least three and possibly four adsorbed states, one of which is probably associated with hydrogen spillover. It has been argued [4] that some
833
30% of the adsorbed hydrogen may have been taken up in hydrogenolysis of Pt-O-Si bonds formed between metal particles and the support during pretreatment and that only about 55% of the total is conventionally chemisorbed. Thus of the 178 ~mol g-l adsorbed at 40 Torr pressure, only about 98 ~mol g-l is strictly relevant, and this value corresponds (assuming H:Pts = 1:1) to a dispersion of about 62%, in harmony with the electron microscopy results. We are forced to conclude that at least with this system there is no foundation in theory for using hydrogen chemisorption to estimate dispersion or mean particle size, and that conventional application of volumetric measurements leads to a gross over-estimation of dispersion. We cannot of course know whether these difficulties are present with other metals and other supports, but clearly it is necessary to investigate whether they are or not before the method can be used with confidence. The validity of the concept of turnover frequencies estimated from hydrogen chemisorption measurements must also be called into question. Other studies of hydrogen chemisorption on EuroPt-l have been reported ~o]. CHEMISORPTION OF CARBON MONOXIDE AND OF OXYGEN Limitations of space do not allow a detailed presentation of the results obtained [4J under this heading. Four reliable volumetric measurements, employing various reduction and evacuation schedules, of carbon monoxide chemisorption afforded isotherms which could be extrapolated to zero pressure to give intercepts of 185-198 ~mol g-l. This is slightly greater than the corresponding value for hydrogen. Extensive measurements of the infra-red spectra of chemisorbed carbon monoxide show that most of it is in the linear form, with only minor contributions from other states. Of the eighteen volumetric measurements of oxygen chemisorption at ambient temperature, sixteen gave isotherms which on extrapolation to zero pressure showed intercepts of between 80 and 91 ~mol g-l, corresponding to an O:Pts ratio of about unity [4J. However the amount adsorbed was found to be a function of the temperature of measurement, and so once again by accident the most convenient conditions provide an empirically useful means of estimating dispersion. CATALYTIC MEASUREMENTS Measurements of the catalytic properties of EuroPt-l have received lesser emphasis than its physical characterisation as reported above, but a certain amount of work has been carried out, particularly on hydrocarbon transformations. The results are still being analysed, but it is hoped that they may be published shortly.
834
As an indication of the kind of results obtained, Table 1 gives the turnover frequencies at 573 K for the hydrogenolysis of C2 to C6 alkanes, together with activation energies. These results were obtained in three different laboratories as indicated, and employed various pretreatment schedules. Alkane partial pressures were in the range 0.04-0.08 atm and hydrogen partial pressures between 0.83 and 0.94 atm. The hydrogen monolayer volume was taken as 166 ~mol g-l for calculating turnover frequencies. Some of this work has already been published [llJ. TABLE 1 Kinetic parameters for alkane hydrogenolysis over EuroPt-l Alkane
T range/K
E/kJ mol-l
C2H6 C3H8 n-C 4H10 n-C 4H10 n-C 5H12 n-C 6H14
613-633 578-608 578-608 523-603 484-518 484-518
210 189 129 121 109 117
TOF/mo1ec. h-l Pt -1* s 0.22 1.35 3.46 4.4 6.75 9.21
Laboratory UKl UKl UK1 F4 N3 N3
*At 573 K CONCLUSIONS This study has revealed a number of unsuspected problems in characterising an apparently simple supported metal catalyst. (1) Small platinum particles are apparently readily oxidised upon exposure to the atmosphere; re-reduction is therefore essential, but is easily effected. (2) Hydrogen chemisorption is an unexpectedly complex phenomenon: there are at least three states distinguishable by TPD, not all of which are associated with chemisorbed hydrogen atoms. (3) Volumetrically measured hydrogen isotherms show no well-defined saturation converage below atmospheric pressure, and hydrogen uptake at 40 Torr equilibrium pressure corresponds to a H:Pttot ratio of greater than unity, whereas particle size estimates from TEM give a dispersion (i .e. H:Pttot ratio) of about 0.6. This discrepancy may be explained by the use of some part of the hydrogen consumed in breaking Pt-O-Si bonds, although we do not know whether this is a general phenomenon.
835
ACKNOWLEDGEMENTS We are indebted to the Council of Europe, and in particular to the officers of the Committee on Science and Technology, for their unstinted support of this project. The following former members of the Research Group on Catalysis also contributed to the project in its early stages: J.J. Fripiat (Fl, Orleans); the late F.G. Gault (F3, Strasbourg); J.K.A. Clark (11, Dublin); D.L.Trimm (NW1, Trondheim); S. Friberg (S2, Stockholm); R.W. Joyner (UK2, Bradford); M.W. Roberts (UK3 Cardiff, formerly Bradford); C. Kemball (UK5, Edinburgh). REFERENCES 1 G.C. Bond and P.B. Wells, Applied Catalysis, 18 (1985) 221. 2 G.C. Bond and P.B. Wells, Applied Catalysis, 18 (1985) 225. 3 J.W. Geus and P.B. Wells, Applied Catalysis, 18 (1985) 231. 4 A. Frennet and P.B. Wells, Applied Catalysis, 18 (1985) 243. 5 P.B. Wells, Applied Catalysis, 18 (1985) 259. 6 R. Bacaud, G. Blanchard,H. Charcosset and L. Tournayan, React. Kin. Catal. Lett., 12 (1979) 357. 7 R.W. Joyner, J. Chem. Soc. Faraday Trans. I, 76 (1980) 357. 8 J.L. Lemaitre, P.G. Menon and F. Delannay in 'Characterization of Heterogeneous Catalysts', (ed. F. Delannay, Dekker, New York/Basel, 1984), p.343. 9 A. Crucq, L. Degols, G. Lienard and A. Frennet, Acta Chim. Acad. Sci. Hung., 111 (1982) 547. 10 A.R. Berzins, M.S.W. Lau Vong, P.A. Sermon and A.T. Wurie, Adsorption Science and Technology, 1 (1984) 51. 11 G.C. Bond and Xu Yide, J. Chern. Soc. Faraday Trans. I, 80 (1984) 969.
The diagrams in this paper are reproduced from Applied Catalysis (references 3 and 4) with permission of Elsevier Science Publishers BV.
