COMPREHENSIVE CHEMICAL KINETICS
COMPREHENSIVE Section 1. THE PRACTICE AND THEORY OF KINETICS Volume 1 Volume 2 Volume 3
The Practice of Kinetics The Theory of Kinetics The Formation and Decay of Excited Species Section 2. HOMOGENEOUS DECOMPOSITION AND ISOMERISATION REACTIONS
Volume 4 Volume 5
Decomposition of Inorganic and Organometallic Compounds Decomposition and Isomerisation of Organic Compounds Section 3. INORGANIC REACTIONS
Volume 6 Volume 7
Reactions of Non-metallic Inorganic Compounds Reactions of Metallic Salts and Complexes, and Organometallic Compounds Section 4. ORGANIC REACTIONS (6 volumes)
Volume 8 Volume 9 Volume I0 Volume 12 Volume 13
Proton Transfer Addition and Elimination Reactions of Aliphatic Compounds Ester Formation and Hydrolysis and Related Reactions Electrophilic Substitution at a Saturated Carbon Atom Reactions of Aromatic Compounds Section 5. POLYMERISATION REACTIONS (3 volumes)
Volume 14 Volume 14A Volume 15
Degradation of Polymers Free-radical Polymerisation Non-radical Polymerisation Section 6. OXIDATION AND COMBUSTION REACTIONS (2 volumes)
Volume 17
Gas-phase Combustion Section 7. SELECTED ELEMENTARY REACTIONS (1volume)
Volume 18
Selected Elementary Reactions Section 8. HETEROGENEOUS REACTIONS (4volumes)
Volume 19 Volume 20 Volume 21 Volume 22
Simple Processes a t the Gas-Solid Interface Complex Catalytic Processes Reactions of Solids with Gases Reactions in the Solid State Additional Section KINETICS AND TECHNOLOGICAL PROCESSES
CHEMICAL KINETIC: EDITED BY
C.H. BAMFORD M.A.,Ph.D., Sc.D. (Cantab.), F.R.I.C., F.R.S. Campbell-Brown Professor o f Industrial Chemistry, University of Liverpool AND
C.F.H. TIPPER Ph.D. (Bristol), D.Sc. (Edinburgh) Senior Lecturer in Physicaf Chemistry, University of Liverpool
VOLUME 20
COMPLEX CATALYTIC PROCESSES
ELSEVIER Amsterdam - Oxford - New York 1978
- Tokyo
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. 655, Avenue of the Americas New York, NY 10010,U.S.A. First edition 1978 Second impression 1991
Library of Congress CaIaloging in Publication Data
Barnford, C H 'Complex c a t a l y t i c processes. (Their Comprehensive chemical kinetics ; v. 20) Bibliography: p. Includes index. 1. Catalysis. I. Tipper, Charles Frank Howlett, j o i n t author. 11. T i t l e . QD501.€!242 vol.20 [QDSOSI 541l.39~ rs+1'.3951 78-4165 ISBN 0-W-41651-X
ISBN 0-444-41 651-X with 79 illustrations and 109 tables 0 Elsevier Science Publishers B.V., 1978
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./ Academic Publishing 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. (CCC 1, 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 publisher. No responsibility is assumed by the Publisher for any injury and/or damage to persons or property as a matter of products liability, negligence or otherwise, or from any use or operation of any methods, products, instructions or ideas contained in the material herein.
This book is printed on acid-free paper. Printed in The Netherlands
COMPREHENSIVE CHEMICAL KINETICS
ADVISORY BOARD Professor S.W. BENSON Professor SIR FREDERICK DAINTON Professor G. GEE the late Professor P. GOLDFINGER Professor G.S. HAMMOND Professor W. JOST Professor G.B. KISTIAKOWSKY Professor V.N. KONDRATIEV Professor K.J. LAIDLER Professor M. MAGAT Professor SIR HARRY MELVILLE Professor G. NATTA Professor R.G.W. NORRISH Professor S. OKAMURA the late Professor SIR ERIC RIDEAL Professor N.N. SEMENOV Professor Z.G. SZABO Professor 0. WICHTERLE
Contributors to Volume 20 L. BERANEK
Institute of Chemical Process Fundamentals, Czechoslovak Academy of Sciences, 165 0 2 Praha 6 - Suchdol, Czechoslovakia
M. KRAUS
Institute of Chemical Process Fundamentals, Czechoslovak Academy of Sciences, 165 0 2 Praha 6 - Suchdol, Czechoslovakia
P.J. VAN DEN BERG Department of Chemical Technology, Delft University of Technology, Delft, The Netherlands K. VAN DER WIELE
Department of Organic Products, Akzo Zout Chemie Nederland bv Research, Hengelo, The Netherlands
G. WEBB
Chemistry Department, The University, Glasgow, Scotland
Section 8 deals with reactions which occur at gassolid and solidsolid interfaces, other than the degradation of solid polymers which has already been reviewed in Volume 14A. Reactions at the liquidsolid interface (and corrosion) involving electrochemical processes outside the coverage of this series are not considered. With respect to chemical processes at gassolid interfaces, it has been necessary to discuss surface structure and adsorption as a lead-in to the consideration of the kinetics and mechanism of catalytic reactions. In Volume 20, complex processes catalysed by solids are covered. Chapter 1 deals with hydrogenation. After consideration of the nature of the metal catalysts, general aspects of the kinetics and alternative reaction pathways, the hydrogenation of olefins, alkynes, dienes and cyclic molecules are dealt with in detail. Finally, the relationship between catalyst structure and hydrogenation activity is discussed. Chapter 2 is concerned with heterogeneous oxidation processes. The oxidation of ethylene and propene, so important industrially, is considered at length and then higher olefins and aromatic hydrocarbons; the influence of ammonia (ammoxidation) is also discussed. There is a section on the oxidation of methanol, ammonia and sulphur dioxide and, to conclude, the role of the catalysts is considered. Elimination, addition and substitution processes occurring on solid acid-base catalysts are covered in the last chapter. These reactions include dehydration, deamination, dehydrohalogenation, dealkylation by cracking, dehydrosulphidation, hydration, hydrohalogenation, alkylation by olefins, aldol condensation, esterification and hydrolysis. The editors are very grateful for much invaluable advice from their colleague Professor D.A. King.
Liverpool January, 1978
C.H. Bamford C.F.H. Tipper
This Page Intentionally Left Blank
Contents ................................................. Chapter 1 (G . Webb) Catalytic hydrogenation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2 . General principles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1 Variables in the system . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Preface
2.1.1 Variables associated with the substrate . . . . . . . . . . . . . . . 2.1.2 Variables associated with the catalyst . . . . . . . . . . . . . . . . 2.2 Kinetics and the derivation of rate expressions . . . . . . . . . . . . . . . 2.2.1 Rate expressions for bimolecular surface reactions . . . . . . . 2.3 Selectivity and the concept of alternative reaction paths . . . . . . . . 2.4 Application of absolute rate theory t o bimolecular surface reactions 3. The hydrogenation of olefins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.2 Adsorbed states of olefins . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3 Possible reaction mechanisms . . . . . . . . . . . . . . . . . . . . . . . . . . 3.4 Treatment of experimental results . . . . . . . . . . . . . . . . . . . . . . . 3.5 Hydrogenation of ethylene . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.6 Hydrogenation of propene . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.7 Reactions of the n-butenes . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.8 Reactions of higher aliphatic olefins . . . . . . . . . . . . . . . . . . . . . . 4 . The hydrogenation of alkynes and alkadienes . . . . . . . . . . . . . . . . . . . . 4.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2 Nature of the adsorbed state of alkynes and alkadienes . . . . . . . . . 4.3 Possible reaction mechanisms . . . . . . . . . . . . . . . . . . . . . . . . . . 4.4 Treatment of experimental results . . . . . . . . . . . . . . . . . . . . . . . 4.5 The hydrogenation of acetylene . . . . . . . . . . . . . . . . . . . . . . . . 4.6 Hydrogenation of monoalkylacetylenes . . . . . . . . . . . . . . . . . . . 4.7 Hydrogenation of dialkylacetylenes . . . . . . . . . . . . . . . . . . . . . . 4.8 The hydrogenation of alka-l:2-dienes . . . . . . . . . . . . . . . . . . . . . 4.9 The hydrogenation of conjugated alkadienes . . . . . . . . . . . . . . . . 5 . The hydrogenation of cyclic molecules . . . . . . . . . . . . . . . . . . . . . . . . 5.1 The hydrogenation of alicyclio alkenes . . . . . . . . . . . . . . . . . . . . 5.2 The hydrogenation of cyclopropane . . . . . . . . . . . . . . . . . . . . . . 6 . Catalyst structure and hydrogenation activity . . . . . . . . . . . . . . . . . . . 6.1 Geometric factors in catalysis . . . . . . . . . . . . . . . . . . . . . . . . . . 6.2 Electronic factors in catalysis . . . . . . . . . . . . . . . . . . . . . . . . . . 6.3 Surface migration and the influence of catalyst supports . . . . . . . . 7 . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
vii
1 1 2 2 2 3 4 6 8 13 16 16 16 23 27 29 37 38 48 50 50 50 55 57 58 68 71 74 81 94 94 100 103 103 106 109 112 114
Chapter 2 (K . van der Wiele and P.J. van den Berg)
..............................
123
1. Scope of the chapter . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2 . Oxidation processes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1 Ethylene oxidation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
123 124 126
Heterogeneous oxidation processes
2.1.1 Ethylene oxide production . . . . . . . . . . . . . . . . . . . . . . . 2.1.2 Acetaldehyde and acetic acid production . . . . . . . . . . . . . 2.2 Propene oxidation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2.1 Propene oxide production . . . . . . . . . . . . . . . . . . . . . . . 2.2.2 Acrolein production . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2.3 Acrylic acid production . . . . . . . . . . . . . . . . . . . . . . . . . 2.2.4 Dimerization and aromatization . . . . . . . . . . . . . . . . . . . 2.2.5 Acetone production . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2.6 Ammoxidation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3 Butenes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3.1 Isobutene to methacrolein and methacrylanitrile . . . . . . . . 2.3.2 Normal butenes to butadiene, furan and maleic anhydride . . 2.3.3 Dimerization and aromatization of iso- and n-butenes . . . . . 2.3.4 Oxyhydration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.4 Higher olefins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5 Aromatic hydrocarbons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5.1 Benzene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5.2 Toluene and xylene . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5.3 Ortho-xylene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5.4 Naphthalene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5.5 Anthracene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5.6 Other aromatic hydrocarbons . . . . . . . . . . . . . . . . . . . . . 2.5.7 Ammoxidation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.6 Methanol . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.6.1 The silver process . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.6.2 Iron molybdate and other metal oxide catalysts . . . . . . . . . 2.7 Ammonia . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.7.1 The production of NO . . . . . . . . . . . . . . . . . . . . . . . . . . 2.7.2 The formation of N2 and NzO . . . . . . . . . . . . . . . . . . . . . 2.8 Sulphur dioxide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 . Role of the catalyst . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1 Thermodynamic considerations . . . . . . . . . . . . . . . . . . . . . . . . . 3.2 Metal-oxygen bond strength . . . . . . . . . . . . . . . . . . . . . . . . . . 3.2.1 Heat of formation of metal oxides, AHf . . . . . . . . . . . . . . 3.2.2 Heat of oxygen desorption . . . . . . . . . . . . . . . . . . . . . . . 3.2.3 I8O2 isotope exchange . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3 Oxygen transfer . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3.1 Participation of lattice oxygen . . . . . . . . . . . . . . . . . . . . 3.3.2 Role of Me=O type of oxygen . . . . . . . . . . . . . . . . . . . . . 3.3.3 Significance of 0; and 0- radicals . . . . . . . . . . . . . . . . . . 3.4 Aspects of charge transfer . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.4.1 Bulk electrical properties . . . . . . . . . . . . . . . . . . . . . . . . 3.4.2 Charge transfer on an atomic scale . . . . . . . . . . . . . . . . . . 3.5 Nature of the active sites . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5.1 Acid-base properties . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5.2 Bifunctionality . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.6 Adsorption and reaction complexes on the catalytic surface . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
126 133 135 136 137 159 160 162 164 174 175 179 194 195 195 196 197 204 210 217 218 219 221 224 224 225 227 228 228 230 231 231 233 233 234 234 235 236 239 241 242 243 244 247 248 250 251 253
Chapter 3 (L . Berinek and M . Kraus) Heterogeneous eliminations, additions and substitutions . . . . . . . . . . . . . . . . 263 1. General features . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.1 Correspondence between homogeneous and heterogeneous reactions
263 263
Nature of the catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2.1 Silica . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2.2 Alumina . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2.3 Aluminosilicates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2.4 Metal salts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2.5 Ion exchange resins . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2.6 The working surface . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.3 Type of kinetics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2 . Elimination reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1 Common features of heterogeneous catalytic eliminations . . . . . . . 2.1.1 Mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1.2 Orientation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1.3 Kinetics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2 Dehydration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2.1 Types of dehydration reactions . . . . . . . . . . . . . . . . . . . . 2.2.2 Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2.3 Experimental kinetic results . . . . . . . . . . . . . . . . . . . . . . 2.2.4 Mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3 Deamination . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3.1 Types of deamination reactions . . . . . . . . . . . . . . . . . . . . 2.3.2 Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3.3 Experimental kinetic results . . . . . . . . . . . . . . . . . . . . . . 2.3.4 Mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.4 Dehydrohalogenation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.4.1 Types of dehydrohalogenation reactions . . . . . . . . . . . . . . 2.4.2 Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.4.3 Experimental kinetic results . . . . . . . . . . . . . . . . . . . . . . 2.4.4 Mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5 Dealkylation by cracking . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5.1 Types of cracking reactions . . . . . . . . . . . . . . . . . . . . . . 2.5.2 Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5.3 Experimental kinetic results . . . . . . . . . . . . . . . . . . . . . . 2.5.4 Mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.6 Dehydrosulphidation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.6.1 Types of dehydrosulphidation reactions and catalysts . . . . . 2.6.2 Experimental kinetic results . . . . . . . . . . . . . . . . . . . . . . 2.6.3 Mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 . Addition reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1 Hydration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1.1 Hydration of olefins to alcohols . . . . . . . . . . . . . . . . . . . 3.1.2 Hydration of acetylene to acetaldehyde . . . . . . . . . . . . . . 3.1.3 Hydration of alkene oxides to glycols . . . . . . . . . . . . . . . . 3.2 Hydrohalogenation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.2.1 Types of hydrohalogenation reactions and catalysts . . . . . . 3.2.2 Experimental kinetic results . . . . . . . . . . . . . . . . . . . . . . 3.2.3 Mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3 Alkylation by olefins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3.1 Types of alkylation reactions and catalysts . . . . . . . . . . . . 3.3.2 Experimental kinetic results . . . . . . . . . . . . . . . . . . . . . . 3.3.3 Mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.4 Addition of alcohols to alkenes . . . . . . . . . . . . . . . . . . . . . . . . . 3.5 Aldol condensation and related reactions . . . . . . . . . . . . . . . . . . 3.5.1 Types of reaction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5.2 Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2
264 264 266 268 269 270 271 272 274 275 275 277 280 281 281 282 282 290 295 295 296 296 298 300 300 300 301 308 309 309 310 311 315 318 318 319 319 320 321 321 327 329 332 332 332 333 334 334 335 336 336 337 337 340
3.5.3 Experimental kinetic results . . . . . . . . . . . . . . . . . . . . . . 3.5.4 Mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4 . Substitution reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.1 Esterification and transesterification . . . . . . . . . . . . . . . . . . . . . 4.1.1 Types of reactions and catalysts . . . . . . . . . . . . . . . . . . . 4.1.2 Reactions catalysed by inorganic catalysts . . . . . . . . . . . . . 4.1.3 Reactions catalysed by organic polymer-based cation exchangers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2 Hydrolysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2.1 Hydrolysis of esters . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2.2 Other hydrolyses . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Index
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342 345 348 348 348 351 356 371 371 380 385 399
1 Chapter 1
Catalytic Hydrogenation GEOFFREY WEBB
1.Introduction Since the early reports of Sabatier et al. [l] of the catalysed hydrogenation of unsaturated hydrocarbons and aldehydes, a vast number of studies of catalytic hydrogenation have been reported. The extent of this literature is a reflection, in part, of the rich variety of systems and the wealth of attainable detailed information and, in part, of the great complexities which face the catalytic chemist. Selwood [ 21 in 1962 remarked: “No problems in surface chemistry have been more hotly debated than the adsorption and hydrogenation mechanisms for ethylene; and few debates have resulted in such meagre conclusions”. Some 15 years later, the subject remains one in which there is still considerable dispute and debate regarding detailed mechanisms, although over the past few years some measure of order has begun to appear. It is against this background of complexity and uncertainty that this chapter sets out to review the present status of the problem of heterogeneous catalytic hydrogenation from the standpoint of (a) the kinetics and TABLE 1 Types of catalysed hydrogenation reactions Reactant a
Possible reaction products
a
>c=c<
-c*-
>c=c=c< >C=CH-CH=C< >CH . (CH2)n+l . CH=C< ; >CH(CH?),+&H<
Alicyclic rings R . CHO R . CO . R’ RCOOH R . NO2 ___
a
~~
R = H, alkyl or aryl.
References pp. 1 1 4-1 21
Linear olefin; alkane R . CH2OH R . C O ( 0 H ) . R’ R . CHO; RCHzOH RNH2
2 mechanism of hydrogenation reactions and (b) the properties of the catalyst which give rise t o hydrogenation activity. In general, reactions involving the addition or abstraction of hydrogen are most efficiently catalysed by metals [3]. Those which have been found to possess hydrogenation activity are the nine Group VIII metals, of which nickel, palladium and platinum have been most extensively studied, together with rhenium, copper, gold and tungsten. Metal oxides have also been used as hydrogenation catalysts [ 4 ] although less commonly than metals. The principal types of catalysed hydrogenation reactions, together with the likely reaction products, are shown in Table 1.Of the reactions listed, the hydrogenation of monoolefins and diunsaturated hydrocarbons acetylenes and diolefins - have attracted most attention and, for this reason, much of the ensuing discussion will be concerned with these particular reactions,
2. General principles Before undertaking a detailed examination of individual catalysed hydrogenation reactions, we consider some of the more general aspects of the subject and highlight some of the problems associated with heterogeneous reactions in general, and catalytic hydrogenation in particular. 2.1 VARIABLES IN THE SYSTEM
Catalytic hydrogenation is, by definition, a bimolecular reaction in which it is now generally agreed that a t least one of the reactants, the unsaturated molecule, is adsokbed on the surface. The role and reactive state of hydrogen is less clear and, as we shall see later, hydrogen may either react from an atomically adsorbed state or as a physically adsorbed molecular state. The introduction of the surface as an active partner in the reaction gives rise t o complexities not present in homogenous systems and makes the establishment of the mechanism of a heterogeneous reaction an incomparably more difficult task than for a homogeneous reaction. The process of chemisorption of the substrate molecule a t the catalyst surface involves a chemical interaction between the substrate and an active site on the surface. The introduction of the adsorption process makes the heterogeneous system susceptible to a large number of variables, some of which are readily controllable, others less easily so. These variables may be basically divided into two classes, (1) those associated with the substrate and (2) those associated with the catalyst itself.
2.1 .l Variables associated with the substrate Using the hydrogenation of the olefinic double bond as an example, the double bond may occur in a wide variety of environments depending upon
3 the other substituents in the molecule. Thus, the number and chain length of the substituent alkyl groups may be varied; the double bond may be located in a molecule containing any number of other functional groups, some of which may themselves be susceptible t o hydrogenation, or the double bond may be part of two alicyclic rings. Such variations may give rise t o either or both of two effects. First, an electronic effect due to the electron-releasing or -withdrawing properties of the substituent may be apparent. This will result in variations in electron density within the double bond with a consequential effect upon the strength of interaction of the olefin with the surface and the energetics of subsequent surface processes. Secondly, there may be a geometric effect due t o the size of the substituents. Bulky groups may reduce the number of molecules which may be accommodated on the surface. Thus with molecules such as, for example, 2,2,5,5-tetramethylhex-3-ene, (CH3)3C. CH=CH * C(CH3)3, the packing of molecules on the surface is likely t o be such that hydrogen can adsorb non-competitively, whereas with other substrates such as ethylene, the chemisorption of hydrogen would occur competitively. The size of the substituents may also affect the distance of approach of the olefin t o the surface and hence the strength of the chemisorption bond. While such variables are readily controlled, they nevertheless give rise t o significant differences in kinetics and mechanism from one substrate t o another. 2.1.2 Variables associated with the catalyst Metal hydrogenation catalysts may be employed in any one of a variety of forms: (a) macroscopic forms as wires, foils or granules; (b) microscopic forms as powders obtained by chemical reduction, colloidal suspensions, blacks or evaporated metal films; (c) supported catalysts where varying concentrations of metal are dispersed to a varying degree on a carrier such as alumina, silica or carbon. Clearly, there is an infinite number of variations of the physical form of the catalyst which may be employed. One of the major problems is the production of a reproducible metal surface. Irreproducibility may be due t o any one of a number of factors, the following being some of the more important ones: (a) variations in the degree of cleanliness and state of reduction of the surface; (b) variation in the degree of exposure of certain crystallographic planes; (c) variation in the concentration of surface defects and (d) variation in the distribution of particle sizes. Most of these factors are not readily controllable and may not be without effect upon the rate and mechanism of the reaction being catalysed; it is important, therefore, that the effects of each of the variables is assessed independently. Metal single crystals and evaporated metal films have been used in an attempt t o overcome some of the difficulties of irreproducibility. HowReferences PP. 114-121
4 ever, these suffer from a number of disadvantages. Whereas single crystal surfaces can be prepared with relatively large areas of exposed and welldefined crystallographic planes and can be obtained in an extremely clean state, the actual surface area exposed for catalytic reaction is too small to be of practical use, unless special techniques are used [ 5-71. Similarly, evaporated metal films of relatively high surface area, which can be relatively easily prepared in a clean state, are susceptible to good characterisation. However, their inherent activity may lead to poisoning since the reactants may be strongly and irreversibly adsorbed. Thus, for example, acetylene cannot readily be hydrogenated on evaporated metal films [S] but is easily hydrogenated on supported metals [ 91. A further disadvantage of using metal films is that their high activity means that either (a) hydrogenation reactions performed at or about room temperature are so rapid as t o prevent the detection of intermediate products or (b) reactions must be studied a t low temperatures, thereby limiting the range of molecules studied or (c) flow systems must be used in order to minimise contact times [ 101. Supported metal catalysts are much easier to employ and have obvious attractions for industrial use from their ease of handling and economic considerations of obtaining maximum utilisation of the catalytically active metal, by using very small particles with a high surface-to-volume ratio, which are stable on the support and not susceptible to sintering. In spite of the inherent difficulties of variable activity, kinetics and activation energies [ 111 associated with their use, supported metals have been extensively used as hydrogenation catalysts. 2.2 KINETICS AND THE DERIVATION O F RATE EXPRESSIONS
A knowledge of the kinetic parameters and, in particular, the orders of reaction of a catalysed reaction is important to the accurate definition of the reaction mechanism. However, catalytic hydrogenation reactions proceed through a series of elementary steps, only one of which may be ratedetermining. In consequence, the observed rate expressions give little or no direct information about most of the steps involved and kinetics alone are not sufficient for a precise description of the mechanism. The general principles of the kinetics of surface-catalysed reactions have been discussed in detail elsewhere [ 12,131 and in this section we shall confine our attention to a discussion of the kinetics and rate expressions which are applicable to hydrogenation reactions. While in homogeneous systems the reaction is occurring throughout the entire volume of the reaction vessel and the partial pressures (concentrations) of the species participating in the rate-controlling step are often directly observable, the same is not true for heterogeneous systems. Here, reaction is confined to a monomolecular layer at the surface, around of the total volume of the reaction system, and the concentrations of
5 adsorbed species are not directly observable unless appropriate measurements can be made by suitable methods [ 14,151. In heterogeneous systems, the rate expressions have to be developed on the basis of (a) a relation between the rate and concentrations of the adsorbed species involved in the rate-determining step and (b) a relation between the latter and the directly observable concentrations or partial pressures in the gas phase. In consequence, t o obtain adequate kinetic rate expressions it is necessary t o have a knowledge of the reaction mechanism, and an accurate means of relating gas phase and surface concentrations through appropriate adsorption isotherms. The nature and types of adsorption isotherm appropriate t o chemisorption processes have been discussed in detail elsewhere [16,17] and will not be discussed further except t o note that, in spite of its severe theoretical limitations, the Langmuir isotherm is almost invariably used for kinetic interpretations of surface hydrogenation reactions. The appropriate equations are (a) For the adsorption of a single species A ha
A(gas) + A(ads.) hd
where eA is the fractional surface coverage of A , b A the adsorption coefficient (=ka/k,) and PA the pressure of A in equilibrium with the surface. (b) For the dissociative adsorption of a single species ha
A,(gas) + n A(ads.) kd
(c) For the competitive adsorption of two species on the same surface sites
In a catalytic reaction, the overall rate depends upon the relative rates of five distinct processes, any of which may be rate-determining. These are (i) mass transport of reactants t o the catalyst surface; (ii) adsorption of reactants, or a t least one of them, a t the surface; (iii) reaction between adsorbed species on the surface (this may involve more than one elementary step); (iv) desorption of products from the surface; (v) mass transport of the reactants away from the surface. References p p . 1 1 4-1 21
6 In general, experimental conditions are such that mass transport of reactants and products is not rate-limiting and the observed rate expressions refer to the true chemical processes in steps (ii)-(iv), The diffusion limitation is likely to be important in liquid phase hydrogenation reactions, particularly when hydrogen has a limited solubility in the liquid phase, and in gas phase hydrogenation where the catalyst is porous and the reaction occurs within the catalyst pores. As noted above, a necessary prerequisite to the derivation of an adequate kinetic rate expression is a knowledge of the mechanism of the reaction. In catalytic hydrogenation, two types of mechanism have been proposed. According t o the original Langmuir postulate [18] adopted by Hinshelwood [19] and others, both of the reactants are adsorbed and reaction occurs between adjacently adsorbed species. The postulated alternative, due originally to Rideal [ 20,211 and adopted by Eley [ 221, Twigg [23]and others, supposes that it is only necessary for one of the reactants to be adsorbed on the surface, and that reaction occurs on collision of the second molecule, coming from the gas phase, with the surface. In this mechanism, it is usually considered, following the original proposals of Rideal, that the unsaturated substrate molecule is chemisorbed as an immobile species and that the hydrogen enters the reaction via a “physically” adsorbed molecular state involving single sites left vacant following the adsorption of the substrate. The alternative situation of chemisorbed hydrogen atoms reacting with “physically” adsorbed substrate molecules has also been considered in certain cases [ 24,251.
2.2.1 Rate expressions for bimolecular surface reactions Throughout this section, it will be assumed that the mass transfer of reactants and products and the adsorption of reactants or desorption of products are not rate-limiting. For a reaction proceeding via a Langmuir-Hinshelwood mechanism A(ads.) + B(ads.) + Products the rate of reaction will be proportional to the probability that the species A and B are adsorbed on neighbouring sites and this is proportional to the fractions covered by A and B. Thus Rate = keA8B
(4)
and from eqn. (3)
If the pressure of B is kept constant and P A varied, the rate will change
7
V
3
Reactant pressure
Fig. 1. Variation of rate with reactant pressure for a bimolecular surface reaction proceeding by a Rideal-Eley or a Langmuir-Hinshelwood mechanism,
in accordance with Fig. 1. A similar variation will occur with PB at constant P A . Such a situation has been found t o occur in the hydrogenation of cyclopropane over pumice-supported Group VIII metals [ 261. A number of special cases of eqn. (5) can be considered. If the surface is only sparsely covered with the reactants, then bAPA + b B P B << 1and eqn. (5) reduces to
(6) Rate = kbAPA bBPB The reaction becomes first order in each reactant and second order overall. The hydrogenation of ethylene over copper [ 271 under certain conditions obeys such a rate law. If reactant A is weakly adsorbed, then b AP A << 1 + bBPB and the relevant rate equation is
and if reactant B is sufficiently strongly adsorbed that b B P B >> 1,then
~ ~ A P A Rate = ~BPB Examples of the applicability of eqns. (7) and (8) are, respectively, the hydrogenation of carbon dioxide over platinum [ 281 and the hydrogenation of ethylene over copper [29] and platinum [30]under certain conditions. There are, of course, an infinite variety of intermediate cases lying between the extreme conditions considered above. Rate expressions for a Rideal-Eley mechanism may be readily formulated. Assuming A is not adsorbed at all
References PP. 114-121
8 This equation differs from eqn. (5) in that the rate can only attain a limiting value and cannot pass through a maximum as the pressure of B is increased (Fig. 1). Whilst, in principle, kinetic measurements should allow a differentiation between the two possible mechanisms, it must be noted that in catalytic hydrogenation reactions relatively few examples are sufficiently clear cut t o allow this differentiation to be made. Thus, for example, it is quite commonly found that the experimentally observed orders of reaction are zero in the unsaturated substrate A and unity in hydrogen. Such results are readily interpreted by the “adjacent-site” mechanism by assuming A to be much more strongly adsorbed than hydrogen or by the RidealEley type of mechanism. Clearly, kinetic measurements alone are insufficient for the establishment of mechanism. Before leaving the discussion of kinetics, two points concerning the experimental determination of reaction orders should be noted. First, the kinetics of surface reactions, in contrast to those of homogeneous systems, are temperature-dependent. This must be the case since the relative surface coverages of the reactants A and B are
where 6 A 0 , is the difference between the free energies of adsorption of the reactants. Secondly, it is commonly found that the orders of reaction with respect to each reactant, as determined by the frequently used initial rate method, are not in agreement with those derived by following the change in rate as a function of time [31,32]. The reasons for such discrepancies are not altogether clear and it is likely that several factors may be responsible. One possible explanation is that reversible poisoning of the surface occurs during the course of the hydrogenation leading to a greater decrease in rate than expected from the observed pressure dependencies of the rate. 2.3 SELECTIVITY A N D THE CONCEPT OF ALTERNATIVE REACTION PATHS
One of the characteristics of many catalytic hydrogenation reactions is the ability of the catalyst t o promote the formation of more than one reaction product. Thus, for example, in the hydrogenation of acetylene, ethylene may be formed as an intermediate in the production of ethane and may be the major product in the initial stages of the reaction
Similarly, the hydrogenation of an unsaturated aldehyde, e.g. cinnamaldehyde t o 3-phenylpropanol, may proceed by two routes as shown in Fig. 2.
9 CH=CH * CH,OH I
Fig. 2. Alternative reaction pathways for the hydrogenation of cinnamaldehyde.
Other examples of alternative reaction pathways are shown in Table 2. It is clear that an understanding of the factors which influence the preference of a catalyst for one product are paramount t o our understanding of the mechanisms of catalytic hydrogenation and t o the industrial applications of hydrogenation catalysis. Thus, for example, the removal of trace amounts of acetylene from ethylene feedstocks depends upon the ability of the catalyst t o hydrogenate the acetylene t o ethylene without hydrogenating the latter to ethane [ 331. Similarly, the isomerisation activity of a catalyst may have an important bearing upon its use for the hydrogenation of monoolefinic compounds such as glyceride oils. It has also been suggested that, by studying these more complex systems, involving more than one reaction pathway, the elucidation of reaction mechanisms is less dependent upon such factors as variable catalyst activity, since both reactions are equally affected [ 341. The concept of alternative reaction pathways leads t o the concepts of selectivity and stereospecificity in hydrogenation catalysis. The general features of selectivity in catalytic systems have been considered from a kinetic standpoint by Wheeler [ 351 and by Waterman et al. [36,37]. According t o these workers, three types of selectivity can be distinguished. Type I selectivity arises when two reactants A and C are initially present
References pp. 11 4-1 21
TABLE 2 Alternative reaction paths in catalytic hydrogenation [ 34 ] Substrate
Other
Intermediate product
Final product
Designation of alternative path
Acetylene Buta-l:3-diene Ethylene But-1-ene
H2 H2 DZ HZ
Ethylene n-Butenes Deuteroethylenes But-Benes
Ethane n-Butane Deuteroethanes n-Butane
Selective hydrogenation Selective hydrogenation Olefin exchange Double bond migration
11 Assuming that A and C react according t o first-order kinetics and that the rate constants k l and k 2 refer t o unit surface area, the selectivity factor, u, is given by the ratio k , / k , and the relative fractions, a , of A and C reacted at any given time are given by the equation a* = 1- (1- a , )
(10)
Type I1 selectivity arises with the simultaneous production of two products from a single reactant
Defining as before, we have off, (11) Type I11 selectivity concerns the consecutive reaction of a product B to yield a further product C (Yg
=
The relative rates of change of the concentrations of A and B are given by 1 __ [BI d[Bl - 1 _ _ d[AI (3 [A1 from which it follows that
where a A is the fraction of A reacted and a B is the fraction of B formed. The concentration profiles for the fractions of A converted into B ( a ~ ) and into C ( a C )are shown in Fig. 3. I t is clear that all three types of selectivity are relevant t o catalytic hydrogenation reactions and from a consideration of the reaction scheme for alkyne hydrogenation (Fig. 4), it can be deduced that all three factors may be operative simultaneously. Clearly, the selectivity for the formation of the alkene relative to alkane will depend upon a number of factors. If both the alkene and the alkane are formed during one residence of the parent molecule on the surface, the selectivity will depend upon the relative values of k l and k 2 (Type I1 selectivity) and upon the ratio k l / k 4 (Type I1 selectivity). Since both of these depend upon the specific properties of the catalyst, they have been termed the mechanistic selectivity factor [38]. Once the alkene is produced, the system contains another potential adsorbate and Type I selectivity must be taken into account. It References p p . 1 I4--121
12
I
,
n "
1 CC
Fig. 3. Variations of t h e fractional yeilds of B and C with time for various type I11 selectivity factors. The arrows denote t h e direction of increasing time.
has been suggested that this selectivity corresponds t o a thermodynamic factor [38,39] on the basis that the surface coverages of alkyne (0,) and alkene (0, ) are given by
8,
=
bXPX (1+ bxPx + byP,)
and
Thus
Fig. 4 . Reaction scheme for t h e hydrogenation of a diunsaturated hydrocarbon.
13 where 6AG is the difference in free energies of adsorption of X and Y. A difference of only a few kilojoules per mole in the free energies of adsorption would result, therefore, in a very high surface coverage of the more strongly adsorbed species. This approach makes the tacit assumption that the same surface sites are responsible for both alkene and alkyne adsorption. Recent studies with supported metal hydrogenation catalysts using 14C-tracers [40,41] show that this assumption may not be completely justified. These studies show that ethylene and acetylene are adsorbed on different sites and that the adsorptions occur non-competitively. Type I selectivity, as observed in hydrogenation reactions, may thus find its origins in the inherent properties of the catalyst, rather than in the thermodynamics of the competitive adsorption of the two reactants. 2.4 APPLICATION OF ABSOLUTE RATE THEORY TO BIMOLECULAR SUR-
FACE REACTIONS
The formulation of rate expressions for surface reactions has been described in detail by Laidler [ 421.In this section, we confine our attention to its application t o the derivation of rate expressions relevant t o hydrogenation reactions. Again, it is necessary t o specify the mechanism in order t o derive relevant rate expressions. For a bimolecular reaction proceeding by an adjacent site (LangmuirHinshelwood) mechanism*
A
+ B + **
T~
AB* + Products
**
the rate of reaction, v, is given by
where f * , f,, , FA and FB are the partition functions for the species denoted by the subscript. If n, is the total number of surface sites and s is the number of sites adjacent t o any chosen site, it follows that the concentration of pairs of vacant sites [**I is equal to 1 s[ *]*/nS. Further, since [
*
I
-- 1+
n, bAPA + bBPB
where b A and bB are the adsorption coefficients for A and B respectively, it follows that
Equation (15) is of the same general form as that derived from the simpler Throughout this chapter References p p . 1 1 4-1 21
* will be used to represent a surface adsorption site.