836
APPENDIX Members of the Research Group on Catalysis as at October 1984 Country Laboratory code AUSTRIA
Al
BELGIUM
Bl B2 B3 B4
FRANCE
F2 F3 F4 F5
GERMANY
Gl
THE Nl NETHERLANDS N2 N3 N4 SPAIN
SPl
SWEDEN
Sl S2
UNITED KINGDOM
UKl UK4 UK5
Member and affiliation
H.L. Gruber Institut fur Physikalische Chemie, INNSBRUCK B. Delmon Universite Catholique de Louvain, LOUVAIN-LA-NEUVE E.G. Derouane Facultes Universitaires de Namur, NAMUR A. Frennet Universite Libre de Bruxelles, BRUXELLES P.A. Jacobs Katholieke Universiteit Leuven, LEUVEN B. Imelik (until 1984), R. Maurel (from 1984), C. Naccache, J.C. Vedrine Institut de Recherches sur la Catalyse (CNRS), VILLEURBANNE G. Maire Universite Louis Pasteur, STRASBOURG J. Barbier Universite de Poitiers, POITIERS G. Leclercq Universite de Lille, LILLE G. Ertl, H. Knozinger Institut fur Physikalische Chemie, MUNCHEN J.W.E. Coenen Katholieke Universiteit, NIJMEGEN J.W. Geus Rijksuniversiteit, UTRECHT V. Ponec Rijksuniversiteit, LEI DEN J.H.C. van Hooff, R. Prins Technische Hogeschool, EINDHOVEN G. Munuera Universidad de Sevilla, SEVILLA R. Larsson University of Lund, LUND P. Stenius Institute of Surface Chemistry, STOCKHOLM G.C. Bond Brunel University, UXBRIDGE P.B. Wells University of Hull, HULL D.A. Whan University of Edinburgh, EDINBURGH
837
DISCUSSION G.R. LESTER: Could you comment further on the statement in the paper on the poorer relation accuracy of the single-point surface area measurements ? G.C. BOND: I think it is self-evident that a surface area based on a full isotherm determination will be more accurate than a single point determination, simply because of the larger number of data points involved. I am less clear as to why the single point procedure gives significantly higher values, but can only assume this is inherent in the approximations introduced in the theoretical analysis. B. NOTARI : The question is raised how to detect Pt particles in the case -they are in the range 1 + 20 A and escape detection at the TEM. I would like to suggest the measurement of catalytic activity towards a test reaction. In my experience this is a very sensitive parameter of the degree of dispersion of Pt. For instance, the temperature at which CO is oxidized on Pt. Under well defined conditions (gas composition, gas flow, etc ... ) the temperature of the reactor is increased with a linear program: the temperature at which the reaction starts - as detected by thermal and chemical analysis - is lower the better the degree of dispersion. No effects other than the dimensions of the Pt particles have ever been detected. A correlation can thus be drawn between the temperature at which CO is oxidized and the degree of dispersion of Pt. G.C. BOND : Your suggestion is a most interesting one and should certainly be followed up. However, it would appear necessary to have some kind of calibration curve of activity or light-off temperature versus dispersion, and the construction of such a curve would require activity measurements on a number of catalysts of known dispersion. I therefore cannot see how the procedure you advocate can be used without having an independent method for determining dispersion in the first instance.
J. KIWI: How small would a Pt-cluster stoichiometry as a minimum? In other for the metal cluster to show metallic or number of atoms of Pt (per cluster)
have to be to have a 1:1 chemisorption words, the metallic character necessary behaviour would start at which dimension ?
G.C. BOND : There is probably no single critical minimum size for the appearance of metallic behaviour. The answer will depend upon what parameter is selected as indicative of the true metallic state, and there is more likely to be a gradual transition than an abrupt change. Since surface metal atoms are different from those in the bulk in a variety of ways, metallic character is achieved when the particle size is such that surface atoms no longer influence significantly the properties of the whole particle. Theoretical calculations seem to indicate that the band structure is fully developed in particles containing 200 atoms, i.e. at a particle size of about 1.8 nm and a dispersion of about 60%. A. FRENNET : In spite of the fact that I certainly agree with the difficulties you mentioned about the used of H2 measurements to characterize a metal surface, I am convinced that the comparison between dispersion so determined and derived from TEM is not easy when one deals with highly dispersed metals, because many TEM's do not see the very small particles and because the stoichiometry of chemisorption on highly dispersed metals may be larger than 1:1. The comparison between the informations from both TEM and H2 chemisorption would probably have been more significant in a first step working with a much less dispersed catalyst. G.C. BOND: In seeking to calibrate hydrogen chemisorption against surface area determined by TEM, one faces a dilemma: the measurement of chemisorption becomes more accurate as the dispersion increases, whereas a TEM particle size
838
becomes less accurately measurement. It is not clear where the position of best compromise lies, but it may indeed be at a dispersion lower than 60%. Another uncertainty is whether the several forms of chemisorbed hydrogen retain their relative or absolute amounts as the particle size is changed. A. CRUCQ : There is a good agreement from surface science studies that on well developed crystal faces, the ratio HIM is close to 1. For very small particles (such as in supported catalysts) this is not true. If the chemisorption site is the center of a square defined by 4 metal atoms on a 100 face, (or in the center of a triangle on a 111 face) some metal atoms on corners and edge may contribute twice to the definition of a chemisorption site and one can calculate for particles in the range 10-20 Athat the HIM ratio should be in the range 1.5-2.0. High value of HIM around 2.0 can thus be easily understood: that was previously mentioned in a paper several years ago by people from the Institute of Catalysis in Villeurbanne. B.C. BOND: I believe there is a good deal of force in your comment. I would add the proviso that in all probability the HIM ratio of unity is achieved with small particles at rather low pressures, but without the appearance of a welldefined plateau: and that hydrogen which is taken up at higher pressures, corresponding to the region in which the slope of the isotherm is low, and which is weakly and reversibly adsorbed, is that which is held as MH2 on atoms of low coordination number. This idea is indeed not new: you will find that my paper in the Fourth International Congress on Catalysis (Moscow, 1968) makes reference to it. A. FARO : Would not the use of dynamic chemisorption methods obviate the problem of pumping speed associated with volumetric methods. Was there any comparison made between static and dynamic methods as far as this catalyst is concerned ? G.C. BOND: Dynamic methods have indeed much to commend them as rapid and empirical procedures for estimating dispersion. I am less convinced of their value in fundamental studies, particularly in cases where (as with hydrogen) there is a multiplicity of adsorbed states of various strengths. J.M. HERRMANN: You have found on Europt-1 catalyst an average size of ~1.8 nm. We have found the same value for thermally reduced Pt catalysts deposited on titania, with metal contents ranging between 0.5 and 10 wt Pt% and whose size distribution is given in Nouv. J. Chim. 6, 1982. Don't you think that this particle size of 1.8 nm could be considered as a textural characteristics of divided platinum, whatever its support? G.C. BOND: I regret that I do not recall your results, which frankly I find rather surprising, as the mean size is generally found to increase with increased metal loading, and there are several models which account for such behaviour. There is no difficulty in producing even more highly dispersed metal, down to sizes as small as 1 nm or less, and it therefore seems more likely that the agreement between the particle size in Europt-1 and in your Pt/Ti02 catalysts is due to a coincidence. J.W. GEUS : The reason we are using silica as a support is that this support is not eXhibiting a diffraction pattern. As a result any heavy metal atom and at any rate small clusters of metal atoms can be distinguished rather easily using modern electron microscopes. G.C. BOND: No comment. A. VANNICE : Have you measured isotherms to distinguish irreversibly adsorbed hydrogen at 300K from reversible adsorbed hydrogen? If so, how do the irreversible uptakes compare to irreversibly adsorbed CO and to oxygen
839
chemisorption? We have repeatedly found in our laboratory that 40-50% of the total H2 uptake can be removed by evacuation at 300K (for 1h), thus the irreversible uptake is much closer to the O2 uptake. G.C. BOND: I am not convinced that there is a clear and meaningful distinction between reversibly and irreversibly held hydrogen, although the existence of at least two forms is not in question. If there is any truth in the response I made to Dr. Crucq's comment, I would expect the fraction of "reversibly" adsorbed hydrogen to decrease with increasing particle size, i.e. as the mean surface coordination number increases. I can see no likelihood that by adjustment of pressure alone, either in adsorption or desorption, will lead to the isolation of any particular adsorbed state. P.G. MENON: The discrepancy in H-chemisorption data on Europt-1 catalyst resulted from the different pretreatments (calcination and reduction, reduction temperature) given by different laboratories. But, where the pretreatments were quite comparable (Universities of Innsbruck and Leiden, Chalmers University of Technology, and Micromeretics Laboratory in USA), the chemisorption value for H2 was practically the same, 134 t 2 ~mole H2/g catalyst. G.C. BOND: The point you make is a very valid one, but the value for the monolayer volume which you quote seems to me to be rather low in comparison with those given in our paper (see Fig. 2). I presume the isotherms were measured in a low pressure region. A. CRUCQ : The fact that the extrapolation to zero pressure gives different results following the pressure range considered is not peculiar to the Europt-1 catalyst. It has been observed on other metals, (Rh, Ni, Ru) on alloys, either supported or unsupported. G.C. BOND: I am sure you are correct. This must be true for any system exhibiting an isotherm which does not have a plateau region in which the slope is close to zero. C.J.G. van der GRIFT : Is IR spectroscopy used during H2-chemisorption? It might give evidence for the formation of SiOH groups during chemisorption. G.C. BOND : We have not tried to use IR spectroscopy to study chemisorbed hydrogen. although of course other workers have used it successfully.
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841
LIST OF PARTICIPANTS
AMBACH
E.
BASF Aktiengese11schaft, Zak/1 6700 Ludwigshafen F.R.G.
AMPHLETT J.C.
Royal Military College of Canada Department of Chemistry and Chemical Engineering Kingston, Ontario K7K 5LO CANADA
ANDERSSON B.
Chalmers University of Technology Department of Chemical Reaction Engineering 41296 GBteborg SWEDEN
ANDERSSON L.H.
Chemical Technology Chemical Center P.O. Box 124 22001 Lund SWEDEN
ANDERSSON U.
EKA Nobel AB 44501 Surte
ANTOS G.
U.O.P. Drawer C Riverside, Illinois 60546
SWEDEN
U.S.A.
ARNTZ D.
Degussa AG, Abt. FC-O Postfach 1345 6450 Hanau 1 F.R.G.
BABIC D.
I.P.L.A.S. Kemicna Industrija Koper o.sd.o. Krpanova 2 66000 Koper YUGOSLAVIA
BACHEllER J.
Universite de Caen Laboratoire de Cata1yse Esplanade de 1a Paix 14000 Caen FRANCE
BAlKER A.
ETH-Zentrum Technical Chemical Laboratory 8092 ZUrich SWITZERLAND
BAKER B.G.
Flinders University Bedford Park 5042 South Australia
AUSTRALIA
BALKENENDE A.R.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
BARRAULT J.
Cata1yse en Chimie Organique U.A. CNRS 350 Avenue du Recteur Pineau, 40 86022 Poitiers FRANCE
BARTEN H.
Kema NV Utrechtseweg 310 6812 AR Arnhem
THE NETHERLANDS
842
CHRZASZCZ J.
Institute of Industrial Chemistry Ul. Rydygiera. 8 01-793 Warsaw POLAND
CORDIER G.
Rhone-Poul enc Avenue des Freres Peret. 85 69190 St Fons FRANCE
CORMACK B.
Johnson-Matthey Catalytic Systems Division Orchard Road Royston. Herdfordshire SG8 5HE
COURT J.
Universite de Grenoble I LEDSS I - U.S.T.M. B.P. 68 38402 St Martin d'Heres Cedex
CRUCQ A.
U.L.B. Laboratoirede Catalyse Campus de la Plaine. B.P. 243 1050 Bruxelles BELGIUM
DALLONS J. L.
U.C.B. - Secteur Chimique 1620 Drogenbos BELGIUM
DANNER A.
Universitat Mainz Institut fUr Anorganische Chemie Becherweg 24 6500 Mainz F.R.G.
DE CLIPPELEIR G.
Labofina Zoning Industriel 6520 Feluy BELGIUM
DE KEYSER F.
Monsanto Rue Laid Burniat 1348 Louvain-la-Neuve
ENGLAND
FRANCE
BELGIUM
DELLER K.
Degussa Wolfgang AC-AT 3-CK Postfach 1345 6450 Hanau F.R.G.
DELMAS H.
CNRS. U.A. 192 Laboratoire de Genie Chimique Chemin de la Loge 31078 Toulouse Cedex FRANCE
DELMON B.
U.C.L. "- Groupe de Physico-Chimie Minerale et de Catalyse Place Croix du Sud. 1 1348 Louvain-la-Neuve BELGIUM
DER KINDEREN
VEG-GAS Instituut Wilmersdorf 50 - Postbus 137 7300 AC Apeldoorn THE NETHERLANDS
DERKS L.J.G.M.
Dow Chemi ca1 (Ned.) BV Postbus 48 4530 AA Terneuzen THE NETHERLANDS
843
BYRNE J.W.
Engelhard Corporation Menlo Park U.S.A. Edison, NJ 08818
CAHEN R.
Labofina SA Zone Industrielle 6520 Feluy BELGIUM
CAILLOD M.
RhOne-Poulenc Recherches Rue des Cardinaux, 12 93306 Aubervilliers FRANCE
CAMPELO PEREZ J.M. Departamento de QUlmica Organica Facultad de Ciencias 14004 Cordoba SPAIN CANDLIN J. P.
ICI, Petrochemical and Plastics Division Research Department wilton, Middlesborough, Cleveland ENGLAND
CANDY J.P.
C.N.R.S., I.R.C. Avenue Einstein, 2 69626 Villeurbanne
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CANESSON P.
Universite de Poitiers Avenue du Recteur Pineau, 40 86022 Poitiers FRANCE
CARRUTHERS J.D.
American Cyanamid 1937, W. Main Street - P.O. Box 60 Stamford, Connecticut 06904 U.S.A.
CARTLIDGE S.
Grace GmbH Postfach 449 6520 Worms
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CARVALHO FONSECA M.
Petroflex Industria e Commercio Rua Parana Campos Eliseos Duque de Caxias Rio de Janeiro BRAZIL
CAVAN I F.
Enichem Sintesi Via Luini 241 20099 Sesto S. Giovanni
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CENTI G.
Istituto Tecnologie Chimiche Speciali Viale Risorgimento 4 40136 Bologna ITALY
CERBONI M.
Ch imet SpA
Via dei Laghi 31/33 52041 Badia Al Pino (AR)
ITALY
CHAUMETTE P.
Institut Fran~ais du Petrole B.P. 311 92506 Rueil Malmaison FRANCE
CHEN T.-N.
Chemical Society of China Research Institute of Nanjing Da-Chang-Zhen Nanjing, Jiangsu CHINA
844
BONNELLE J.-P.
Universite des Sciences et Techniques Li1le 1 Laboratoire de Catalyse Heterogene et Homogene Bat. C3 59655 Villeneuve d'Ascq Cedex FRANCE
BONNIER J.-M.
L.E.D.S.S. - Bat. de Chimie U.S.M.G. B.P. 68 38402 St Martin d'Heres Cedex
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BOOM A.Q.M.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
BOSCH H.
Twente University Box 217 7500 AE Enschede
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BOSSCHAERTS M.
Metallurgie Hoboken-Overpelt A. Greinerstraat 14 2710 Hoboken BELGIUM
BOURDILLON
E1 f France Centre de Recherche Elf Solaize B.P. 22 96360 St Symphorien FRANCE
BOURNONVILLE J.P.
Institut Fran~ais du Petro1e B.P. 311 92506 Ruei1 Ma1maison FRANCE
BRAHMA N.
State University of Utrecht Meentweg 99 3454 AR DE MEERN THE NETHERLANDS
BRANDU J.
Chemical Technology Chemical Center P.O. Box 124 22001 Lund SWEDEN
BRASSER C.D.
Harshaw Chemie BV Strijkviertel 67 - P.O. Box 19 3454 ZG DE MEERN THE NETHERLANDS
BREUKELAAR J.
Kon./Shell Laboratorium Badhuisweg 3 1031 CM Amsterdam THE NETHERLANDS
BROOKS C.S.
Recycle Metals Baldwin Lane, 41 Glastonbury, Connecticut 06033
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Rhone-Poulenc Recherches Rue des Cardinaux, 12 93306 Aubervilliers FRANCE
BUJADOUX K.
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Center for Industrial Research P.O. Box 350, Blindern 0314 Oslo 3 NORWAY
BAUWENS J.
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BELLUSSI G.
BELGIUf
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BENEDETTI F.
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BIBBY D.M.