14 treatment, eqn. (5), and the same limiting conditions, depending upon the strengths of adsorption of A and B, can be applied. If both reactants are weakly adsorbed, bA bB << 1 and eqn. (15) becomes
-
f,
kT u = isn,[A] [B] h
exp
FAFBf,,
(- RT 3)
i.e. the reaction will be overall second order. If reactant A is strongly adsorbed and B weakly adsorbed such that bA >> 1>> bB , eqn. (15) simplifies to
The adsorption coefficient can now be written in terms of partition functions
where fA , F A and f, are the partition functions for adsorbed A, gas phase A and surface sites, respectively, and E is the energy difference between adsorbed A and gas phase A. Thus eqn. (17) becomes
Assuming that the adsorbed state of A and the transition state AB , are immobile on the surface, the partition functions f t , f , , fA and ?:* can be considered to be unity and thus
Rate expressions may similarly be derived for a reaction proceeding via a Rideal -Eley mechanism A(ads.) + B(gas) * AB*
**
+
Products
kT f* u = [A][B] - FBfa
Assuming that both A and B are adsorbed, although adsorbed B does not
15 take part in the reaction directly
and
A virtually identical rate expression t o that given in eqn. (20) is obtained for the alternative postulate of gaseous A reacting with adsorbed B. It may also be noted that the maximum limiting rate predicted by eqn. (20) for the Rideal-Eley mechanism is very close t o the maximum rates predicted using eqn. (15) for a Langmuir-Hinshelwood mechanism. Thus quantitative calculations do not easily differentiate between the two mechanisms. The application of Absolute Rate Theory t o the interpretation of catalytic hydrogenation reactions has received relatively little attention and, even when applied, has only achieved moderate success. This is, in part, due to the necessity t o formulate precise mechanisms in order to derive appropriate rate expressions [ 431 and, in part, due t o the necessity to make various assumptions with regard t o such factors as the number of surface sites per unit area of the catalyst, usually assumed t o be 10'' cm-2, the activity of the surface and the immobility or otherwise of the transition state. In spite of these difficulties, it has been shown that satisfactory agreement between observed and calculated rates can be obtained in the case of the nickel-catalysed hydrogenation of ethylene (Table 3), and between the observed and calculated apparent activation energies for the
TABLE 3 Absolute rates of ethylene hydrogenation Catalyst
Temp. ("C)
PHz (torr)
Surface ( cm2)
Rate X (molecules cm-' sec-') Obs.
Calc
2.11 2.6 142.0 150.0
16.3 14.3 320.0 250.0
_____~Ni Ni Ni Nisilica
117 120 156 156
R e f e r e n c e s p p . 114-121
12.5 14.0 29.9 760.0
1.88 0.41 2.51 0.04
Ref.
~
45 46 41 48
16 hydrogenation of benzene over supported platinum, palladium and nickel catalysts [ 4 4 ] . In this latter study, it was assumed that 5-10 exposed metal atoms constituted a site for benzene adsorption.
3. The hydrogenation of olefins 3.1 INTRODUCTION
It has already been noted (Sect. 2.2) that an analysis of the kinetics of a catalysed hydrogenation reaction is insufficient for a precise description of the reaction mechanism. Whereas from the kinetics, it may be possible to determine values for the adsorption coefficients and hence the relative surface coverages of the reactants, such measurements reveal little or nothing about the intervening processes between the adsorption of the reactants and the appearance of products. That the mechanism is not simply the straight addition of two hydrogen atoms across the olefinic double bond is abundantly clear when one examines (a) the products from the reaction of hydrogen with higher olefins (C, and above) and (b) the products of the reaction of olefins with deuterium. These show that the occurrence of one o r more of the following processes may accompany the hydrogenation of the olefin to the alkane. (a) Olefin exchange; the replacement of H by D in the parent olefin CnH2, -+ CnH,n-xD, (b) Hydrogen exchange; the formation of HD and Hz in the gas phase. (c) Cis-trans isomerisation R
\
IR'
/c=c
H
\
H
+
\
H
I
H
/c=c\
R'
(d) Double bond migration
-
-
RCH=CH2 * CHZ CH2 R'
*R
*
CH2 CH=CH CH2R'
As will become apparent during the ensuing discussion, an understanding of the above processes, in terms of the elementary steps occurring on the catalyst surface, is essential to the specification of precise mechanisms for olefin hydrogenation. 3.2 ADSORBED STATES OF OLEFINS
It is now generally accepted that the chemisorption of an olefin precedes its hydrogenation. Since any mechanism will depend upon the nature of the adsorbed state of the olefin, it is pertinent t o examine the
17 present state of.knowledge of the adsorption of olefins at catalyst surfaces, and t o consider the possible ways in which the adsorbed species may interact with hydrogen-containing species. Information regarding the adsorbed states of olefins has been obtained from three main sources: (i) indirectly from studies of surface processes such as self-hydrogenation, self-poisoning and olefin-deuterium exchange; (ii) from changes in the properties of the metal, for example, magnetic susceptibility, surface structure and work function, during and after adsorption; and (iii) by direct observation using infrared spectroscopy. Most studies have been concerned with ethylene; relatively few studies of higher hydrocarbons have been reported. It is well established that when ethylene is admitted t o a freshly prepared evaporated metal film, “self-hydrogenation” resulting in the rapid production of ethane is observed [50-521. A similar phenomenon is observed when ethylene is adsorbed on supported metal catalysts i 49,533 (see Fig. 5). These observations have been interpreted as indicating that ethylene is first chemisorbed dissociatively, viz. C2H4(gas) C2H4-,(ads.) + x H(ads.) +
The hydrogen atoms thus liberated t o the surface may then react with either gaseous ethylene [ 50,511, or associatively adsorbed ethylene [ 531 or with the surface C2H4-, complex [52]. Volumetric [52,54] and magnetic susceptibility measurements [ 551 suggest that the extent of dissociation is dependent upon the temperature and varies from metal to metal. From the changes in magnetic susceptibility of nickelsilica catalysts during ethylene adsorption at room temperature, Selwood [ 551 has concluded that ethylene exists both as an associatively and a dissociatively adsorbed species. On increasing the temperature, the dissociative adsorption becomes more important. Thus at 100”C, the susceptibility changes are consistent with the formation of six bonds t o the surface for each adsorbed ethylene molecule, suggesting the following process
H
C2H4(gas)+ 6
*
7 \,?-?,7 jH
+
+
+
* * * ** *
Further increase in temperature results in carbon-arbon bond fission and the formation of a surface carbide containing single carbon units. These results are supported by volumetric studies [ 52,541. With both nickel and palladium there is a slow self-hydrogenation at -78°C. This becomes increasingly important as the temperature is increased; a t 0°C the overall hydrogen/carbon ratio in the surface species is 1.5, falling to 1.0 at room temperature. Field emission microscopic (FEM) studies of ethylene adsorption on iridium [56] and tungsten [57] are also satisfactorily References p p . I 1 4-1 21
18
/
0
2
4
6
8
Total molecules in gas phase ( x ~ O - " )
--.
m
5h X
v
4 % W -c
a
3: m C .-
2 m
-Q
3Q
1 5
z
0
2.5 5.0 7.5 Total molecules in gas phase (xlO-'*)
10
0
Fig. 5. Adsorption isotherms and composition of the gas phase for the adsorption of ethylene on (a) rhodiumsilica and (b) palladiumsilica at 2OoC. 0,Total molecules adsorbed;@,ethylene; 0 , ethane.
interpreted in terms of associative adsorption at low temperature, giving way to dissociation and the ultimate formation of surface carbide as the temperature is increased. FEM investigations also indicate that ethylene dimerisation occurs at step edges and on the (111) terraces of nickel [58]. Ethylene dimerisation has also been observed with cobalt catalysts [ 591 and with alumina and silica-supported platinum catalysts [ 601. Low energy electron diffraction (LEED) studies of ethylene adsorption on the (111)face of platinum [61]suggest that adsorbed ethylene occupies four
19
I
I
I
200 400 600 Gas phase count r a t e (mi+)
0
10
I
I
20 30 Time (min)
I
1
40
Fig. 6. Adsorption isotherms for I4C-ethylene and the effect of hydrogen on the adsorbed species for (a) nickelalumina and (b) palladium-alumina at 20°C [ 631.
sites, supporting the postulate of species such as structure B (see Fig. 8). The coexistence of at least two modes of ethylene adsorption has been clearly demonstrated in studies of 14C-ethyleneadsorption on nickel films [62] and various alumina- and silica-supported metals [53,63-65] at ambient temperature and above. When 14C-ethylene is adsorbed on to alumina-supported palladium, platinum, ruthenium, rhodium, nickel and iridium catalysts [63],it is observed that only a fraction of the initially adsorbed ethylene can be removed by molecular exchange with nonradioactive ethylene, by evacuation or during the subsequent hydrogenation of ethylene-hydrogen mixtures (Fig. 6). While the adsorptive capacity of the catalysts decreases in the order Ni > Rh > Ru > Ir > Pt > Pd, the percentage of the initially adsorbed ethylene retained by the surface which was the same for each of the processes, decreased in the order Pd > Ru > Ni > Rh > Ir > Pt 24.0 22.5 16.0 6.5 42.0 63.5 Similar results have been obtained using silica-supported metals [ 53,641, although the shape of the isotherm and the extent of retention appear t o be dependent upon the physical nature of the catalyst (Fig. 7). With alumina-supported palladium, platinum and rhodium and silicasupported platinum [65,66] in the temperature range 20-200"C, no molecular exchange between adsorbed 14C-ethyleneand gaseous ethylene is observed, whilst with hydrogen, small quantities of methane are formed at 100°C and above with platinum and rhodium and at 200°C with palladReferences PP. I 14-1 21
20 I
I
I
I
I
I
C ._
E
v
X
2 0,
+-'
0 I
w
C
Z J
8
c
1
0
'c
L
3 v)
-0 0
I
I
I
1
2
3
Gas phase count -rate f min-' x
)
Fig. 7 . Adsorption isotherms for 14C-ethylene and the effect o f hydrogen on the adsorbed species for palladiumsilica at 2OoC 1531.
ium. These observations have led to the postulate [67] that the retained species in ethylene adsorption arises from the formation of multiply bonded hydrogen-deficient surface complexes of the type shown in Fig. 8.
/H
H \
H
H\
*
c-c I T ? \ ** * * ******
I /c=c \
*
/H /I\
(B)
(A)
/I\
(C)
cII
C I\
H H \I C
* *
*I \ *
(D)
(E)
Fig. 8 . Hydrogen-deficient surface adsorbed states of ethylene.
14C-Propene adsorption on platinumalumina and platinumsilica
[66] differs from ethylene adsorption insofar as a fraction of the initially retained l 4 C-propene is relatively easily exchanged or removed by hydrogen treatment. This suggests less extensive dissociation of the adsorbed propene and a n-ally1 species (structure F) has been proposed in this case, viz .
* (F)
These studies also showed that, on a surface effectively saturated with 14C-propene, ethylene adsorption could still occur, although the amount
21
of ethylene adsorbed was less than on a “clean” catalyst. Such observations suggest that at least part of the adsorbed ethylene occupies sites not available for propene adsorption. The existence of several adsorbed states of an olefin on metal surfaces is shown by infrared spectroscopic studies [68]. This technique has the advantage that it yields direct information regarding the chemical identity of the various adsorbed species. although there are limitations t o its use. One of the main limitations is that the presence of surface intermediates may not be revealed if the appropriate band intensities are too weak [69]. In this context, it has been suggested [70] that the C-H bands associated with carbon atoms which are multiply bonded t o the surface are t o o weak to be observed. Pearce and Sheppard [71] have also proposed the operation of an optical selection rule, similar t o that found with bulk metals [ 721, in determining the bands observed with adsorbed species on supported metal catalysts. In spite of these limitations, however, the infrared approach has contributed significantly t o the understanding of the nature and reactivity of adsorbed hydrocarbons. Results obtained using the infrared technique indicate that the chemical identity of the adsorbed olefin is critically dependent upon the availability or otherwise of surface hydrogen. The types of adsorbed species also depend upon the structure of the olefin itself and, at any given temperature, upon the metal. Adsorption of ethylene on “bare” palladiumsilica produces only weak bands [ 721 which indicate the presence of species with double bond character (structure A, Fig. 8), together with surface methyl and methylene groups. Admission of hydrogen t o an ethylene-precovered surface results in the disappearance of the C=C bands and the appearance of bands ascribable t o adsorbed ethyl groups. On “bare” platinurnsilica at room temperature [ 73-75], adsorption of ethylene gives a spectrum consisting of bands ascribable t o both dissociatively and associatively adsorbed surface complexes, the latter being the predominant species. Admission of hydrogen causes an intensification of the spectrum, which has been interpreted as indicating the formation of surface carbidic residues during the initial adsorption. On increasing the temperature to 95”C, bands attributable to surface n-butyl groups are observed together with n-butane in the gas phase. It has been suggested that the C4-species arise from random polymerisation of dissociatively adsorbed ethylene residues, although thermal desorption data [ 761, and the observation [ 771 that in the reaction of 1-13C-ethylene with hydrogen no carbon-13 redistribution occurs, may be taken as a clear indication that on platinum the retained species is a discrete C2-unit. The spectra observed for ethylene adsorbed on nickelsilica [75,78] show similar features t o thpse found with platinurnsilica. The major difference is the temperature range over which the various bands are observed; with nickelsilica, the surface n-butyl groups are present at room Referencesp p . 1 14-1 21
22
temperature. With nickel, the retained species appear to be predominantly C4-units [ 75,791 rather than C2 -units as found with platinum. In general, the degree of dissociation of adsorbed species is greater with platinum than with nickel, although this metal gives a greater retention. With higher hydrocarbons, the spectra depend upon the structure of the olefin [ 701. With platinumsilica catalysts, linear chain olefins tend to form dehydrogenated surface residues more readily than branched chain olefins, which give predominantly saturated adsorbed alkyl species. Adsorption of hex-l-ene, a mixture of cis- and trans-hex-2-ene, and cis-hex-3-ene on nickelsilica results in identical infrared spectra [ 831. Addition of hydrogen results in an intensification of the spectrum suggesting that the initial spectrum results from dissociatively adsorbed species, a conclusion substantiated by the observation that the gas in equilibrium with the surface during the initial adsorption contains isomerised hexenes. Evacuation of the hydrogen causes a decrease in intensity and the reappearance of the initial spectrum. The effect of preadsorbed hydrogen upon the spectra of adsorbed olefins has been extensively investigated [ 75,78,79]. In general, the spectra are more intense when hydrogen-precovered surfaces are used and they show bands resulting from predominantly saturated species formed by the hydrogenation of associatively adsorbed olefin. From the foregoing discussion it is clear that the interaction of an olefin with a metal surface is a complex process resulting in the formation of both dissociatively adsorbed hydrogen-deficient species and associatively adsorbed surface complexes. The question remains as to the relative importance of these various species in the context of catalytic hydrogenation 1951. The early work of Horiuti and Polanyi [80,81] regarding the metalcatalysed olefin-deuterium exchange reaction led to their postulating that the olefin was chemisorbed as a di-o-bonded species (structure G , 3 g . 9) formed by the rupture of the olefinic n-bond and the formation of two carbon-metal o-bonds. Subsequently it has been generally assumed that the olefin species active in catalytic hydrogenation is associatively bonded. However, there is still considerable debate as to the type of surfaceadsorbate bonding involved in the associatively adsorbed species, and of the role of dissociatively adsorbed species in catalytic hydrogenation [ 82, 951. The existence of a di-o-bonded species is consistent with the observed infrared spectra as discussed above. The formation of such species leads t o the idea of optimum metal-metal distances for olefin adsorption and the geometrical properties of such species have been discussed in detail [ 841. Following the suggestions of Rooney et al. [ 85-87] in the early 1960’s, a so-called n-complex (structure H, Fig. 9) has been assumed for hydrogenation. The major differences between structures G and H lie in the geometries of the two species. Thus, whereas the di-o-bonded species requires two metal atoms at a suitable distance apart and has sp3-hybri-
23 H \ lH H-C-C-H I \
*
*
H \
I
H
F\
H * H (H)
(G)
Fig. 9 . Di-o-bonded and n-bonded associatively adsorbed ethylene.
dised carbon atoms, the n-complex requires only one surface site, although others may be obscured, and retains some double bond character depending upon the extent of the metal-olefin interaction. Surface potential measurements of ethylene chemisorbed on nickel [88] and palladium films [89] are consistent with the n-complex structure H. Recently, Sheppard and co-workers [ 901, using an extremely sensitive infrared interferometry technique, have also obtained evidence for the existence of both a di-o-bonded and a n-bonded species when ethylene is chemisorbed on hydrogen-precovered silica-supported palladium and platinum catalysts. These workers also claim that both species are easily hydrogenated, the n-complex being the more reactive. Although the precise nature of the reactive adsorbed state of an olefin remains a subject for debate and conjecture, in the ensuing discussion it will be assumed that-it is adequately represented by structure H. 3 . 3 POSSIBLE REACTION MECHANISMS
We have already noted (p. 16) that the interaction of an olefin with hydrogen or deuterium may lead t o the occurrence of any of a number of processes. There is much evidence t o suggest that each of these processes may be accounted for by considering a number of elementary steps in which a hydrogen atom, from a meantime unspecified source, is added to or removed from an adsorbed hydrocarbon species. It has been observed [ 23,91,92] that when an unsaturated hydrocarbon is reacted with (a) equilibrated and (b) non-equilibrated hydrogen-deuterium mixtures, the deuteroalkane distributions are identical, Such observations indicate that the direct addition of a hydrogen molecule across the olefinic bond does not occur, and provides strong evidence for the formation of a “half-hydrogenated state”, that is, an adsorbed alkyl radical, first suggested by Horiuti and Polanyi [ 811, as a relatively stable reaction intermediate. The process of hydrogenation may thus be represented as
RCH=CHR‘ I
*
2 RCH-CH~R’ kl
I
*
Unless otherwise stated, the term hydrogen is used in a generic sense to indicate either H o r D. References p p . 114-1 21
24
RCH-CH~R’ I
*
2RCH~-CH~R’
(2)
k2
Under the conditions generally used for olefin hydrogenation, the alkane, once formed, is unreactive [93] and its readsorption on the surface can be neglected. The processes of olefin exchange, double bond migration and cis-trans isomerisation, observed t o occur concomitant with hydrogenation, may be accounted for by considering that the formation of the half-hydrogenated state is reversible. For olefin exchange we can write
RCH=CHR‘ I
*
2 RCHD-CHR’ 3 RCD=CHR’ I
k3
*
I
k3
(3)
*
Similar mechanisms may be written for double bond migration
RCH2CH=CHR’
I
*
kl
RCH,CH-CH2R’ I
*
r5f RCH=CHCH2R’ I k4
(4)
*
and the cis-trans isomerisation
Clearly, if the reversal of alkyl formation is rapid compared with the further hydrogenation of the adsorbed alkyl radical t o alkane, this will result in a build-up on the surface of deuteroolefin and protium atoms in the case of mechanism (3), or isomerised olefin in the case of mechanisms (4) and (5). However, it should be noted that these species must undergo desorption before appearing in the gas phase. Consequently, the extent of olefin exchange or isomerisation which is observed will depend upon both the ratio kJk, or h 4 / h 2 and the ratio k l / k d , kd being the rate coefficient for desorption of the olefin. Since, according to the mechanisms outlined above, double bond migration and cis-trans isomerisation are mutatis mutandis identical with ethylene4euterium exchange, it might be expected that correlations exist between the abilities of various metals to catalyse olefin isomerisation and olefin4euterium exchange. Table 4 shows values for 0, and pi defined as
lo,
=
8. =
Cdeutero-olefin yields Cdeutero-alkane yields Cisomerised butene yield butane yield
25 TABLE 4 Abilities of various metals t o catalyse olefin isomerisation and exchange ___
~-
-~
Metal -~
Ni
__
1.8
PA1-C4Hs)
2.3
Ru
Rh .
~~~
Pe(C2H4) Be(C&)
__
~~~~~
Pd 2.2 1.1 6.9
~-
--____-_______
__
0s
Pt
Ir
0.4
0.13 0.06 0.03
0.03 0.01
-
1.6 1.6 5.7
1.4
0.2
1.1
0.13
for ethylene, propene and but-1-ene respectively over various metals 1941. From these results it can be seen that, in general, good isomerisation catalysts are good exchange catalysts and vice versa, thus supporting the proposed mechanisms. For olefins possessing one o r more a-CH2 groups, it is possible that olefin exchange and isomerisation could occur through the formation of an adsorbed n-allylic intermediate [ 871. The appropriate mechanism for olefin exchange and double bond migration is RCHX-€H=CHR’
RCH,CH=CHR‘ I
2 RCH-CH-CHR’ I
*
*
+P
*I
ARCH=CH-CHXR’
(6)
I
* where X = H or D. Repetition of these steps, with or without intermediate desorption of the olefin, could lead t o complete exchange of all the hydrogen atoms except those attached t o the 2-carbon atom. This contrasts with the alkyl reversal mechanism (3), where all the hydrogen atoms are exchangeable and consequently, it should be possible t o test the occurrence of (6) as the sole mechanism of exchange. Since, in the n-ally1 complex, free rotation is not possible, cis-trans isomerisation can only occur through n-ally1 if it is preceded by double bond migration, as shown in Fig. 10. The source of hydrogen atoms for the above mechanisms is not readily established. There are several possibilities. (a) Dissociative adsorption of molecular hydrogen, which may occur either competitively o r non-competitively with the olefin. H2(gas)+ 2 [ * ] + H + H I I
(7)
* *
(b) Interaction of molecular hydrogen with chemisorbed olefin. RCH=CHR’ + H2 = RCH2-€HR’ + H
I
* References pp. 1 1 4-1 21
I
*
I
*
26
But-l-ene (g)
11
H,C=CH. CH:! . CH3 I
+HJ--H
H,C
*
H H\ / ?IC\ H3C * CH,
H
It
11
cis-but-2-ene (g)
trans-but-2-ene (g)
Fig. 10. Mechanism for the isomerisation of t he n-butenes involving n-allylic intermediates.
(c) Hydrogen transfer between associatively adsorbed hydrocarbon species. RCHyCHR’ + RCH--CHZR’
I
*
I
*
=
RCH,-€HR’ + RCH=CHR’ I I
*
(9)
*
(d) Hydrogen transfer between associatively adsorbed olefin and a dissociatively adsorbed hydrocarbon residue, which may be represented as C,H, 1953.
C,H,(ads.) + RCH=CHR’
I
=
C,H,-,(ads.)
*
+ RCH,-CHR’ I
10)
*
followed by C,H,-I(ads.)
+H I
=
C,H,(ads.) +
[*I
*
C,H,-l(ads.)
+ H,
= C,H,(ads.) + H
I
* To decide between these various possibilities requires detailed analysis of the kinetic data and of the deuteroolefin and deuteroalkane distributions (Sect. 3.5). It should be noted that the situation may be further complicated since it is possible that more than one of the above processes
27 may occur simultaneously, although the extent t o which each contributes to the overall reaction may not be the same. 3.4 TREATMENT O F EXPERIMENTAL RESULTS
In this section, the techniques which have been applied t o the experimental results to obtain detailed information regarding the basic reaction mechanisms are reviewed. As noted in Sect. 2.2 (p. 5) the determination of the kinetic parameters cannot of itself lead t o the specification of a complete mechanism, although any postulated mechanism must be compatible with the kinetics. Furthermore, the interpretation of kinetic parameters in terms of absolute rate theory [ 421 has not, in general, been successful. Attempts t o interpret the mechanism of ethylene hydrogenation over nickel [96-991 and over platinum catalysts [100,101] in terms of a statistical mechanical approach have not met with any substantial success, partly due t o the limitations of the model which must be assumed in order to perform the calculations and partly due t o the complexity of the calculations themselves. Results obtained from the reaction of ethylene with deuterium have been used t o obtain information regarding the probabilities of the various changes which the adsorbed hydrocarbon species may undergo. The procedure, due originally t o Kemball [ 1021 and subsequently used by Bond et al. [ 103-1051 and Wells and co-workers [ 1061, is based upon a steady state analysis of the following general mechanism. +H
C2X4(g)+ C2X4(ads.)+ C2X,(ads.) -X
+X +
C,X6(gas)
where X = H or D, their origins and fates being unspecified. C,X, (ads.) and C2X5(ads.) are specified t o react with the probabilities C,X4(ads.) +
C2X,(ads.) (gas)
c2x4
+
C,X,(ads.) + H C,X,(ads.) + D -+ C,X,(ads.) C,X,(ads.) + X
+
+
+
C2X4H(ads.) C,X4D(ads.) C2X4(ads.)+ X CzX6
P ( 1 + P)-l l(1 + p ) - ' l(1+ q)-I q(1 + 4)-l
r ( 1 + r)-l 1(1+ r)-I
l/(l+ s)-' C,X,(ads.) + H + CzXSH s/(l + s)-l C2X,(ads.) + D + C,X,D Assuming these probabilities are independent of isotopic content, the six isotopic adsorbed ethylenes, denoted A l + A6 and the twelve isotopic ethyl radicals, denoted B1 + BIZ,are related by a series of simultaneous equations of the form
References p p . I J 4-J2J
28 where Q, represents any of the A’s or B’s, pmnis the ratio (the probability of any entity of type in becoming one of type n ) t o (the probability of m leaving the surface) Q1 and film have similar meanings and a, (0) represents the fraction of entity m initially present. a, (0) is zero for all A’s and B’s except CzH4. The eighteen simultaneous equations so obtained are solved for various values of p , q, r and s t o obtain a satisfactory fit between the calculated and experimentally observed deuteroethylene and deuteroethane distributions. Examples of the results obtained are given in Sect. 3.5. More recently, the procedure has been modified t o allow for isotope effects in the rupture of C-H and C-D bonds [107], although further improvement in the agreement between calculated and observed distributions appears to be small. The values of p and r obtained by this procedure give valuable information regarding the relative rates of the surface processes involving hydrocarbon species, while q and s reveal information about the relative surface coverages of H and D relevant to the hydrogenation of adsorbed ethylene and ethyl respectively. In principle, this approach is applicable t o higher olefins although the complexity of analysis increases rapidly with increasing chain length. Thus, for example, in the propene-deuterium reaction, a total of eighty simultaneous equations involving twenty-four propenes and fifty-six propyl radicals are required. Further, the situation may be more complex with higher olefins since there is more than one mechanism possible for olefin exchange (Sect. 3 . 3 ) . To date, no attempts have been made at applying the treatment t o higher olefins. A semi-quantitative treatment based upon the observation that, in the reactions of olefins with deuterium, the yields of deuterated paraffins often decrease logarithmically with increasing deuterium content, that is yield - -~
of CnH2n+2-xDx _____ =o yield of CnH2n+3-,Dx--l
has been proposed by Bond [108]. According t o this treatment, any adsorbed olefin which, having undergone exchange, contains one deuterium atom must necessarily be converted into an alkyld, species. If the fate of the adsorbed alkyl radical is to disproportionate, that is 2 RCH-CHDR‘ I
+
*
RCH=CHR’ + RCHD-CHDR’ I
*
+
RCH=CDR‘ + RCH,--CHDR’ I
*
there is an approximately equal chance of each alkyl radical desorbing as a paraffin-d, or remaining on the surface, undergoing further exchange and ultimately desorbing as a paraffin-d,,,. Accordingly, values of u are expected to be approximately 0.5 and departures from this value may
29 be attributed to (i) unequal occurrence of the exchange reactions, (ii) addition of deuterium by other mechanisms and (iii) zero point energy effects. Thus values of u greater than 0.5 are expected if processes such as
RCH-CHDR' + I
I*]
+
*
H + RCH=CDR' I I
*
*
and
RCH-CHDR' + D + RCH=CDR' + HD I I I
*
*
*
occur, whilst u < 0.5 will be observed if either olefin desorption or the processes
RCH-CHDR' + D + RCHD-CHDR' I I
*
+ 2[*]
*
and RCH-CHDR' I
*
+ Dz
*
+
RCHD-CHDR' + D I
*
are predominant. 3 . 5 . HYDROGENATION OF ETHYLENE
The hydrogenation of ethylene has been extensively studied over a wide variety of metal catalysts. In this section we review some of the results obtained for the kinetics and activation energies and from the use of deuterium as a tracer. Table 5 summarises the kinetics and activation energies observed for the hydrogenation of ethylene over a variety of metals. From these results it can be seen that, although there is some considerable variation in the precise values of the orders with respect to hydrogen and ethylene, in general the order with respect to hydrogen is commonly approximately unity and that for ethylene is zero or slightly negative. This indicates the strong adsorption of ethylene relative t o hydrogen. A second feature of the results is the narrow variation in the values of the activation energy, this is especially true for silica-supported metals, where E , = 35.2 kJ mol-'. Such an observation led Beeck [ 501 and Schuit and van Reijen [ 1091 to suggest that variations in specific activity from metal to metal must be attributed to differences in the temperature-independent term A in the Arrhenius equation. *4ccording to the latter authors, the general rate equation
References p p . 1 1 4-1 21
30 TABLE 5 Kinetics and activation energies for ethylene hydrogenation
Catalyst
X
Fe/SiOz Fe film Co/SiOz Ni film Ni wire Ni powder Ni/SiO2 Ni/A1203 Nilsilica-alumina Cu/SiOz Cu powder
0.91 0.87 0.55 1 1 1 0.67 1 1.09 0.69 1 1 1 0.95 1 0.85 0.66 1 0.9 1
-0.04 -0.6 -0.19 0 0
1.6 1.3 1.2 1.0 0.77 1.0
-0.4 -0.8 -0.5 -0.3 0.25 -0.2
Ru/A1203 Ru/SiO2 Rh/A1203 Rh/SiO2 Pd/SiOz Pd/A1203 Re/SiOz O~/A1203 Ir/SiO2 Ir/Al2O3 Pt foil Pt/A1203 Pt/A1203 Pt/Si02 Pt/SiOZ a
ya
L -0.08
L 0.21 0.06 0 1 -0.2 -0.59 0 4.74 -0.03 0 -0.6 0
Temp. ("C) 30 32
-6@ 11'9 156 99-165 -4 0 70 128 80 100 200 46 -7 6 79 -7 6 -3 0 -1 8 104 50 60 ?
0-18 0 -4 0 60
Ea (kJ mol-') 35.2 30.5 35.2 42.7 58.6 ? 35.2 48.6 50.2 35.2 45.2 29.3 36.4 35.2 50.2 35.2 35.2 47.7 33.5 35.6 35.2 57.8 41.8 41.4 62.8 35.2 87.9
(? temp.) (32-80') (? temp.) (20-150') (60-100") (? temp,) (30-80") (90-135") ( ? temp.) (150-200') (200-250') (32-80") (? temp.)
(73-100") ( ? temp.) ( ? temp.) (50-77") (25-150") (17-47") (? temp.)
(80-120') (0-150') (0-50") (0-40")
(?te mp.) (60-120')
Ref. -~ 109 120 109 51 110 111 109 112 139 109 113 104 109 105 109 109 105 106 104 105 103 114 108 103 109 137
L = Langmuir expression.
may be simplified by assuming x = 1and y = 0 and hence values of A may be calculated. The specific activity, k , , , is given by h,, = A N ; / N , , where N: is the number of molecuies in the reaction vessel at the reaction temperature and N , is the number of sites at the catalyst surface. Table 6 shows values of the specific activities for a series of silica-supported metals and metal films, expressed as log,, [(hsp)M/(ksp)Rh ] , rhodium being the most active metal. The values so obtained will critically depend upon the value of N , ; the radiochemical results discussed in Sect. 3.2 (p. 16) show that there is no simple relationship between the adsorptive capacity of a catalyst and the fraction of the adsorbed molecules active in the hydrogenation. Consequently, the only satisfactory values of N , would be those determined from the actual number of adsorbed molecules undergoing reaction; values for N , based upon total surface areas, as used in Table 6,
eb
TABLE 6 Specific activities of metals for ethylene hydrogenation
Films On SiO? a-See ref. 106.
Fe
co
Ni
Ru
Rh
Pd
Re a
Ir
Pt
Ref.
-3.0 -3.4
-2.1
-2.8 -1.5
-4: 3
0.0 0.0
-4.8 -4.9
-6.2
-2.0
-1.6 -1.5
109
50
32
0
E
1 0
20
30
40
50
60
70
80
5 0 6 0 70 Conversion
80
90
(b)
E I
2 2-
3 VI
-0. 0
1 0
2 0
30
40 O/O
90
100
Fig. 11. Distribution of (a) deuteroethylene yields and (b) deuteroethane yields as a function of conversion observed in the reaction of ethylene with deuterium over a nickel wire at 90°C [91].
may give misleading values of the specific activities. One of the earliest studies of the reaction of C2H4with D2,in which a full mass spectrometric analysis of the products was performed, used a nickel wire as catalyst [115,116]. Some typical results are shown in Fig. 11. These results showed that ethylene exchange was rapid and the deuteroethylenes are probably formed in a stepwise process in which only one deuterium atom is introduced during each residence of the ethylene molecule on the surface, that is there is a high probability of ethylene desorption from the surface. From Fig. l l ( a ) it can also be seen that the major initial products are ethaned, and ethanedl. This is consistent with a mechanism in which hydrogen transfer occurs by the reaction
HZC--CH?D + H*C=CH2 +- H,C=CHD + H*C--CH3 I I I I
*
*
*
*
33 followed by a disproportionation reaction of the type
*
*
*
As the reaction proceeds, the ethyIene becomes progressively more deuterated [Fig. l l ( b ) ] and this is reflected in a progressive growth of the more heavily deuterated ethanes in the latter stages of the reaction. These conclusions are in general agreement with those subsequently obtained for the reaction over nickel films [ 1021 (see Table 7). A systemmatic study of the reaction of ethylene with deuterium over alumina-supported Group VIII metals has been reported by Bond et al. [103-1051. Table 7 shows a selection of typical results together with theoretical distributions calculated as discussed in Sect. 3.4. Steady state analysis of the reaction scheme
shows that ec2x5 = kdPr ec2x4 h ( r + 1)
Thus the values of p and r may be used to give qualitative information regarding the relative surface coverages of adsorbed olefin and alkyl, although, because of a lack of knowledge of the ratio hd/k4, quantitative values can not be derived. The widely differing values of p and r for the different metals suggest that the state of affairs on the various surfaces are very different. Figures 1 2 and 1 3 show, respectively, the percentage
-
80 -
$60-
v
*Q
40-
Pt
-50
0
50 Tern peratu re ("C)
100
150
Fig. 12. Variation of the percentage chance of ethylene desorption @*)with temperature for various alumina-supported metals [ 1051. References p p . I 14-1 21
34 TABLE 7 Observed and calculated deuteroethane and deuteroethylene distributions over various metals Catalyst
Temp. ("C)
D2/olefin
Conversion
-dl
Ni-film
-100
4
3
+4
0
46.7 13.2 2.7 41.4 18.8 6.4
0.7 1.0
7.8 12.5 2.8 9.9
52.1 53.5 45.1 46.0 57.5 52.8
3.9 3.1 4.1 2.4 3.5 3.0
0.0 0.1 0.0 0.1 6.2 0.0
0.0 0.0 0.0 0.0 0.0 0.0
2.6 0.8 0.7 1.0 0.7 4.1
9.9 10.5 11.3 12.6 10.4 10.7
24.8 23.7 15.9, 13.9 30.7 30.8
1.6 2.8 1.1 11.0 1.5 4.5
0.0 0.2 0.0 0.1 0.0 0.5
0.0 0.0 0.0 0.0 0.0 0.0
3.1 1.3 0.0 0.8 0.0 0.1
14.2 16.5 10.6 13.9 12.3 13.9
65.8 66.7 26.2 26.3
7.7 8.4 4.2 6.8
0.6 0.7 0.8 1.2
0.0 0.0 0.2 0.1
8.3 7.5 2.3 9.4 7.9 19.7 4.5 24.1
3.1 3.3 1.1 14.9 11.9
0.4 1.5 0.2 4.3 6.2
0.4 0.4 0.1 1.4 1.7
0.0 0.0 0.1 0.4 0.2
0.0 3.0 0.0 1.7 2.1
32.0 30.2 32.8 21.3 21.6
40.8 41.3 20.6 23.2 45.4
7.9 8.1 5.3 4.0 8.6
2.2 1.0 1.9 0.4 1.1
0.8 0.1 0.7 0.0 0.0
21.6 22.3 17.2 18.7 22.0
17.6 19.7 27.5 32.1 16.8
9.2 1.2 11.9 6.2 36.5 8.6 32.0 20.3 44.0 11.4 41.4 16.9
0.2 1.7 3.9 6.1 4.7 3.8
0.0 0.2 2.1 0.7 1.0 0.4
19.3 13.8 4.9 3.9 8.2 4.2
28.2 26.7 11.2 11.2 13.0 13.2
2.7 0.8
0.0
53
0.54
7
Ru-Al203
53
1.65
7
Ru-Al203
80
1.0
7
Os-A1203
24
0.56
16
Os-A1203
24
4.0
16
Os-A1203
47
1.o
15
Rh-Al203
76
0.26
5
-18
1.0
5
Ir-A1203
-1 6
0.26
2.5
IrA1203 IrA1203
-1 6 124
4.74 1.00
5 5
Pd-Al203
-1 6
0.24
5
Pd-Al203
-1 6
4.94
5
Pd-AI203
67
1.0
5
Pt-Al203
54
1.0
10
Pt-AI203
150
1.0
10
Pt-Al203
150
0.3
5
Pt-pumice
25
03
--dZ
3
Ru-Al203
Rh- A12
Ethanes
Ethylenes
("/.I
1.17
50
7.4
19.2 21.2 ~
a M = mean deuterium number of ethane, i.e. x in C2HexDx.
chances of olefin desorption (p*) and of ethyl reversal (r*) as a function of temperature for the metals studied [105].Clearly, the chances of olefin desorption are least with platinum and iridium and highest with ruthenium and osmium, whilst the chances of alkyl reversal are lowest
Ma +2
4
3
4
4
4
5
4
P
q
r
S
H in D2
Ref.