D.S.I.R., Chemistry Division Private Bag Petone NEW ZEALAND
BICKLEY R. I.
University of Bradford School of Chemistry Bradford BD7 1DP ENGLAND
BIRKE
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BLEKKAN LA.
SINTEF Division of Applied' Chemistry 7034 Trondheim-NTH NORWAY
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BOCK W.
Degussa AG, Abt. FC-O Postfach 1345 6450 Hanau F.R.G.
BODART-RAVET I.
Facultes Univ. Notre-Dame de la Paix Laboratoire de Catalyse Rue de Bruxelles, 61 5000 Namur BELGIUM
BOELLAARD E.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
BOND G.C.
Brunel University Department of Chemistry Uxbridge UB8 3PH ENGLAND
BONGAARTS J.
Dow Chemical (Nederland) BV P.O. Box 48 4530 AA Terneuzen THE NETHERLANDS
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DEROUANE E.G.
Facultes Univ. Notre-Dame de la Paix Laboratoire de Catalyse Rue de Bruxelles, 61 5000 Namur BELGIUM
DESCHRYVER P.
Interox, c/o Solvay &Cie Rue de Ransbeek, 310 1120 Bruxelles BELGIUM
DES COURIERES T.
ELF France Centre de Recherche Elf Solaize B.P. 22 69360 St Symphorien FRANCE
DESTRIAU M.
Universite de Bordeaux Rue de Loustalot, 75 33170 Gradignan FRANCE
DI RENZO F.
Politecnico di Milano Dipartimento di Chimica Industriale e Ingegneria Chimica "G. Natta" Piazza L. da Vinci, 32 20133 Milano ITALY
DOGSON I.L.
Johnson Matthey Chemicals Orchard Road Royston, Herts. SG8 5HE
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DOESBURG E.B.M.
Technische Hogeschool Delft Lab. voor Anorganische and Fysische Chemie Julianalaan 136 2628 BL Delft THE NETHERLANDS
DOMESLE R.
Degussa AG, ZN Wolfgang Dept. FC-PH-AK P.O. Box 1345 6450 Hanau 1 F.R.G.
DOTSCH H.
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DROSTE
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DYPVIK T.
Technical University of Trondheim 7034 Trondheim-NTH NORWAY
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I.C.I. Plc Plastics and Petrochemicals Division The Heat Runcorn, Cheshire WA7 4QD ENGLAND
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GUTEC GmbH Seligmannallee 1 - Postfach 220 3000 Hannover 1 F.R.G.
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C.S.I.C. - Instituto de Catalisis Serrano 119 28006 Madrid SPAIN
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Regional Technical College Kil kenny Road Carlow IRELAND
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Hills AG GB BC/Anorganische Abt. Postfach 13 20 4370 Marl F.R.G.
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Aluminum Company of America Alcoa Technical Center Alcoa Center, PA 15069 U.S.A.
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Universita di Milano Dipartimento di Fisica ed Elettr. Via C. Golgi, 19 20133 Milano ITALY
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Lonza Ltd. C.F. Meyerstrasse 4059 Basel SWITZERLAND
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Lonza Ltd. Milnchensteinerstrasse 38 SWITZERLAND 4002 Basle
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Universite Louis Pasteur - LCCS Laboratoire de Catalyse - CNRS Rue Blaise Pascal, 4 67000 Strasbourg FRANCE
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State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
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F.R.G.
GRANGE P.
U.C.L. Groupe de Physico-chimie Minerale et de Catalyse Place Croix du Sud, 1 1348 Louvain-la-Neuve BELGIUM
GREMMELMAIER C.
CIBA-GEIGY Werke Schweizerhalle AG 2093.4.14 4133 Schweizerhalle SWITZERLAND
GRIFFIN K.G.
Johnson Matthey Chemicals Orchard Road Royston, Herts. SG8 5HE
GRITZ E.V.M.
GROENENDAAL W.
Sasol One Ltd. P.O. Box 1 Sasolburg 9570 G.C. I. Gooland 13 2631 CP Nootdorp
ENGLAND
SOUTH AFRICA
THE NETHERLANDS
GROOT C. K.
Dow Chemical (Ned.) BV P.O. Box 48 4530 AA Terneuzen THE NETHERLANDS
GROS J.
C.N.R.S. L.T.P.C.M. - ENSEEG - DU B.P. 75 38402 St Martin d'Heres
FRANCE
GUCZI L.
Institute of Isotopes of the Hungarian Academy of Sciences P.O. Box 77 1525 Budapest HUNGARY
GUISNET M.
Universite de Poitiers Lab. de Chimie VII "Catalyse Organique" Avenue du Recteur Pineau, 40 86022 Poitiers Cedex FRANCE
HAGENAARS R.
Harshaw Chemie BV Postbus 19 3454 ZG De Meern
THE NETHERLANDS
849
HALVORSEN S.
Center for Industrial Research P.O. Box 350, Blindern 0314 Oslo 3 NORWAY
HANNAN W. K.
University of Wollongong Northfields Avenue Wollongong NSW 2500 AUSTRALIA
HANSEN J.H. Bak
Haldor Tops_e A/S Nym_llevej 55 2800 Lyngby DENMARK
HANSON C.B.
I.C.I. - Plastics and Petrochemicals Division R. and T. Department Wilton, Middlesborough, Cleveland ENGLAND
HATTORI T.
Nagoya University Department of Synthetic Chemistry Chikusa Nagoya 464 JAPAN
HAVIL
Laboratoire de Reactivite de Surface Tour 54-55 - 2eme etage Place Jussieu, 4 75252 Paris FRANCE
HECQUET G.
CdF Chimie C.R.N.C. B.P. 57 62 Mazingarbe
FRANCE
HEINRICH L.
Hills AG - Ethyl enabteil ung Pos tfach 1320 4370 Marl F.R.G.
HELLEBORG S.
Norsk Hydro A/S Research CTR. Porsgrunn Fabr. P.O. Box 110 3901 Porsgrunn NORWAY
HERCULES C.M.
University of Pittsburgh Department of Chemistry Pittsburgh. PA 15260
HERMANN J.M.
Ecole Centrale de Lyon photocatalyse B.P. 163 69131 Ecully Cedex
U.S.A.
FRANCE
HIGASHIO Y.
Sumitomo Chemical Co. Ltd. Chiba Research Laboratory Anesaki-Kaigan 3-1 JAPAN Ichihara-City. Chiba
HILAIRE L.
Universite Louis Pasteur Laboratoire de Catalyse et Chimie des Surfaces Rue Blaise Pascal. 4 67000 Strasbourg FRANCE
850
HOEK A.
Kon./Shell Laboratorium Postbus 3003 1003-AA Amsterdam THE NETHERLANDS
HOFMANN H.
Universit~t Erlangen-NUrnberg Egerlandstrasse 3 852 Erlangen F.R.G.
HOLMEN A.
University of Trondheim Department of Industrial Chemistry 7034 Trondheim NORWAY
HOLMES M.W.
I.C. I. Plc. West Bradford Road Clitheroe, Lancashire
ENGLAND
BP Research Centre Chertsey Road Sunbury-on-Thames, Middlesex
ENGLAND
HOLT A.
HOOGSTOEL M.
Reilly Chemicals S.A. Rue de Villerot 7350 Hautrage BELGIUM
HORN G.
Ruhrchemie AG Katalyseforschung Postfach 13 01 35 4200 Oberhausen 11
HOWDEN M.G.
IMAI T.
CSIR - CERG P.O. Box 395 Pretoria 0001
F.R.G.
SOUTH AFRICA
Signal Research Center, Inc. 50 E. Algonquin Road Des Plaines, Illinois 60017-5016
U.S.A.
IRANDOUST S.
Chalmers University of Technology Department of Chemical Reaction Engineering 41296 Gtlteborg SWEDEN
IRGANG M.
BASF Aktiengesellschaft 6700 Ludwigshafen F.R.G.
JACOBS P.A.
K.U.L. Laboratorium voor Oppervlaktechemie Kardinaal Mercierlaan 92 3030 Heverlee BELGIUM
JANNES G.
CERIA Service de Chimie Physique IFIMC Avenue Gryson 1 1070 Bruxelles BELGIUM
JANSSEN F.
KEMA - Chem. Dept. P.O. Box 9035 Arnhem THE NETHERLANDS
851
JENKINS J. W.
Johnson Matthey Technology Centre Blountscourt Sonning Convnon Reading RG4 9NH ENGLAND
JENKINS R.J.
Air Products & Chemicals, Inc. Box 2842 Lehigh Valley, PA 18001 U.S.A.
JENSEN E.J.
Haldor Tops~e A/S Nymo 11 evej 55 2800 Lyngby DENMARK
JENSEN P.E.
Haldor Tops¢e A/S Nymollevej 55 2800 Lyngby DENMARK
JOCKERS R.
Bergbau-Forschung GmbH Franz-Fischer-Weg 61 4300 Essen 13 F.R.G.
JOHANSEN K.
Haldor Tops~e A/S Nymollevej 55 2800 Lyngby DENMARK
JOHNSON J.W.
EXXON Research and Engineering Annandale, NJ 08801 U.S.A.
JOUD J.Ch.
LN.P.G. - CNRS - L.T.P.C.M. Domaine Universitaire B.P. 75 38402 St Martin d'Heres FRANCE
KAMEYAMA M.
Idemi tsu Kosan 25 St. James' Street London SW1A 1HA ENGLAND
KERA Y.
Kinki University 3-4-1 Kowakae, Higashiosaka Osaka 577 JAPAN
KHARE G.P.
Engelhard Corporation Menlo Park CN 28 Edison, NJ 08818 U.S.A.
KIRCHEN R.
Syncrude Canada Ltd. P.O. Box 5790 - Station L 17 Street, 10120 Edmonton, Alberta T6C 4G3
CANADA
KIWI J.
LP.F.L. Institut de Chimie Physique 1015 Lausanne SWITZERLAND
KNUUTTILA P.
Neste Oy Technology Center 06805 Kulloo FINLAND
KONOK F.
Dev. Compo for Org. Chern. Ind. Stahly u, 13 1085 Budapest HUNGARY
852
KRABETZ R.
BASF Aktiengesellschaft Abt. Zak 15 6700 Ludwigshafen F.R.G.
KRAUSE R.
Degussa AG - AC-GKA-CK P.O. Box 11 05 33 6000 Frankfurt 11 F.R.G.
LA GINESTRA A.
University of Rome Dipartimento di Chimica Piazza A. Moro, 5 Rome ITALY
LANCASTER P.
Dycat International 381 Fulwood Road Sheffield 5103GB
ENGLAND
LANSINK ROTGERINK H.G.J.
Twente University of Technology P.O. Box 217 7500 AE Enschede THE NETHERLANDS
LAPUYADE
Atochem Usines de Maurienne 73130 La Chambre
FRANCE
LEDOUX M.J.
CNRS - LCCS Universite L. Pasteur Rue Blaise Pascal, 4 67000 Strasbourg FRANCE
LEHMANN U.
Condea Chemie GmbH Mittelweg 13 2000 Hamburg 13
LEIGH D. LESTER G.R.
F.R.G.
Engelhard Corporation Vall ey Road Cinderford, Glos. GL14 2PB
ENGLAND
Allied-Signal Engin. Materials Research Center 50 E. Algonquin Road - Box 5016 Des Plaines, Illinois 60017-5016
U.S.A.
LI Dadong
Research Institute of Petroleum Processing P.O. Box 914 Beijing CHINA
LIGORATI F.
Enichem Sintesi Via Luini 241 20099 Sesto S. Giovanni
ITALY
LINDQVIST G.
Chemical Technology Chemical Center P.O. Box 124 22100 Lund SWEDEN
LIU Fu
Hebei Chemical Engineering Institute Department of Chemical Engineering Catalysts Laboratory CHINA Shijiazhuang
853
LIVRELL J.
EKA Nobel AB 44501 Surte
LOEWENSTEIN M.A.
Engelhard Corporation 429 Delancy Street Newark, NJ 07.105 U.S.A.
LOVINK H.J.
Ketjen Catalysts - Akzo Chemie P.O. Box 247 3800 AE Amersfoort THE NETHERLANDS
LUNA MARTINEZ D.
Departamento de Quimica Organica Facultad de Ciencias 14004 Cordoba SPAIN
LYCOURGHIOTIS A.
University of Patras Chemistry Department Patras GREECE
MAIRE G.
Universite L. Pasteur Laboratoire de Catalyse - U.A. 423 CNRS Rue Blaise Pascal, 4 67000 Strasbourg FRANCE
MARCQ J.P.
Societe Fran~aise des Petroles BP Centre de Recherches de Dunkerque B.P. 4519 59381 Dunkerque Cedex 1 FRANCE
MARENGO S.
Stazione Sperimentale Combustibili Viale A. De Gasperi. 3 20097 San Donato Milanese ITALY
MARGITFALVI J.L.
Central Research Institute for Chemistry P.O. Box 17 1025 Budapest HUNGARY
MARINAS RUBIO J.M.
Departamento de Quimica Organica Facultad de Ciencias 14004 Cordoba SPAIN
MARSH C.
Engelhard Corporation 429 Delancy Street Newark, NJ 07105 U.S.A.
MARTENS J.
K.U.L. Laboratorium voor Oppervlaktechemie Kardinaal Mercierlaan 92 3030 Leuven BELGIUM
MARTENS L.R.M.
K.U.L. Laboratorium voor Oppervlaktechemie Kardinaal Mercierlaan 92 3030 Leuven BELGIUM
MARTIN F.-G.
Institut fUr Techn. Chemie I Egerlandstrasse 3 8520 Erlangen F.R.G.
SWEDEN
854
MARTINS FERREIRA J.
Petroflex Industria e Commercio SA Rua Parana s/m Campos Eliseos duque de Caxia Ri 0 de Janei ro BRAZIL
MATTHIEU S.
CERIA / IFIMC Service de Chimie Physique Avenue E. Gryson, 1 1070 Bruxelles BELGIUM
McNAIR R.
Johnson Matthey Inc. 2001 Nolte Drive West Deptford, NJ 08066
U.S.A.
MEIMA G.R.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
MELE P.
Dycat International 381 Fulwood Road Sheffield S10 3GB
ENGLAND
MELO F.V.
Instituto de Catalisis Serrano 119 28006 Madrid SPAIN
MENDIOROZ E.
C.S. I.C. Serrano 119 28006 Madrid
SPAIN
MENON P.G.
Dow Chemical (Nederland) BV P.O. Box 48 4530 AA Terneuzen THE NETHERLANDS
MENZER R.
KFA JUlich - IRB Postfach 1913 5170 JUlich F.R.G.
MEURIS T.
Shell Belgium Passagierstraat 100 9000 Gent BELGIUM
MICHEL J.
E.I. Du Pont de Nemours Experimental Station E356 Wilmington, DE 19898 U.S.A.
MIZUKAMI F.
National Chemistry Laboratory for Industry Catalysis Division Yatabe Ibaraki 305 JAPAN
MOFFAT J.B.
University of Waterloo Department of Chemistry Waterloo. Ontario N2L 3G1
MONE R.
Akzo Chemie Postbus 15 1000 AA Amsterdam
CANADA
THE NETHERLANDS
855
MONTI D.
Ciba-Geigy AG R-1055 615 Postfach 4002 Basel
SWITZERLAND
MORETTI G.