(%)
6
9.5 4.0 2.0 11.1 5.7 2.2
0.8 0.6
0.2 0.1
1.55 1.90
29.1 29.6 35.7 35.5 26.5 25.9
2.5 2.3 3.1 2.4 2.1 3.3
0.0 0.1 0.0 0.1 0.1 0.0
0.0 0.0 0.0 0.0 0.0 0.0
0.0 0.0 0.0 0.0 0.0 0.0
1.70 1.76 1.79 1.777 1.64 1.63
46.8 47.6 62.8 63.7 44.4 40.6
6.1 7.1 6.0 6.1 7.0 7.5
2.5 0.6 0.8 0.1 1.8 0.9 0.5 0.5 3.0 0.6 1.1 0.1
1.6 0.0 0.9 0.2 0.6 0.0
1.63 1.93 2.10 2.06 1.99 1.91
9.1 10.6 26.5 26.3
0.9 1.7 8.1 8.4
0.1 0.2 4.4 2.0
0.0 0.0 1.8 0.3
0.0 0.0 0.2 0.0
1.12 1.51 1.83 1.70
37.5 34.8 40.0 28.9 28.8
13.8 19.1 14.4 11.0 18.7
9.1 6.5 8.2 9.3 7.2
3.3 1.2 2.7 5.8 1.5
0.4 0.1 0.4 1.4 0.1
2.12 2.00 2.07 2.35 2.18
7.3 6.1 17.1 17.0 5.4
1.2 1.2 5.2 3.4 0.6
0.5 0.2 3.0 0.5 0.1
0.1 0.0 1.4 0.1 0.0
0.0 0.0 0.1 0.0 0.0
0.79 0.73 1.36 1.10 0.66
25.6 11.9 23.9 13.1 14.6 12.2
8.7 3.7 4.3 8.6 2.2 5.8
4.9 0.6 2.9 3.3 0.5 1.2
4.9 0.1 1.7 0.7 0.3 0.3
2.3 0.1 0.2 0.1 0.1 0.0
1.55 1.59 1.90 1.95 1.36 1.70
23.2 10.5 6.4
3.0
0.5
1.68
3
2
0.2
9
2.5
4
0.2
9
1.8
4
0.2
9
3
4
0.8
9
1
4
0.8
9
0.5
9
0.8
9
1.5
4
0.5
2
9
a
12
102
1
‘4.5 1.8 4.5 8.5 0.5
104
5.2 15.6 -
105
0.4 3 99
1
4
4
0.5
9
9 0.05 0.9
49
0.5
19
2.5
0.5
9
0.1
3
0.5
4
1
103
9 0.2 0.05
105
0.01 0.3
49
0.5
19
0.5
49
0.5
99
1
9
0.5
29
2
8.4 16.7 1.2
108
with ruthenium and osmium and high with the other four metals. The low exchange activity of platinum and iridium may, therefore, be accredited t o the high stability of the adsorbed ethylene, whereas the high exchange activity of, for example, palladium is due t o a balance between a high References p p . 1 1 4-1 21
'oo 80 -
-
*c v 0
-
60 -
40
I
f
1
I
Fig. 13. Variation of the percentage chance of adsorbed ethyl reverting to adsorbed ethylene ( r * ) with temperature for various alumina-supported metals [ 1051.
probability of alkyl reversal and of olefin desorption. A more quantitative interpretation would require a knowledge of the actual surface eoverages of H and D atoms during the reaction. The parameters q and s give some indication of the origins and fates of the H and D atoms. There appears to be an inverse relationship between TABLE 8 Distribution of products from the reaction of propene with deuterium Conversion a
Propenes
Catalyst
Temp. ("C)
Rh-AI203 Rh-AI203 Rh-A12 O3
90 90 128
0.7 3.1 3.0
8.7 9.9 12.7
76.8 17.8 4.1 1.1 0.3 61.2 27.1 8.7 2.4 0.5 77.5 16.0 4.9 1.3 0.3
II-AI~O~ Ir-A1203 II--AI~O~
16 78 78
0.83 0.83 4.28
53.0 49.0 56.0
97.6 2.0 96.7 2.2 99.3 0.6
Pd-Al203 Pd-AI203 Pd-Al203
-20 20 20
3.0 3.0 0.75
10.5 17.1 15.8
86.0 10.9 2.5 0.6 82.3 12.0 4.1 1.3 78.7 16.0 4.2 1.0
0.1 0.3 0.2
0.0 0.0 0.0 0.0 0.0 0.0
Pt-AI203 Pt-AI203
38 75
0.83 0.83
84.0 83.0
54.3 24.4 12.0 5.8 83.4 10.6 3.6 1.4
2.5 0.3
0.9 0.2
0.2 0.1
Pt-pumice Pt-pumice
18 75
0.03 1.0
97.0 3.0
0.0
0.0
0.0
a
100 100
Based upon reactant which is in deficiency. Mean deuterium number of propanes, i.e. x in C3H,+.xDx.
0.4 0.0 0.7 0.3 0.1 0.0
0.0 0.0
0.1 0.0 0.0 0.0 0.1 0.0
0.0 0.0 0.0 0.1 0.0 0.0 0.0 0.0 0.0
37 the values of r* and q * , these being the percentage chances of ethyl reversal and of adsorbed ethylene acquiring a D atom, respectively. As r* tends towards loo%, q* tends t o a low value (33%)and conversely as r* approaches 50%, q* tends t o 100% [ 1051. This has been interpreted as indicating that the reversal process occurs via a redistribution reaction C2X4(a)+ C2Xs(a) = C2Xs(a)+ C2X4(a) Low values of q* found for palladium, platinum and iridium (Table 7), together with negligible amounts of hydrogen exchange, indicating a low stationary concentration of adsorbed H and D atoms, suggest that the above reaction may be of particular importance over these metals. 3.6 HYDROGENATION O F PROPENE
Compared with the extensive studies of the kinetics of ethylene hydrogenation, the kinetics of the hydrogenation of propene have received little attention. Over platinum-pumice catalysts at 18"C, the kinetic rate law was observed to be
P"G;
Rate = kPC'&
while the activation energy has a value of 26.3 kJ mol-'. [92]. Over palladium-alumina, the apparent activation energy is 48.9 kJ mol-' [119]. More recently, Mann and Lien [ 1171 have reported that for pumice-sup-
Mb
Propanes 4
0
--dl ~~
4
2
4
3
4
4
4
5
4
6
-dl
4
u
H in D2
Ref.
("/.I
8
~
0.0 0.0 0.0 0.0 0.0 0.0
1.06 1.51 1.91
0.22 0.36 0.33
12.1 4.6 2.8
15.0 28.5 26.1 15.4 8.5 3.9 1.9 13.4 23.1 22.7 17.5 11.2 7.1 3.4 3.2 12.1 17.9 16.1 13.4 11.0 10.4
0.7 0.1 1.5 0.2 5.0 3.83
1.97 2.34 3.83
0.52 0.65
?
?
12.2
61.9 26.6 9.1 51.7 28.0 12.9 63.9 25.5 7.9
36.6 31.1 24.9 5.5 23.5 27.0 31.8 11.8 11.8 18.4 45.0 17.7
1.2 4.4 6.0
0.7 1.5 1.1
0.0 0.0 0.0
2.1 5.3 2.3
0.4 1.8 0.4
0.0 0.3 0.0
0.0 0.0 0.0
0.0 0.52 0.0 0.0 0.0 0.0
0.52 0.18 0.50
0.38 0.41 0.32
16.4 30.9 23.8 14.4 19.3 26.7 22.2 13.8
7.5 9.5
4.1 5.0
2.0 2.8
0.7 0.1 1.0 0.2
1.90 1.99
0.54 0.60
45.2 28.4 15.8 5.6 18.7 26.1 22.0 15.0
2.4 8.1
1.4 5.2
0.8 2.8
0.3 0.1 1.2 0.3
1.01
2.04
0.43 0.63
Referencespp. 114-121
105
1.9 103 0.2 0.2 0.2 ? ?
105 103 92
38 ported Ni, Fe, Co, Pt, Pd, Rh, Ir, Ru andOs the kinetic rate law is Rate = kP;, P&H6 where the values of x are unity or slightly less and of y are zero or slightly negative; the respective values of the activation energies are 54.4, 41.9, 33.9, 67.0, 46.0, 54.4, 62.8, 27.2 and 31.0 kJ mol-’. Based upon the initial rate of hydrogenation per unit weight of catalyst, the activity sequence was Rh > Ir > Ru > Pt > Pd > Ni > Fe > Co > 0 s . However, no attempts were made t o allow for possible variations in surface area or of the fractions of the surface which was catalytically active and hence the significance of the above sequence is of doubtful value. The reaction of propene with deuterium has been studied over a variety of catalysts [ 92,103,105,1181. Some typical deuteropropane and deuteropropene distributions are shown in Table 8. Values of Pe as defined in Sect. 3.3 and u (Sect. 3.4) are also quoted. Comparison of these results with those presented for ethylene in Table 7 show that, for each catalyst, the two reactions show close similarities suggesting that the general features of the mechanisms are the same for both reactions. There are some important differences in detail between the two reactions, particularly with palladium and iridium. With palladium, olefin exchange occurs more readily with propene than with ethylene. This may be due t o easier desorption of propene or possibly due to an alternative mechanism such as H3C * CH=CH2 I
*
3 H2C-CH-CH, 3 H2DC I
*
CH=CH2 I
*
contributing t o the olefin exchange. With iridium, increasing deuterium pressure has a pronounced effect on the deuteropropane distributions, whereas with ethylene no such effect is apparent. This has been attributed to differing temperature dependencies of the alkyl disproportionation reaction for the two olefins [103], although the same effect could be achieved by considering that the relative surface coverage of deuterium is greater with propene than with ethylene. 3.7 REACTIONS OF THE n-BUTENES
Whilst the use of deuterium allows a deeper insight into the mechanism of catalytic reactions than was previously possible, it nevertheless does not allow an absolutely rigorous analysis to be made. One of the major problems in ethylene4euterium and propenedeuterium studies is that there is no method whereby the true fraction of olefin which has undergone an “olefinalkyl-olefin” cycle and reappeared in the gas phase as olefindo can be determined. This is especially true for reactions on metals such as palladium, ruthenium and rhodium where the olefin exchange results sug-
39 gest a ready redistribution of H and D atoms between adsorbed alkyl radicals and olefin molecules (Sect. 3.5). A deeper analysis of this problem may be, in principle, carried o u t by studying the reactions of the n-butenes, since there is the possibility that in undergoing an ‘‘olefin-lkylolefin” cycle the original olefin will reappear as an isomerised product, e.g. +H
-H
HzC=CH * CH2 * CH3 -r C4H9(ads.) += H3C CH=CH. CH3 I --H +H I
*
*
Relatively few reports of the catalysed reactions of n-butenes with hydrogen were extant up t o the early 1960’s. Those studies which had been performed were mainly concerned with nickel as catalyst. The major problem was the difficulty of chemical analysis of the reaction products. However, with the advent of gas chromatography as a general analytical technique, the analysis of reaction products has become a relatively simple task and, accordingly, over the last 15 years the hydrogenation of higher olefins has received considerable attention. One of the earliest studies of n-butene hydrogenation was that reported by Twigg [ 1211 who observed that, for the reaction of 1;butene with hydrogen over a nickel wire between 76 and 126”C, both hydrogenation and double-bond migration occurred. Hydrogenation and double-bond migration followed the same kinetic rate law, namely Rate a P i ; X Pf$, whilst the activation energies were 10.5 kJ mol-’ (hydrogenation) and 24.7 kJ mol-’ (double-bond migration). Subsequently, Taylor and Diebler [ 1221 studied the reactions of all three n-butenes with hydrogen and deuterium over a nickel wire catalyst. Rates of hydrogenation were in the order b u t - l e n e > cis-but-2-ene = trans-but-2-ene. With the but-2-enes, c i s t r u n s isomerisation was 4-5 times faster than hydrogenation a t 75°C. The kinetics for hydrogenation and double-bond migration were similar and were dependent upon both the total pressure and the hydrogedbut-1< 100 torr the rate expression was of the ene ratio; when PH2 form
-
Rate = h Pki Pf?B but with large excesses of hydrogen, the rate was independent of the hydrogen pressure and proportional to the but-1-ene pressure. Activation energies for the various reactions are shown in Table 9. The results were interpreted in terms of a concerted hydrogen switch mechanism involving an intermediate complex which was common t o all three butenes. The reactions of the n-butenes with hydrogen and with deuterium catalysed by supported noble Group VIII metals have been extensively studied [ 103,123-1271. The variation of the butene composition with the extent References p p . 114-121
40 TABLE 9 Activation energies for butene hydrogenation and isomerisation over a nickel wire catalyst [ 122 ] Olefin
E (hydrogenation) (kJ mol-')
E (isomerisation) (kJ mol-')
But-1-ene cis-But-2-ene trans-But-2-ene
8.4 14.7 14.7
20.9 22.2 20.1
of hydrogenation for typical reactions over palladium, ruthenium and rhodium catalysts are shown in Figs. 14-16, respectively. From these results it can be seen that, as the reactions proceed, the butenes attain their thermodynamic equilibrium proportions [ 1281, although in the early stages of the reaction the cis/truns ratio in the but-1-ene yield is in excess of that expected from the thermodynamic equilibrium. In contrast to these results, those obtained with platinum [lo31 and iridium [126] show that little isomerisation occurs with these metals. Osmium [123] exhibits behaviour intermediate between ruthenium and platinum. The sequence of hydrogenation activity at 100°C is Ni = Rh > Pd > Ru > 0 s > Pt > Ir. Rates of isomerisation may be calculated from the variation of the product distributions with time either directly or as the ratio of the rate constants for isomerisation and hydrogenation [ 123,1251. The direct method, based upon that originally described by Twigg [121] for double bond migration assumes that the but-1-ene and but-2-enes have similar
' " " _.
0
20
40 60 80 n -Butane (%)
1 0
Fig. 14. Distribution of the n-butenes as a function of extent of hydrogenation of butl e n e over palladiumalumina at 37OC [ 1241. 0, But-1-ene; @, trans-but-2-ene; @, cis-but-2ene. Dotted lines indicate thermodynamic equilibrium yields.
41
-0
20 40 600 n-Butane (%)
20 40 0 n-Butane (%)
20 n-Butane
40
(n)
Fig. 15. Distribution of t h e n-butenes as a function of t h e extent of hydrogenation of ( a ) but-1-ene a t 33'C, ( b ) trans-but-2ene a t 33'C and ( c) cis-but-2-ene a t 49OC over r u t h e n i u m a l u m i n a catalysts [ 1231. 0,But-1-ene; 0 , trans-but-2-ene; cis-but-2ene.
*,
adsorption coefficients, in which case ex~(--c~t) where yes is the equilibrium fraction of isomerised olefin, y is the fraction of isornerised olefin at time t, a = k f n ( P B P H 2 ) / & q ( P B ) O , k is the rate coefficient for isomerisation and (PB),is the initial pressure of butene. Thus (Yeq
- Y ) = Yeq
log(Ye q - Y )
=
fft log Y e q - 2.30~
and the rate of isomerisation, ri, is given by
100
-
2 e 75 m Q,
C
$50 3 m 2s 0 0
50 100 /?-Butane (%)
0
50 n -Butane
100 (O/d
Fig. 1 6 . Distribution o f t h e n-butenes as a function of t h e extent of hydrogenation of but-l-ene a t 25OC over (a) rhodium-alumina and (b) r h o d i u m s i l i c a [ 1 2 5 ] . 0 , But-1ene; 3,t r a n s - b u t - l e n e ; f cis-but-2-ene. References p p . 1 1 4-1 21
TABLE 10 Kinetics of hydrogenation and isomerisation of the n-butenes over various metal catalysts h' = kh&zPk;
'dpm = kdpm%2q-:_B;
'ct = k c ,
a
b
m
n
0 60
1.0 1.0
0.0 0.0
0.5
0.0
78.5 66 54
1 .o 1.0 1.0
0.0 4.05
0.5 0.0 0.0
0 0
1.0 1.0
0.0 0.0
80 120
0.6 1.0
? ?
0.6 1.1 a
? ?
0
0.08
0
Catalyst
Olefin
Temp. ("C)
Ru-Al203
1 -B C-2-B
OS-AIZO~ Rh-AI203 Rh-Si02
1-B 1-B 1-B
IrA1203
1-B C-2-B 1-B 1-B
P1-A1203
%zp:-~
-0.1
C-2-B t-2-B
0 0
0.3 0.5 0.3
0 0 0
1-B C-2-B
63 60
1 .o 1 .o
0 0
~~
Valid at low pressure; as P,,
increased, m decreased.
X
Y
1.0
0.0
-0.25 -0.4
Ref.
123 123 125 125
-0.4
103
126 4.43 4.05
? 0.4
124 103
43 The second method, derived by Hamilton and Burwell El271 for the hydroisomerisation of cis-but-2-ene over palladiumalumina, requires the assumption that the isomerisation and hydrogenation reactions have identical kinetic form and that the rates of hydrogenation of the butenes are identical. It was shown that, under these conditions
where x is the fraction of n-butene in total hydrocarbon product, y is the fraction of isomerised olefin in total hydrocarbon product, h h , hi are the rate coefficients for hydrogenation and isomerisation respectively, and K = y e q / ( l- y e s ) = equilibrium constant for isomerisation. The initial rate kinetics for hydrogenation, double bond migration and c i s t r a n s isomerisation are shown in Table 10. The activation energies for isomerisation and hydrogenation are shown in Table 11. The reactions of the n-butenes with deuterium have been studied over alumina-supported platinum and iridium [ 1031 and palladium [ 1241. In general, the results obtained are similar to those discussed above for ethylenedeuterium and propene-deuterium reactions. A comparison of the deuteroalkane distributions over platinum is shown in Fig. 17. TABLE 11 Activation energies for hydrogenation and isomerisation over various metal catalysts
~
_
~
~
Temp. range ("C)
_
_
_-
~
Ei Ei-Eh Ref. (kJmol-') (kJmol-') (kJmol-')
Eh
Catalyst
Olefin
Ru-Al203
1-B C-2-B t-2-B
12-4 5 29.66 32-64
60.7 60.7 41.9
39.8 25.1 27.2
-20.9 -35.6 -14.7
123
Os-A1203
1-B C-2-B t-2-B
62-122 94-137 100-123
35.6 39.8 35.6
54.4 27.2 29.3
18.8 -12.6 -6.3
123
Rh-AI203
1-B
-18 to +27.5 56-105
25.1 20.9
41.9 40.6
16.8 19.7
125
Rh-Si02
1-B
-18 to +18 41-86
38.5 35.6
46.1 41.9
7.6 6.3
125
Ir-AI203
1-B
82-150
39.8
46.1
6.3
126
Pd-Al203
1 -B
29.3 a 29.3 27.2 a 15.5 0
67-85
52.3 a 55.4 69.0 a 53.1 83.7 21.3
-~
-
C-2-B t-2-B
P t- A1 2
03
C-2-B
References p p . 1 1 4--121
-23.0 -26.1 -41.8 -37.6 -83.7
124
103
44
O “ -r
Number of deuterium atoms
Fig. 1 7 . Deuteroalkane distributions observed in the reactions of deuterium with ethylene at 54OC ( O ) , propene at 73OC (@) and but-1-ene at 67OC ( 0 )over platinumalumina [103].
With palladium, the extent of olefin exchange in both the reactant olefin and the isomerised olefin is small, olefin-d,, being by far the major product. This fact, together with the observation that the amount of hydrogen exchange is negligible, strongly suggests the occurrence of intramolecular hydrogen transfer processes. The results for the three n-butenes are broadly similar as shown in Table 12. The broad similarities between the reactions of ethylene, propene and the n-butenes over the noble Group VIII metals may be taken as an indication that the general features of the mechanisms of the three reactions are similar. Accordingly, the isomerisation of all three n-butenes over these metals has, in general, been interpreted in terms of a hydrogen additionabstraction mechanism [mechanism (5), p. 241 rather than an abstraction-addition mechanism involving r-ally1 intermediates [mechanism (6), p. 251. A similar conclusion was drawn by Holbrook and Wise [ 1381 from studies of the isomerisation of butenes over palladium microspheres and palladium supported on a-alumina at low hydrogen pressures. The observed differences in kinetics and activation energies between the various reactions may be accounted for by assuming a common mechanism but different rate-determining steps [ 1231. In the palladium-catalysed reactions, it has been suggested [ 1241 that the reaction involves dissolved hydrogen atoms rather than surface-adsorbed hydrogen. The preferential formation of cis-but-2-ene from but-1-ene, particularly noticeable with rhodi-
-3
; n
1 b
P
," 4
1
TABLE 12 Product distributions from the reaction of the n-butenes with deuterium over palladiumalumina catalysts [ 1241
)-r
2
(J'R)o = (Pc~H&o = 60 torr _____________~~_
Reactant
Conversion
--
Temp. f"C)
_
_ ~ - ---
Product
Deuteroproduct distribution (%)
-do
-dl
4
__-
But- 1-ene
cis-But-2-ene
trans-But-2-ene
4
3
X b 4
__ _ _
4
-ds -
Composition
_~__-
10.2
17
But-1-ene trans-But-2-ene cis-But-2-ene n-Butane
99.3 68.7 67.7 49.9
0.5 19.5 20.9 29.2
0.2 7.5 7.6 13.8
0.0 3.1 2.9 5.3
0.0 1.0 0.8 1.5
0.0 0.2 0.1 0.3
0.01 0.49 0.49 0.80
77.5 8.3 4.0 10.2
9.3
19
cis-But-2-ene But-1-ene trans-But-2-ene n-Butane
94.6 72.2 74.7 44.5
4.0 21.3 19.2 38.6
1.1 5.1 4.6 12.6
0.3 1.3 1.3 3.6
0.0 0.1 0.2 0.7
0.0
0.0 0.0 0.0
0.07 0.36 0.33 0.77
59.1 2.0 29.6 9.3
trans-But-2-ene But-1-ene cis-But-2-ene n-Butane
92.3 68.2 18.0 47.7
5.9 24.6 18.3 46.5
1.0 5.6 3.0 5.2
0.7 1.4 0.5 0.6
0.1 0.3 0.2 0.0
0.10
78.4 1.9 13.6 6.1
6.0
36
a
2
a
No species above - d ~were observed. x in C4HB,Dx or C4HIo7Dx % Composition as given by GLC analysis.
0.0 0.0
0.0 0.0
0.41 0.27 0.58
46
%4
Fig. 18. The effect of mercury coverage ( 6 ~upon ~ ) t h e rates of hydrogenation olefin exchange (@) and isomerisation ( 8 ) of but-l-ene over r h o d i u m s i l i c a at 4 [129].
um at low temperatures (see Fig. 16 and ref. 125) and with iridium [ 1261, has been interpreted as being due t o a particularly restrictive ordering of the but-l-ene molecules in the chemisorbed layer at high surface coverages [ 126,1301, whilst the more usual cis/trans ratios of 0.5-1 .O, observed with the other noble Group VIII metals, may be explained by supposing that the conformations of the adsorbed but-2-yl group immediately before H-atom abstraction are important [ 1311. In an attempt t o understand further the mechanism of double bond migration, the reaction of but-l-ene with deuterium over mercury-poisoned rhodiumsilica catalysts was investigated [ 1291. Figure 18 shows the variation of the rates of hydrogenation ( r h ) ,isomerisation ( T i ) and butl e n e exchange ( r o e ) with mercury coverage. It was also observed that increasing mercury coverage had little effect upon the deuterobut-2-ene distributions, although the hydrogen exchange reaction was markedly suppressed. It was concluded that hydrogenation and isomerisation occurred independently of each other; whereas but-l-ene hydrogenation and exchange occurred on the metal, the isomerisation reaction involved migration of the adsorbed but-l-ene from the metal t o the support, followed by isomerisation on the silica.
47 In an attempt to obtain detailed information regarding the nature of the hydrocarbon intermediate involved in olefin isomerisation, Wells et al. [130,132,133] and Ragaini e t al. [134,135] have studied the reactions of the n-butenes over various catalysts in the absence of molecular hydrogen. Double bond migration and cis-trans isomerisation occur in the absence of molecular hydrogen over cobalt wire and cobaltalumina catalysts [ 1321 via an abstractionaddition mechanism involving I-methyl-r-allyl intermediates [mechanism (6), p. 251. Alumina-supported nickel, ruthenium, rhodium, osmium and platinum [ 1331, iridium [ 130, 1331 and palladium [133,134] are also active for but-1-ene isomerisation at 100°C in the absence of molecular hydrogen. Unsupported osmium, iridium and platinum powders were found t o be inactive under similar conditions, although platinum black is active for both but-1-ene and cisbut-2-ene isomerisation at 135°C in the absence of molecular hydrogen [ 1351. Over the alumina-supported metals, the activity sequences for isomerisation and hydroisomerisation were observed t o be similar. With these catalysts, the results were interpreted in terms of a H-atom additionabstraction mechanism involving an adsorbed but-2-yl intermediate; the importance of the alumina support as a source of hydrogen atoms to initiate the isomerisation was stressed [ 130,1331. I t was suggested that, with platinum black, isomerisation occurred by both an additionabstraction and an abstractionaddition mechanism, the relative contributions of each depending upon the experimental conditions [ 1351. Metallic gold has been reported t o catalyse the hydrogenation of but-lene at 110°C provided hydrogen atoms are provided t o the gold surface 11361. The rate of hydrogenation was proportional to t h e second power of the hydrogen atom concentration and the first power of the but-1-ene pressure. Gault et al. have studied the isomerization and deuterium exchange of the n-butenes over evaporated iron [140] and nickel films [141]. With iron films it was observed that, whereas the but-2-enes formed from but1-ene (& p isomerisation) contained no deuterium, the but-1-ene formed from cis- or trans-but-2-ene ( p a isomerisation) was extensively exchanged. But-1-ene underwent exchange by a stepwise process: the rate of exchange was some 20 times more rapid than OL fl isomerisation. The rates of isomerisation of cis- and trans-but-2-ene were approximately the same as that for a + p isomerisation of but-1-ene, although but-2-ene exchange was negligible. The trans-but-2-ene formed from cis-but-2-ene was little exchanged, in contrast t o the but-1-ene formed by 0 Q isomerisation. A similar result was obtained for cis-but-2-ene formed from trans-but-2-ene. These observations were interpreted in terms of different mechanisms for Q + p isomerisation, but-1-ene exchange and cis-trans isomerisation. It was suggested that but-1-ene exchanges via a vinylic dissociative adsorption mechanism involving the preferential formation of an adsorbed but-2-enyl species, viz. -+
-+
-+
-+
References p p , 114-1 21
48
H,C=CH
+D
--K
*
CH2 . CH3 + H,C=C. CH, . CH3 + H,C=CDCH, I
*
CH3
*
a -+ /3 double bond migration was postulated as occurring by an intramolecular hydrogen transfer mechanism, in which a symmetrical intermediate involving a bridged hydrogen atom between the a - and y-carbon atoms, as originally suggested by Smith and Swoap [142], was formed (Fig. 19).
Fig. 19. Mechanism of double-bond migration by intramolecular hydrogen transfer.
Cis-trans isomerisation was thought t o occur by an addition-abstraction mechanism, involving a single interconversion between adsorbed but2-yl and adsorbed but-2-ene, although it was suggested that a direct process, which does not involve the breaking or formation of any C-H bonds, was also operative. It was further suggested that the mechanisms occurred on different types of site, which were distinguishable by the differing strengths of adsorption of the a - and 0-olefins. Results obtained with nickel films [ 1411 were similar t o those observed with iron, although minor detailed differences were found. The preferential formation of adsorbed but-2-enyl in the vinylic dissociation of but-lene was confirmed by microwave spectroscopic examination of the exchanged but-1-ene. This preferential loss of the 0-hydrogen atom has also been noted in propene-deuterium exchange [ 1431 on nickel catalysts, where propene-2-d, is the most favoured product. As with iron, it was concluded that a + /3 double bond migration occurred by an intramolecular hydrogen transfer, although both a /3 isomerisation and the /3 a migration were thought t o occur on the same sites with nickel. Cis-trans isomerisation was claimed t o occur by both a direct route and via a nonrepetitive olefin-alkyl interconversion. --f
+
3.8 REACTIONS O F HIGHER ALIPHATIC OLEFINS
The metal-catalysed hydrogenation of the higher olefins exhibit general features which are similar t o those observed with the n-butenes. Thus, for example, the hydrogenation of hex-1-ene over Adams platinum catalyst [144] is accompanied by very low amounts of double-bond migration; the relative rates of isomerisation and hydrogenation are in the ratio 0.03 : 1. Similarly, in the liquid phase hydrogenation of the n-pentenes over platinum-charcoal and iridium-charcoal [ 1451, little or no isomerisation
49
," 80 e, I
0 % Hydrogenation
Fig. 20. Percentage yields of trans- and cis-pent-2-ene observed in the reaction of pent1-ene with hydrogen over palladium--charcoal in methanol (0,@),glacial acetic acid (0,B)and benzene (A+) solution [ 1451.
of the reactant olefin is observed. In contrast, the paUadium--charcoalcatalysed hydrogenation of the n-pentenes [ 1441 is accompanied by extensive isomerisation (Fig. 20). Furthermore, the relative rates of isomerisation appear t o be independent of the nature of the catalyst support and of the solvent used. Comparison of the initial cis/trans ratios for pent1-ene and but-1-ene shows that it is appreciably higher in the former case. The rhodium-charcoal-catalysed hydrogenation of the n-pentenes at room temperature exhibits similar results t o those found with platinum [ 1451. Whilst at first sight this might appear t o be unusual, the apparent lack of isomerisation can be readily accounted for by assuming that, as with the n-butenes [ 1251, isomerisation only becomes the predominant reaction a t relatively high temperatures, (Ei - Eh ) being positive. In the isomerisation of the tetra-substituted olefin 3,4-dimethylhex-3ene over palladiumalumina [146], it has been shown that double bond migration is a necessary precursor to cis-trans isomerisation. This has been interpreted as showing that the mechanism involves a series of elementary steps, each of which is stereospecific, although no definite conclusions were drawn as t o whether an additionabstraction or an abstract i o n a d d i t i o n mechanism was involved. Conclusive evidence for the participation of .rr-allylic intermediates in double bond migration has been obtained from a study of the nickelcatalysed hydrogenation of the isomeric olefinic esters methyl oleate and methyl elaidate using tritium as a tracer [ 1471. It was also concluded that in this system c i r t r a n s isomerisation occurred by an addition-abstraction mechanism. References p p . 1 1 4-1 21
50 4. The hydrogenation of alkynes and alkadienes 4.1. INTRODUCTION
The metal-catalysed hydrogenation of multiply unsaturated hydrocarbons is, of necessity, more complex than that of monoolefins. The problems encountered in alkyne and alkadiene hydrogenation are essentially similar and it is appropriate, therefore, that the two systems are considered together. It has already been noted in Sect. 2.3 (p. 8) that, in the hydrogenation of a diunsaturated hydrocarbon, the corresponding monoolefin may be formed as an intermediate product along with the alkane; indeed, the olefin may be the major product in the initial stages of the reaction. This gives rise t o selectivity in catalytic hydrogenation, as noted in Sect. 2.3, and an understanding of the factors which govern the selectivity in any system is of prime importance t o the specification of detailed mechanisms. A further problem arises when one considers the hydrogenation of alkynes and alkadienes containing four or more carbon atoms. In such cases, it is possible that the intermediate olefin may be formed in more than one isomeric form. Hence, the stereospecificity of the hydrogenation must be considered and attempts made t o explain the observed stereospecificity in terms of the mechanism (see Sect. 4.3). 4.2 NATURE OF THE ADSORBED STATE OF ALKYNES AND ALKADIENES
As with monoolefins, it is generally considered that, before undergoing reaction, the hydrocarbon is adsorbed at the metal surface. However, compared with monoolefins, the adsorption of alkynes and alkadienes has been little studied. Acetylene, when adsorbed on active nickel catalysts, undergoes self-hydrogenation with the production of ethylene [ 911, although the extent of this process is less than with ethylene. Similar behaviour has been observed with alumina- and silica-supported palladium and rhodium [ 531, although with both of these metals ethane is the sole self-hydrogenation product; some typical results for rhodiumsilica are shown in Fig. 21. 14C-Tracer studies of acetylene adsorption on alumina- and silica-supported palladium [ 53,651, platinum [ 661 and rhodium [ 531 show the coexistence of at least two adsorbed states, one of which is retained on the surface, the other being reactive undergoing molecular exchange and reaction with hydrogen. Acetylene adsorption exhibits the same general characteristics as those observed with ethylene (see Sect. 3.2). However, there are important differences. The extent of adsorption and retention is substantially greater with acetylene than with ethylene. Furthermore, the amounts of acetylene retained by “clean” and ethylene-precovered sur-
51
10101
molecules in gas pnose t x iu
")
on equilibrium equilibrium with Fig. 21. Adsorption isotherm and composition of the gas phase on 2OoC. 0, 0, Total molethe surface for the adsorption of acetylene o n rhodiumsilica at 2OoC. Total cules adsorbed; s ~gas , phase acetylene; ethane.
*,
faces is identical. The results have been interpreted as showing that the adsorbed states of acetylene involved in retention and hydrogenation are different and that the sites involved in acetylene and ethylene adsorption are different. This latter conclusion is substantiated by field emission microscopic studies of acetylene and ethylene adsorbed on iridium and tungsten [ 56,571 which have shown that, while acetylene is most readily adsorbed on the (110) planes, ethylene is preferentially adsorbed on the (111)planes. The existence of several adsorbed states of acetylene on palladium has been demonstrated in a particularly elegant study by Inoue and Yasumori [ 1481. In this study, it was demonstrated that the catalytic activity of a cold-worked palladium foil showed a marked variation with the annealing temperature. I4C-Acetylene adsorption demonstrated the existence of four types of adsorbed acetylene: A, which underwent desorption on evacuation; B, which was removed from the surface during hydrogenation; C, which was not removed during hydrogenation, but could be removed by treatment with hydrogen a t 150°C; and D, which was retained even after reduction at 150°C. The variation of the relative amounts of A, B, C and D with annealing temperature are shown in Fig. 22. It was also concluded that the type B species could occupy two types of site, (I) and (11). Site (I) was identified with lattice imperfections in the (110) planes, which disappeared on annealing a t 200--300°C. Site (11) was correlated with lattice planes or boundaries preferentially developed during the disappearance of the (110) planes and the growth of (111) planes a t annealing temperatures of around 600°C. References p p . 1 1 4-1 21
52
Annealing temperature ("C)
Fig. 22. Variation in the fractions of various forms of adsorbed acetylene on palladium foil with temperature of annealing [ 1481.
Infrared spectra of acetylene adsorbed on silica-supported nickel, palladium and platinum [ 1491 in the absence of hydrogen show bands ascribable t o an olefinic species (J)
H
IH
\
*
,c=c\
*
(J)
and to surface alkyl groups. Subsequent admission of hydrogen resulted in and intensification of these bands and the appearance of new bands ascribable t o surface alkyl groups of average structure CH3(CH2),, where n = 3 with nickel and n > 4 with platinum. The proportion of olefinic C-H bonds was somewhat higher with palladium than with platinum or nickel. Similar spectra to those observed with the silica-supported metals have been observed with palladiumalumina [ 150,1511, platinum on y-alumina [ 1521 and explosively dispersed palladium and nickel [ 1531. The greater amount of retention observed with acetylene than with ethylene has been ascribed t o the ability of the former t o polymerise extensively. The existence of surface polymers following acetylene adsorption on alumina- and silica-supported platinum [ 601, evaporated palladium films [ 1541 and silica-supported rhodium [ 671 has been demonstrated by thermal desorption studies. Almost no studies of the adsorption of diolefins and higher alkynes
53
have been reported. Avery [70] has found that when buta-l:3-diene was adsorbed on a “clean” palladium-silica surface, no infrared spectrum was observed. Subsequent addition of hydrogen, however, yielded an intense spectrum containing bands similar t o those observed from adsorption of monoolefins on hydrogen-covered surfaces. It was also observed that subsequent heating t o 300°C in hydrogen did not remove the adsorbed hydrocarbon. These observations were interpreted in terms of the formation of (C4)npolymer units which were multiply bonded t o the surface. Khulbe and Mann [155] have obtained infrared spectra of allene adsorbed on silica-supported cobalt, nickel, palladium, platinum and rhodium. The spectra were similar for all the metals, although variations in band intensity from metal t o metal were observed. Addition of hydrogen to the allene-precovered surface resulted in similar spectra t o those found for chemisorbed and hydrogenated propene in which the surface species was thought t o be an adsorbed prop-1-yl group. The authors concluded that the initial allene spectrum was consistent with the adsorbed species being a 1:2di-o-bonded allene (structure K) H2C+=CH2 I I
* *
(K)
The effects of hydrogen on the infrared spectra of adsorbed acetylene together with evidence from mechanistic studies of alkyne hydrogenation has led to the general conclusion that the acetylenic species active in hydrogenation is associatively bonded t o the surface. However, as with monoolefins, there is still doubt as t o the precise formulation of the surfacealkyne bonding. In the early work [ 1561, it was assumed that the associatively adsorbed complex was adequately represented as a di-abonded olefin, which adopted a cis-configuration.