Centro CNR (SACSO) Dip. Chimica Universita la Sapienza 60185 Roma ITALY
MOSER M.D.
Signal Eng. Mat. Res. CTR 50 E. Algonquin Road Des Plaines, Illinois 60017-5016
U.S.A.
MOUREAUX P.
Shell Recherche SA Centre de Recherches de Grand Couronne 76530 Grand Couronne FRANCE
MRMOLJA N.
I.P.L.A.S. - Kemicna Industrija Koper o. sd. o. Krpanova 2 66000 Koper YUGOSLAVIA
NAGY J.-B.
Facultes Univ. Notre-Dame de la Paix Rue de Bruxelles, 61 5000 Namur BELGIUM
NAKASHIMA C.
Mitsubishi Chemical Industries Yokkaishi Plant 1 - Toho-cho, Yokkaichi Mie 510 JAPAN
NEWSON E.
Alusuisse R&D 8212 Neuhausen
NOACK K.
F. Hoffmann-La Roche & Co, AG Grenzacherstrasse 124 4002 Basel SWITZERLAND
NOTARI B.
Eni Coordinamento Ricerca Innovazione Via Pisticci, 1 20097 San Donato Milanese ITALY
NOWAK W.
Flelt AG Dufourstrasse 101 - Postfach 393 8034 ZUrich SWITZERLAND
NOWECK K.
Condea Chemie GmbH Fritz-Staiger-Strasse 15 2212 BrunsbUttel 1 F.R.G.
ODENBRAND I.
Chemical Technology Chemical Center P.O. Box 124 22001 Lund SWEDEN
OGINO Y.
Tohoku University - Faculty Eng. Department of Chemical Engineering Aoba Aramaki Sendai 980 JAPAN
SWITZERLAND
856
OKUHARA T.
University of Tokyo Department of Synthetic Chemistry Faculty of Engineering Tokyo 113 JAPAN
ONO Y.
Tokyo Institute of Technology Department of Chemical Engineering Meguro-ku Tokyo 152 JAPAN
ONUMA K.
Mitsubishi Chemical Industries Research Center 1000 Kamoshida-cho, Midori-ku Yokohama 227 JAPAN
ORMANIEC W.
Institute of Industrial Chemistry ul. Rydygiera, 8 1-793 Warsaw POLAND
OTTERSTEDT J.E.
Chalmers University of Technology Department of Chemical Engineering 41296 Gllteborg SWEDEN
OUDEJANS J.C.
Unilever Research Olivier van Noortlaan 120 3133 AT Vlaardingen THE NETHERLANDS
PACKET D.
K.U.L. Laboratorium voor Oppervlaktescheikunde Kardinaal Mercierlaan 92 3030 Leuven BELGIUM
PAIS DA SILVA M.I.
COPENE - Polo Petroquimico Rua Eteno 1561 40000 Camacari - Bahia
BRAZIL
PAJONK G.M.
Universite Cl. Bernard - Lyon Laboratoire associe CNRS 231 Bld. du 11 Novembre 1918, 43 69622 Villeurbanne Cedex FRANCE
PALMISANO L.
Universita di Palermo Istituto di Ingegneria Chimica Viale delle Scienze 90128 Palermo ITALY
PATRONO P.
C.N.R. Via Salaria Km 29300 Monte Rotondo Scalo, C.P. 10 Roma ITALY
PAUCOT A.
Solvay & Cie SA Rue de Ransbeek, 310 1120 Bruxelles BELGIUM
PEDERSEN L.A.
Aluminum Company of America Alcoa Technical Center Alcoa Center, PA 15069 U.S.A.
857
PERNICONE N.
Montedison Spa Istituto Guido Donegani Via Fauser 4 28100 Novara ITALY
PERRICHON V.
Institut de Catalyse Avenue A. Einstein, 2 60626 Villeurbanne Cedex
PHILP J.
FRANCE
Engelhard Corporation Valley Road Cinderford, Gloucestershire
ENGLAND
PINNA F.
Universita di Venezia Dipartimento di Chimica Calle Larga S. Marta. 2137 30123 Venezia ITALY
PLATZER O.
Laboratoire de Physico-chimie des Rayonnements Bat. 350 91405 Orsay Cedex FRANCE
POElS E.K.
Unilever Research Olivier van Noortlaan 120 3133 AT Vlaardingen THE NETHERLANDS
POHl J.
Henkel KGAA Henkelstrasse 67 4000 DUsseldorf
F.R.G.
PONCELET G.
U.C.l. Groupe de Physico-chimie Minerale et de Catalyse Place Croix du Sud, 1 1348 Louvain-la-Neuve BELGIUM
POPOV T.S.
Higher Institute of Agriculture "V. Kolarov" Plovdiv BULGARIA
PREUSS U.
Ruhr-UniversitKt Bochum Lehrstuhl fUr Technische Chemie Postfach 10 21 48 4630 Bochum F.R.G.
PRIGENT M.
Institut Fran~ais du Petrole 1-4 avenue Bois Preau - B.P. 311 92506 Rueil Malmaison FRANCE
PUTTEMANS J.P.
CERIA Service de Chimie Physique - IFIMC Avenue E. Gryson. 1 1070 Bruxelles BELGIUM
RAJARAM P.
Alchemie Research Centre P.O. Box 155 Thane 400 601. Maharashtra
RAJBENBACH-RAVON L.A.
E.P.F.L. - Ecublens Institut de Chimie Physique 1015 Lausanne SWITZERLAND
INDIA
858
RASMUSSEN B.S.
Haldor Tops~e A/S Nymollevej 55 28000 Lyngby DENMARK
REINALDA D.
Kon./Shell Laboratorium Dept. CRD Postbus 3003 1003 AA Amsterdam THE NETHERLANDS
RESASCO D.E.
Yale University 9 Hillhouse Avenue New Haven, CT 06520
U.S.A.
RICHARD D.
CNRS - Institut de Recherches sur la Catalyse Avenue A. Einstein, 2 69626 Villeurbanne FRANCE
RICHARDSON J.T.
University of Houston Department of Chemical Engineering University Park Houston, Texas 77004 U.S.A.
RIESER K.
Grace GmbH P.O. Box 449 6520 Worms
F.R.G.
ROOS J.A.
Twente University of Technology P.O. Box 217 7500 AE Enschede THE NETHERLANDS
ROSS J.R.H.
Twente University of Technology P.O. Box 217 7500 AE Enschede THE NETHERLANDS
RUIZ P.
U.C.L. Groupe de Physico-chimie Minerale et de Catalyse Place Croix du Sud, 1 1348 Louvain-la-Neuve BELGIUM
SCHANKE D.
University of Trondheim Department of Industrial Chemistry 7034 Trondheim-NTH NORWAY
SCHAPER H.
Kon./Shell Laboratorium Badhuisweg 3 1031 CM Amsterdam THE NETHERLANDS
SCHEVE J.
Zentral Institut fur Physikalische Chemie RUdower Chaussee 5 1199 Berlin-Adershof D.D.R.
SCHMIDT L.D.
University of Minnesota Department of Chemical Engineering and Materials Science Minneapolis, Minnesota 55455 U.S.A.
SCHNE IDER P.
Institute of Chemical Process Fundamentals Czechoslovak Academy of Sciences Rozvojova 135 16502 Praha 6 CZECHOSLOVAKIA
859
SCHNEIDER G.
Unilever Research Laboratory P.O. Box 114 3130 AC Vlaardingen THE NETHERLANDS
SCHOLTEN J.J.F.
Central Laboratory DSM Catalysis Department Geleen THE NETHERLANDS
SCHUBERT P.J.
Degussa AG, ZN Wolfgang Dept. FC-PH-AK P.0. Box 13 45 6450 Hanau 1 F.R.G.
SCHUTZ A.
University of Pittsburgh School of Engineering Chemical and Petroleum Engineering Dept. 1249 Benedum Hall Pittsburgh, PA 15261 U.S.A.
SCHWARTZ J.T.
E.I. du Pont de Nemours C &P Department Jackson Laboratory Wilmington, DE 19898
SCHWARZ J.A.
SERMON P.A.
Syracuse University Chemical Engineering 320 Hinds Hall Syracuse, NY 13244
U.S.A.
U.S.A.
Brunel University Department of Chemistry Uxbridge, Middlesex
ENGLAND
SHIBATA M.
Idemitsu Kosan Co Ltd. Central Research Laboratory 1280 Kamiizumi, Sodegaura, Kimitu-gun Chiba 299-02 JAPAN
STAAL L.H.
Unilever Research Olivier van Noortlaan 120 3133 AT Vlaardingen THE NETHERLANDS
STEPHENS H.P.
Sandia National Laboratories Division 6254 Albuquerque, NM 87185 U.S.A.
STERTE J.
Chalmers University of Technology Department of Chemical Engineering 41296 GBteborg SWEDEN
STOBBE D.E.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
STRAUSZ D.P.
University of Alberta Department of Chemistry Edmonton, Alberta T6G 2G2
STRINGARO J.-P.
Sulzer Bros. Ltd. Kost 0637 - Werk OW 8401 Winterthur
CANADA
SWITZERLAND
860
SUN G.
Goethe-Institut Ifflandstrasse 2-6 6800 Mannheim
F.R.G.
SVENSSON B.
Chemical Technology Chemical Center P.O. Box 124 22001 Lund SWEDEN
TAMBLYN W.H.
Johnson Matthey Inc. 2001 Nolte Drive West Deptford, NJ 08066
TAMHANKAR S.S.
TAUSZIK G.R.
THORSTEINSON E.
U.S.A.
The BOC Group Inc. Group Technical Center 100 Mountain Avenue Murray Hill, NJ 07974 Montedipe Spa Via San Pietro 50 20021 Bollate Mi
U.S.A.
ITALY
Union Carbide Corp. P.O. Box 8361 South Charleston, W. Virginia 25303
TIJBURG I. LM.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
TOKARZ M.
EKA Nobel AB 44501 Surte
TRAVERS Ch.
Institut Fran~ais du Petrole B.P. 311 92506 Rueil Malmaison FRANCE
TRONSTAD O.
University of Trondheim Laboratory of Industrial Chemistry 7034 Trondheim-NTH NORWAY
TWIGG M.V.
I.C. I. Plc. Agricultural Division Billingham, Cleveland TS23 1LB
ULRICH R.
Grace GmbH P.O. Box 449 6520 Worms
U.S.A.
SWEDEN
F.R.G.
UYTDEWILLIGEN D.
Metallurgie Hoboken-Overpelt Greinerstraat 1 Hoboken BELGIUM
VACCARI A.
Istituto Tecnologie Chimiche Speciali Viale Risorgimento 4 40136 Bologna ITALY
VAN BERGE P.C.
Rand Afrikaans University P.O. Box 524 Johannesburg SOUTH AFRICA
ENGLAND
861
van BEYNUM J.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
van der GRIFT C.J.G.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
van DIJK A.
Kon./Shell Laboratorium Dept. HCP Postbus 3003 1003 AA Amsterdam THE NETHERLANDS
van DILLEN A.J.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
van HENGSTUM A.J.
Akzo Chemie P.O. Box 15 1000 AA Amsterdam
van LEEUWEN W.F.
E.C.N. P.O. Box 1 1755 ZG Petten
THE NETHERLANDS
THE NETHERLANDS
VANNICE A.
Penn. State University Department of Chemical Engineering University Park, PA 16802 U.S.A.
van OMMEN J.G.
Twente University of Technology P.O. Box 217 7500 AE Enschede THE NETHERLANDS
van REISEN C.A.M.
Akzo Chemie (Nederland) BV P.O. Box 25 7550 GC Hengelo (0) THE NETHERLANDS
van STIPHOUT P.C.M.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
VAN'T HOF H.A.
Akzo Zout Chemie P.O. Box 25 7550 GC Hengelo
VERMAIRE D.
Sasol One P.O. Box 1 Sasolburg 9570
THE NETHERLANDS
SOUTH AFRICA
VERSLUIS F.
Harshaw Chemie BV Strijkviertel 67 - P.O. Box 19 3454 ZG De Meern THE NETHERLANDS
VILLA P.
Politecnico di Milano Piazza L. da Vinci, 32 20133 Milano ITALY
VOGT E.T.C.
State University of Utrecht Croesestraat 77A 3522 AD Utrecht THE NETHERLANDS
862
WADDILOVE A.
WAGNER B.E.
British Petroleum BP Research Centre Chertsey Road Sunbury-on-Thames, Middl. TW16 7LN Union Carbide Corporation P.O. Box 670, 200-2-19 Bound Brook, NJ 08805
ENGLAND
U.S.A.
WAL W. K.
Institute of Industrial Chemistry ul. Rydygiera 8 01-793 Warsaw POLAND
WALZ B.
Institute of Physics Klingelbergstrasse 82 4056 Basel SWITZERLAND
WARD J.W.
Unocal Corporation 3765 Valencia Avenue Brea, California 92621
U.S.A.
WESSEL J.
Consortium fUr Elektrochem. Ind. GmbH Zielstattstrasse 20 8000 MUnchen 70 F.R.G.
WETSTEIN H.
Nepera Inc. Fine Chemicals Division Harriman, NY 10926 U.S.A.
WIECZOREK F.
Bergbau-Forschung Essen Franz-Fischer-Weg 61 Essen F.R.G.
WIELERS A.F.H.
Kon./Shell Laboratorium Badhuisweg 3 1031 CM Amsterdam THE NETHERLANDS
WIESENHAAN H.
Unilever Research Olivier van Noortlaan 120 3133 AT Vlaardingen THE NETHERLANDS
WILHELM F.C.
Air Products &Chemicals Inc. Box 2842 - P.O. Box 538 Lehigh Valley, PA 18001 U.S.A.
WILLS G.B.
Virginia Polytecnic Institute Department of Chemical Engineering Blacksburg, Virginia 24061-6496
WILMS M.
KFA JUlich - IRB Postfach 1913 5170 JUlich F.R.G.
WOLD J.
Norsk Hydro A/S Research CTR, Porsgrunn Fabr. P.O. Box 110 3901 Porsgrunn NORWAY
U.S.A.
863
WUNDE C.
HUls AG Ethylenabteilung Postfach 1320 4370 Marl F.R.G.
WUNDER F.
Hoechst AG Fo. Aliphat. Zwipro. Postfach 80 03 20 6230 Frankfurt 80
F.R.G.
XU Xiaoding
Technische Hogeschool Delft Laboratorium voor Anorg. en Fys. Chemie Julianalaan 136 2628 BL Delft THE NETHERLANDS
YANG Y.
Dalian Institute of Chemical Physics CHINA Dalian
YIN X.Z.
Research Institute of Petroleum Processing P.O. Box 914 Beijing CHINA
ZHAO Jiusheng
Tianjin University Chemical Engineering Department Tianjin CHINA
ADDENDUM EKA Nobel AB 445 01 Surte ~1ALETZ
G.
SWEDEN
SUd-Chemie AG Waldheimer Strasse 15 8206 BruckmUhl-Heufeld
F.R.G.
MAZZOCCHIA C.
Politecnico di Milano Dipartimento di Chimica Industriale e Ing. Chimica Piazza L. da Vinci, 32 20133 Milano ITALY
TEMPEST! E.