R RC-CR’+
*
\ I /C=C \
R’
*
This postulate led t o various conclusions regarding (a) the stereochemistry of acetylene hydrogenation (see Sect. 4.3) and (b) the geometrical requirements of the adsorbed state. Thus, it was concluded that, for the face-centred cubic metals, adsorption of acetylene could only occur with minimum strain on those faces showing the longer interatomic specings, namely the (100) and (110) faces [ 1571. It was also concluded that the (111)faces would be inactive in acetylene adsorption, and the absence of suitable spacings on the close-packed hexagonal metals, ruthenium and osmium, was claimed t o be responsible for the alleged non-activity of these metals in acetylene hydrogenation [ 1581. Subsequently, it has been
6
References p p . 1 1 4-1 21
54
Fig. 23. Possible modes of adsorption of acetylene on a Pt (111) surface [ 1601.
shown that both metals possess appreciable activity and that suitable fi spacings exist on the (3054)planes [ 1591. In a recent study of the adsorption of acetylene on platinum single crystals by low energy electron diffraction [ 1601, it has been shown that acetylene adsorbs on the (111)planes. These results show that, on a clean pt (111)surface, acetylene adsorbs at a distance of 1.95 A above the topmost plane of platinum atoms, either in the C2 or, less likely, the B, mode shown in Fig. 23. No evidence was found for adsorption in the A, or A2 modes, which corresponds to a n-complex structure or for the B2 mode corresponding to a di-a-complex, although it was stated that such structures may be possible with a less stable overlayer which had been observed. Following the proposals of Rooney et al. [85-871, it has generally been assumed that, as with monoolefins, the adsorbed state of an alkyne active in hydrogenation is a n-complex formed by the interaction of the n-orbitals of the acetylenic bond with two metal atoms. The n-complexed alkyne may be represented as structure L. R-CZEC-R’ I
**
(L)
This postulate has several implications regarding the mechanism of alkyne hydrogenation; these will be discussed in Sect. 4.3. It should be noted, however, that there is as yet little or no direct evidence for structure L, although analogous structures are known t o exist with organometallic complexes [161]. Such a structure is also consistent with the positive surface potentials observed for acetylene adsorption on evaporated nickel films [ 881. In the ensuing discussion, it will be assumed that structure L is the relevant species in alkyne hydrogenation, and that the catalytically active adsorbed state of an alkadiene can be represented as a n-olefin complex in which either one or both olefinic bonds interact with the surface.
55 4 . 3 POSSIBLE REACTION MECHANISMS
As noted in Sect. 4.1,in the hydrogenation of diunsaturated hydrocarbons it is generally observed that both the corresponding monoolefin and alkane are formed in the initial stages of the reaction, the former product generally predominating. Further, in the reactions of alkynes and diolefins containing more than three carbon atoms, a distribution of isomeric olefins is usually observed. It is generally agreed that the kinetics and the distributions of deuterated products from the reactions of alkynes or alkadienes with deuterium are satisfactorily interpreted in terms of the consecutive addition of two hydrogen atoms, of unspecified origin, t o the adsorbed hydrocarbon to yield the monoolefin. The identity of the distributions of deuteroethylenes from the reaction of acetylene with equilibrated and non-equilibrated hydrogen-deuterium mixtures also provides strong evidence for such a mechanism [ 911. The interpretation of the experimentally .observed selectivity and stereoselectivity depends t o some extent on the assumed nature of the adsorbed alkyne and alkadiene. If the alkyne is adsorbed as a di-n-complex (structure L), the product olefin will be formed as an adsorbed species, which must, therefore, undergo desorption before appearing in the gas phase (see Fig. 4). Consequently, the selectivity defined as
will, in part, be dependent upon the relative abilities of the olefin t o desorb, which may be aided by displacement by the more strongly adsorbed acetylene, or undergo further hydrogenation. Alternatively, if, as was originally envisaged [157] the alkyne is adsorbed as a di-o-bonded complex (structure J), hydrogenation will lead t o the direct formation of olefin in the gas phase.
R
*
R
R
c=cI +-,H c\= C/ I \ I \
\
*
H
R'
R
+ H ---f
*
R'
\ I /C=C (gas) \ H H
Further hydrogenation of the monoolefin, and hence the selectivity, would depend upon the readsorption of the olefin in competition with the alkyne. In either of the above cases, the selectivity will also depend upon the existence, or otherwise, and the relative importance of a direct route from the diunsaturated hydrocarbon to the alkane, not involving the monoolefin as an intermediate. For example, a possible route t o the direct formaReferencesp p . l 14-1 21
56
H-GC-H4 I +H
**
H C=C + C=C-H 1 1 \ +H /I \ * * H ** H
H\
I
I
* Fig. 24. Possible reaction scheme for the direct formation of ethane from adsorbed acetylene.
tion of ethane from acetylene could be envisaged as shown in Fig. 24. Similar mechanisms can be envisaged for other alkynes and diolefins. The stereoselectivity observed in alkyne hydrogenation may depend upon two factors. It is generally agreed that the consecutive addition of two hydrogen atoms to adsorbed alkyne will yield the monoolefin in the cis-configuration. This is the case since, in the di-.rr-adsorbed alkyne the substituent groups will adopt a cis-configuration [ 1621 and hydrogen atom addition can then only occur from below the plane of unsaturation.
*
*
On the other hand, it is possible that, if the di-.rr-adsorbed alkyne reacts with molecular hydrogen, for which there is kinetic evidence [ 91, addition above the plane of unsaturation with the formation of the trans-olefin could be envisaged, viz.
The second factor which may influence the stereoselectivity is the
anti-alka-1:3-diene
syn-alka-l:3diene
Fig. 25. Conformations of an alka-l:3-diene.
57
simultaneous occurrence of hydrogenation and isomerisation of either (a) the reactant alkyne or (b) the olefin product, either before desorption or following readsorption. This factor is also of importance in the hydrogenation of diolefins. However, with conjugated diolefins, a further factor must also be taken into consideration, namely the existence of more than one conformation of the reactant (Fig. 25). 1 : 4-addition of hydrogen to the anti-conformer would be expected t o yield the trans-olefin, whereas the syn-conformer would yield the cis-olefin. 1 : 2-addition of hydrogen to both conformers would yield the a-olefin. 4.4 TREATMENT O F EXPERIMENTAL RESULTS
The interaction of an alkyne or alkadiene with deuterium leads t o the formation of deuteroalkenes whose isotopic composition yields valuable information regarding possible reaction mechanisms. In an attempt to interpret in detail the deuteroalkene distributions, two approaches have been used. The first, due to Bond [ 1631, is a simplified version of the general theory proposed by Kemball for the hydrogenation of ethylene (see Sect. 3.4) and has been used to interpret the results of the reaction of acetylene with deuterium [ 163-1651. The method comprises a steady state analysis of the reaction scheme
3 C,X,(a)
C,X,(ads.) : + C2X3(ads.) -X
(where X = H or D) in which the possible deuteroacetylenes and deuteroethylenes are related by a series of simultaneous equations involving three parameters p , q and s where p is the percentage chance that adsorbed vinyl, once formed will react further to form ethylene rather than revert to adsorbed acetylene, q is the percentage chance that adsorbed acetylene will react with a deuterium atom rather than a hydrogen atom and s is the percentage chance that adsorbed vinyl will react with a deuterium atom rather than a hydrogen atom. The simultaneous equations so obtained are solved for various values of p , q and s until a satisfactory fit between calculated and observed deuteroethylene distributions is obtained. The second approach is that developed t o interpret the products of the reactions of octalins with deuterium [144] and is equally applicable t o the reactions of mono- or di-unsaturated hydrocarbons with deuterium. Smith and Burwell [144] pointed out that, whereas the experimental deuterohydrocarbon distributions are obtained in terms of the number of deuterium atoms in the product hydrocarbon, the quantities of fundamental importance to the discussion of the mechanisms of catalytic reactions are the fractions of the hydrocarbon sample which have equilibrated with the surface deuterium-hydrogen pool. Thus, for example, in the reaction of buta-1: 3diene with deuterium, the product butenes consist of a series of species, butene-(h, d ) 2 ,-(h,d ) 3 ,..., -(h, d), in which 2,3 ..., n positions References p p . 1 14-1 21
58 have been equilibrated. The fractions of such species are designated N 2 , N 3 ..., N,, the complete series being termed an “N-profile”. Butene-(h, d ) , will thus appear as a random distribution of the possible deuterated species, butene-d,, - d l ..., d,. If d,(n) is the fraction of the species butene-(h, d), which has gained rn deuterium atoms and s hydrogen atoms, (rn + s) = n, then
(alb)“ n ! d m ( n ) = (1 + a/b)”rn!s! where a/b = surface D-atoms/surface H-atoms. The fraction of molecules containing rn deuterium atoms, d, ,in the total butene is given hv d, = C d , ( n ) N , I
for n > rn. The N-profile is then calculated by assigning trial values to N o , N , ..., N,, and to the ratio a/b such that the calculated deuterium distribution, derived from the N-profile, is in good agreement with the experimentally observed distribution. This method has been used by Wells and co-workers [ 166,1671 t o interpret the deuterobutene distributions observed in the catalysed hydrogenation of buta-l:3-diene (see Sect. 4.9). 4.5 THE HYDROGENATION OF ACETYLENE
Among the early systemmatic studies of the metal-catalysed hydrogenation of acetylene were those of Sheridan et al. [158,168-1701 who investigated the kinetics and product distributions over pumice-supported metals. Subsequently, the reaction has been extensively studied by Bond et al. [ 9,165,171-1751 over pumice- and alumina-supported metals and metal powders. The reaction of acetylene with deuterium over nickel [ 91, 1631 and alumina-supported noble Group VIII metals [ 164,1651 has also been investigated. For reactions carried out in a constant volume reactor, the shapes of the pressure fall against time curves are dependent upon the initial hydrogen:acetylene ratio and upon the metal. Figure 26 shows typical pressure fall against time curves. Analysis of the reaction products shows that with those metals which exhibit a high selectivity for ethylene formation, the pressure fall against time curves are of types B and C in Fig. 26, the reaction occurring in two distinct stages, the onset of the second stage being accompanied by an increase in rate. Conversely, those metals which exhibited a low selectivity show pressure-time curves with either an “acceleration-point” occurring very late in the reaction, type D (Ir), or no “acceleration” at all, type E (0s). The overall order of reaction as determined from the pressure fall
Fig. 26. Forms of pressure fall against time curves observed in the hydrogenation of acetylene over noble Group VIII metal catalysts.
against time curves was observed t o be dependent upon the initial hydrogen : acetylene ratio and upon the metal. For alumina-supported metals [9,165], ruthenium, osmium and iridium gave rise to pressure fall-time curves in which the order changed continuously throughout the reaction; with rhodium, palladium and platinum, the rate of pressure fall, r, during the first stage of the hydrogenation, obeyed the expression r = h ( P H , ) q , where (PH2)1 is the instantaneous pressure of hydrogen at any point in the reaction. For initial (H2/C2H,) ratios >2, values of x were 1 . 5 (fresh Rh) decreasing t o -1.0 (well-used Rh) and 0 (Pd and Pt). For ( H 2 / C 2 H 2 )ratios of unity, values of x were 0.5 (Pd) and 1.0 (Pt). Similar behaviour has been observed with silica-supported rhodium and palladium using (H,/ C,H,) ratios of 3, although in both cases the value of x was observed t o be unity [ 531. The kinetics, determined from initial rates, and the activation energies for acetylene hydrogenation over a variety of metal catalysts are shown in Table 13. One of the characteristic features of the metal-catalysed reaction of acetylene with hydrogen is that, in addition t o ethylene and ethane, hydrocarbons containing more than two carbon atoms are frequently observed in appreciable yields. The hydropolymerisation of acetylene over nickel-pumice catalysts was investigated in some detail by Sheridan [ 1691 who found that, between 200 and 2 5 0 ° C , extensive polymerisation to yield predominantly C4 - and C6 -polymers occurred, although small amounts of all polymers u p t o C , , where n > 31, were also observed. It was also shown that the polymeric products were aliphatic hydrocarbons, although subsequent studies with nickelalumina [ 1761 revealed that, whilst the main products were aliphatic hydrocarbons, small amounts of cyclohexene, cyclohexane and aromatic hydrocarbons were also formed. The extent of polymerisation appears t o be greater with the first row metals, iron, cobalt, nickel and copper, where up t o 60% of the acetylene may polymerise, than with the second and third row noble Group VIII metals. With alumina-supported noble metals, the polymerisation prodReferences p p . 1 1 4-1 21
60 TABLE 1 3 Initial rate kinetics and activation energies for acetylene hydrogenation Rate
=
k P " H ~ P at ~ T ~ H C~
Catalyst
Temp. (" C)
X
136
1.4 1.0 1.0 1.0 1.0
Fe-pumice Fe-powder Co-pumice Co-powder Ni-pumice
- 135 - 130 - 135
Ni-powder Cu-pumice Ru-A12 O3 Rh-pumice Rh-A12 O3 Rh-A12 O3 Pd-pumice Pd-A12 O3 Pd-AI203 Pd-SiO2 Os-A1203 Irpumice Ir-A1203 Pt-pumice Pt- A12 03 Pt-A12 O3
30 150 112 85 130 130 49 20 0 114 165 175 130 73 105 110
79
-0
0
1.0 Variable 1.0 -1 -1 1.O-1 .5 -1 1.4 1.0 1.0 1.0 -1 1.0 1.2 1.5 1.5
-
Temp. range ("C)
Ref.
(kJ mol-') 64.0 29.7 17.2 37.7 50.7
20-200 108-162 107-157 111-156 0-126
58.6 79.5 44.0 64.9 44.4 31.7 49.8
16-57 150-195 92-145 17-110 132-162 115-165 0-120
45.7 71.2 33.5
0-30 114-160 144-202
31.8 50.2 73.7 38.9
115-180 0-120 77-161 40-96
168 172 172 172 168, 178 178 172 165 158 173 9 168 173 9 180 165 158 9 170 173 9
Ea
Y
0.5 0.3 0.0 -0 0 -0.5 --0.5 0 0 0 -0.3 -0.7 -0.7
ucts have been observed to be almost exclusively C4 -hydrocarbons [ 9, 1651; typical distributions of the polymer products are shown in Table
14. The mechanism of the hydropolymerisation of acetylene is not too clear. It has been suggested [9,169] that in the hydrogenation of acetylene to ethylene, the half-hydrogenated state, an adsorbed vinyl species, may exist in either a normal or free radical form, viz.
H-CrC-H I
+ I$HC=CH2 or HC=CH2 +
I1
**
I
**
*
Interaction of the free radical with either an adsorbed acetylene or a normal vinyl radical could then lead to polymerisation.
HC=CH2 + HCcCH I I
*
+
**
HC=CH2 + HC=CH2 I I I
*
**
H,C=CH--CH=CH I I
*
+
*
H2C=CH-&I=CH2 I I
*
*
U
P
TABLE 14 Percentage composition of C4 -products from acetylene hydrogenation over alumina-supported Group VIII metals Initial C2H2/H2 Metal
Ru Rh Pd 0s Ir Pt a
a
- 1.
Temp. ("C)
Cz yield
166 135 16 123 135 136
92 85-88 63 84 85-88 72
Butadiene
1-Butene
trans-But-2-ene
cis-But-2-ene
0.0 4.0 0.0 0.0 8.0
42.5 54.0
7.3 24.0 Trace amounts 1.5 22.0
34.3 14.0
Butane
(%I
Initial HZ/C2H2 = 2. Initial H2/C2H2 = 4. 8.7% Isobutene also formed.
35.4 47.0
17.9 20.0
7.2 4.0 > 95 45.2 3.0
Whether an adsorbed species can exist as a free radical on a metal surface is open t o some debate. It seems unlikely that the vinyl free radical will exist as a relatively stable surface intermediate although it could be envisaged as a transition state in the formation of adsorbed vinyl from acetylene. An alternative mechanism involving the direct insertion of a vinyl group into an acetylene
HC-CH + HC=CH2 + HC=CH-CH=CH, I I I I I I
**
**
*
**
similar to that envisaged in Zeigler-Natta polymerisation [ 1 7 7J seems to have received little attention t o date, although it would appear t o satisfy the experimental observations. The general theoretical approach t o the selectivity observed in the hydrogenation of acetylene has been discussed in Sect. 2.3, where it was noted that the observed selectivity may be dependent upon both thermodynamic and mechanistic factors. A possible explanation of the operation of a mechanistic factor has been discussed in Sect. 4.3. The selectivity values, defined as S = P C 2 H 4 / ( P C 2 H 4 + PCZH6), observed for various metal catalysts are shown in Table 15. Selectivities have been observed to TABLE 1 5 Selectivity observed in acetylene hydrogenation over supported metals Catalyst
H ' 2 lpC 2 H2
Temp. ("C)
Selectivity
Ref.
Fe-pumice Co-pumice Ni-pumice
1 1 1
Cu-pumice Ru-AI2 O3 Rh-pumice Rh-AI203 Rh-AI2 O3 Rh-Si02 Pd-pumice Pd-Al2 0 3 Pd-AI203 Pd-SiO2 Pd-SiO2 0s-A12 O3 Ir-pumice Ir-AI2 0 3 Pt-pumice Pt-AIz (33 Pt-AI, 0 3
1 2 1 2 2 3 2 2 2 3 3.75 5 1 2 1 2 2
156 197 80 125 200 135 85 133 150 20 36 22 0 20 181 123 175 130 163 105 110
0.91 0.90 0.83 0.91 0.91 0.80 0.86 0.92 0.90 0.75 0.92 0.95 0.97 0.96 0.97 0.54 0.30 0.55 0.82 0.86 0.90
168 168 169 168 168 165 158 173 9 53 168 173 9 53 180 165 168 9
_ . _ . _ ~ _ _ _ _ _ _
1
173 9
63 decrease with increasing hydrogen pressure and t o increase with increasing temperature. The shapes of the pressure-time curves together with the observation that the selectivity remains constant or nearly so until the acceleration point is reached has been taken t o indicate that the thermodynamic factor is high, that is the presence of acetylene effectively prevents the readsorption of ethylene from the gas phase and also aids the desorption of ethylene (Fig. 4). Such a conclusion makes the implicit assumption that the Same sites are involved in acetylene and ethylene adsorption. Recent studies using I4C-tracers [ 531 have shown that, with alumina- and silica-supported palladium and rhodium, the admission of acetylene t o I4C-ethylene-precovered surfaces results in the displacement of a small fraction of the l 4 C-ethylene, although a further fraction will undergo hydrogenation to I4C-ethane. Furthermore, the addition of 14C-ethylene t o an acetylene-hydrogen reaction mixture shows that the added 14C-ethyleneundergoes hydrogenation independently of the hydrogenation of acetylene. By the selective use of I4C-tracers, it has also been observed that the adsorbed acetylene which comprises the primary region of the adsorption isotherm (Fig. 27) yields only ethane during an hydrogenation reaction, while the acetylene adsorbed on the secondary region yields both ethylene and ethane. These results show that earlier interpretations of selectivity [ 9, 1571 in terms of a thermodynamic factor which governs the ratio O C Z H 2 / 6 C 2 H 4 and a mechanistic factor, which arises from the ratio of the rates of ethylene desorption and further hydrogenation (p. 13), are inadequate. A full understanding of the factors which influence the selectivity must
I
I
2 2
3 3
I
+ Primary region a,
U
0 L
i3 n0uJ
0
e 1 -1-
-
1
4 4
1
Gas phase count-rate (min-' x lo-')
Fig. 2 7 . ''C-acetylene adsorption isotherm on palladiumsilica at 20°C (0) and the I4C-ethylene adsorption isotherm on an acetylene-precovered palladiumsilica catalyst at 20°C ( 0 ) . References p p . 1 1 4-1 21
64
include the following: (i) the relative amounts of acetylene and ethylene adsorbed on independent sites; (ii) the relative amounts of different forms of adsorbed intermediates which can only lead to the formation of the alkane (Fig. 24, Sect. 4.3) The results obtained using carbon-14 tracers are in general agreement with those obtained in a recent study of the hydrogenation of acetylene in the presence of excess ethylene over palladium-alumina catalysts [ 1791. These show that at least two types of site exist on the catalyst surface. On type X sites, the hydrogenation of both acetylene and ethylene can occur, although acetylene is adsorbed some 2200 times more strongly than ethylene at 20°C. Type Y sites, which are easily poisoned by carbon monoxide, can hydrogenate ethylene in the presence of acetylene, but are inactive for the hydrogenation of acetylene. From the 14C-tracer results, a general reaction scheme for acetylene hydrogenation involving three types of surface sites can be envisaged (Fig. 28). Type I sites are those responsible for the acetylene adsorbed on the secondary region; type I1 sites are those responsible for the primary adsorbed acetylene species and type I11 sites are identified with the type Y sites discussed above. The type I and I1 sites may together be identified with the type X sites mentioned above. The reaction of acetylene with deuterium has been studied over alumina-supported noble Group VIII metals [ 164,1651, whilst over nickelpumice catalysts the reaction of perdeuteroacetylene with hydrogen has been investigated [ 1631. In both of these studies, the deuteroethylene distributions have been interpreted in terms of the steady state analysis discussed in Sect. 4.4. Typical deuteroethylene distributions together with the values of p , q and s are shown in Table 16. N o acetylene exchange was observed with Rh, Pd, Ir or Pt, although the steady state analysis showed that -10% (Pd, Pt) or -30% (Ir, Rh) of the adsorbed acetylene was either C2HD or C2Dz.Thus acetylene adsorp-
a 0 3
2 b
P
I
L.
2
TABLE16 Observed and calculated deuteroethylene distributions observed in reaction of C2H2 with deuterium Catalyst
Ru-Al203
Temp. ("C) 144
D2/C2H2
3.36
(Calc.) Rh-AI203 Pd-Al203
134
(Calc .) 15
2.2 2.0
(Calc.) O S - A ~ ~ O ~ 148 (Calc.) I ~ - A I ~ O ~ 42 (Calc.) Pt-Al203 110 (Calc.)
2.63 4.0 2.0
Ethylene distribution (%) C2H4
C2H3D
CzH2D2 C2HD3 C2D4
1.5 2.0 1.2 1.0 2.3 1.9 10.5 6.2 1.8 1.3 2.5 1.7
10.0 15.2 11.7 11.2 22.0 22.0 28.2 28.2 15.2 14.6 19.5 20.2
37.7 37.2 36.7 36.7 65.2 65.2 38.9 39.9 51.2 51.2 63.1 63.1
30.4 30.0 27.7 32.5 8.8 10.0 18.2 20.6 22.5 25.4 11.9 13.3
20.4 15.5 22.7 18.9 1.8 0.8 5.0 5.0 9.3 7.8 3.0 1.5
x in
H i n D2
CZH4-xDx
(%)
2.60 2.42 2.59 2.58 1.86 1.86 1.81 1.91 2.23 2.24 1.93 1.92
Ref.
p
q
s
85
80
70
82
82
82
39
85
85
80
60
70
63
86
86
164
37
86
86
164
7 5.5 0
14 2.5 0
165 164 164 165
66 tion was irreversible. A similar conclusion was reached for Ni, where acetylene exchange wsls extremely slow [ 1631, and for the Ru- and Os-catalysed reactions, although the results were complicated due to rapid exchange of the acetylene with the alumina support [ 1651. The steady state analysis shows that the abilities of the metals to promote vinyl reversal, p , were in the same order as those observed for ethyl reversal, p * , in ethylene deuteration (Sect. 3.5). The values of the parameters p , q and s obtained from the steady state analysis together with the observed kinetics enable the establishment of fairly precise mechanisms for acetylene hydrogenation. Bond and Wells [ 341 have proposed three mechanisms for the reaction over alumina-supported Group VIII noble metals and nickel. Mechanism (A) occurs with Ru, Os, Ir and Rh. With these metals, the occurrence of hydrogen exchange shows the hydrogen adsorption is reversible. The high values of p observed with these metals suggests that the ratedetermining step is probably the addition of hydrogen t o adsorbed vinyl. The following mechanism is consistent with the experimental obervations. H2(g) ?+ 2 H(a)
CZH4(a)
+
(192)
C2H4(g)
Steady state analysis shows that
and the rate of ethylene production, u , is given by
From the observed kinetics, the value of p is high. Thus
eCZH2
+
1 and O H
a
P i ; , k 5 >> k 6 e ~since
The order of unity in hydrogen is thus accomodated. Mechanism (B) was postulated to explain an order of unity observed with palladium and nickel, where hydrogen adsorption was @eversible and p was low. The relevant steps for mechanism (B) are (l),(3), (4), (7)
67 and the ratedetermining step 2 C2Hda) -,C2H4(a)+ C2H2(a)
(8)
The rate of formation of pairs of vinyls will be proportional t o the hydrogen pressure and hence the rate of formation of ethylene by step (8) will be first order in hydrogen, An alternative mechanism for the palladium-catalysed reaction involving the reaction of adsorbed acetylene with molecular hydrogen has been proposed [ 91. Mechanism (C) was proposed for the platinum-catalysed reaction to explain the observed order of 1.5 in hydrogen. The mechanism comprises steps (l), (3)-(6), (9) and the reaction C2H2(a) + H2
CZH3(a) + H(a)
Steady state analysis with respect t o CzH3(a)shows that
The rate of ethylene formation, u, assuming step (6) t o be rate-determining, is given by u = k6eC2H38H
- k6eH(k
l0eC2HzPH2
+ k4eC2H2eH)
k5 + k6eH It was suggested that an order of 1.5 could result if k 5 k 6 d H . The stereochemistry of acetylene hydrogenation has been examined by determining the relative yields of cis-, trans- and asyrn.-C2H2D2 by infrared spectroscopy [ 163-1651. Some typical results are shown in Table 17. Cis-C2H,Dz will result from the reaction of either adsorbed
-
TABLE 17 Typical distributions of cis-, trans and asyrn.-dideuteroethylenefrom the reaction of C2H2 D2 +
Metal
Temp.
cis-CzHzD2
trans-C2HzDz
nsy m -C2H2D2
2
Ni a Pd Pt
41 15 50
I7 83
21 15
85
14
Rh Ir
135
48
63
40 30
12
120
42
42
16
56
33
11
Ru 0s a
167 153
Reaction of CzDz with H2.
References p p . 1 1 4-1 21
2 1
I
68
D \
H I
?TC\
P
H * *
D\
c=c
1 I \D H *
lH
c=c * *
1 I iD
D--C=C-H I
**
+H I
Fig. 29. Routes t o the formation of truns- and usyrn.-CzH2D2in the reaction of CzH2 with D2.
C2H2 with two D-atoms or of adsorbed C2D2 with two H-atoms, while truns-C2H2D2and usym.-C2H2D2can only result from the addition of one D-atom and one H-atom to adsorbed C2HD as shown in Fig. 29. The yields of truns- and usym.-C2H2D2 would therefore be expected t o be the same, both being less than that of cis-C2H2D2.The formation of more truns-C2H2D2 than expected has been attributed t o the existence of a free radical form of adsorbed vinyl [ 1631, viz. H H H \ I \. IH ,C=C * C-C-De
D H \ I ,,C=C
*
*
I
\
D
*
\
*
\
H
4.6 HYDROGENATION OF MONOALKYLACETYLENES
In comparison with the extensive studies of acetylene itself, the hydrogenation of monoalkylacetylenes has received much less attention. The hydrogenation of methylacetylene over pumice-supported metals [ 170, 181,184,1851 and over metal powders [181-1851, has been studied. No studies of t h e reaction of propyne with deuterium have been reported. Table 18 shows the kinetics, activation energies and selectivity observed over various metal catalysts. These results show a close resemblance t o those obtained with acetylene and, by analogy, similar mechanisms to those proposed for acetylene have been invoked. In the absence of deuteration studies, unequivocal mechanisms cannot readily be obtained. Like acetylene, some polymerisation of methylacetylene has been observed
TABLE 18 Kinetics, activation energies and selectivity observed in the hydrogenation of methylacetylene
~ T(^C);S Rate = h PX,, P & H at
= P C ~ H J ( P C+~ PHc~~
Catalyst
Temp ("C)
X
Fe-powder Co-powder Ni-powder Ni-powder Ni-pumice Ni-pumice Ni-Kieselguhr Cu-powder Rh-pumice Rh-powder Pd-powder Pd-pumice Pd-pumice Ir-powder Ir-pumice Pt-powder Pt-pumice Pt-pumice Ru-pumice Ru-powder
131-220 30-100 70-130 46-70 20-110 0-133 5-64 181-220 47-76 73-111 12-45 16-4 1 47-89 40-70 85-121 95-135 72-135 47-89
1.0 1.0 1.1 1.1 1.14 1.0 1.77 0.97 1.0 1.0 1.25 1.25 1.0 1.03 1.03 1.01 1.01 1
H~) Y
0 -0.16 -0.12 0
0 0 0 0.49 0 0 -0.24 -0.52 0 0 -0.2 -0.3 -0.4 0
Ea ( k J mol-') (at T"C)
5
30.6 (130-220) 33.9 ( 30-100) 51.1 ( 70-130) 72.0 ( 46-70) 70.3 ( 20-110) 59.4 ( 0-133) 58.6 ( 5-64) 176.6 (180-220) 44.8 ( 47-76) 49.0 ( 73-111) 39.8 ( 12-45) 44.0 ( 16-41) 69.1 ( 78-198) 26.0 ( 40-70) 36.0 ( 85-121) 61.5 ( 99-135) 51.9 ( 70-135) 72.4 ( 65-180)
0.99 (218) 0.90 ( 98) 0.88 (115)
Ref.
(at T"C )
?
0.75 ( 97) 0.93 ( 91) ?
1.00 (173) 0.93 ( 72) 0.52 (106) 1.00 ( 37) 0.97 ( 40) 0.79 ( 80) 0.71 ( 95) 0.79 ( 95) 0.94 ( 95) 0.92 ( 95) 0.89 ( 75) 0.88 (154) 0.88 (140)
182 182 182 181 181 170 181 183 185 185 185 185 170 184,185 184,185 184,185 184,185 170 185 185
70 TABLE 19 Selectivities observed in the hydrogenation of but-l-yne over iron, cobalt and nickel catalysts (187,1881 at
= ‘ C ~ H S / ( ~ C ~+H‘C4H10) ~
Catalyst
Initial pH21pC&
Fe-powder Fe-pumice Co-powder Co-pumice Ni-powder Ni-pumice Ni-pumice
2.14 3.33 3.13 2.74 2.73 2.63 1.00
pc Temp. (“C)
s
130 117 35
0.88 0.84 0.78 0.82 0.62 0.70 0.74
50 40 50 60
although, in general, the yields of polymers are much lower with methylacetylene. The hydrogenation of but-l-yne in alcoholic solution using palladium on barium sulphate yielded only but-l-ene (98%) and n-butane [186]. But-lene and n-butane were the sole initial products from the hydrogenation of but-l-yne over pumice-supported iron, cobalt and nickel and unsupported powders of the same metals [ 187,1881 , although in the later stages of the reaction, when but-l-ene underwent further hydrogenation t o n-butane, small amounts of truns- and cis-but-2-ene were observed. The selectivity observed with these catalysts are shown in Table 19. Small amounts of polymer were also observed, although, in general, less than 4% of the but-l-yne was consumed in this manner. The initial rate orders were invariably unity in hydrogen and zero or slightly negative in hydrocarbon over all the catalysts. With palladium-alumina, the products of the reaction of but-l-yne with deuterium [ 1891 were but-l-ene, 99.1%;truns-but-2-eneY0.2%; cisbut-2-eneY0.2%; n-butane, 0.5%, until at least 75% of the but-l-yne had reacted. But-l-ene hydrogenation and hydroisomerisation were observed t o occur when all the but-l-yne had reacted. The formation of but-2-ene as an initial product was postulated as being the result of a slow isomerisation of but-l-yne t o absorbed buta-l:2-diene
CH3 CH2 - C-CH 9
I I
**
‘tH CH3 - CH2 -H
*
C=CH2 I I
**
3 CH3 - CH=C=CH2 I I * *
which on hydrogenation yielded both but-l-ene and the but-2-enes. Analysis of the but-l-ene and but-l-yne by mass spectrometry showed small amounts of but-l-yne-d, and of all deuterobutenes up t o - d 3 . Furthermore NMR analysis showed that the ethyl group of the but-l-ene contained no deuterium.
71
J C2HSCX=CH + C H C=CHX I I sll
**
I
**
C2 H 5 CX=CHX I
*
1
But-1-ene-d,, -dl, -do
C2HSCX=CD + C H C=CDX sll
**
I
**
C2H SCX=CDX I
*
1
But-1-ene-d, -dz, -dl
(X = H or D) Fig. 30. Reaction scheme for t h e hydrogenation of but-1-yne over palladium-alumina [ 1891.
From these results, the reaction scheme shown in Fig. 30 was proposed [ 1891. It was shown that the rate of formation of but-l-ene-d3 was some twelve times faster than the rate of but-1-yne-d equilibration, suggesting that the interconversion of but-1-yne-do and -dl occurred without desoption of the but-1-yne from the surface. An alternative route to but-l-yned l involving reaction of adsorbed but-1-yne with molecular deuterium and reversal of the formation of adsorbed but-2-enyl has been suggested [ 341. Reaction of pent-1-yne and hex-1-yne with hydrogen over nickel catalysts leads t o the selective formation of the corresponding alk-1-ene and the alkane [ 190,1911, in an approximate 3 : 1ratio. 4.7 HYDROGENATION OF DIALKYLACETYLENES
The hydrogenations of dialkylacetylenes are of particular interest from a stereochemical viewpoint since, as noted in Sect. 4.3, the adsorbed state of the acetylene is expected t o adopt a cis-configuration and, consequently, upon hydrogenation t o yield the cis-olefin. Wide use of this fact has been made in preparative organic chemistry as noted by Burwell [ 1921 and by Campbell and Campbell [ 1931. Although early studies of the catalytic hydrogenation of disubstituted acetylenes [ 194-1961 revealed the formation of trans- as well as cis-olefins, it was generally assumed that the trans-isomer was formed by isomerisation of the cis-olefin. However, more recent studies have shown that this view may have References p p . 1 1 4-1 21
72 TABLE 20 Initial distribution of products observed in the hydrogenation of but-2-yne over alumina-supported Group VIII metals [ 199,2001
(“(2
Initial cis-But P H ~ I P c ~ H -2-ene ~
Ni
200 140 152
2 : l 2 : l 2 : l
Ru Rh
90 154
0s Ir Pt
120 160 158
Metal Fe
co
Temp.
trans -But -2-ene
But-l-ene
Selectivity
76 88 95
20 5 1
1.00 1.00 1.00
2 : l 4 : l
79 85
16 7
0.97 0.99
2 : l 10 : 1 7 : l
74 87 87
22 5 5
0.90 0.96 0.97
been erroneous and that the trans-olefin may be formed directly from the adsorbed alkyne. Studies of the palladiumalumina-catalysed hydrogenation of but-2yne in a flow system [127,197] have revealed that the reaction is highly stereoselective for the formation of cis-but-2-ene in the temperature range 14-58” C; mere traces of truns-but-2-ene, but-l-ene and n-butane were observed as long as unreacted but-2-yne remained in the system. In the reaction with deuterium, the expected product, cis-but-2-ene-2,3-d2, accounted for 99% of the total initial product yield. Reaction of but-2yne with an equimolar hydrogen-deuterium mixture led to a random deutero-cis-but-2-ene product distribution. Thus it was concluded that the “hydrogen” was dissociatively adsorbed and that the mechanism involved the consecutive addition of adsorbed hydrogen atoms to adsorbed but-2yne and but-2-enyl. Similar stereoselective hydrogenation has been observed with evaporated palladium films [ 1273, wires [ 1981 and palladium on BaSO, in alcoholic solution [ 1991. The hydrogenation of but-Zyne in the gas phase has been investigated using alumina-supported Group VIII metals, other than palladium, and over copperalumina [ 200,2011. With the exception of copper, which was 100% stereoselective for cis-but-2-ene formation, the distribution of the initial reaction products, as shown in Table 20, are more complex than was observed with palladium. Over all the metals studied, except cobalt, nickel and copper, the selectivity and stereoselectivity decreased slightly as the reaction proceeded. In addition to the products shown in Table 20, in the rhodium- and iridiumcatalysed reactions small yields (2-3%) of buta-1: 2diene were also observed. For all the catalysts, except rhodium, iridium and platinum, which were not investigated, the initial rate kinetic orders were unity in hydrogen and zero or slightly negative (Ni) in but-2-yne.
73 The reaction of but-2-yne with deuterium was studied over ruthenium, rhodium and osmium. Typical deuterobutene distributions are shown in Table 2 1 for each catalyst. In all cases, but-2-yne exchange was absent and the extent of the hydrogen exchange reaction was small. The close similarities in the deuterocis- and deutero-trans-but-2-ene distributions, together with the observation that the trans-but-2-ene:butI-ene ratios for but-2-yne hydrogenation are significantly less than the corresponding ratios in cis-but-2-ene hydroisomerisation, suggest that the trans-but-2-ene and but-2-ene are formed directly from but-2-yne rather than by the subsequent isomerisation of cis-but-2-ene. The trans-addition of hydrogen t o adsorbed but-2-yne may be envisaged as involving the interaction of the adsorbed hydrocarbon with molecular hydrogen (Sect. 4.3, p. 55). However, if this was the case, then, since the cis-but-2ene was considered to arise from reaction of but-2-yne with adsorbed hydrogen atoms, the cidtrans ratio would be expected t o be hydrogen-pressure dependent. This was not observed. Furthermore, Smith [ 2021 has commented that the results of Bream e t al. [ 2031 concerning the reaction of hydrogen with )-D-homogonene show that the reaction of molecular hydrogen with adsorbed species does not occur. An alternative mechanism for the direct formation of trans-but-2-ene from but-2-yne was proposed. In this mechanism, it was envisaged that a free radical was a transitory intermediate of high potential energy and short lifetime in the formation of but-2-enyl from adsorbed but-2-yne and dissociatively adsorbed hydrogen. The formation of but-2-ene and the more highly deuterated cis- and trans-but-2-enes (-d3 and above) were considered to arise from the hydrogenation of adsorbed buta-1: 2-diene formed by isomerisation of but2-yne. In this context, the observation of small yields of buta-l:2-diene in the rhodium and iridium-catalysed reactions is significazt. Meyer and Burwell [ 1891 have observed the reverse process of isomerisation of buta-1:3diene t o but-2-yne over palladium-alumina catalysts. Gold wires have been observed t o catalyse the partial hydrogenation of but2-yne in the temperature range 335-395°C [198]. The reaction became completely poisoned after ca. 40% hydrogenation of the but-2yne. The predominant hydrogenation product was cis-but-2-ene, although trans-but-2-ene and but-1-ene were also observed; traces of buta-l:3diene (0.2%) were also detected. Above 395"C, hydrogenolysis of the but-2-yne was observed along with hydrogenation and the poisoning of the catalyst became more extensive. The results observed in the hydrogenation of higher dialkylalkynes are in general accord with those observed with but-2-yne. Thus, for example, the selectivity and stereoselectivity observed in the liquid phase hydrogenation of pent-2-yne over carbon-supported palladium, rhodium, ruthenium and platinum and iridium-alumina [145], as shown in Table 22, show a similar pattern t o that observed in but-2-yne hydrogenation. Similarly, the References p p . 1 1 4-1 2 1
74 TABLE 2 1 Typical deuterobutene distributions observed in the reaction but-2-yne with deuterium over alumina-supported metals ~ _ _ _ _ _ _ _ _
Catalyst
Temp. ("C)
Rh
165
Ru
0s
-
Initial DZ/C4H6
Deuterobutene distribution
(a)
-do
-d,
-d?
-d3
2
But-1-ene cis-But-2-ene trans-But-2-ene
2.0 1.3 2.2
10.7 23.7 22.4
43.8 73.1 60.5
36.1 1.7 8.1
85
2.1
But-1-ene cis-But-2-ene trans-But-2-ene
0.0 1.1 2.2
5.9 18.0 18.3
40.4 79.0 71.9
38.3 1.9 6.0
100
2.0
But-1-ene cis-But-2-ene trans-But-2-ene
0.5 2.3 2.2
8.2 21.0 16.5
21.2 64.8 56.7
40.1 5.0 7.8
nickel-catalysed hydrogenation of hex-2-yne and hex-3-yne produces the cis-olefin as the principal product with a selectivity of about 0.97 [ 1911. 4.8 THE HYDROGENATION O F ALKA-1: 2-DIENES
Studies of the hydrogenation of propadiene have been reported for all of the Group VIII metals either when supported on pumice [ 170,2041, or silica [155] or alumina [94], or as metal powders [205,206]. The results closely resemble those observed in alkyne hydrogenation. Thus, with all the Group VIII metals, the pressure-fall against time curves for reactions with initial hydrogen: propadiene ratios of two or greater exhibit two distinct stages; with iron, cobalt apd nickel, the rate of the second stage is less than that of the first stage, whereas with the noble Group VIII metals, the onset of the second stage is accompanied by an increase in rate [ 155, 1701. Over all catalysts, the initial rate orders of reaction are zero in propadiene and approximately unity in hydrogen. The selectivity values observed in the first stage of the reaction, summarked in Table 23, show similar trends from metal to metal t o those observed in alkyne hydrogenation. Again, palladium is almost completely TABLE 22 Selectivity and stereoselectivity observed in the liquid phase hydrogenation of pent-2yne [ 1451 _
Catalyst: Selectivity: cis-Pent-2-ene (%):
Pd-C 0.99 (8) 97.8
Rh-C 0.81 95.6
~
-
Pt-C 0.90 92.5
- ~ _ - - _ _ _ Ru< 0.86 91.7
I~A1203 0.60 91.5
75
5.1 0.2 3.2
2.1 1.8
0.2 0.0 0.1
0.0 0.0 0.0
0.0 0.0 0.0
2.41 1.78 1.94
0.0 0.0 0.0
2.68 1.82 1.87
2.2 0.4 1.2
3.27 2.00 2.41
0.0
10.2 0.0 1.6
5.2
0.0
0.0 0.0
0.0
0.0
0.0 0.0 0.0
12.1 3.0 6.2
8.2 2.4 5.8
4.7 0.6 2.0
2.8 0.5 1.6
selective for olefin formation, whilst iridium is characterised by a very low selectivity. Over each metal the selectivity was observed t o decrease with increasing initial hydrogen pressure and decrease in temperature. Propadiene polymerisation during hydrogenation has been observed t o occur, particularly with iron, cobalt and nickel. Over these metals, up to approximately 25% of the propadiene has been observed to polymerise [ 2041, although the chemical identity of the polymeric products was not established. Detailed studies of the mechanism of propadiene hydrogenation using deuterium as tracer have not been reported. By analogy with the mechanisms proposed for propyne and acetylene hydrogenation (see Sect. 4.7), propadiene hydrogenation has been postulated as occurring by the mechanism shown in Fig. 31. The hydropolymerisation of propadiene has been explained by postulating that the adsorbed propenyl half-hydrogenated
Fig. 31. Reaction scheme for the hydrogenation of propadiene. R e f e r e n c e s p p . 1 1 4-1 2 1
76 TABLE 23 Selectivities (S) observed in the hydrogenation of propadiene [94,155,170,205] Catalyst
S
Fe-powder
0.93
42
Ru-Al203
0.84
51
Os-Al203
0.73
130
Co-powder
0.87
40
Rh-Al203 II-A~~O~
0.92 0.36
61 28
Temp. ("C)
Catalyst
S
Temp. ("C)
Ni-powder Ni-pumice
0.97 0.93
105
Pd-Al203 Pd-pumice
0.99 1.00
19 116
Pt-Al203 Pt-pumice Pt-SiOz
0.89 0.80 0.80
79 89 100
73
state can exist in both a normal and a free radical form [ 2061 HzC=C--CH3 6 HzC-C--CH3 I 1 I 1
**
* *
In the platinum-catalysed reaction, it has been observed [ 1551 that. the effect of increasing the hydrogenlpropadiene reactant ratio was t o increase the yield of propane without affecting the yield of propene. This has been interpreted as showing that propane may be formed directly from adsorbed propadiene by a mechanism which does not involve adsorbed propene as an intermediate. However, no conclusions were drawn regarding the nature of the surface intermediates involved in such an interconversion. For a series of silica-supported metals, the specific activity, calculated as the rate coefficient at 100°C referred t o unit area of metal, has been found [155] t o decrease in the order Pd > Rh > Pt > Ni > Co > Ru > I r > > Fe, 0 s = 0. For pumice-supported metals, the corresponding
(C)
Fig. 32. Conformations of adsorbed buta-l:2diene after Grant et al. [ 2071
77 sequence is [204] Pt > Pd > Rh > Ir > Ni > Co > Fe > Ru > 0 s . The hydrogenation of buta-1: 2-diene appears to have received relatively little attention. Over palladium-alumina a t room temperature, the products of the gas phase hydrogenation were cis-but-2-ene, 52%; but-1-ene, 40%; trans-but-2-ene, 7% and n-butane, 1%[ 1891. Some isomerisation of the buta-1: 2diene t o but-2-yne (10%) together with traces of but-1-yne and buta-1: 3diene was also observed. A similar butene distribution (namely, cis-but-2-ene 52%, but-1-ene 45% and trans-but-2-ene 3%)was observed in the liquid phase hydrogenation over palladium [ 1861. In the gas phase deuteration of buta-l:2-diene over palladiumalumina, Meyer and Bunvell [189] showed that the deutero-but-1-ene and deutero-cis-but-2-ene distributions contained all isotopic isomers up t o - d 5 ; both distributions showed a pronounced maximum a t - d z . At 8 ” C , the surface deuterium/hydrogen atom ratio was calculated t o be 15.6 : 1, and it was concluded that 95% of the products were formed by simple cis-1,2 or 2,3-addition of two “hydrogen” atoms t o adsorbed buta-1: 2diene. This conclusion was substantiated by NMR analysis of the but-lene-d, and cis-but-2-ene-d2,which were found t o be H3C-€HD-CD=CHz and H3C - C H = CD-CH2D, respectively. The marked preference for the formation of the thermodynamically less stable cis-isomer in the but-2ene yield was attributed t o the operation of steric factors in the adsorption of buta-1: 2diene. The formation of trans-but-2-ene by the cis- addition of two hydrogen atoms required the formation of adsorbed buta-l:2diene with the methyl group directed towards the surface. On steric grounds, such a species was considered t o be unlikely compared with the other possible conformations shown in Fig. 32. The virtual absence of deuterium in the product but-2-yne was taken as an indication that the isomerisation of buta-1: 2diene occurred by an intramolecular hydrogen transfer mechanism, rather than by the loss of a hydrogen atom from adsorbed but-2-enyl. Grant et al. [207] have recently reported a detailed study of buta-l:2diene hydrogenation and deuteration over various nickel catalysts. Two types of catalyst behaviour, designated types A and B, were observed depending upon the method of preparation of the catalyst. The chief difference between the two types of catalyst was the distribution of the isomeric n-butenes observed in the first stage of the reaction. Over each type of catalyst, approximately half of the buta-1: 2-diene was converted to n-butenes, the other half appearing as C8 and higher hydrocarbons, although detailed chemical analysis of the polymeric products was not performed. The distribution of C4 -products was virtually independent of the extent of hydrogenation and of the reaction conditions; typical distributions after 20% hydrogenation are shown in Table 24. A slow isomerisation of buta-1: 2diene t o buta-1: 3-diene was also observed. R e f e r e n c e s P P . 1 1 4 - 1 21
78 TABLE 24 Distribution of butenes observed in the nickel-catalysed hydrogenation of b u t a - l : 2 diene [ 207 ] Catalyst Temp. Butene distribution (%) -. ("C) But-1-ene trans-ButCis-But2-ene 2-ene
N i- p ow d e r
75-110
38
5
57
2
62
26
44
(Type A )
Ni-AI203 (Type B)
4 7-7 5
30
TABLE 2 5 Distribution of deuterium in t h e products of the reaction of buta-1 :Z-diene with deuterium over nickel catalysts [ 2 0 7 ] Catalyst
Hydrocarbon
X
1
2
(%)
0.9 0.0 3.1
26.9 0.0 26.4
60.5 82.5 58.8
But-1-ene
obs. C4H8-xDx (%) Nx calc. C4H8,D, (%)
4.2 0.0 1.3
17.2 0.0 18.1
60.6 77.8 62.6
trans-But-2-ene
obs. C4H8-xDx (%) calc. C4H8-xDx (%)
3.9 0.0 1.9
19.4 0.0 20.0
54.0 70.4 55.4
obs. C4H6-xDJ (%) obs. C4H6-xDx (%)
71.5 98.5
17.0 0.5
1.8 0.2
1.9 0.0 0.5
11.9 0.0 12.3
81.3 94.3 82.3
calc. C4H8-xDx (%)
1.6 0.0 0.3
10.6 0.0 10.4
80.3 91.5 81.7
obs. C4Hs-.vDx Nx calc. C4H8-.xDx (%)
2.3 0.0 0.6
13.1 0.0 13.5
76.8 91.1 78.4
obs. c4H6-.yDx ( y o ) obs. C4H6-xDx (%)
63.6 99.3
24.3
2.9 0.1
0 __ C4H8-xDx (%) .-
Ni-powder (Type A) a t 75°C ('D2/'CqHg)O = 3.0-5.5
cis-But-2-ene
obs.
N X
calc. C4H8,Dx
N X
cis-But-2-ene
obs. C,Hs-XD,
(%)
N X
calc. C ~ H S - ~ (%) D~ But-1-ene
obs. C4H8,Dx
(%)
N X
trans-But-2-ene
a
D.N. = mean number of D atoms per molecule.
0.4
79 In order to obtain detailed information about the reaction mechanism, the results obtained from the reaction of buta-l:2-diene with deuterium were used t o calculate N-profiles and hence theoretical deuterobutene distributions as described in Sect. 4.4.The distributions of deuterium in the n-butenes, buta-1: 3diene and buta-1: 2diene together with the butene N-profiles and calculated deuterobutene distributions are shown in Table 25. In developing the mechanism, it was assumed that the adsorbed state of buta-1: 2diene active in hydrogenation was the di-7r-adsorbed species C in Fig. 32. The close similarity in the N-profiles for each butene formed in a given reaction suggests that each was formed as a primary product and the general mechanism shown in Fig. 33 was proposed. The difference in behaviour of the type A and B catalysts was explained by proposing that, a t the type A surface, a-bonded and .rr-olefinic species are of importance as
~ ~_ _ _ .
3
4
5
6
7
7.4 11.1 7.4
2.2 2.7 2.2
1.1 1.8 1.1
0.5 1.0 0.5
0.3 0.0 0.3
8 0.2 0.9 0.2
9.7 10.2 9.7
5.5 8.0 5.5
1.6 2.2 1.6
0.7 0.8 0.7
0.4 0.7 0.4
0.1 0.3 0.1
2.05
11.6 11.8 11.6
5.0 5.3 5.0
3.1 3.1
1.5 1.9 1.5
1.o 0.7 1.0
0.5 2.1 0.5
2.18
4.8
2.7 0.1
6.0 0.6
0.7 0.1
0.2 0.0
3.5 4.0 3.5
0.8 1.0
0.1 0.1 0.1
0.1 0.1 0.1
0.1 0.2 0.1
1.92
0.8
0.3 0.3 0.3
5.3 5.8 5.3
1.6 1.9 1.6
0.3 0.4 0.3
0.2 0.2 0.2
0.1 0.1 0.1
0.1 0.1 0.1
1.97
4.a 5.6 1.8
1.4 1.7 1.4
0.6 0.8 0.6
0.3 0.3 0.3
0.2 0.1 0.2
0.2 0.4 0.2
2.4 0.0
5.8 0.2
0.7 0.0
0.2 0.0
-~ ~
_
_
~-
References p p . 1 1 4-1 21
1.91 81/19
87/13 2.10 84/16
2.21 0.58 0.05 9317 1.94 9416 2.00 1.95 9218 1.98 0.65 0.01
~
80
cis-But-2-ene
But-1-ene
trans-But-2-ene
Fig. 33. Reaction scheme for the hydrogenation of buta-l:2-diene [207].
intermediates, whereas at the type B surface, the formation of 7r-allylic intermediates is favoured. The detailed mechanisms are shown in Fig. 34. From the deutero-buta-1: 3diene distributions, it was concluded that the isomerisation occurred predominantly by intramolecular hydrogen atom transfer as shown in Fig. 35. The subsidiary maximum in the distribution
c-2-b
1-b TYPE A
1-b
t-2-b
TYPE B
Fig. 34. Mechanisms for buta-l:Z-diene hydrogenation operative over Types A and B nickel catalysts [207]. (Numbers refer to species in Fig. 33.)
81
* Fig. 35. Reaction scheme for the isomerisation of buta-l:2-diene.
at C4HZD4was not accompanied by any maximum at C4HZD6in the n-butene distributions and it was concluded that different sites were involved in buta-1: 2diene exchange and isomerisation on the one hand, and buta-l:2diene hydrogenation on the other. It was also concluded that direct dimerisation of buta-1: 2diene was responsible for the polymerisation reaction 2 H2C=C=CHCH3 2 HzC=&CHCH3 + polymers I I I I +
* *
*
*
4.9 THE HYDROGENATION OF CONJUGATED ALKADIENES
The gas phase hydrogenation of buta-1: 3-diene catalysed by aluminasupported Group VIII metals and metal powders and wires has been TABLE 26 Kinetics and activation energies observed in t h e hydrogenation of buta-l:3-diene over various metals Rate = k P"H~P g 4 ~ at 6 T"C Catalyst
Temp.
x
Y
? 1.5 1.0 1.0 1.0 1.0 1.0 1.0 -1.0
? -0.5 0.0 0.0 -0.2 0.0 0.0 0.1 0
1.7
-0.7 ? 0 0.0 0.0 -0.5 0.0 0.0
("C)
Fe-AI2O3 Co-A12 O3 Co-Al203 Co-powder Ni-Al203 Cu-Al2 O3 Ru-AI203 Rh-AI 2 03 Rh-wire Pd-AI2 O3 Pd-film Pd-wire Os-A12 O3 I r A 1 2O3 Pb-AI2 03 Au-AI203 Au-Boehmite
148 111 106 87 105 14 0 ? 16
?
?
70 0 16 200 200
References pp. I 1 4-1 21
-1.0 1.0 0.8 1.3 1.0 1.0
E , (kJ mol-I) at (TOC)
Ref.
46.5 (200-260) 51.1 ( 94-153) 36.0 (107-184) ? 62.8 ( 84-110) 54.8 ( 82-130) 52.3( 0-48) 46.5 ( 16-82) 49.8 (100-198) 54.8 (175-231) 69.9 ( 20-42) 24.7 ( 38-206) 103.8 (134-225) 42.7 ( 24-70) 18.8 ( 0-62) 81.6( 0-152) 36.5 (200-260) 36.5 (170-260)
166 166 166 166 166 166 208 208 209 208 209 209 208 208 208 2 11 211
TABLE 27 Distributions of butenes observed in the hydrogenation of buta-l:3-diene over various metal crystals Catalyst
Hz/CzH6
ma (torr)
Temp. ("C)
Butene composition (%) 1-
trans-2-
cis-2-
-
trans-2-Butene
Sb
Ref.
cis-2-Butene
2.0
18
198 259
23 27
45 41
32 32
1.41 1.28
0.980 0.984
166
4.0
13
94 152
53 50
32 33
15 17
2.13 1.94
0.968 0.980
166
2.0
14
75 130
28 22
64 70
8 8
8.00 8.75
1.000 1.000
166
1-7 2.0
17 15
77 110
49 55
35 30
16 15
2.19 2.00
0.988 0.998
166
1.0
10
100
27
63
10
6.30
2.0
10
60 130
87 83
6 8
7 9
0.86 0.89
1.000 1.000
166
3.1
12
0 49
69 61
19 23
12 16
1.58 1.43
0.736 0,835
208
3.1
20
16 82
51 48
32 32
17 20
1.88 1.60
0.743 0.906
208
>0.99
207
Rh-wire
3.0
24
215
37.9
36.5
25.6
1.43
0.987
209
Pd-AIz
0 3
3.1
25
0 43
64.4 59.4
33.2 36.8
2.4 3.8
13.83 9.68
1.000 1.000
208
Pd-wire
3.0
30
21 5
46.8
40.1
13.1
3.06
1.000
209
OS-AI 2 0 3
5.0
15
24 70
65 57
19 21
16 22
1.19 0.95
0.431 0.630
208
IrA1203
3.1
16
24 75
59 32
19 34
22 34
0.86 1.00
0.251 0.384
208
Ir-wire
3.0
10
190
57.9
21.7
20.4
1.06
0.869
209
3.1
25
0 152
72 38
18 36
10 26
1.80 1.78
0.501 0.920
208
Pt-wire
3.0
23
260
42
30
28
1.07
0.985
209
Au-Boehmite
2.0
15
200 260
50.5 51.9
17.6 20.6
31.9 27.5
0.55 0.75
1.000 1.000
211
Au-Al203
2.0
15
170 260
58.4 59.7
12.8 15.4
27.7 24.9
0.46 0.62
1.000 1.000
2 11
Pt-AI
a
2 03
(PCqHJO = 50.0 torr throughout. = PCqHg/(PCqH~
+
‘C4H
10).
84
--
80
$
0' 40
.-
+ ._ $
Q
g 80o Q,
C Q,
+ 3 m
40
0 0
50
100 0 Pressure fall ( t o r r )
100
Fig. 36. Variation of the yields of but-1-ene ( O ) , trans-but-2ene (a) and cis-but-Z-ene ( 0 ) with pressure fall in the hydrogenation of buta-l:3-diene over (a) Rh at 18OC, ( b ) Pd at O°C, (c) Ir at -2OOC and ( d ) Pt at 104OC. P c ~ H = 50, ~ P H =~ 155 torr [ 2 0 8 ] .
extensively studied by Wells et al. [ 166,167,208-2101. Meyer and Burwell [ 1891 have also studied the gas phase deuteration of buta-l: 3diene over alumina-supported palladium, whilst Buchanan and Webb [ 2111 have reported a similar study using y-alumina- and boehmite-supported gold catalysts. In the liquid phase, the hydrogenation of buta-1: 3diene in ethanolic solution has been investigated using nickel, palladium and platinum catalysts [ 2121. For the catalysed reaction in the gas phase, the initial rate orders of TABLE 28 Buta-l:3-diene and hydrogen exchange over alumina-supported metals [ 166,167 ] Catalyst
Temp. ("C)
D2/C4H6
co
101
2
Ni cu
68 120
2 2
Rh Pd Pt
100 17 20
1 1 1.5
a
Reaction
a
Deuterobuta-l:3diene distribution (%)
(%)
10 90 10 10 80 16 15 15
Based upon pressure of buta-l:3-diene.
C4H6
C4HsD
C4H4D2
90.0 30.1 98.8 83.4 11.0 99.1 98.3 99.6
7.9 36.0 0.9 14.4 27.8 0.9 0.9 0.4
0.6 19.4 0.2 2.0 34.9 0.0 0.4 0.0
85 reaction, selectivities for olefin formation and the shapes of the pressure fall against time curves were similar t o those observed in alkyne and alka1:2diene hydrogenation for a particular metal. The kinetics and activation energies are summarised in Table 26. One of the characteristic features of the hydrogenation of buta-1: 3diene is that, over all metals, all three isomeric n-butenes are observed as initial products. In general, but-1-ene is the major product, although both the but-lene/but-2-ene ratio and the relative yields of trans- and cisbut-2-ene vary widely from metal t o metal. Figure 36 shows some typical plots of the variation of the butene distribution as a function of the extent of hydrogenation; the initial butene distributions observed with the Group VIII metals, copper and gold, together with the selectivity values are shown in Table 27. The selectivity was observed t o increase with increasing temperature and t o decrease with increasing initial hydrogen pressure over all catalysts. In the nickel- and cobalt-catalysed reactions [ 166,2071 it was observed that the butene distribution depended upon the temperature of reduction of the catalyst. For both powders and alumina-supported catalysts prepared by reduction of the oxides, reduction a t temperatures below ca. 330"C gave catalysts which exhibited so-called Type 'A behaviour where but-2-ene was the major product and the truns-but-2-ene/cis-but-2-ene ratio was around unity. Reduction above 360°C (Ni) or 440°C (Co) yielded catalysts which gave trans-but-2-ene as the major product (Type B behaviour). It is of interest to note that the yield of cis-but-Zene was not significantly dependent upon the catalyst reduction temperature with either metal. The formation of but-2-ene as an initial product raises some important questions regarding the mechanism of hydrogenation. In principle, two
'AH3D3 0.4 1.9 0.1 0.2 20.5
Trace 0.2 0.0
x in
y in
C4H6-A
HyD2y 0.008 0.078 0.002
C4H2D4
C4HD5
C4D6
0.2 4.6 0.0
0.0 1.6 0.0
0.0 0.0
0.11 1.28 0.02
0.0
0.0
0.0
0.19
0.009
5.2
0.1
0.1
0.067
Trace 0.1 0.0
0.0 0.1
0.0 Trace
1.83 0.01
0.0
0.0
References p p . 1 1 4-1 21
0.4
0.03 0.004
routes leading t o the formation of but-2-ene exist; (a) direct 1 : 4-addition of two hydrogen atoms t o adsorbed buta-1 :3-diene, and ( b ) 1 : 2-addition of two hydrogen atoms t o yield but-l-ene which, before desorption, undergoes isomerisation. In order t o differentiate clearly between these two possibilities, it is necessary to consider the results obtained from the use of deuterium as a tracer. The reaction of buta-1: 3diene with deuterium has been studied over several metals [ 166,167,1891 and the deuterated product distributions used t o obtain N-profiles as described in Sect. 4.3. Over all the metals studied, some buta-1: 3-diene-deuterium exchange was observed as shown in Table 28. This exchange was particularly marked with cobalt and copper, where the extent of exchange increased noticeably as the reaction proceeded; with the other metals, the extent of exchange was small. Hydrogen exchange was also observed with cobalt, nickel and copper, as shown in Table 28, but was absent in the rhodium-, palladium- and platinum-catalysed reactions indicating that, with these metals, the surface coverages of hydrogen and deuterium atoms were very low under the reaction conditions. The observed deuterobutene distributions together with the calculated N-profiles, for those metals where unique N-profiles could be obtained, the surface D/H ratio and the calculated deuterobutene distributions are shown in Table 29. One of the major features of these results is that, over all the metals studied, the trans- and cis-but-2-ene profiles show pronounced maxima a t - N 2 .This clearly shows that the predominant route t o the formation of but-2-ene was direct 1:4-addition of two hydrogen atoms t o adsorbed buta-1 :3-diene. 1 : 2-Addition of hydrogen t o yield but-l-eneN 2 followed by isomerisation would have led t o a zero value for but-2ene-N2 and a maximum a t but-2-ene-N, or higher depending upon the number of butene-butyl interconversions before desorption of the but-2ene. The detailed interpretation of the N-profiles has been discussed fully by Wells and co-workers [166,167] who have proposed the two mechanisms shown in Fig. 37. Mechanism A is a generalised mechanism which was proposed for those metals where the trans-but-2-ene : cis-but-2-ene ratio was around unity. This mechanism contains a variety of reversible steps which permit the conformational interconversion of the diadsorbed buta-1: 3-diene. Consequently, the trans : cis ratio will depend upon the relative rates of these reversible steps and the ratio may be much lower than would be expected if the relative surface concentrations of anti- and syndiadsorbed buta-1:3diene, species I and 111, respectively, in Fig. 37, were similar to the relative amounts of anti- and syn-buta-l:3-diene in the gas phase. It was also suggested that the relative importance of the various steps in mechanism A may be different for different metals. Thus, for example, the type A behaviour of nickel and cobalt catalysts, as deduced from the butene distributions and a detailed examination of the butene N-profiles [ 1661, was
87
Mechanism A ~
Buta-1,3-diene (g)
ll
+HI 1-H
+H]I-H
L trans-but-2-ene (g)
1
cis-but-2-ene (g)
Mechanism B Buta-1,3-diene (g)
CH2%H F)CH
CH2=$H L'CH-CH*
*,:I CH2 [ I I I A ]
[IAI
+H/I--H
1
trans-but-2-ene (g)
But-1-ene(g)
+H] 1-H
cis-but-2ene (g)
Fig. 37. Reaction schemes for the hydrogenation of buta-l:3-diene. References p p . 1 14-1 2 1
J
TABLE 29 Deuterobutene distributions observed in the reaction of buta-l:3-diene with deuterium ow alumina-supported metals [ 166,1671 Catalyst
co
Temp. ("C) 101
DZ -
Reac, tion
C4H6
(X)
2
74
co
Ni
68
2
54
Hydrocarbon
X
0 ____
But-1-ene
C ~ H S - ~ (obs.) D~ Nx C4H8-xDx (calc.)
8.4 0.0 6.5
trans-But-2-ene
C~HS-~D (obs.) ~ NX C4Hs7Dx (calc.)
7.8 0.0 8.2
cis-But-2-ene
C4Hs,Dx NX C4Ha7Dx
(calc.)
7.6 0.0 7.3
But-1-ene
C4HS-xD.v (obs.) NX C4H8-xDx (calc.)
9.6 0.0 7.4
trans-But-2-ene
C4Hs,Dx NX C4Hs,Dx
(calc.)
9.0 0.0 7.6
C4H8-xDx (calc.)
10.2 0.0 10.1 30.6 26.6 27.4
(obs.)
(obs.)
cis-But-2-ene
cu
120
2
80
But-1-ene trans-But-2-ene cis-But-2-ene
C4Hs7Dx (obs.) C4H8--sDx (obs.) C4Hs,Dx (obs.)
Rh
100
1
80
But-1-ene
C4Hs,Dx NX C4Hs7DX
(obs.)
trans-But-2-ene
C4Hs7Dx NX C4H8,Dx
(obs.)
cis-But-2-ene
C4Hs,DX NX C4Hs7DX
(obs.) (calc.)
7.5 0.0 5.4
But-1-ene
C ~ H S - ~ (obs.) D~ NX C4HS7Dx (calc.) C4Hs7Dx (obs.) NX C4Hs7Dx (calc.)
13.5 0.0 9.2 9.2 0.0 8.4
C4Hs,DX (obs.) NX C~HS-~D (calc.) ~
13.9 0.0 9.2
trans-But-2-ene
cis-But-2-ene
(calc.)
(calc.)
6.9 0.0 5.2 6.9 0.0 5.5
1
2 __
~
26.8 0.0 28.9
3
4
5
18.7 35.5 18.6
5.0 17.7 5.0
~-
40.4 41.8 40.1
0.7 4.6 0.7
1.88
0.0
0.0 0.0 0.0 0.0 0.0 0.0
1.69
1.80
<0.1 0.5 0.1
0.0
0.0
1.89
35.9 0.0 34.4
40.5 60.5 41.6
12.5 26.0 12.5
2.9 11.0 '1.9
0.4 2.0 0.4
<0.1 0.5 0.1
0.0 0.0 0.0
31.1 0.0 31.6
41.0 57.4 40.8
15.6 30.1 15.6
4.0 14.7 4.0
0.6 2.8 0.6
0.1 0.1
0.0 0.0 0.0
0.0 0.0 0.0
27.8 0.0 30.7
37.2 55.0 37.0
13.1 15.0 13.3
7.4 11.0 7.4
3.2 10.0 3.1
1.5 5.0 0.9
0.1 1.0 0.2
0.0 1.0 0.0
31.4 0.0 32.1
38.7 63.0 38.9
11.1 15.5 11.1
5.7 9.0 5.8
3.0 2.5 3.0
1.1 10.0 0.9
0.0 0.0 0.0
0.0 0.0 0.0
1.88
35.3 0.0 36.4
38.8 65.7 37.5
9.5 18.8 9.6
3.9 5.3 3.9
1.8 4.7 1.8
0.5 4.5 0.5
0.1 1.0 0.1
0.0 0.0 0.0
1.71
29.8 28.8 29.4
21.6 23.8 23.4
12.0 13.4 13.3
4.8 5.8 5.3
1.1 1.5 1.3
0.1 0.1 0.1
0.0 0.0 0.0
0.0 0.0 0.0
1.35 1.48 1.34
26.1 0.0 28.1
12.7 59.9 42.6
13.3 19.1 13.3
5.8 10.9 5.8
2.8 2.0 2.8
1.5 5.9 1.5
0.5 0.0 0.5
0.2 2.2 0.2
2.03
27.8 0.0 29.5
43.7 64.6 43.4
11.6 16.2 11.6
5.2 10.0 5.2
2.6 1.1 2.6
1.5 5.9 1.5
0.5 0.0 0.5
0.2 2.2 0.2
1.98
27.9 0.0 29.2
42.4 63.4 43.2
11.9 17.5 11.9
5.1 9.3 5.1
2.7 1.9 2.7
1.6 4.9 1.6
0.7 0.0 0.7
0.3 3.0 0.2
1.98
29.5 0.0 32.4
34.4 58.7 34.4
12.6 1.0 14.0
8.2 30.5 7.7
1.6 1.7 1.8
0.5 8.1 0.5
0.0 0.0 0.0
0.0 0.0 0.0
1.80
27.2 0.0 28.7
34.3 53.5 32.3
14.6 1.5 14.6
9.1 27.0 9.5
3.9 0.5 3.9
1.4 15.0 1.4
0.2
0.0 1.5 0.0
2.05
31.2 0.0 33.3
34.0 50.9 36.4
12.7 8.5 12.8
4.8 8.5 4.8
2.4 0.0 2.5
1.0 10.0 0.9
0.1 0.5 0.1
0.0
1.75
R e f e r e n c e s p p . 114-121
1.0
1.0 0.2
1.69
1.80
66 : 34
66 : 34
1.98 1.97
1.87
1.74
67 : 33
67 : 3 3
67 : 33
72 : 28 2.04 72 : 28 1.99 72 : 28 2.01 62 : 38
1.86 62 : 38 2.04
1.0 0.0
66 : 34
62 : 38 1.82
TABLE 29 (continued) Catalyst
Temp. ("C)
D2/ C4H6
Reaction
Hydrocarbon
X
0
Pt
20
1.5
66
But-1-ene
tmns-But-2-ene
C4H8-xDx (calc.)
5.4
C4Ha,Dx
7.0 0.0 5.5
(obs.)
N X
C4Ha-,yD, (calc.) cis-But-2-ene
C4 Ha -.yD.y
( obs. I
N... .
C4H8-.\.Ds (calc.)
a
8.8 0.0 4.9
D.N. = Mean number of D atoms per molecule. a ; b = surface D : surface H.
concluded t o be consistent with a restricted form of mechanism A shown in Fig, 38. Mechanism B was originally proposed for the palladium-catalysed reaction [208],which is characterised by a high trans : cis ratio. In this mechanism, conformational interconversion of adsorbed species does not occur and the trans : cis ratio is a direct reflection of the relative surface concentrations of the anti- and syn-diadsorbed buta-l:3-diene, which are dependent upon the nature of the sites available a t the surface and upon the relative stabilities of the two conformers. In addition t o palladium, mechanism B has been proposed, along with a contribution from mechanism A , TABLE 30 Deuterobutane distributions observed in the reaction of buta-l:3-diene with deuterium over rhodium and platinum [ 1 6 7 ] Catalyst
D2/C4H6
Rh (100°C)
1.0
X ~-
C4H(10-x)Dx (Obs.) C4H(Io-xfDx (calc.) N X
Pt (20°C)
1.5
-_ a
C4(H10--xIDx (obs.1 C~H(IO-,)D, (ca1c.I Nx -
___
D.N. = Mean number of D atoms per molecule. a . b = surface D : surface H.