Politecnico di Milano Dip. di Chimica Industriale e Ingegneria Chimica Piazza L. da Vinci, 32 20133 Milano ITALY
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865
AUTHOR INDEX
AGOSTINI. G. AKRATOPULO, A. ARMBRUSTER. E. ARNTZ. D. BAlKER, A. BAKER. B.G. BARTEN. H. BAYENSE. C.R. BECHADERGUE-LABICHE. C. BERRY. F.J. BETTEGA DE PAULI, L.C. BICKLEY. R. I. BIELANSKA. E. BIELANSKI, A. BILOEN. P. BLANCHARD. M. BOND. G.C. BONNELLE. J.P. BONGAARTS. J.E. BOSCH, H. BOURNONVILLE. J.P. BOUTROIS. J.P. BROOKS. C.S.
569 309 389 137 389 455 103 55 725 467 713 297 217 217 165 725 827 581 83 151 713 581 375
DANNER. A. DAVEY, R.J. DEKKER. B.G. DE KONING. B. DELCOUR. K. DELMON. B. DEL PIERO. G. DENG, M. DEROUANE. E.G. DESTRIAU. M. DI RENZO. F. DISDIER. J. DOESBURG. E.B.M. DONKER, H. DOSCH. R.G. DU. H. DURAND. D.
343 15 83 767 83 591. 605 753 701 505 809 175 285 767 55 271 467 739
ELLIOTT. P.
15
FENOGLIO. R.J. FERRETTI. O.A. FIERRO. J. L. G. FOLGADO. M.A.
125 713 591,605 619
CAMPELO, J.M. CAMRA. J. CANDY. J.P. CANESSON, P. CARDEW. P.T. CAVALCANTI. F.A.P. CAVANI, F. CENTI. G. CHAUMETTE. P. CHRZASZCZ, J. CLARK. N.J. CONNER, W.C. CORMA, A. COURTY, P.
199 217 713 725 15 165 227 227 739 217 455 323 557 739
GALLEZOT. P. GAO, S. GARCIA. A. GELLINGS. P.J. GEOFFROY, G.L. GEUS. J.W. GIUFFRE. L. G8BOdiS, A. GONZALEZ-CARRENO, T. GONZALEZ-VELASCO. J.R. GRANGE, P. GREB, H. GRIMBLOT. J.
71 421 199 151 479 55. 83. 103, 401 175 689 297 619 591. 605 493 581
866
GROBET, P.J. GRONCHI, P. GUENTHERODT, H.J. GUO, X. GUSI, S. GUTTIEREZ-ORTIZ, J.I. GUTTIEREZ-ORTIZ, M.A.
531 175 389 701 753 619 619
HALLER, G.L. HANSEN, J.B. HATTORI, T. HEGEDUS, M. HEISE, M.S. HERCULES, D.M. HERRMANN, J.M. HILAIRE, L. HOJLUND-NIELSEN, P. E. HOPPENER, R.H. HOUALLA, M.
125 785 113, 815 689 1 433 285 569 785 767 431
IMAI, S. IMAMURA, J. INOUE, T. ISHIHAMA, Y. ITOH, S. IVANOV, K.I.
45 45 259 259 113 191
JACOBS, P.A. JACOBSON, A.J. JANSSEN, F. JOCKERS, R. JOHNSON, J.W. JOHNSTON, D.C.
531 181 103 493 181 181
KAMADA, T. KAMINSKY, M. KAWASHIMA, T. KERA, Y. KERN-TALAS, E. KIKUCHI, N.
259 479 259 259 689 519
KLEIN, J. KLISSURSKI, D.G. KOEBRUGGE, G.W. KOUTSOUKOS, P.G.
493 191 401 309
LACROIX-REPELLIN, M. LANSINK-ROTGERINK, H.G.J. LECLERCQ, C. LEDFORD, J.S. LEDOUX, M.J. LEE, C. LEFERINK, R. LI, H. LIANG, D. LIN Liwu LIN, Y.J. LIN Zhiyi n LIU Fu LIU, Z. LUCAT, C. LUNA, D. LYCOURGHIOTIS, A.
333 795 285 433 569 421 103 701 467 467 125 701 443 649, 659 809 199 309
MABILLON, G. MAILLE, S. MAIRE, G. MARGITFALVI, J. MARINAS, J.M. MARTENS, L.R.M. MARTIN LUENGO, M.A. MATSUMOTO, H. MASUMOTO, T. MATSUKAZE, Y. MAZZOCCHIA, C. MEIMA, G.R. MELO, F. V. MENDIOROZ, S. MENIL, F. MESTERS, C.M.A.M. MICHEL, J.B. MIFSUD, A.
713 725 569 689 199 531 29, 619 815 353 259 175 83 557 619 809 401 669 557
867
MIZUKAMI, F. MOFFAT, J.B. MURAKAMI, Y.
45 241 113, 815
NAGY, J.B. NAJBAR, M. NIENOW, A.W. NIWA, S.
505 217 15 45
ONO, Y. ONUMA, K. ORMANIEC, W. OTTERSTEDT, J.E.
519 543 217 631
PAJARES, J.A. PAJONK, G.f4. PALMISANO, L. PARRINELLO, F. PEDERSEN, L.A. PESHEVA, J. PETERS, U. PETRAKIS, L. PICHAT, P. PIZZOLI, F. POPOV, T.S. PORTIER, J. PRADA SILVY, R. PRESCHER, G.
619 333 297 227 313 191 493 433 285 753 191 89 605 137
RASMUSSEN, B.S. RAVET, I. RESASCO, D.E. RICHARD, D. ROSS, J. R. H.
785 505 115 71 151, 795
SALARDENNE, J. SANZ, E. SCHLOEGL, R. SCHMIDT, L.D. SCHOLTEN, J.J.F. SCHUTZ, A. SCHWARTZ, J.T.
809 557 389 421 767 165 669
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1 29 353 45 29 271 631 659 649, 659 689 739
TAGAWA, T. TANG, R. TEICHNER, S.J. TEMPESTI , E. TOBA, M. TOULHOAT, H. TOURNIER, G. TRAVERS, C. TRIFIRO, F. TSUCHIYA, T.
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UNGER, K.K.
343
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753 83 103 401 83, 103 725 151, 795 55 479 479 531 55 785 103 309
WAL, W.
217
868
WALZ, B. WAMBEKE, A. WANG, C. WATANABE, H. WEIST, E.L. WELLS, P.B. WIELERS, A.F.H. WIGMAN, J.M.
389 581 467 519 323 827 401 83
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15 767
YAN, F. YANG, Y. ZHANG, Su
467
Also available: Preparation ofCatalysts I, II and III Scientific Bases for the Preparation ofHeterogeneous Catalysts Preparation of Catalysts I Proceedings of the International Symposium, Brussels, Belgium, 14-17 Oct. 1975
B. Delmon, P.A. Jacobsand G. Poncelet(editors) (Studies in Surface Science and Catalysis, 1) 1976 3rd repro 1987 xvi + 706 pages ISBN 0-444-41428-2
Preparation of Catalysts II Proceedings of the 2nd International Symposium, Louvain la Neuve, sept. 4-7,1978 B. Delmon, P. Grange, P. Jacobsand G. Poncelet (editors) (Studies in Surface Science and Catalysis, 3) 1979. 2nd repro 1987 xiv + 762 pages ISBN 0-444-41733-8
Preparation of Catalysts III Proceedings of the 3rd International Symposium, Louvain-Ia-Neuve, 6-9 Sept. 1982
G. Poncelet, P. Grangeand P.A. Jacobs (editors) (Studies in Surface Science and Catalysis, 16) 1983 xvi + 854 pages ISBN 0-444-42184-X For details write to:
ELSEVIER SCIENCE PUBLISHERS P. O. Box 211 - 1000 AH Amsterdam - The Netherlands
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