2
3
0
1
7.1 8.5 0.0
5.7 7.5 0.0
9.3 8.2 0.0
18.9 16.7 0.0
1.9
2.4 2.0 0.0
10.9 10.3 0.0
21.3 25.1
0.2
0.0 ~
~
0.0 ~-
91
1
2
3
4
28.2 0.0 29.4
42.9 64.1 43.6
12.4 16.0 12.4
5.8 11.0 5.8
2.3 4.3 2.3
28.4 0.0 29.3
42.4 65.4 42.8
11.3 12.3 11.3
6.3 11.9 6.3
4.0
24.9 0.0 26.6
38.0 57.1 40.2
13.3 16.2 13.3
7.6 9.7 7.6
_
_
_
5
6
7
8
0.9 2.6 0.9
0.2 2.0 0.2
0.0 0.0 0.0
0.3 0.0 0.3
0.1 1.4 0.1
1.98
3.0
1.4 7.2 1.4
4.2 7.3 4.2
2.0 5.9 2.0
0.9 0.0 0.9
0.3 3.8 0.3
2.12
1.8
~
1.93 72 : 28 1.95 1 2 : 28 2.00 72 : 28 2.19
to account for the type B behaviour of cobalt and nickel. The formation of but-2-ene by direct 1:4-addition was confirmed by NMR analysis of the trans-but-2-ene. This showed that 90% (Pd) or 97% (Co) of the deuterium was located on the terminal carbon atoms. In the rhodium- and platinum-catalysed reactions [ 1671, it is of significant interest that the N-profiles calculated from the distributions of deuterium in the n-butane (Table 30) appear t o bear no clear relationship to the N-profiles of the n-butenes formed simultaneously. This observation has been interpreted as indicating that either the butene which undergoes further hydrogenation never desorbs as butene, or that the sites responsi-
D.N. 5
4
6
8
7
a
a: bb
10
9
-
23.0 23.6 30.0
13.1 13.6 15.5
8.0 8.0 14.0
5.0 5.0 2.5
4.5 3.0 7.5
2.3 3.0 0.0
2.3 1.4 10.0
4.08 3.96
29.1 29.5 40.0
16.4 16.9 20.0
9.0 9.6 24.0
4.2 4.5 0.0
1.8 1.8 15.7
0.1 0.1 0.0
trace trace trace
3.94 4.05
References p p . 1 1 4-1 21
8 2 : 18 (80%) 1 4 : 8 6 (20%)
76 : 24
92
-
Buta-1,3-diene (g)
Buta-l,3-diene (g)
11 (1)-
-It
(IV)(XIII)
(11)
lt
(V)(VI)
/
(111)
It
\
(XIII)(7
Mechanism B
Mechanism A
Fig. 38. Predominant routes t o t h e formation of butenes over Type A (mechanism A ) and Type B (mechanism A + mechanism B) nickel and cobalt catalysts after Phillipson e t al. [166]. (Figures refer t o species in Fig. 37 except (XIII) which is S H 2 - C H 2 C H 4 H 2 .)
I *
ble for n-butane formation are distinct from those which catalyse the formation and desorption of butene. This latter conclusion is also supported by the observation that the surface D/H ratios for n-butane formation are
TABLE 31 Initial distributions of products observed in the hydrogenation of penta-l:a-diene over various alumina-supported metals [ 2151 ______ . -HZ I C S H8 Temp. Penta-l:3-diene comp. (%) Catalyst Reactant isomer (“C) frans-CSH8 cis-C5H8 ~
trans cis
2 2
81 81
98 11
2 88
trans
2 2
98 98
81 71
19 29
4 2
70 70
99 16
1 84
cis
2 2
116 116
83 72
17 28
trans cis
2 2
20 20
0 87
cis
2
94
100 13 -
cis
trans cis (95%) -
trans (5%) trans
~
~.
-
93 not the same as those operative for butene formation. Wells and Wilson [215] have reported a study of the gas phase hydrogenation of cis- and trans-penta-1: 3-diene catalysed by alumina-supported cobalt, nickel, copper, palladium and platinum. The general features of the reaction resembled closely those of buta-l:3-diene hydrogenation and it was concluded that the introduction of the substituent methyl group did not alter the basic features of the reaction. Reactant isomerisation was observed t o be rapid over copper and one form of cobalt, but was relatively slow over the other catalysts studied. Some typical product distributions are shown in Table 31, from which it can be seen that all three isomeric n-pentenes were formed over each catalyst, although the pentene distribution varied considerably from metal t o metal. From the pentene distributions, the relative amounts of 1 : 2-, 3 : 4and 1 : 4-addition of hydrogen were calculated (Table 31). The results show that over all catalysts the 3 : 4-addition process is more important than the 1 : 2-addition process and that the fraction of product formed by 1 : 4-addition in penta-1:3-diene hydrogenation and in buta-l:3-diene hydrogenation is similar, confirming the operation of similar mechanisms in the two cases. A number of reports of the liquid phase hydrogenation of substituted buta-1: 3dienes are extant. Although in these studies product distributions have been quoted, precise mechanisms have not been established. The situation is confused since the product distributions appear t o be dependent upon the solvent used and also upon the nature of the catalyst sup-
Relative % addition process
Pentene composition (%) -.
1-C5HI0
trans-C5Hlo
c ~ s -Hi C ~0
1:2
3:4
1: 4
31 39
55 36
14 25
19 10
31 40
50 50
5 6
90 88
6 5
25
5
70
29 24
60 56
11 20
20 18
30 23
50 59
69 70
19 18
l12 2
14
69
17
10 38
52
a
44
18
10 16
40 38
46
52
19
29
16
52
32
References P P . 1 14-1 21
1
1
50
94 port. In general, it would appear that the introduction of a substituent at an olefinic carbon atom renders that carbon atom less vulnerable to hydrogen atom attack. Thus in the hydrogenation of 2-methylbuta-l:3diene over Raney nickel and palladium and platinum blacks, more 2-methylbut-1-ene than 3-methylbut-1-ene is formed [ 2131. Similarly, pent-2ene is the major product in the hydrogenation of penta-l:3-diene catalysed by carbon-supported rhodium, palladium and platinum, whilst the 1 : 4-addition products, 2,5-dimethylhex-3-enes, are only formed in minor amounts during the hydrogenation of 2,5-dimethylhexa-2 :4-diene [214]. In these reactions, palladium and nickel showed a higher selectivity for olefin formation than platinum in agreement with results obtained for buta-1: 3-diene itself.
5. The hydrogenation of cyclic molecules The metal-catalysed hydrogenation of cyclic hydrocarbons is of considerable interest from the stereochemical viewpoint and numerous reports of the hydrogenation of both alicyclic and aromatic hydrocarbons are extant. The chemisorption and hydrogenation of aromatic compounds has been extensively reviewed by Moyes and Wells [216] and by Garnett and Sollich-Baumgartner [217,218], whilst Weitkamp [ 2191 has provided a very comprehensive review of the stereochemistry of the hydrogenation of substituted naphthalenes. In this section, attention will be confined t o the hydrogenation of alicyclic molecules. In this context, in the metal-catalysed reactions of cyclopropane with hydrogen, cyclopropane shows properties intermediate between those of an alkene and an alkane. The hydrogen-cyclopropane reaction is, therefore, classified as an hydrogenation reaction, whereas the reactions of cyclobutane, which exhibits much less alkene-like behaviour, and higher cycloalkanes with hydrogen are more accurately described as hydrogenolysis reactions. 5.1 T HE HYDROGENATION OF ALICYCLIC ALKENES
The metal-catalysed hydrogenation of alicyclic alkenes differs from that of linear alkenes in that the hydrogenated products have a definite stereochemical configuration. Accordingly, from the simple Horiuti-Polanyi mechanism involving the consecutive addition of two hydrogen atoms t o the adsorbed alkene, the hydrogenation of a 1,2-dialkylcycloalkene is expected t o yield only the cis-1:2-dialkylcycloalkane, whereas hydrogenation of a 2,3-dialkylcycloalkene may yield either cis- or trans-dialkylcycloalkane depending upon the configuration of the adsorbed alkene [ 2201. Stereochemical analysis of the products of the hydrogenation of cycloalkenes may, therefore, give valuable information regarding the conformations of adsorbed species or the stereochemistry of the hydrogena-
TABLE 32 The yields of cis-dialkylcycloalkanes observed in the hydrogenation of various cyclohexenes over palladium and platinum catalysts a t atmospheric pressure and room temperature ____~
~ ~ _ _ -~ _ _ ~_ . ~
Cyclohexene
Catalyst
cisCycloalkane
Ref.
("/.I 1,2-Dimethyl
1,3-Dimethyl
Pt Pd Pt Pd Pt Pd Pt Fd Pt
2,4-Dimethyl
Pt
2,3-Dimethyl 2-Meth ylmethylene
A9~'G-OctaIin
82 25 75 24 70 30 51 10 80 74 69
220 222 220 222 220 222 231 223 221 23 1 221
tion processes. An understanding of such processes is of obvious value in the elucidation of reaction mechanisms [ 192,2211 and for this reason, a considerable number of reports of the hydrogenation of cycloalkenes are extant. Most of the studies t o date have employed either palladium [ 222-2291 or platinum [ 220,224,226,228-235], commonly as Adams reduced platinum oxide, although nickel [ 228,236,2371, rhodium [ 238,2391, ruthenium [ 2391, iridium [ 2391, iron [ 2371 and tungsten [ 2371 have also been used. Many of these studies have been concerned with the stereochemistry of the hydrogenation of disubstituted cycloalkenes. Table 32 shows some typical results for the platinum- and palladium-catalysed hydrogenation of disubstituted cyclohexenes. Table 33 shows comparative results for the hydrogenation of 1,4-dialkylcyclohexenes over palladium, platinum and rhodium catalysts. From these results it can be seen that in the hydrogenation of 1,2dialkylcyclohexenes the expected cis-l,2-dialkylcyclohexaneis not the sole product. Similarly, in the hydrogenation of the 1,4-dialkylcyclohexenes where both the cis- and trans-cyclohexanes are expected, the transisomer being the thermodynamically more stable, the stereoselectivity varies from metal t o metal. Thus with palladium, the cisltrans ratio approaches the equilibrium composition, whereas with platinum and rhodium, the equilibrium composition is never approached. It is also instructive t o note that in the palladium-catalysed reactions, hydrogenation is accompanied by extensive alkene isomerisation [ 220-2231, whereas with rhodium and platinum, little or no isomerisation is observed [220, References p p . I 1 4-1 21
96 TABLE 33 Yields of cis-isomer in the hydrogenation of 1,4-dialkylcyclohexenes over rhodium, palladium and platinum catalysts at atmospheric pressure and 25°C in acetic acid solution Cyclohexene
1,4-Dimethyl 1 -Methyl-.lethyl 1 -Ethyl-4-methyl 1,4-Diethyl 1 -Methyl-4-i-propyl 1 -i-Propyl-4-methyl a
cis-Dialkylcyclohexane (%) over Rha
Pd
Pt
61 55 55 54 52 54
28 24 a 27 a 25 a 26 21 b
57 48 58 49 a 43 58
ref. 238. ref. 223. ref. 231.
235,2381. Nickel has been reported t o show behaviour similar t o platinum [ 2361. Further, in the reactions of cycloalkenes with deuterium, the product cycloalkanes are much more extensively exchanged over palladium than over nickel or platinum [ 2361. Such behaviour is not unexpected by comparison with the results obtained in the hydrogenation of linear alkenes (Sect. 3, p. 25). A number of mechanisms have been proposed in attempts t o explain the observed stereoselectivity in cycloalkene hydrogenation. Siege1 et al. [ 220-222,224,2321 have postulated that the stereochemical results are explicable in terms of a simple Horiuti-Polanyi mechanism involving the interconversion of associatively adsorbed alkene and monoadsorbed alkyl species (Sect. 3.3 p. 23), the different stages of the mechanism taking place with retention of configuration at the adsorbed carbon atom. Thus only cis-addition of hydrogen is allowed during a single sojourn of the molecule at the catalyst surface. The formation of truns-l,2-dialkylcycloalkane, observed in the hydrogenation of 1,2-dimethylcyclohexene [ 2201 and 1,2dimethylcyclopentene [ 2321 is explained by supposing that the 1,2dimethylcycloalkane is first isomerised to the thermodynamically less stable exo- or the 2,3-isomer, which undergoes a desorption-readsorption cycle before being hydrogenated t o a mixture of the cis- and truns-saturated product (see Fig. 39). The appearance of small, but significant, yields of 2,3-dimethylcyclopentene during the platinum-catalysed hydrogenation of 1,2dimethylcyclopentene [ 2321 and of A'q9-octalin in the hydrogenation of A9,'o-octalin over Adams platinum [144] have been cited as evidence for this mechanism. The key step in this mechanism, leading t o inversion at one of the substituted carbon atoms and hence t o the formation of the trans-cycloalkane, is the desorption and readsorption of the
97
*
H3C Hs;CH
3*
H3C HQCH2-H *
1 H3C HO CH H 3 iH3CQCH3
H *
*
H *
H
1
~
H
3*
~
C
H H~ *H
*,
~
C
~H H C
H *
Fig. 39. Reaction scheme f o r the hydrogenation of 1,2-dimethylcyclohexene by a Horiuti-Polanyi mechanism as proposed by Siege1 [ 2211.
isomerised cycloalkene. However, extrapolation from the results obtained with linear alkenes (Sect. 3.5 p. 33), where it was shown that the chances of alkene desorption is extremely small with platinum, suggests that, even though isomerisation may occur, the extent of alkene desorption is unlikely t o be sufficient to account for the observed yields of the transcycloalkanes. A similar view has been expressed by Hussey et al. [223, 234,2351, who have suggested from kinetic studies that the chemisorption of most cycloalkenes on the hydrogenation sites of platinum surfaces is irreversible, a conclusion substantiated by their observations [ 2351 that on platinum catalysts n o isomerisation of thermodynamically unstable cycloalkenes t o the more stable and less readily hydrogenated isomers occurred. These workers have proposed a mechanism in which the interconversion of adsorbed cis- and trans-cycloalkenes occurs through a stereochemically symmetrical intermediate cis-cycloalkene (ads.) + symmetrical X (ads.) =+trans-cycloalkene (ads.) although the chemical identity of this intermediate was not discussed. Once formed, the adsorbed cis- and trans-cycloalkenes were postulated as reacting by a simple Horiuti-Polanyi mechanism. A similar type of mechanism was suggested by Smith and Burwell [ 1441 who observed that, References p p . I 14-1 21
~
~
H
98
c:3=a=cG:: 1
cis-Decalin
I
*
1
cis- + trans-Decalin
Fig. 40. Reaction scheme for the hydrogenation of A9*lo-octalin[ 1441.
in the reaction of A9”O-octalin with deuterium over platinum, the cis- and truns-decalins contained, on average, three deuterium atoms per molecule. They proposed that the A9’ l o -0ctalin was dissociatively adsorbed and that the trans-decalin was formed by isomerisation t o A’ *9-octalinthrough the dissociatively adsorbed intermediate as shown in Fig. 40.A similar “dissociative mechanism” was proposed to explain the extensive deuterium exchange of cyclopentene and cyclohexene over iron films [ 2371. Although the chemical nature of the symmetrical adsorbed intermediate envisaged by Hussey et al. was not discussed, it is possible that it may be identified as a n-allylic adsorbed intermediate as originally suggested by Rooney et al. [85-871. From the exchange of cycloalkanes with deuterium, it was suggested that n-allylic species, in which the carbon atoms have a planar configuration, may react with hydrogen at both their upper and lower faces [ 871. Consequently, the interconversion of adsorbed cycloalkene and an adsorbed n-allylic complex may cause inversion of the carbon atoms and hence provide a route to the formation of cis- and transcycloalkane. Hilaire and Gault [ 2271 have concluded that n-allylic species are of importance in the palladium-catalysed deuteration of Cs -cycloalkanes, although, as expected from cycloalkane4euterium exchange studies [85], they d o not appear t o play an important role in the platinum- or nickel-catalysed reactions [ 2361. Clearly, the situation with regard t o the precise mechanism of cycloalkene hydrogenation, particularly over platinum, is still rather confused. Pecque and Maurel [228], from studies of the stereospecificity of the hydrogenation of 2,3-dimethylbicyclo[ 2,2,2] oct-2-ene (compound A) and 3-methyl-2-methylenebicyclo[ 2,2,2] octane (compound B) over Raney nickel, Adams palladium and Adams platinum catalysts, have suggested that direct truns-addition occurs with tetrasubstituted alkenes over platinum. In this study it was observed that the extent of truns-hydrogenation of compound A increased in the order Ni < Pd < Pt, whereas with compound B, which would be expected to yield more truns-saturated product than compound A, the order was Pt < Pd < Ni. Furthermore, the yield of truns-cycloalkane from hydrogenation of compound A was greater than from compound B over platinum; the reverse was found for nickel. It was concluded that the stereochemistry observed with nickel was consistent with the occurrence of a simple Horiuti-Polanyi mechanism, the transproduct arising from double-bond migration, as shown in Fig. 41,i.e. for
99
v-trans-
::;: cisFig. 4 1 . The hydrogenation of 2,3dimethylbicyclo-2,2,2-oct-2ene.
nickel ha = 0; with platinum, however, ha # 0 and direct trans-hydrogenation occurred. No deductions regarding the stereochemistry of the palladium-catalysed reaction could be made because of the rapid isomerisation and equilibration of compounds A and B. It should be noted, however, that in this study ethanol was used as solvent and, as observed by Phillipson and Burwell [ 2401 in the deuteration of cycloalkenes over platinumalumina, polar hydrogen-bonding solvents are not inert but are intimately involved in the hydrogenation mechanism. The latter conclusion is also substantiated by the studies of van Rantwijk et al. [229] who observed that the yield of trans-cycloalkane in the palladium-catalysed hydrogenation of dimethylbicyclo[ 2,2,2]oct-2-ene-2,3-dicarboxylate (Fig. 42) was dependent upon the solvent, although in the platinum-catalysed reaction, sohent participation was claimed not to be involved, the ca. 10% transproduct yield arising by direct trans-addition. The competitive hydrogenation of pairs of cycloalkenes over metal catalysts has been studied by a number or workers [ 234,235,239,2411in an attempt t o establish the mechanism of hydrogenation. The results of these studies have been discussed in detail in a recent review [239] and will not be discussed further here. COOMe
/”
e C O O M e
//
COOMe
Fig. 42. The hydrogenation of dimethylbicyclo-2,2,2-oct-2-ene-2,3-dicarboxylate. References p p . 1 1 4--121
100 5.2 THE HYDROGENATION OF CYCLOPROPANE
The metal-catalysed hydrogenation of cyclopropane has been extensively studied. Although the reaction was first reported in 1907 [ 2421, it was not until some 50 years later that the first kinetic studies were reported by Bond et al. [ 26,243-2451 who used pumice-supported nickel, rhodium, palladium, iridium and platinum, by Hayes and Taylor [246] who used K20-promoted iron catalysts, and by Benson and Kwan [247] who used nickel on silicaalumina. From these studies, it was concluded that the behaviour of cyclopropane was intermediate between that of alkenes and alkanes. With iron and nickel catalysts, the initial rate law is values of rn and n being respectively zero and -1 (Fe at 101-212°C); 0.3 and -1 (Nisilica/alumina at 75°C) and 1 and 0 (Ni-pumice at 170°C). With the pumice-supported Group VIII metals [ 261, cyclopropane orders between 0.2 and unity were found depending upon the temperature and the initial fixed hydrogen pressure. With increasing hydrogen pressure, the rate rose to a maximum value and then either remained constant or decreased. The results were interpreted in terms of a Langmuir-Hinshelwood mechanism in which hydrogen was more strongly adsorbed than cyclopropane and the slow step was the reaction between adsorbed cyclopropane with adsorbed hydrogen atoms. The activation energies observed in these and subsequent studies are summarised in Table 34. In the reaction of cyclopropane with deuterium over pumice-supported rhodium, palladium, iridium and platinum, it was observed [ 244,2451 that the deuteropropanes were extensively exchanged, consisting predominantly of propaned, with varying amounts of all other deuteropropane isomers. Only very small amounts (<2%) of cyclopropane exchange were observed. It was also observed that the deuteropropane distributions obtained from cyclopropane-deuterium and propane-deuterium exchange over the same catalysts were closely similar [ 2451, suggesting that the two reactions occurred by a similar mechanism, differing only in the initiating steps. These observations, together with the kinetic evidence and the observation that methylcyclopropane was more strongly adsorbed than cyclopropane itself [26], led to the postulate that the cyclopropane was adsorbed as a n-complex rather than as a 1,3-diadsorbed species or a dissociatively adsorbed - C 3 H 5 complex. Throughout these studies, no product other than propane was observed. However, subsequent studies by Sinfelt et al. [ 249-2511 using silica-supported Group VIII metals (Co, Ni, Cu, Ru, Os, Rh, Ir, Pd and Pt) have shown that, in addition to hydrogenation, hydrocracking to ethane and methane occurs with cobalt, nickel, ruthenium and osmium, but not with the other metals studied. From the metal surface areas determined by hydrogen and carbon monoxide chemisorption, the specific activities of
TABLE 34 Activation energies observed in the hydrogenation of cyclopropane over various metals Catalyst
Temp. range ("C)
Ea a (kJ mol-')
Ref.
Fe-powder Fe-powder
24 8-3 20 100-212
121.3 56.5
252 24 6
Co-SO2 Ni-Si02 /A1 O 3 Ni-AIz0, Ni-S i 0 /A1 0 Ni-Si 0 Ni-pumice Cu-SiO
93-158 56-100 30-70 30-90 3 2-4 2 130-200 113-207
44.8 (78.3) 63.6 55.0 63.2 (67.0) 54.4 (67.0) 44.4 45.6
250 247 256 256 250 24 3 250
Re-powder
70-128
52.3
252
Ru-powder Ru-Si 02
50-160 0-80
46.0 50.2 (50.2)
252 251
Rh-pumice Rh--Si02
49-124 -35 t o -10
41.9 46.0
26 251
Pd-pumice Pd-Si02
95-215 -10 to +25
41.9 67.0
26 251
0s-powder 0s-Si 0 2
19-50 0-6'0
50.2 54.4 (54.4)
252 251
Ir-pumire It-pumice Ir-Si02
0-60 55-177 0-3 0
48.1 41.0 54.4
26 245 251
Pt-pumice Pt-pumice Pt-Si 0
16-126 50-201 -20 to +30
35.2 37.3 46.0
26 24 3 251
-
a
Values in parenthesis are for the hydrocracking reaction.
the silica-supported noble Group VIII metals [251] were found to decrease in the order Rh > Pt > Pd > Ir > 0 s > Ru. Wallace and Hayes [252] also reported the occurrence of hydrocracking to ethane and methane during the hydrogenation of cyclopropane over rhenium, iron, ruthenium and osmium powders. The extent of hydrocracking, which was severe over iron, decreased in the order Fe > Re 2 Ru > 0 s . Under comparable experimental conditions, no hydrocracking of propene, propane or ethane was observed and consequently the hydrocracking, which has a similar activation energy t o that for hydrogenation [251], must involve an intermediate which is readily formed by cyclopropane but not by propene or propane. On this basis, it was proposed that the initial step in the hydrocracking and hydrogenation of cyclopropane was the formation of a 1,3-diadsorbed C3H6 species (HzTCHz%Hz). A similar conclusion was reached by Merta and Ponec [253, References p p . 1 1 4-1 21
102
2541 from their studies of cyclopropane hydrogenation and hydrocracking over evaporated nickel films and nickel powders, and by McKee [255] for the platinum-catalysed reaction. In a detailed kinetic study, Sridhar and Ruthven [256], using nickel supported on Kieselghur (58%Ni), alumina (14% and 40% Ni) and silicaalumina (5% Ni), showed that over all four catalysts the rates of both hydrogenation and hydrocracking could be correlated according t o the power rate law equation
r = k P "H2P & 3
6
For all four catalysts and for both hydrogenation and hydrocracking, it was found that rn 0 and n 0.6. Furthermore, although the relative rates of hydrogenation and hydrocracking varied from catalyst t o catalyst, the total specific reaction rates per unit area of nickel were similar for the different catalysts. It was concluded that the rates of both reactions were governed by a common ratedetermining step and a similar mechansim (Fig. 43) t o that proposed by Addy and Bond [245] and Anderson and Avery 12573, from deuterium exchange studies, was postulated. It was suggested that the cyclopropane is initially adsorbed as a weakly bound n-complex (A), which then undergoes ring fission t o form the 1,3diadsorbed species (B). The formation of (B) was suggested t o be the ratedetermining step in hydrogenation, hydrocracking and deuterium exchange. Relatively few studies of the hydrogenation of substituted cyclopropanes are extant. From the studies which have been reported [26,248, 2581, it would appear that, with alkyl-substituted cyclopropanes, ring cleavage occurs by rupture of the bond opposite the carbon atom carrying the greatest number of substituent groups. Thus in the platinum on
* * (B)
(A)
3 C3H8(g) [Hydrogenation] +21! CH,(g)
+ C2H6(g)[Hydrocracking]
* Fig. 43. Reaction scheme for the hydrogenation, hydrocracking and deuterium exchange of cyclopropane over metal catalysts.
103 pumice-catalysed hydrogenation of methylcyclopropane [ 261 the products are 95% iso-butane and 5% n-butane at ambient temperature and 70% iso-butane and 30% n-butane a t 250°C. In contrast, the hydrogena[ 2481 at room temperature yields tion of methylene-cyclopropane n-butane and methylcyclopropane, but little iso-butane. Other alkenylcyclopropanes show similar behaviour [ 2591.
6. Catalyst structure and hydrogenation activity In the foregoing sections of this review, attention has been largely confined to a discussion of the interaction of species once they are adsorbed at the catalyst surface. It has been implicitly assumed that, for the processes of chemisorption and reaction, sites possessing the necessary characteristics are available at the catalyst surface. In this last section, consideration is given t o the physical and chemical properties of the catalyst which give rise to hydrogenation activity. 6.1 GEOMETRIC FACTORS IN CATALYSIS
It is now some fifty years since Taylor [260] first proposed his theory of “active sites”. He considered that the amount of surface which is catalytically active depends upon the reaction which is being catalysed and that a “geometric factor” existed in catalysis. The existence of a “geometric factor” was also the basis of Balandin’s Multiplet Theory of Catalysis [261]. In the ensuing years, much research has been directed towards an understanding of this geometric factor by attempting t o obtain answers to such questions as: Is there an optimum metal-metal distance for catalytic activity? Does the catalytic activity vary with the size of the metal particles? Is a minimum lattice exhibiting bulk metal properties required for normal catalytic behaviour? It is in the past fifteen or so years that the most significant advances in the understanding of such factors have been made. Such understanding as currently exists has come with the development of such techniques as electron microscopy [2622641, X-ray line broadening [ 2651 and selective chemisorption [ 266-2691 for the accurate characterisation of small metal particles and the use of LEED and Auger spectroscopy for the study of single crystal metal surfaces [ 5-7,61,160,270-2741. Early attempts t o establish the existence or otherwise of a geometric factor were based upon the assumption that the surfaces of metal particles consisted of extensive arrays of atoms arranged in welldefined low index planes; the optimum metal-metal distances for the strain-free adsorption of the reactant hydrocarbon were calculated [84,157]. As noted in Sect. 4.2 (p. 50) such an approach led to the conclusion that only certain crystal planes should be active in alkene and alkyne hydrogenation References p p . I 14-1 2 1
104 [ 1581. Similarly, such considerations, which formed the basis of the Multiplet Hypothesis [ 2611, led t o the conclusion that the chemisorption and hydrogenation of benzene required a hexagonal array of six suitably spaced atoms such as found with the (111) planes of the face-centred cubic metals and the corresponding planes of the close-packed hexagonal metals. Thus metals having either of these structures should be more active than metals having a body-centred cubic structure. Recent studies [ 160,2741 with metal single crystals have suggested that such an approach is invalid. Thus Dalmai-Imelik et al. [ 274-2771 have shown that the activities of the (loo), (110) and (111) planes of nickel are approximately the same, even though the benzene is adsorbed flat on the (111) plane and differently (probably on edge) on the other two faces. An alternative approach t o an understanding of the role of surface geometry is to consider the geometrical consequences of the aggregation of metal atoms t o give microcrystallites and, in particular, to consider the various coordination numbers of the surface atoms at edges, corners, surface steps and within the faces and their variation with crystallite site. Such an approach was used by Poltorak et al. [278-2801 in an attempt t o specify, in terms of the minimum number of metal atoms, the geometry of a surface site for various catalytic reactions. These workers determined the distributions of surface atoms of different coordination numbers as a function of crystallite size for a face-centred cubic metal, which was assumed to form octahedral crystallites. It was concluded that, for crystallites in the range 10-20 A, a large number of atoms of abnormally low coordination number (<7) are present, whilst in the range 2050 A , there is a preponderance of atoms with coordination numbers corresponding t o crystal edges (7) and faces (9). Thus, by determining the specific activities, for a variety of metal-catalysed reactions, of catalysts with particle size distributions in the range 10-50 A diameter, it should be possible to establish relationships between activity and average coordination number of a surface metal atom. A similar approach and similar conclusions were reached by van Hardeveld et al. [281,282] and by Bond [283], both of whom examined the nature of the sites present at the surface of complete and incomplete cubo-octahedral crystallites. They were especially concerned with the concentration of the five coordinate B5 sites which were concluded t o be present in appreciable numbers only with crystallites of diameter <70 A . It has recently been stated [ 2741 that the situation may be even more complex than was assumed by the above workers, since it is claimed that, with crystallites containing, for example, 13 atoms, the stable crystallographic structure is an icosahedron [284,2851 and, in consequence, the exposed planes of small crystallites are not perfect, but deformed (111) planes. Experimental determination of the specific activities of small metal particles have revealed the surprising result that, in most hydrogenation systems, there is little variation of the activity with particle size. Thus, the
105 TABLE 35 Specific activities of silica-supported nickel catalysts for benzene hydrogenation (kH) and exchange ( k E ) a t 298 K [39,282] Metal content (wt.% Ni) CrystaIIite size ( A ) k H (mol h-' m-') X k E (mol h-' m-2) x
lo5 lo5
6.8
- 200
4.7-5.3 17-90
7.4
< 70 11.0 20
23.3
<50 9.0-1 2.5 0.7-3.8
25.3
<50 9.5 0.28
-
specific activities of a series of platinumsilica catalysts for hex-1-ene hydrogenation at 198 K were observed t o be independent of particle size [286] over a range in which the number of surface platinum atoms varied from 100% ( < l o A diameter) t o 30% (ca. 40 A diameter). A similar result has been obtained by Dorling and Moss [287] for the hydrogenation of benzene over platinumsilica catalysts; over an increase in particle size of two orders of magnitude, the specific activity remained constant. Boudart e t al. 12881 have also observed that, in the hydrogenation of cyclopropane at O"C, the difference in activity between highly dispersed platinum7-alimina and sintered platinum-q-alumina or platinum foil catalysts was only a twofold change in specific activity, while the platinum specific surface area varied by more than four orders of magnitude. Such observations led Boudart [ 2891 to classify metal-catalysed reactions as facile (structure-insensitive) and demanding (structure-sensitive) and t o conclude that hydrogenation reactions were facile. From the results of van Hardeveld and Hartog [282], it would appear that, for nickelsilica, benzene hydrogenation at 25" C is facile, whereas benzenedeuterium exchange is structure-sensitive (Table 35). With iridium, however, a particle size effect was apparent both in the hydrogenation and exchange of benzene. The lack of dependence of the specific rate of hydrogenation on particle size has also been observed with nickel catalysts by Ljubarskii [290], Nikolajenko e t al. [291,292] and Aben e t al. [44]. Other examples of facile hydrogenation are the platinum-catalysed hydrogenation of cyclohexene and ally1 alcohol [ 2791 and the nickel-silicacatalysed hydrogenation of buta-l:3-diene [ 2931. Bond [ 2941 used comparisons between homogeneously and heterogeneously catalysed interconversions of unsaturated hydrocarbons t o deduce that the reactive state of an adsorbed hydrocarbon may reasonably be assumed t o be a r-complex (see Sect. 3.2, p. 22). On this assumption, a molecular orbital model appropriate t o a face-centred cubic metal was deveIoped. By considering the direction of emergence and degree of oceupation of the metal atomic orbitals a t the (loo), (110) and (111)faces, assuming that the atomic orbitals on the surface keep the same orientation as in the bulk metal, which may not be valid [295], he concluded that the (111)planes were least suited t o the adsorption requirements of References p p . 1 1 4-1 21
106 alkenes, alkadienes and alkynes. Recent work with nickel single crystals suggests, however, that in ethylene hydrogenation the (100) pianes are inactive and that the activities of the (111)and (110) planes differ by about a factor of two, the (111)planes being the most active [284], although these results are in sharp contrast with the lack of dependence of specific activity on particle size noted above. It is clear that the influence of surface geometry upon catalytic activity is extremely complex and many more studies are required before any definitive relationship between catalytic activity and metal particle size can be established. Such studies will require t o take cognisance of such factors as the perturbation of surface structure due t o the formation of carbidic residues, as noted by Boudart [289] and by Thomson and Webb [95], and by the modification of catalytic properties on adsorption, as noted by Izumi et al. [296--2981 and by Groenewegen and Sachtler [299] in studies of the modification of nickel catalysts for enantioselective hydrogenation. Possible effects of the support, as will be discussed in Sect. 6.3, must also be taken into account. It is possible that future studies could lead t o the preparation of supported metal catalysts in which each cluster of metal atoms is simply that which constitutes a surface site, as suggested by Kobosev [ 3001. In this context, it is of interest t o note that Prestridge and Yates [301] have prepared catalysts containing atomic dispersions of rhodium on silica in which some 5-atom clusters were observed. (Rh), and (Rh)6 clusters have been made on a phosphinated polystyrene polymer support [302] using Rh4(C0)12and Rh6(C0)16as starting materials; the (Rh)6 catalysed the hydrogenation of aromatics at 25°C and 1 atm pressure, although the (Rh), was apparently inactive. Webb and co-workers [ 303,3041 have also prepared silica-supported metal catalysts from Ru,( CO) 2 , Os,( CO) and RhzCoz(CO),,, which appear t o contain highly dispersed metal and which are active for alkene hydrogenation and hydroisomerisation. 6 . 2 ELECTRONIC FACTORS IN CATALYSIS
In the early 1 9 5 0 ’ ~the ~ geometric approach t o the interpretation of catalytic activity was largely abandoned in favour of the so-called “elecTABLE 36 Relative values for log h s p for ethylene hydrogenation and for 6 for various metals Metal -
~
Ta
W
Cr
Fe
co
Ni
RU
log 4, 1% k , ,
-4.4
-4.0
-4.2
6
39
-3.0 -3.4 39.5
-2.1 39.7
-2.6 -1.5 40
-0.3 50
(%)
43
39
107 tronic factor”, in which the catalytic activity was related to the electronic properties of the bulk metal [305]. Beeck [50], in studies of the hydrogenation of ethylene over evaporated metal films, suggested that the specific activity could be correlated with the percentage d-character (6) of the metal bond as calculated by Pauling [306]. Thus, as shown in Table 36, the specific activity increased as the value of 6 increased. A similar conclusion was reached by Schuit and van Reijen [ l o g ] for ethylene hydrogenation over silica-supported metals, although these workers suggested that correlation should be sought between the activity and the product of the d-character and the metallic valency, rather than with 6 alone. The concept of an electronic factor in catalysis led to numerous studies of the catalytic behaviour of metal alloys. According t o the Band Theory of metals [307], the transition metals contain vacancies in the d-band which can be filled by being alloyed with, for example, a Group IB metal. Studies of catalysis by alloys were prompted by the expectation that abrupt changes in catalytic behaviour would be observed at an alloy composition corresponding t o the complete filling of the d-band; the presence of d-band vacancies having been suggested as being a necessary prerequisite for catalytic activity [ 308,3091. Thus, for example, in the palladium-old alloy system, changes in activity might be expected at a composition of 60 at.% gold [ 3101. Most of the early studies were carried o u t using alloy powders [311,312], foils [309,313,314] or wires [198,315], alloys between the Group VIII metals, nickel, palladium and platinum, and the Group IB metals, copper, silver, and gold, being used in order to minimise changes in surface geometry with alloy composition. More recently, evaporated metal films, which are the subject of an extensive review by Moss and Whalley [316], have been employed. Attention has been centred upon changes in specific hydrogenation activity with alloy composition. However, in many cases the existence of a compensation effect between the activation energy and pre-exponential factor necessitates the consideration of the variation of these two parameters with alloy composition. In sharp contrast t o the results obtained in, for example, the palladium-gold-catalysed para-hydrogenation conversion [ 3171, formic acid
Ref. Rh
Pd
0 0 50
-0.8
-0.9 46
References P P . 114--121
Ir
Pt
cu
-2.0 49
-1.65 -1.5 44
-4.1 36
50 109
108 At.
-
:
-1
0 1201
10
20
I
I
1
0
10
20
O/.
Pd in Pd/Au alloy ( ( 3 )
60
30
40
50
1
I
I
I
1
1
1
30
70
80
1
I
50 60 70 80 A t . % Ni in Ni/Cu alloy ( 0 , O )
40
90
100
I
I
90
100
Fig. 44. Variation of the apparent activation energy with alloy composition for the hydrogenation of ethylene ( 0 ) [ 3201 and acetylene (0)[ 311 ] over Cu-Ni alloys and for but-2-yne hydrogenation (a)over Pd-Au alloys [ 1981.
decomposition [ 3181 and oxidation of carbon monoxide [ 3191, where, as expected, a sharp change in activation energy is observed at approximately 60 at.% gold, alloy-catalysed hydrogenation reactions provide little evidence t o support the d-band theory. Thus, for example, as shown in Fig. 44,the activation energy for acetylene hydrogenation over nickelcopper alloy powders [311] passes through a maximum at ca. 80% copper, whilst in the nickel-copper alloy-catalysed hydrogenation of ethylene [ 3201 and the hydrogenation of but-2-yne over palladium-old wires [198], the activation energy appears to be almost invariant with alloy composition. Similarly, in the hydrogenation of buta-1: 3diene over pumice-supported palladium-gold alloys [ 2101 the activation energy remained constant at -44 kJ mol-' from 0 t o 60 at.% gold, showed a maximum of 6 3 kJ mol-' at 65-70 at.% gold and decreased to 37.5 kJ mol-' for 9 5 at.% gold. It should be noted that, in the interpretation of activity patterns of alloy catalysts, extreme care is needed to ensure that the surface composition is known. It has been shown [321,322] with copper-nickel alloys, which show two phases in the composition range 2-80% copper, that, within this miscibility gap, the surface composition remains constant at 80% Cu-20% Ni, independent of the nominal bulk composition. Furthermore, the surface composition may vary depending upon the catalyst pretreatment [ 3221. No miscibility gap occurs with palladium-old or palladiumsilver alloys [ 3231. In the apparent absence of a correlation between hydrogenation activity and the bulk electronic properties of alloys, Sachtler et al. [321,322,
109 3241 have attempted to explain their results for benzene hydrogenation over copper-nickel alloys in terms of a surface atom model. In this model, it is envisaged that chemisorption involves the interaction of the adsorbate with localised d-orbitals on individual surface nickel atoms, the properties of which may be modified by the presence of neighbouring copper atoms. An essentially similar conclusion was reached by Clarke et al. [325,326] who claimed that their results for the hydrogenation of but-1-ene and of buta-1: 3-diene over nickel-copper alloy films were consistent with a model in which the alloys were acting as mixtures of active nickel centres in copper. Rushford and Whan [198] also concluded that, in the palladium-gold alloy-catalysed hydrogenation of but-2-yne, the changes in activity, which were due to changes in the pre-exponental factor rather than the activation energy, were consistent with a model in which the active sites were islands of four adjacent palladium atoms. Clarke [347], in a recent review, has discussed both the collective electron and localised atomic site models. From the evidence available a t present, it would appear that, in catalytic hydrogenolysis, the influence of the electronic structure of the metal may be best explained by considering the electron population of the localised d-orbitals of individual surface sites, which may be an individual metal atom or a cluster of metal atoms as discussed in the previous section. If this is the case, it follows that geometric and electronic factors will be interdependent and indeed it may be impossible t o separate the effect of each upon the hydrogenation activity. The localised orbital model also provides a rational basis for considering possible correlations between the metal-catalysed hydrogenation of unsaturated hydrocarbons and the organometallic chemistry of the catalytically active metals, as discussed by Rooney and Webb [87] and Bond and Wells [131]. However, such an approach may be of limited value in understanding the origins of the hydrogenation activity of metals, since it does not take into account the perturbation of the electron population of the surface orbitals arising from such effects as the formation of irreversibly adsorbed carbonaceous residues, or of possible electronic interactions between the metal and support in supported-metal catalysts [ 3271, unless it is assumed that these effects are the same as those of the other ligands present in the organometallic compounds. 6.3 SURFACE MIGRATION AND THE INFLUENCE O F CATALYST SUPPORTS
The type of catalyst most commonly used in studies of the hydrogenation of unsaturated hydrocarbons is one in which the metal is dispersed upon a support such as alumina or silica. Although for many years it was generally thought that the support was catalytically inert, studies over the past few years have suggested that this may not be the case and there is a growing body of evidence t o suggest that the support may influence the References P P . I 1 4-1 21
110 catalytic activity in any of several ways. Thomson and Harvey [327] suggested that a chemical interaction between metal and support was responsible for the modifications of electrical conductivity when a metal is dispersed on different supports. A similar reason has been invoked t o explain the differences in activities and selectivities for benzene hydrogenation of platinum when supported on polyamides, molecular sieves and alumina [ 3281, and the differences in activity for buta-1: 3-diene hydrogenation over y-alumina, boehmite and silica-supported gold [ 2111. Figueras et al. [329] have observed that, in the hydrogenation of benzene over supported palladium, the activity per surface metal atom increased with increasing Lewis acidity of the support. This was interpreted in terms of the modification of the electronic state of the palladium due to metaloxidising site interactions. A similar effect has also been observed for benzene hydrogenation over platinum-amorphous alumina [ 3311. Another way in which the support may modify the catalytic behaviour of a supported metal is by acting as a source of reactive intermediates or as a seat of reaction, through the migration of adsorbed species between metal and support. Since the observations by Sinfelt and Lucchesi [330] that the activity of platinumsilica catalysts for ethylene hydrogenation could be increased by mixing the catalyst with alumina, the problems of surface migration, which has been termed “spill-over” [ 3321 has attracted the attention of a number of workers and has been the subject of a review [ 3331. Most of these studies have been concerned with the migration of hydrogen between metal and support, although Webb e t al. [ 40,53,6467,1251 using a radiotracer approach have obtained evidence for the migration of hydrogen and of olefins and acetylene. Boudart e t al. [332,334] obtained direct evidence for spill-over of hydrogen by following the reduction of W 0 3 to H,W03 with hydrogen at room temperature using mechanical mixtures of WO, with supported or unsupported platinum catalysts. Sancier [ 3351 has claimed that, in the hydrogenation of benzene over palladiumalumina/alumina mixtures, chemically reactive hydrogen can migrate over distances of up t o 0.5 mm, part of the benzene being hydrogenated on the palladium and part on the alumina. However, this result has been disputed by Vannice and Neikam [ 3361, who claim that the specific activity of palladiumalumina, as determined from the turnover number (molecules of benzene reacted per second per surface Pd atom), is independent of the dilution of the catalyst with alumina. Teichner and co-workers [337], in a particularly elegant study, demonstrated hydrogen spillover from a nickelalumina catalyst to alumina by hydrogenating ethylene on alumina which had previously been treated with hydrogen in the presence of the nickel-alumina. These workers [ 331,3381 have also obtained evidence for hydrogen spill-over in the hydrogenation of cyclopropane and methylcyclopropane over platinum supported o n amorphous alumina, and in the hydrogenation of ethylene over 6-alumina-supported platinum.
111 In the studies mentioned above, it has generally been considered that the hydrogen spill-over occurs from metal t o support and that the subsequent reaction occurs on the support. Evidence for the back migration of hydrogen from support t o metal is less well documented. Altham and Webb [ 60,661 showed that, when platinurnalumina or platinumsilica catalysts were pretreated with tritiated hydrogen and subsequently titrated with ethylene or acetylene, hydrogen equivalent t o a hydrogen: ) of up to 66 could be reacted. It was also total platinum ( n H / n p tratio shown that the support alone was not active under the conditions used and hence it was concluded that the tritium, which was suggested to be associated with the support hydroxyl groups, was capable of migrating from the support to the metal. A similar conclusion was drawn by Sermon and Bond [ 3331, who used an alkene titration method t o show that platinum-H,WO, catalysts at 100°C yielded hydrogen equivalent t o a nH/npt ratio of around 277. However, subsequent studies by Webb e t al. [53,64, 66,671 suggested that the migration of ethylene and acetylene from metal to support could occur, a conclusion in agreement with the observation that, in the reaction of but-1-ene with hydrogen over mercury-poisoned rhodiumsilica, but-1-ene isomerisation occurs on the support [ 1251. Such migration could possibly explain the alkene titration results cited above. Indisputable evidence for the occurrence of hydrogen spill-over exists although the mechanism whereby it occurs is still not clear. Benson e t al. [ 3391 have suggested that the migration occurs by a surface rather than a gas phase transport of hydrogen species, whilst Neikam and Vannice [340] claim that organic residues, arising from the preparation of the catalyst, from carriers or greased stopcocks, may play an important part by acting as bridges for the migration; it has also been reported that the apparent absence of hydrogen spill-over is pertinent t o relatively clean catalyst systems [ 341,3421. Thus, far too little attention appears to have been paid to the possible role of impurities such as Fe in the support, which have been shown to play an important part in determining the hydrogenation activity of alumina [ 3431, or of the presence of ions of the catalytically active metal in the support matrix [ 3441. Clearly, further research is required before the mechanism of spill-over is established, and the role, if any, of spill-over in catalytic hydrogenation by supported metals is deduced. Such studies must be made on well-characterised catalysts under carefully controlled conditions t o minimise the effects of impurities and contaminants; they must also take into account such effects as the occlusion of hydrogen by metals as shown by the studies of Paal and Thomson [345] with platinum blacks and Babenkova e t al. [346] for pall ad iu m blacks .
R c f e r e f i c c sp p , 1 1 4-1 21
112
7. Conclusions In the foregoing sections, the various aspects of the metal-catalysed hydrogenation of unsaturated hydrocarbons have been discussed. In this final section, it is appropriate to attempt t o summarise, in a general way, the current state of knowledge of the subject and t o make suggestions as t o possible future developments. From the wealth of information which has accrued over the past two decades, it is clear that a measure of order in the subject has appeared. Thus, it is clear that, for example, the hydroisomerisation and deuterium exchange of alkenes are a property of the metal rather than of the alkene being hydrogenated (Sect. 3). Similarly, in the hydrogenation of diunsaturated hydrocarbons (Sect. 4) experimental variables such as the selectivity for olefin formation and the stereospecificity are inherent properties of the metal rather than the alkyne or alkadiene under consideration. On the other hand, for a given hydrocarbon, the activation energy appears t o vary only within narrow limits from metal t o metal, activity changes being ascribed t o variations in the pre-exponential factor. Hydrocarbon hydrogenation does not appear to depend upon the surface topography or upon the bulk electronic properties of the metal (Sect. 6). In spite of the vast literature and plethora of information which is available, it is apparent that the subject of the metal-catalysed hydrogenation of hydrocarbons still lacks a basic unifying theory. One of the possible reasons for this is that, as noted in Sect. 2, the study of the actual surface reaction is inherently difficult. However, it is also apparent that many studies have tended to concentrate on single aspects of the subject. Thus, in studies of the kinetics of catalytic hydrogenation, although one finds innumerable statements on pretreatments and variations in activity, in the interpretation of the observed kinetics little or no attention is paid to the possible effects of the pretreatment upon the surface properties. Similarly, it is well established from, for example, spectroscopic and radiochemical studies that much, if not all, of the surface is covered with retained hydrocarbondeficient carbonaceous residues. LEED studies of single crystals also suggest the formation of hydrocarbon overlayers. However, as has been done in this review, the interpretation of reaction mechanisms have been based upon the assumption that the hydrocarbon, and usually the hydrogen as well, are adsorbed directly upon the metal. It has been stated that the results obtained from adsorption studies on single metal crystals and in the absence of hydrogen are of doubtful relevance t o the problem of the catalytically active form of an adsorbed hydrocarbon. Correlations have been sought between catalytic activity and electronic praperties of the pure metal rather than with the electronic properties of the metal, be it a single surface atom or an atom in a surface array, which has almost certainly been modified by the presence of non-participating hydrocarbon species.
113 The message seems t o be clear, In order t o obtain a fundamental understanding of catalytic hydrogenation it will be necessary t o consider together all of the various facets of the adsorption and reaction of hydrocarbons on metals. In one such attempt, Thomson and Webb [95] have recently suggested that the various features of the hydrogenation of unsaturated hydrocarbons, such as those discussed above, are consistent with a general mechanism, in which hydrogenation on metals should be interpreted as hydrogen transfer between an adsorbed hydrogen-deficient hydrocarbon species, M-C,H, , and adsorbed alkene or alkyne, rather than the direct addition of hydrogen t o an associatively adsorbed hydrocarbon. Such a proposal has the merit that it makes the metal only of secondary importance, in agreement with the experimental observations regarding geometric and electronic factors. Correlation between activity and metal surface area would be expected since the concentration of active M<,H, sites will depend upon the latter. Further, the necessity for cacalyst pretreatment t o obtain reproducible activity and the observed retention of hydrocarbons can be considered t o represent the formation and stabilisation of active sites giving a steady state concentration of permanently retained M--C,H, species. Similarly, in the reactions of unsaturated hydrocarbons with deuterium, the formation of isomerised and hydrogenated products which, in the initial stages of reaction, are predominantly the -do species can be readily interpreted in terms of hydrogen atom transfer between the active centre M--C,H, and the associatively adsorbed hydrocarbon. A similar, although more restricted mechanism has also been suggested for ethylene hydrogenation over metals by Gardner and Hansen [348], whilst surface carbiding has been suggested as being responsible for creating active sites on palladium-gold alloys [ 3491. This latter observation is equally well explained by considering the active site to be M--C,H,, . The proof, or otherwise, of such proposals can only come from further research studies. It is apparent that in such studies it will be necessary to use catalysts which have been thoroughly characterised in terms of metal areas, particle size distributions and impurity levels. More important, however, is the obvious need t o characterise the catalyst surface in terms of the reaction being catalysed and, in particular, t o determine the effects on the surface of the permanently retained hydrocarbon species. In this context, the use of techniques such as infrared spectroscopy and the application of radiotracer methods t o monitor the surface under reaction conditions are particularly useful. The possible extension of studies of adsorption and reactions currently being carried out using welldefined metal single crystals to examine the applicability of the results obtained with single crystals t o the more usual forms of catalyst together with further studies of catalysts containing metal clusters of known structure are worthy of further investigation. Such an approach, albeit experimentally difficult, could provide rich rewards in terms of our further understanding of hydrogenation catalysis. References p p . 1 1 4-1 21
114 References 1 P. Sabatier and J.B. Senderens, Compt. Rend., 135 (1902) 8 7 ; 137 (1903) 301; P. Sabatier and B. Kubota, Compt. Rend., 172 (1921) 173. 2 P.W. Selwood, Adsorption and Collective Paramagnetism, Academic Press, New York, 1962. 3 See e.g. G.C. Bond, Catalysis by Metals, Academic Press, New York, 1962, p. 2. 4 R.J. Kokes, Catalysis, 1 (1973) 1. 5 B. Lang, R. Joyner and G.A. Somorjai, J. Catal., 27 (1972) 405. 6 S.L. Bernacek and G.A. Somorjai, J. Chem. Phys., 62 (1975) 3149. 7 D.W. Blakely and G.A. Somorjai, J. Catal., 42 (1976) 181. 8 0. Beeck, Discuss. Faraday SOC.,8 (1950) 118. 9 See e.g. G.C. Bond and P.B. Wells, J. Catal., 4 (1965) 211; 5 (1966) 65, 419. 1 0 W.F. Madden and C. Kemball, J . Chem. SOC.,(1961) 302. 11 G.C. Bond and J . Newham, Trans. Faraday SOC.,56 (1960) 1501. 12 K.J. Laidler, Catalysis, 1 (1954) 75. 1 3 O.A. Hougen and K.M. Watson, Chemical Process Principles, Vol. 3, Wiley, New York, 1955, pp. 908 e t seq. 14 K. Tamaru, Adv. Catal., 1 5 (1964) 65. 15 S.J. Thomson and G. Webb, Platinum Met. Rev., 11 (1967) 46. 1 6 J. M. Thomas and W.J. Thomas, Introduction t o t h e Principles of Heterogeneous Catalysis, Academic Press, New York, 1967, Chap. 2. 17 M. Bowker and D.A. King, in C.H. Bamford and C.F.H. Tipper (Eds.), Comprehensive Chemical Kinetics, Vol. 19, Elsevier, Amsterdam, in press. 18 I. Langmuir, Trans. Faraday SOC.,17 (1921) 621. 1 9 C.N.Hinshelwood, Kinetics of Chemical Change, Oxford University Press, O x ford, 1926, p. 145. 20 E.K. Rideal, Proc. Cambridge Philos. SOC.,35 (1939) 130. 21 E.K. Rideal, Chem. Ind. (London), 62 (1943) 335. 22 D.D. Eley and E.K. Rideal, Proc. R. SOC.London, Ser. A, 178 (1941) 429. 23 G.H. Twigg, Discuss. Faraday SOC.,8 (1950) 152. 24 0. Beeck, Discuss. Faraday SOC.,8 (1950) 118. 25 G.I. Jenkins and E.K. Rideal, J. Chem. SOC., (1955) 2490, 2496. 26 G.C. Bond and J. Newham, Trans. Faraday SOC.,56 (1960) 1501. 27 V. Grassi, Nuovo Cimento, 11 (1916) 147. 25 G.M. Schwab and K. Naicker, Z. Elektrochem., 42 (1936) 670. 29 E.N. Pease, J. Am. Chem. SOC., 4 5 (1923) 1196. 30 A. Farkas and L. Farkas, J. Am. Chem. SOC., 60 (1938) 22. 31 N. Thon and H.A. Taylor, J. Am. Chem. SOC.,7 5 (1953) 2747. 32 G.C. Bond, G. Webb and P.B. Wells, J. Catal., 1 2 (1968) 157. 3 3 H.C. Andersen, A.J. Haley and W. Egbert, Ind. Eng. Chem., 52 (1960) 901. 34 G.C. Bond and P.B. Wells, Adv. Catal., 1 5 (1963) 91. 35 A. Wheeler, Adv. Catal., 4 (1951) 250. 36 H.I. Waterman and A.B.R. Weber, J. Inst. Pet., 43 (1957) 315; H. Breimen, H.I. Waterman and A.B.R. Weber, J . Inst. Pet., 4 3 (1957) 299. 37 H.I. Waterman, C. Boelhouwer and J. Cornelissen, Anal. Chim. Acta, 18 (1958) 497. 38 G.C. Bond, D.A. Dowden and N. Mackenzie, Trans. Faraday SOC., 54 (1958) 1537. 39 P.B. Wells, Surf. Defect. Prop. Solids, 1 (1972) 245. 40 G.F. Taylor, S.J. Thomson and G . Webb, J. Catal., 1 2 (1968) 150; 12 (1968) 157. 41 A S . Al-Ammar and G. Webb, J. Chem. SOC.Faraday Trans. I, 74 (1978) 195.
115 42 K.J. Laidler, Catalysis, 1 (1954) 105. 43 K.J. Laidler, Catalysis, 1 (1954) 219. 44 P.C. Aben, J.C. Platteeuw and B. Stonthamen, Proc. Fourth Int. Congr. Catalysis, Akademiai Kiado, Budapest, 1 (1971) 395. 45 0. Toyama, Rev. Phys. Chem. Jpn., 12 (1938) 115. 46 A. Farkas, L. Farkas and E.K. Rideal, Proc. R. Soc. London, Ser. A, 146 (1934) 630. 47 G.H. Twigg and E.K. Rideal, Proc. R. SOC.London, Ser. A, 1 7 1 (1939) 55. 48 G.M. Schwab and H. Zorn, Z. Phys. Chem., Abt. B, 32 (1936) 169. 49 J.U. Reid, S.J. Thomson and G.Webb, J. Catal., 30 (1973) 433. 50 0. Beeck, Discuss. Faraday SOC.,8 (1950) 118. 51 G.I. Jenkins and E.K. Rideal, J. Chem. SOC.,(1955) 2490,2496. 52 S.J. Stephens, J. Phys. Chem., 6 3 (1959) 512. 53 A.S. Al-Ammar, S.J. Thomson and G. Webb, unpublished results. 54 D.W. McKee, J. Am. Chem. SOC.,8 4 (1962) 1109. 55 P.W. Selwood, J. Am. Chem. Soc., 83 (1961) 2853. 56 J.R. Arthur and R.S. Hansen, J. Chem. Phys., 36 (1962) 2062. 57 R.S. Hansen and N.C. Gardner, J. Chem. Phys., 74 (1970) 3646. 58 L. Whalley, B.J. Davis and R.L. Moss, Trans. Faraday Soc., 66 (1970) 3143. 59 R.J. Kokes, J. Catal., 14 (1969) 83, 60 J.A. Altham, Ph.D. Thesis, University of Glasgow, 1969. 61 D.L. Smith and R.P. Merril, J. Chem. Phys., 52 (1970) 5861. 62 S.J. Thomson and J.L. Wishlade, Trans. Faraday SOC.,58 (1962) 1170. 63 D. Cormack, S.J. Thomson and G. Webb, J. Catal., 5 (1966) 284. 64 J.U. Reid, S.J. Thomson and G. Webb, J. Catal., 29 (1973) 421. 65 G.F. Taylor, S.J. Thomson and G. Webb, J. Catal., 1 2 (1968) 191. 66 J.A. Altham and G. Webb, J. Catal., 18 (1970) 133. 67 J.U. Reid, S.J. Thomson and G. Webb, J. Catal., 29 (1973) 433. 68 L.H. Little, Infrared Spectra of Adsorbed Species, Academic Press, London and New York, 1966. 69 N. Sheppard, Discuss. Faraday Soc., 41 (1966) 254. 70 N.R. Avery, J. Catal., 19 (1970) 15. 7 1 H.A. Pearce and N. Sheppard, Surf. Sci., 59 (1976) 205. 72 L.H. Little, Infrared Spectra of Adsorbed Species, Academic Press, London and New York, 1966. 73 B.A. Morrow and N. Sheppard, J. Phys. Chem., 70 (1966) 2406. 74 N. Sheppard, Discuss. Faraday Soc., 41 (1966) 171. 75 B.A. Morrow and N. Sheppard, Proc. R. Soc., London, Ser. A, 311 (1969) 391. 76 R. Komers, Y. Amenoyima and R.J. Cventanovic, J. Catal., 1 5 (1969) 293. 77 B.A. Morrow, J. Catal., 1 4 (1969) 279. 78 R.P. Eischens and W.A. Pliskin, Adv. Catal., 10 (1958) 1. 79 J.B. Peri, Discuss. Faraday SOC.,41 (1966) 121. 80 J. Horiuti and M. Polanyi, Nature (London), 132 (1933) 931. 81 J. Horiuti and M. Polanyi, Trans. Faraday Soc., 30 (1934) 1164. 82 J. Erkelens and T h . J . Liefkens, J. Catal., 27 (1972) 165. 8 3 R.S. Hansen and N.C. Gardner, J. Phys. Chem., 74 (1970) 3298. 84 G.C. Bond, Catalysis by Metals. Academic Press, New York, 1962, p. 234. 85 J.J. Rooney, F.G. Gault and C. Kemball, J. Catal., l ( 1 9 6 2 ) 255. 8 6 J.J. Rooney, J. Catal., 2 (1963) 53. 87 J.J. Rooney and G. Webb, J. Catal., 3 (1964) 488. 88 J.C.P. Mignolet, Discuss. Faraday Soc., 8 (1950) 105. 89 R. Dus, Surf. Sci., 50 (1975) 241. 90 J.D. Prentice, A. Lesiunas and N. Sheppard, J. Chem. SOC., Chem. Commun., (1976) 76.
116 91 92 93 94 95 96 97 98 99 100 101 102 103
J.E. Douglas and B.S. Rabinovich, J. Am. Chem. SOC.,74 (1952) 2486. G.C. Bond and J. Turkevich, Trans. Faraday SOC.,49 (1953) 281. C. Kemball, Adv. Catal., 11 (1959) 223. G.C. Bond, G. Webb, P.B. Wells and J.M. Winterbottom, J. Catal., 1 (1962) 74. S.J. Thomson and G. Webb, J. Chem. SOC.Chem. Commun., (1976) 526. J. Horiuti, J. Res. Inst. Catal., Hokkaido Univ., 6 (1958) 187. J. Horiuti and I. Matsuzaki, J. Res. Inst. Catal., Hokkaido Univ., 6 (1958) 250. J. Horiuti, Proc. 2nd Int. Congr. Catal., Technip, Paris, 1961, p. 1191. T. Keii, J . Chem. Phys., 22 (1954) 144. T. Keii, J. Chem. Phys., 23 (1955) 210. T. Keii, J. Chem. Phys., 25 (1956) 364. C. Kemball, J. Chem. SOC.,(1956) 735. G.C. Bond, J.J. Phillipson, P.B. Wells and J.M. Winterbottom, Trans. Faraday SOC.,60 (1964) 1847. 104 G.C. Bond, G. Webb and P.B. Wells, Trans. Faraday SOC.,6 1 (1965) 999. 105 G.C. Bond, J.J. Phillipson, P.B. Wells and J.M. Winterbottom, Trans. Faraday SOC.,62 (1966) 443. 106 J, Grant, R.B. Moyes and P.B. Wells, Trans. Faraday SOC.,69 (1973) 1779. 107 C. Kemball and P.B. Wells, J. Chem. SOC.A, (1968) 444. 108 G.C. Bond, Trans. Faraday SOC.,52 (1956) 1235. 109 G.C.A. Schuit and L.L. van Reijen, Adv. Catal., 1 0 (1958) 242. 110 E.K. Rideal and G.H. Twigg, Proc. Roy. SOC.London, Ser. A, 1 7 1 (1939) 55. 111 W.K. Hall and P.H. Emmett, J. Phys. Chem., 6 3 (1959) 1102. 112 L.A. Wanninger and J.M. Smith, Chem. Weekbl., 56 (1960) 273. 1 1 3 R.N. Pease, J. Am. Chem. SOC., 45 (1923) 1196, 2297; R.N. Pease and C.A. Harris, J. Am. Chem. SOC.,57 (1935) 1147. 114 A. Farkas and L. Farkas, J. Am. Chem. SOC.,60 (1938) 22. 115 J. Turkevich, F. Bonner, D.O. Schissler and A.P. Irsa, Discuss. Faraday SOC.,8 (1950) 352. 116 J. Turkevich, D.O. Schissler and A.P. Irsa, J. Phys. Colloid Chem., 55 (1951) 1078. 117 R.S. Mann and T.R. Lien, J. Catal., 1 5 (1969) 1 . 118 J. Addy and G.C. Bond, Trans. Faraday SOC.,5 3 (1957) 377. 119 R.F. Kayser and H.E. Hoelscher, Chem. Eng. Prog. Symp., Ser. 10, 50 (1954) 109. 1 2 0 K.J. Laidler and R.E. Townshend, Trans. Faraday SOC.,57 (1961) 1590. 121 G.H. Twigg, Proc. R. SOC.London, Ser. A, 178 (1941) 106. 122 T.I. Taylor and V.H. Diebler, J. Phys. Colloid Chem., 55 (1951) 1036. 123 G.C. Bond, G. Webb and P.B. Wells, Trans. Faraday SOC.,64 (1968) 3077. 124 G.C. Bond and J.M. Winterbottom, Trans. Faraday SOC.,65 (1969) 2794. 125 J.I. Macnab and G. Webb, J. Catal., 1 0 (1968) 19. 126 S.D. Mellor and P.B. Wells, Trans. Faraday SOC.,65 (1969) 1883. 127 W.M. Hamilton and R.L. Burwell, Proc. 2nd Int. Congr. Catal., Technip, Paris, 1961, p. 987. 128 J.E. Kilpatrick, E.J. Prosen, K.S. Pitzer and F.D. Rossini, J. Res. Nat. Bur. Stand., 36 (1946) 590. 129 G. Webb and J.I. Macnab, J . Catal., 26 (1972) 226; see also J.I. Macnab, Ph.D. Thesis, University of Glasgow, 1969. 130 S.D. Mellor and P.B. Wells, Trans. Faraday SOC.,65 (1969) 1873. 1 3 1 G.C. Bond and P.B. Wells, Adv. Catal., 1 5 (1964) 109. 132 J.J. Phillipson and P.B. Wells, Proc. Chem. SOC.,(1964) 222. 133 P.B. Wells and G.R. Wilson, J. Catal., 9 (1967) 70. 134 S. Carra and V. Ragaini, J. Catal., 1 0 (1968) 230. 135 V. Ragaini, J. Catal., 34 (1974) 1.
117 136 137 138 139 140 141
142 143 144 145 146 147 148 149 150
R.S. Yolles, B.J. Wood and H. Wise, J. Catal., 2 1 (1971) 66. D. Briggs and J . Dewing, J . Catal., 28 (1973) 338. L. Holbrook and H. Wise, J. Catal., 24 (1972) 315. Ho-Peng Koh and R. Hughes, J. Catal., 33 (1974) 7. R. Touroude and F.G. Gault, J. Catal., 32 (1974) 288,294. F.G. Gault, M. Ledoux, J.J. Masini and G. Roussy, Proc. 6th. Int. Congr. Catal., Vol. I, The Chemical Society, London, 1977, p. 469. G.V. Smith and J.R. Swoap, J. Org. Chem., 3 1 (1966) 3904. K. Hirota and Y. Hironaka, J . Catal., 4 (1965) 602. G.V. Smith and R.L. Bunvell, J. Am. Chem. SOC.,84 (1962) 925. G.C. Bond and J.S. Rank, Proc. 3rd Int. Congr. Catal., Vol. 2, Elsevier, Amsterdam, 1965, p . 1225. R. Maurel, M. Guisnet and G. Perot, J. Catal., 22 (1971) 151. I. Heertje, G.K. Joch and W.J. Wosten, J. Catal., 32 (1974) 337. Y. Inoue and I. Yasumori, J . Phys. Chem., 73 (1969) 1618. N. Sheppard and J.W. Ward, J. Catal., 15 (1969) 50. L.H. Little, N. Sheppard and D.J.C. Yates, Proc. R. SOC. London, Ser. A, 259
(1960) 242. 1 5 1 N.P. Sokolova, L.A. Kazakova and L.A. Borisenko, Russ. J . Phys. Chem., 44 (1970) 1515. 152 S.S. Randhava and A. Rehmat, Trans. Faraday SOC.,66 (1970) 235. 153 C.P. Nash and R.P. DeSieno, J. Phys. Chem., 69 (1965) 2139. . 154 J.J. McCarroll and S.J. Thomson, J. Catal., 19 (1970) 144. 155 C.P. Khulbe and R.S. Mann, Proc. 6 t h Int. Congr. Catal., Vol. I , The Chemical Society, London, 1977, p. 447, 156 G.C. Bond, Catalysis, 3 (1954) 109. 1 5 7 G.C. Bond, Catalysis by Metals, Academic Press, London and New York, 1962, p. 283. 158 J. Sheridan and W.D. Reid, J. Chem. SOC.,(1952) 2962. 159 G.C. Bond and G. Webb, Platinum Met. Rev., 6 (1962) 12. 160 L.L. Kesmodel, P.C. Stair, R.C. Baetzold and G.A. Somorjai, Phys. Rev. Lett., 36 (1976) 1316. 161 R.G. Guy and B.L. Shae, Adv. Inorg. Chem. Radiochem., 4 (1962) 78. 162 F.A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, InterscienceWiley, New York, 1972, 3rd edn. p . 749. 163 G.C. Bond, J . Chem. SOC.,(1958) 4288. 164 G.C. Bond and P.B. Wells, J. Catal., 6 (1966) 397. 165 G.C. Bond, G . Webb and P.B. Wells, J. Catal., 1 2 (1968) 157. 166 J.J. Phillipson, P.B. Wells and G.R. Wilson, J. Chem. SOC.A , (1969) 1351. 167 A.J. Bates, Z.K. Leszczynski, J.J. Phillipson, P.B. Wells and G.R. Wilson, J. Chem. Soc. A, (1970) 2435. 168 J . Sheridan, J . Chem. SOC.,(1944) 373, 476. 169 J. Sheridan, J . Chem. SOC.,(1945) 133,301. 170 G.C. Bond and J. Sheridan, Trans. Faraday SOC.,48 (1952) 651. 1 7 1 G.C. Bond, J . Chem. SOC.,(1953) 2705. 172 G.C. Bond and R.S. Mann, J . Chem. SOC.,(1959) 3566. 173 G.C. Bond, D.A. Dowden and N. Mackenzie, Trans. Faraday SOC.,54 (1958) 1537. 174 G.C. Bond and R.S. Mann, J. Chem. Soc., (1958) 4738. 175 G.C. Bond and P.B. Wells, Proc. 2nd. Int. Congr. Catal., Technip, Paris, 1961, pp. 1135-11 39. 176 G.A. Vassiliev, Bull. SOC.Chim. Fr., 15 (1948) 381. 177 P. Cossee, J. Catal., 3 (1964) 80. 178 F. d e Pam0 and J.C. Jungers, Bull. SOC.Chim. Belg., 57 (1948) 618.
118 179 W.T. McGown, C. Kemball, D.A. Whan and M.S. Sciirrell, J. Chem. SOC., Faraday Trans. I, 7 3 (1977) 632. 180 A. Grignon-Dumoulin and C. Thonon, Rev. Inst., Fr. Pet., 1 4 (1959) 214. 181 R.S. Mann and S.C. Naik, Can. J. Chem., 4 5 (1967) 1023. 182 R.S. Mann and K.C. Khulbe, Can. J. Chem., 45 (1967) 2755. 183 R.S. Mann and K.C. Khulbe, Can. J. Chem., 46 (1968) 623. 184 R.S. Mann and K.C. Khulbe, J. Phys. Chem., 7 3 (1969) 2104. 185 R.S. Mann and K.C. Khulbe, Can. J. Chem., 47 (1969) 215. 1 8 6 A. Rieche, A. Grimmm and H. Albrecht, Brennst. Chem., 42 (1961) 5. 187 R.S. Mann and K.C. Khulbe, J. Catal., 1 0 (1968) 401. 188 R.S. Mann and K.C. Khulbe, J. Catal., 1 3 (1969) 25. 189 E.F. Meyer and R.L. Burwell, J. Am. Chem. SOC.,8 5 (1963) 2881. 190 L.K. Freidlin and Y.Y. Kaup, Neftekhimiya, 2 (1962) 154. 191 L.K. Freidlin, Y.Y. Kaup, E.F. Litvin and T.I. Llomets, Dokl. Akad. Nauk SSSR, 1 4 3 (1962) 883. 192 R.L. Burwell, Chem. Rev., 57 (1957) 895. 1 9 3 K.N. Campbell and B.K. Campbell, Chem. Rev., 31 (1942) 77. 194 A.L. Henne and K.W. Greenlee, J. Am. Chem. SOC.,65 (1943) 2020. 195 N.A. Khan, J. Am. Chem. SOC., 74 (1952) 3018. 196 R. Romanet, Compt. Rend., 236 (1953) 1044,1677. 197 E.F. Meyer and R.L. Burwell, J. Am. Chem. SOC.,8 5 (1963) 2877. 198 H.G. Rushford and D.A. Whan, Trans. Faraday SOC., 67 (19711 3577. 199 N. Dinh-Nguyen and R. Rhyage, J. Res. Inst. Catal. Hokkaido Univ., 8 (1960) 73. 200 G. Webb and P.B. Wells, Trans. Faraday SOC.,61 (1965) 1232. 201 J.J. Phillipson, P.B. Wells and D.W. Gray, Proc. 3rd Int. Congr. Catal., NorthHolland Amsterdam, Vol. 2, 1964 p. 1250. 202 G.V. Smith, Proc. 3rd. Int. Congr. Catal., North-Holland, Amsterdam, Vol. 2, 1965, p. 1262. 203 J.B. Bream, D.C. Eaton and H.B. Herbert, J. Chem. SOC., (1957) 1974. 204 R.S. Mann and D.E. T o , Can. J. Chem. 46 (1968) 161. 205 R.S. Mann and D.E. Tiu, Can. J. Chem., 46 (1968) 3249. 206 R.S. Mann and A.M. Shah, Can. J. Chem., 50 (1972) 1973. 207 J. Grant, R.B. Moyes, R.G. Oliver and P.B. Wells, J. Catal., 42 (1976) 213; R.G. Oliver and P.B. Wells, Catalysis, 1 (1973) 659. 208 G.C. Bond, G. Webb, P.B. Wells and J.M. Winterbottom, J. Chem, SOC., (1965) 3218. 209 P.B. Wells and A.J. Bates, J. Chem. SOC. A, (1968) 3064. 210 B.J. Joice, J.J. Rooney, P.B. Wells and G.R. Wilson, Discuss. Faraday SOC., 41 (1966) 223. 211 D.A. Buchanan and G. Webb, J. Chem. SOC.Faraday Trans. I, 7 1 (1975) 134. 212 W.G. Young, R.L. Meier, J. Vinograd, H. Bollinger, L. Kaplan and S.L. Linden, J. Am. Chem. Soc., 69 (1947) 2046. 213 S.V. Lebed’ev and A.O. Yabubchik, J. Chem. Soc., (1928) 2190. 214 B.A. Kazanskii and N.I. Propova, Izv. Akad. Nauk SSSR Otd. Khim Nauk, (1952) 442. 215 P.B. Wells and G.R. Wilson, J. Chem. SOC.A, (1970) 2442. 216 R.B. Moyes and P.B. Wells, Adv. Catal., 2 3 (1974) 121. 217 J.L. Garnett and W.A. Sollich-Baumgartner, Adv. Catal., 1 6 (1966) 95. 218 J.L. Garnett, Catal. Rev., 5 (1972) 229. 219 A.W. Weitkamp, Adv. Catal., 18 (1968) 1. 220 S. Siege1 and G.V. Smith, J. Am. Chem. SOC.,8 2 (1960) 6082. 221 S. Siegel, Adv. Catal., 1 6 (1966) 124. 222 S. Siegel and G.V. Smith, J. Am. Chem. SOC.,8 2 (1960) 6087.
119 223 224 225 226 227 228 229 230 231 232 233 234 235 236 237 238 239 240 241 242 243 244 245 246 247 248 249 250 251 252 253 254 255 256 257 258 259 260 261 262 263 264 265 266 267
J.I. Sauvage, R.H. Baker and A.S. Hussey, J. Am. Chem. SOC.,8 3 (1961) 3874. S. Siegel and B. Dmuchovsky, J . Am. Chem. Soc., 8 4 (1962) 3132. A S . Hussey and G.P. Nowack, J. Org. Chem., 34 (1969) 439. S. Siegel and B. Dmuchovsky, J. Am. Chem. SOC.,8 6 (1964) 2192. L. Hilaire and F.G. Gault, J . Catal., 20 (1971) 267. M. Pecque and R. Maurel, J. Catal., 1 9 (1970) 360. F. van Rantwijk, A.P.G. Kieboom and H. van Bekkum, Catalysis: Heterogeneous and Homogeneous, Elsevier, Amsterdam, 1975, p. 53. S. Siegel and G.S. McCaleb, J. Am. Chem. SOC.,81 (1959) 3655. J.I. Sauvage, R.H. Baker and A.S. Hussey, J. Am. Chem. SOC.,8 2 (1960) 6090. S. Siegel, P.A. Thomas and J.T. Holt, J. Catal., 4 (1965) 73. S. Siegel, M. Dunkel, G.V. Smith, W. Halpern and J. Cozort, J. Org. Chem., 3 1 (1966) 2082. A.S. Hussey, G.W. Keulks, P.G. Norwack and R.H. Baker, J. Org. Chem., 33 (1968) 610. A.S. Hussey, R.H. Baker and G.W. Keulks, J. Catal., 1 0 (1968) 258. C. Lethuiller, D. Cornet and F.G. Gault, Bull. SOC.Chim. Fr., (1963) 1424. J. Erkelens, A.K. Galwey and C. Kemball, Proc. R. SOC.London, Ser. A, 260 (1961) 273. A.S. Hussey, T.A. Schenach and R.H. Baker, J. Org. Chem., 3 3 (1968) 3258. J.K.A. Clarke and J.J. Rooney, Adv. Catal., 25 (1976) 125. J.J. Phillipson and R.L. Burwell, J . Am. Chem. SOC., 9 2 (1970) 6125. R. Maurel and J. Tellier, Bull. SOC.Chim. Fr., (1968) 4191. R. Willstatter and J . Bruce, Chem. Ber., 40 (1907) 4456. G.C. Bond and J. Sheridan, Trans. Faraday SOC.,48 (1952) 713. G.C. Bond and J. Turkevich, Trans. Faraday SOC.,50 (1954) 1335. J. Addy and G.C. Bond, Trans. Faraday SOC.,5 3 (1957) 368, 383,388. K.E. Hayes and H.S. Taylor, Z. Phys. Chem. (Frankfurt), 1 5 (1958) 127. J.E. Benson and T. Kwan, J. Phys. Chem., 60 (1956) 1601. G.C. Bond and J . Newham, Trans. Faraday SOC.,56 (1960) 1851. J.H. Sinfelt, D.J.C. Yates and W.F. Taylor, J . Phys. Chem., 69 (1965) 1877. W.F. Taylor, D.J.C. Yates and J.H. Sinfelt, J. Catal., 4 (1965) 374. R.A. Dalla-Betta, J.A. Cusumano and J.H. Sinfelt, J. Catal., 19 (1970) 343. H.F. Wallace and K.E. Hayes, J . Catal., 1 9 (1970) 343; 1 9 (1973) 83. R. Merta and V. Ponec, Proc. 4th Int. Congr. Catal., Akademiai Kiado, Budapest, Vol. 2, 1973, p. 53. R. Merta and V. Ponec, J. Catal., 17 (1970) 79. D.W. McKee, J. Phys. Chem., 67 (1963) 1336. T.S. Sridhar and D.M. Ruthven, J. Catal., 1 6 (1970) 363; 24 (1972) 153. J.R. Anderson and N.R. Avery, J. Catal., 8 (1967) 48. E.F. Ullman, J. Am. Chem. SOC.,81 (1959) 5386. G.C. Bond, Catalysis by Metals, Academic Press, New York, 1962, p. 276. H.S. Taylor, Proc. R. SOC.London, Ser. A, 108 (1925) 105. A.A. Balandin, Zh. Fiz. Khim. Obshchestva, Chast’Khim., 61 (1929) 909; Z. Phys. Chem., Abt. B, 2 (1929) 289; Adv. Catal., 1 0 (1958) 96; 19 (1969) 1. C.R. Adarns, H.A. Benesi, R.M. Curtis and R.G. Meisenheimer, J. Catal., 1 (1962) 336. T.A. Dorling and R.L. Moss, J. Catal., 7 (1967) 378. J.M. Thomas, Adv. Catal., 1 9 (1969) 293. R.A. van Nordstrand, A.J. Lincoln and A. Carnevale, Anal. Chem., 36 (1964) 819. G.K. Boreskov and A.P. Karnaukhov, Zh. Fiz. Khim., 26 (1952) 1814. H.L. Gruber, J. Phys. Chem., 6 6 (1962) 48.
120 268 T.R. Hughes, R.J. Houston and R.P. Sieg, Ind. Eng. Chem., Process Des. Dev., 1 (1962) 96. 269 J.E. Benson and M. Boudart, J. Catal., 4 (1965) 704. 270 L.H. Germer, Adv. Catal., 13 (1962) 192. 271 N.E. Farnsworth, Adv. Catal., 15 (1964) 31. 272 S.J. Thomson, in C. Kemball (Ed.), Catalysis, Specialist Periodical Reports, The Chemical Society, London, 1977, in press. 273 R.W. Joyner and M.W. Roberts, in M.W. Roberts and J.M. Thomas (Eds.), Surface
274 275 276 277 278 279 280 281
283 284 285 286 287 288
and Defect Properties of Solids, Specialist Periodical Reports, The Chemical Socie t y , London, 1975, p. 68. G. Dalmai-Imelik and J. Massardier, Proc. 6th Int. Congr. Catal., London, 1976, Paper A l . G. Dalmai-Imelik and J.C. Bertolini, J. Vac. Sci. Techno]., 270C (1970) 1079. G. Dalmai-Imelik, J. Rousseau and J.C. Bertolini, Avires 2, Colloq. Int. Appl. Sci. Techno]. Vide Revetements Etats Surface, C.R. 2nd. (1972) 36. G. Dalmai-Imelik and J.C. Bertolini, J . Appl. Phys. Jpn., 2 (1974) 205. O.M. Poltorak and V.S.Boronin, Russ. J. Phys. Chem., 39 (1965) 1329. O.M. Poltorak and V.S. Boronin, Russ. J. Phys. Chem., 40 (1966) 1436. O.M. Poltorak, V.S. Boronin and A.N. Mitrofanova, Proc. 4th Int. Congr. Catal., Vol. 2, Akademiai Kiado, Budapest, 1971, p. 276. R. van Hardeveld and A. van Montfoort, Surf. Sci., 4 (1966) 396; 17 (1969) 90. R. van Hardeveld and F. Hartog, Proc. 4th Int. Congr. Catal., Vol. 2, Akademiai Kiado, Budapest, 1971, p. 295;Adv. Catal., 22 (1972) 75. G.C. Bond, Proc. rth Int. Congr. Catal., Vol. 2, Akademiai Kiado, Budapest, 1971, p . 266. M.R. Hoare and P. Pal, Nature (London), 236 (1972) 35. J.J. Burton, Catal. Rev., 9 (1974) 209. O.M. Poltorak and V.S. Boronin, Russ. J. Phys. Chem., 39 (1965) 781. T.A. Dorling and R.L. Moss, J . Catal., 5 (1966) 111. M. Boudart, A. Aldag, J.E. Benson, N.A. Dougharty and C.G. Harkins, J. Catal., 6 (1969) 92.
M. Boudart, Adv. Catal., 20 (1969) 153. C.D. Ljubarskii, Proc. 2nd. Int. Congr. Catal., Technip, Paris, 1961, p. 1559. V. Nikolajenko, V. Bosacek and V. Danes, J. Catal., 2 (1963) 127. V. Nikolajenko, V. Danes and M. Krivanek, Kinet. Katal., 7 (1966) 816. R.G. Oliver and P.B. Wells, see ref. 17 in P.B. Wells, in M.W. Roberts and J.M. Thomas (Eds.), Surface and Defect Properties of Solids, Specialist Periodical Reports, The Chemical Society, London, 1971, p. 236. 294 G.C. Bond, Discuss. Faraday SOC.,4 1 (1966) 200. 295 0. Kahn and L. Salem, Proc. 6 th Int. Congr. Catal., Vol. 1, The Chemical Society, London, 1977, p. 101. 296 Y. Izumi, M. Imaida, H. Fukawa and S. Akabori, Bull. Chem. SOC.Jpn., 36 289 290 291 292 293
(1963) 21. 297 Y. Izumi, Angew. Chem. Int. Ed. Engl., 1 0 (1971) 871. 298 T. Harada, Y. Hiraki, Y. Izumi, J. Muraoka, H. Ozaki and A. Tai, Proc. 6th Int. Congr. Catal., London, 1976, Paper B41. 299 J.A. Groenewegen and W.M.H. Sachtler, Proc. 6th Int. Congr. Catal., Vol. 2, The Chemical Society, London, 1977, p. 1014. 300 N.I. Kobosev, Zh. Fiz, Khim., 1 3 (1939) 1; Russ. J. Phys. Chem., 19 (1945) 48, 142; 2 1 (1947) 1413; Usp. Khim., 25 (1956) 545. 301 E.B. Prestridge and D.J.C. Yates, Nature (London) 234 (1971) 345. 302 J.P. Collman, L.S. Hegedus, M.P. Cooke, J.R. Norton, G. Dolcetti, and D.N. Marquardt, J. Am. Chem. SOC.,94 (1972) 1789. 303 J. Robertson and G. Webb, Proc. R. SOC.London, Ser. A, 342 (1974) 383; see also J. Robertson, Ph.D. Thesis, University of Glasgow, 1974.
121 304 305 306 307 308 309 310 311 312 313 314 315 316 317 318 319 320 321 322 323 324 325 326 327 328 329 330 331 332 333 334 335 336 337 338 339 340 341 342 343 344 345 346 347 348 349
G.E. Morgan and G. Webb, unpublished results. M. Boudart, J. Am. Chem. SOC.,72 (1950) 1040. L. Pauling, Proc. R. SOC.London, Ser. A, 1 9 6 (1949) 343. See e.g. N.F. Mott and H. Jones, Theory of t h e Properties of Metals and Alloys, Oxford University Press, Oxford, 1936. D.A. Dowden, J. Chem. SOC.,(1950) 242. D.A. Dowden and P.W. Reynolds, Discuss. Faraday SOC.,8 (1950) 184. D.D. Eley, J. Res. Inst. Catal., Hokkaido Univ., 1 6 (1968) 101. G.C. Bond and R.S. Mann, J. Chem. SOC.,(1959) 3566. W.K. Hall and P.H. Emmet, J . Phys. Chem., 6 3 (1959) 1102. G. Reinacker and E.A. Bommer, Z. Anorg. Chem. 242 (1939) 302. G. Reinacker, E. Muller and R. Burmann, Z. Anorg. Chem., 251 (1943) 55. G. Reinacker and E.A. Bommer, Z. anorg. Chem., 236 (1938) 263. R.L. Moss and L. 'A'halley, Adv. Catal., 22 (1972) 115. A. Couper and D.D. EIey, Discuss, Faraday SOC.,8 (1950) 172. D.D. Eley and P. Luctic, Trans. Faraday SOC.,5 3 (1957) 1483. A.G. Daglish and D.D. Eley, Proc. 2nd Int. Congr. Catal., Technip, Paris, 1960, p. 1615. R.J. Best and W.W. Russell, J. Am. Chem. SOC.,7 6 (1954) 838. W.M.H. Sachtler and G. Dorgelo, J. Catal., 4 (1965) 654. W.M.H. Sachtler and R. Jongepier, J. Catal., 4 (1965) 665. E.G. Allison and G.C. Bond, Catal., Rev., 7 (1973) 233. W.M.H. Sachtler and P. van der Plank, Surf. Sci., 18 (1969) 62. J.K;A. Clarke and J.J. Byrne, Nature (London), 214 (1967) 1109. P.F. Carr and J.K.A. Clarke, J. Chem. SOC.A, (1971) 985. S.J. Thomson and G.A. Harvey, J . Catal., 22 (1971) 359. P. Dini, D. Dones, S. Montelatchi and N. Giordano, J. Catal., 30 (1973). F. Figueras, R. Gomez and M. Primet, Adv. Chem. Ser., 1 2 1 (1973) 480. J.H. Sinfelt and P.J. Lucchesi, J. Am. Chem. SOC.,8 5 (1963) 3365. P.A. Compagnon, C. Hoang-Van and S.J. Teichner, Proc. 6th Int. Congr. Catal., Vol. 1, The Chemical Society, London, 1977, p. 117. M. Boudart, M.A. Vannice and J.E. Benson, Z. Phys. Chem., 64 (1969) 171. P.A. Sermon and G.C. Bond, Catal. Rev., 8 (1974) 211. R.B. Levy and M. Boudart, J. Catal., 32 (1974) 304. K.M. Sancier, J . Catal., 20 (1971) 1 0 6 ; 23 (1971) 404. M.A. Vannice and W.C. Neikam, J. Catal., 2 3 (1971) 401. G.E.E. Gardes, G.M. Pajonk and S.J. Teichner, J. Catal., 3 3 (1974) 145. G.M. Pajonk and S.J. Teichner, Bull. SOC. Chim. Fr., (1971) 3847. J.E. Benson, H.W. Kohn and M. Boudart, J. Catal., 5 (1966) 307. W.C. Neikam and M.A. Vannice, Catalysis, 2 (1973) 609. J.C. Schlatter and M. Boudart, J. Catal., 24 (1972) 482. D. Briggs and J. Dewing, J . Catal., 28 (1973) 338. P.A. Sermon, G.C. Bond and G. Webb, J. Chem. SOC.Chem. Commun., (1974) 417. J.W. Bassi, F.W. Lytle and G. Parravano, J . Catal., 42 (1976) 139. Z. Paal and S.J. Thomson, J. Catal., 30 (1973) 96. L.V. Babenkova, N.M. Popova, D.V. Sokol'skii and V. Solnyserkova, Dokl. Akad. Nauk SSSR, 210 (1973) 888. J.K.A. Clarke, Chem. Rev., 7 5 (1975) 291. N.C. Gardner and R.S. Hansen, J. Phys. Chem., 74 (1970) 3298. R.P. Dessing, V. Ponec and W.M.H. Sachtler, J. Chem. SOC. Chem. Commun., (1972) 880.
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Chapter 2
123
Heterogeneous Oxidation Processes K. VAN DER WIELE and P.J. VAN DEN BERG
1. Scope of the chapter Many substances can be partially oxidized by oxygen if selective catalysts are used. In such a way, oxygen can be introduced in hydrocarbons such as olefins and aromatics t o synthesize aldehydes (e.g. acrolein and benzaldehyde) and acids (e.g. acrylic acid, phthalic acid anhydride). A selective oxidation can also result in a dehydrogenation (butene butadiene) or a dealkylation (toluene + benzene). Other molecules can also be selectively attacked by oxygen. Methanol is oxidized to formaldehyde and ammonia to nitrogen oxides. Olefins and aromatics can be oxidized with oxygen together with ammonia t o nitriles (ammoxidation). The Gibbs free energy change for total combustion of such molecules always has a Iarger negative value than partial oxidation. Hence a catalyst t o control kinetically the oxidation process is an absolute necessity. By choosing the right conditions and the proper type of catalyst, the oxidation process can be directed towards intermediates which d o not react further. Because of the rapid developments in the field of heterogeneous catalysis, the material reviewed here is exclusively dedicated to selective oxidations. No attention is given t o total oxidations or combustion processes (including the problem of automotive exhaust gases). There is one exception, however; the oxidation of sulphur dioxide to trioxide. Work on vanadate catalysts for this reaction is close t o research on selective catalysts and therefore included. The work considered here mainly covers the period 1970-1976 (inclusive) because several reviews appeared before that time. Recent articles for special mention are the book by Hucknall [160] and the review articles by Voge and Adams [343], Germain [134], Butt [67] and Krabetz [181]. Of the available literature, articles have been reviewed which were written or translated into English, German, French, Spanish and Italian. Untranslated publications in Russian or Japanese have not been taken into account. Patents were not considered either because, although they may be relevant for catalytic innovations, they seldom contain kinetic data. f
References p p . 2 5 3 - 2 6 2
2. Oxidation processes Substrates which can undergo partial oxidation are characterized by a n-electron system or unshared electrons: olefins and aromatics contain the first, methanol, ammonia and sulphur dioxide the second. Alkanes do not contain such electrons. Their selective oxidation appears t o demand (thermal or catalytic) dehydrogenation to alkenes as the initial process. In the group of olefins, ethylene takes a special position because it is the only one which can be oxidized with a high efficiency to an epoxide. Only silver is known as a selective catalyst for this purpose. A special mechanism applies in the way oxygen is introduced into the molecule. A second possibility for oxidation is the conversion t o acetaldehyde and acetic acid. This presents a strong analogy with the homogeneous oxidations (Wacker process) and true heterogeneous catalysts containing palladium have been found. Propene can only be oxidized to an epoxide with very low yields. Much more interesting is an allylic oxidation in which propene is oxidized via an allylic intermediate resulting in acrolein, acrylic acid, acrylonitrile and others. Such processes are made possible by catalysts which often consist of combinations of metal oxides, e.g. Bi203-Mo03. With these catalysts, n-butenes and iso-butene can also be converted. Iso-butene can be compared with propene and accordingly can be oxidized to methacrolein. n-Butenes show more possibilities. The relative stability of the resonance structure of butadiene makes it possible t o dehydrogenate n-butenes in a selective oxidation. Several metal oxide combinations known from the propene conversion can be used. It is also possible to increase the extent of oxidation by choosing the right compositions. Then butenes (and butadiene) are selectively oxidized to maleic acid anhydride and other compounds, vanadates being specially suitable. Such compositions are also effective in aromatic oxidations. Either the ring or the side chains are attacked giving rise to phthalic acid anhydride from naphthalene as well as from o-xylene. Several other examples are given. In the partial oxidation of hydrocarbons, the molecules are converted stepwise. Germain [134] remarks that a mechanism with the model of a rake applies Gas Surface
A
B
C
11
I1
11
A
+
B
-,
C
D 11 +
D
The degradation of A on the surface to B is the first step. B can then desorb or react further and so on. The last steps in the rake are oxidation t o combustion products and desorption of carbon oxides and water. It will be clear that the selectivity in each step is determined by the rate of reaction and rate of desorption. If some oxidation steps are very rapid, the rake
125 model will result in a combination of consecutive and parallel reactions
For a large group of metal oxide catalysts, it has been proved that a redox mechanism occurs, as originally proposed by Mars and van Krevelen [ 204 J . The oxidation of the hydrocarbons, methanol, etc. is effected by oxygen supplied by the catalyst, very often by oxygen contained in the crystal lattice; a vacancy results which is reoxidized by oxygen from the gas phase, viz.
A + MeO,,
-+
hleO,-,
+ 0,
0
A 0 + Meon-, +
0
MeO,
The replenishment of the vacancy can be directly from the gas phase or indirectly from the catalyst. In the latter case, the oxygen mobility within the catalyst is so large that bulk oxygen can diffuse t o the vacancy. Then oxygen from the gas phase reoxidizes the lattice on sites which differ from hydrocarbon reaction sites. In a steady state, the rate of catalyst oxidation will be equal t o the rate of reduction by the substrate. The steady state degree of reduction, equivalent t o the surface coverage with oxygen, is determined by the ratio of these two rates. Kinetic models based on these principles are called redox models, for which the simplest mathematical expression is
in which Rhc is the rate of reduction of the oxidized catalyst by the hydrocarbon and R,, is the rate of reoxidation of vacancies by oxygen [ 2041. The redox mechanism applies not only t o allylic oxidation of olefins and t o the oxidation of aromatic hydrocarbons, but also t o the oxidation of methanol and sulphur dioxide, as well as the oxidation of ammonia t o nitrogen. Only in the case of ethylene oxidation and oxyhydration of olefins d o catalysts act according to another mechanism. The latter processes seem t o be always low temperature reactions, occurring below 300" C, whereas redox mechanisms are possible above this temperature (e.g. 400--500°C). A large number of authors describe the oxidation kinetics by Langmuir-Hinshelwood type models. Depending on the particular L-H model selected, the mathematical difference between L-H models and redox models can be very small, although the former always contains more References p p . 253-262
parameters. It must be emphasized, however, that L-H models based on single site adsorption (i.e. competitive adsorption) of the reactants (hydrocarbon and oxygen) are incompatible with a redox mechanism; it would imply that a completely oxidized catalyst is inactive. Langmuir-Hinshelwood models have the advantage that inhibition by reactants and products is easily included in the rate equations. Redox models, on the other hand, have the advantage that reaction orders, varying with respect t o both oxygen and hydrocarbon, according t o the feed composition, can be described with a minimum number of parameters. Furthermore, extended redox models, which include the essential parallel and consecutive side reactions, automatically account for inhibition effects that are caused by competition of the reactants for the same catalyst oxygen: a faster reacting component inhibits the reaction of a slower one because it increases the steady state degree of reduction and hence reduces the oxidizing capacity of the catalyst for competing molecules. Examples of the use of extended redox models are provided in the literature by Boag e t al. [ 551 and by authors of this chapter [ 3481. Many investigators use pulse techniques in which a catalyst reacts with hydrocarbons, oxygen etc. separately in time. This can provide an insight into the nature and significance of the individual reaction and sorption steps, but it should be emphasized that selectivities and other data may be unrepresentative for conditions in a flow reactor. In particular, selectivities may be considerably lower under steady state conditions. If the selectivity differences between pulse and flow experiments are very large, a cyclic mode of operation may be attractive for the practical application of the catalyst concerned. Oxidation and reduction are then separated. 2.1 ETHYLENE OXIDATION
Two selective processes are important in the oxidation of ethylene: the production of ethylene oxide and acetaldehyde. The first process is'specifically catalyzed by silver, the second one by palladium-based catalysts. Silver catalysts are unique and selective for the oxidation of ethylene. No similar situation exists for higher olefins. The effect of palladium catalysts shows a resemblance t o the liquid phase oxidation of ethylene in the Wacker process, in which Pd-C, H4 coordination complexes are involved. The high selectivity of the liquid phase process (95%), however, is not matched by the gas phase route a t present.
2.1.1 Ethylene oxide production In the industrially important epoxidation of ethylene, the main byproducts are carbon dioxide and water. These are formed by parallel combustion of ethylene as well as of ethylene oxide according t o the reaction
127
scheme (>C2H40 C,H4 1(3) (z+CO2 + H 2 0 Yields of 65-75 mol% at 230-270°C and 10-20 atm total pressure are reported with commercial catalysts. These consist of supported silver doped with alkali and earth alkali metals. High selectivities are only obtained if very small amounts (ppm range) of moderator (e.g. chloroalkanes) are added to the reactants. (a) Kinetics
The kinetics of the ethylene oxidation are rather complicated as they depend not only on ethylene and oxygen pressure but also on the concentration of the reaction products. These influence the rate by adsorption competition with the reactants. Moreover, different forms of adsorbed oxygen may occur on the catalyst surface. Consequently, the rate equations proposed in the literature consist of either Langmuir-Hinshelwood and Eley-Rideal types or power rate models with non-integer coefficients. Power rate models are less appropriate as their coefficients inevitably depend on the reaction conditions. According t o the oxidation scheme, three reactions have t o be accounted for. Although the combustion of ethylene oxide is less important than the direct combustion of ethylene, it cannot be neglected. A summary of recently reported rate equations is given in Table 1. Several of them will be discussed in more detail. Ayame et al. [32] studied the effect of higher total pressures, up to 11 atm. In the range 200-265"C, data are given for conversions t o a maximum of 46%. Maximum selectivity is 7 1.7%. Neglecting some inhibition effects, the authors find the power rate equations, which include the total pressure as an individual parameter, as presented in Table 1. Numerical values of the rate coefficients in these equations as a function of temperature are collected in Table 2. Increasing the pressure results in a higher ethylene oxide yield. Ionov et al. [164] confirmed this effect. They reached a selectivity of 78% with a mixture containing 3 vol. % C z H 4 , 20.4% O2 at 240°C and a pressure of 11atm. Their catalyst was promoted with selenium. Metcalf and Harriott [ 2161 reported a maximum in the reaction rate as a function of ethylene pressure, and accordingly their rate equation was based on competitive adsorption of the reactants. Verma and Kaliaguine [ 3391 measured the reaction rate using the pulse technique. They found good agreement with the data of Table 2, obtained for stationary conditions. The authors gave a clear presentation of advantages and disadvantages of the pulse method. R e f e r e n c e s p p . 253--262
128 TABLE 1 Rate equations in ethylene oxidation -. ~
_Alfani and Carberry Ayarne e t al.
Reaction conditions
Rate equations a
Ref.
Authors
__
18
R I = 2.2 X = 5X R 3 = 1.3 X
lo-'
R2
32
lo-*
exp(-7,400/RT) P E ~ " exp(-8,450/RT) PE"~' exp(--11,500/RT) PEO
R 1 = k 1 PTOTo.33 xEXoo.' R2 = k z PT0T0.32X E X O R3 = k 3 P ~ oX ~ E ~~ X' ~~ , ~ O - ~ ' ~
P ~ P 0 ' ~ ~ / [ ( 0 . 0+103. 2 3 6 P ~ PEPo"'(1 + 0 . 6 6 1 f l o ) * ]
250-300" C 1 atm 200-265' C -11 a t m
+ 0.121 \@O)'
220'C 1 atm
R'
1 atrn
Klugherz and Harriott
178
Metcalf and Harriott
216
Spath and Handel
300
Several rate equations based o n different mechanisms. Most probable: dual site
300" C 1 atm
Verma and Kaliaguine
339
R1 = 0.00044 P)?-O.O3 Poo." R2 = 0.00015 PE-o.2 Po1.' R3 = 0 (included in R 2 )
221" c 1atm
Rl
= aoPEPo'.5/(1 + aIPoo.5 + a2Po + + a3~01.5 + a4PE + asPEPoo.S+
+ a6Pco,Poo.S + a7PH,oP00.5)2
a
Indices 1 , 2 and 3 correspond to t h e numbering in the reaction scheme o n p. 1 2 7 ; E = ethylene, 0 = oxygen; EO = ethylene oxide.
Kenson and Lapkin [173] measured the rate of isomerization of the epoxide t o acetaldehyde on a supported silver catalyst and found the rate of epoxide conversion: R E o = 3.9 X PEOat 200°C with an activation energy of 9.8 kcal mol-I . Acetaldehyde may be an intermediate in the combustion of ethylene oxide, but, as such, it is unimportant because it is rapidly oxidized. ( b ) Mechanism
The uniqueness of silver as an epoxidation catalyst has continuously attracted interest and several valuable articles have appeared since the TABLE 2 Effect of temperature o n rate coefficients
T
k l x 10'
k2 X
1.97 4.93 8.76
1.10 5.88 12.8
10'
k, X
("C) 200 240 265
1.37 0.90 1.38
lo3
129 excellent review by Voge and Adams [343]. Although discussion of the mechanism has not ended, some facts are now well established, and the alternatives have become clear. It is clear from the kinetics that both ethylene and oxygen adsorption are important since both compounds appear in the rate equations with non-zero orders. Moreover, it is well known that ethylene is not adsorbed on pure silver, but that it does adsorb on a surface that is partially covered with oxygen. This implies that ethylene is either adsorbed on t o p of pre-adsorbed oxygen or on silver sites that are activated by the presence of oxygen (i.e. by formation of surface oxides, or another form of electron transfer or polarization). Consequently, two different mechanisms arise for the formation of ethylene oxide. The (direct) combustion of ethylene is another point of discussion. Although many favour the idea that different oxygen species are involved, others assume the same oxygen species, but different forms of ethylene adsorption. Preceding the discussion of the two proposed mechanisms of ethylene oxide formation, the different forms of oxygen and their suggested roles in the oxidation process will be reviewed briefly.
(i) Forms of oxygen on the surface. It is clear from the thermodynamics that an oxide of silver, e.g. Ag,O, does not exist as a separate phase under reaction conditions. However, Clarkson and Cirillo [ 881 conclude from electron diffraction measurements that a layer of Ag20 is built up under reaction conditions. The authors suggest that such a A g 2 0 layer has semiconductor properties, facilitating electron transfer from the silver bulk to the silver oxide surface. The situation may be even more complex, as indicated by the work of Temkin and Kulkova [315] who conclude that Ag,03 is formed. At high partial pressures and high temperatures, oxygen may even dissolve sub-surface, and form non-stoichiometric silver oxides. Sato and Seo [277] demonstrated that a silver catalyst continuously emits low-energy electrons. This emission is chemically stimulated when the catalyst produces ethylene oxide. There is a strong correlation between the production rate and the emission rate. The emitting layer is continuously renewed and is apparently silver oxide. Chemisorption measurements have shown that ethylene does not adsorb on pure silver, but only on a silver surface which has been preoxidized [ 3391. Complete coverage with an oxygen monolayer, however, also seems to destroy the capacity t o adsorb ethylene, as was demonstrated by Force and Bell [114,116] (favouring the idea of adsorption on silver). Consequently, partial oxygen coverage seems t o be a necessary condition for catalytic activity. Several studies have been devoted to adsorption of oxygen on silver surfaces. It has become evident that both molecular and atomic species of adsorbed oxygen exist. It is generally considered that fast dissociative adsorption takes place first and leads t o a strongly bonded species, which References PP. 2 5 3 - - 2 6 2
130
is probably 0 2 -although , 0- is not excluded. Kilty et al. [ 1761 measured a very low activation energy for this step ( 3 kcal mol-') at 95°C. Further adsorption leads to a layer of oxygen more loosely bound. This is molecular oxygen as has been proved by Clarkson and Cirillo [ 87,881, Kilty et al. [176,177] and Force and Bell [114,116] using infrared spectroscopic and electron spin resonance methods. The activation energy of this step is 7.9 kcal mol-' . Kilty et al. [ 1761 demonstrated that, at temperatures above 150°C, a third chemisorption occurs with an activation energy of 14 kcal mol-' . This adsorption is again dissociative. It is generally assumed that oxygen adsorbed in a molecular form is 0;. The orientation on the surface could be 0
I
0--0
0
Ag Ag
Ag
The 0; species is confirmed by Herzog [ 1561 and Trifiro [ 3261. Trifiro used a gas chromatographic technique and found two types of oxygen (molecular and atomic). Electron spin resonance spectra also prove that molecular oxygen exists as 0; at the surface [ 87,88,291,314]. The occurrence of 0'- and 0- is not excluded, however. Shimizu et al. [291] detected, for the first time, Agz+with ESR techniques and postulated the existence of 0'-. Working with LEED, Bradshaw et al. [60] favoured the hypothesis that 0- is an active species. Indirect evidence for the presence of 0- at high temperatures (220°C) was obtained from IR spectra by Force and Bell [116]. It is interesting that preadsorption of chlorine, which is a well known moderator, completely destroys activity when it covers 25% of the surface. Kilty et al. [176] suggested that the first adsorption step is coupled to an ensemble of four adjacent silver atoms, viz.
O 2+ 4 Ag(adj.) +.2 02-(ads.) + 4 Ag'(adj.) (ii) Role o f the different oxygen species. Many authors assume that different species of oxygen account for selective oxidation and combustion. Molecular oxygen is believed to take part in the epoxidation while a more strongly bonded form of atomic oxygen is responsible for combustion. Evidence indicating these different roles can be taken from a thermal desorption investigation by Tanaka and Yamashina [ 3141, which indicates that 0;reacts more easily with ethylene than with 0'-.Mikami et al. [218] show, with pulse techniques, that more loosely bound oxygen is responsible for production of the epoxide and more fixed oxygen for combustion. Very recently, Prauser et al. [258] have used these tech-
131 niques in the temperature range 200--320°C. They confirm the existence of an atomic oxygen species, which is irreversibly adsorbed, at low coverage of the catalyst surface. At higher coverage, the second type of oxygen is adsorbed. The temperature influences the ratio of atomic t o molecular oxygen. At higher temperatures, there is relatively more atomic oxygen leading to lower selectivity. The diatomic species not only give ethylene oxide but also some total oxidation. The atomic species can arise by adsorption from the gas phase but it can also have an origin in the formation of ethylene oxide, as suggested by Worbs in 1942 (see ref. 84 in Voge and Adams [343]), viz.
6 C,H, + 6 O 2 = 6 C,H,O + 6 0 6 0 + C2Hs
=
2 CO, + 2 H2O
Excluding the effects of recombination of oxygen atoms and combustion of ethylene oxide, the maximum selectivity in such a coupled system is 85.7%. Such a high value is never reached, which is understandable because ethylene oxide can oxidize further, a process which is also catalyzed by silver. The different roles of atomic and molecular oxygen were also suggested by Herzog [ 1561, who demonstrated that N 2 0 in small concentrations is a precursor of atomic oxygen leading t o combustion. At high N 2 0 concentrations, the formation of molecular oxygen is favoured and an enhanced production of the epoxide is observed. It is not necessarily true that different reactions are caused by distinct oxygen species. Force and Bell [114-1161 take the opposite view, based on their combined infrared and kinetic investigations. They suppose that only the more labile 0 - is responsible for all oxidations.
(iii) The mechanism of ethylene oxide formation. Compelling evidence for an Eley-Rideal type of mechanism, in which ethylene is adsorbed on preadsorbed oxygen, is presented by Kilty e t al. [ 1761. The authors adsorbed ethylene on a silver surface at 95°C after pre-adsorption of oxygen. A band at 870 cm-' was found in the IR spectra and assigned to a peroxidic structure
The spectra show stretching and bending frequencies which are in accordance with such a species. By heating the sample, the spectra change and show a character which can also be observed when ethylene oxide is adsorbed on silver. The authors show that ''0 leads to a shift in the 8 7 0 cm-' band and it could be proved that the oxygen atoms in the peroxidic References p p . 2 5 3 - 2 6 2
132 structure stem from the same original oxygen molecule. Kilty and Sachtler [ 1771 proposed that the Eley-Rideal proceeds as follows. H,C
=
CH,
7(6') 0 (6-1
H,C - CHz \ /
0'
I
0-
-
mechanism
HZC - CH, \ /
0
0
Ag Ag Ag The oxygen molecule is polarized during adsorption and the partial positive charge will interact with the nucleophilic ethylene. There is, however, one important objection against the attractive theory of Kilty et al. Their infrared work was carried out a t a relatively low temperature. It is very likely that the peroxidic structure is not sufficiently stable a t the usual reaction temperatures (200-300°C). Noteworthy evidence for a Langmuir-Hinshelwood mechanism is provided by Force and Bell [ 114,1161. Apart from demonstrating that ethylene does not adsorb on silver completely covered with oxygen, the authors show by IR analysis at 220°C that the adsorbed complex has a structure which is analogous to ethylene, coordinated with isolated silver ions, viz. 0-
CHI
=
CH,
0-
The authors believe that initial adsorption of oxygen takes place on uncharged silver, resulting in negatively charged oxygen ions. The counterbalancing positive charge will distribute between several neighbouring silver atoms. The resulting unoccupied silver sites with a partial charge (as well as uncharged sites) contribute t o subsequent adsorption of ethylene. A Langmuir-Hinshelwood type of mechanism is also supported by the work of Klugherz and Harriott [ 1781 and Spath and Handel [ 3001. These authors, and several others, advocate 0; (and 0 2 - )as the active oxygen species. Carberry et al. [74] and Forzatti et al. [118]contributed t o the evidence for 0; with their results on the effect of promotors. ?-Irradiation of a calcium-promoted silver catalyst enhanced the yield of ethylene oxide. It could be shown that, during irradiation, calcium migrated t o the surface, increasing the 0; concentration there. Force and Bell propose a mechanism based only on 0-. They suggest that 0- can react in two ways
The first reaction complex gives ethylene oxide, the second, eventually, carbon dioxide (via acetaldehyde).
133
It is also possible to formulate these reactions as dual site mechanisms, instead of the single site type given above, viz.
The authors prefer reaction (a) as the main route to ethylene oxide and reaction (d) t o combustion products. Nevertheless, they find that selectivity is not a strong function of temperature, suggesting that there is a common rate-determining step. Recently, they studied the effect of dichloroethane as a moderator and concluded that such a compound decreases the activity because silver sites are occupied by chloride ions. But the selectivity is higher because the single site mechanism is first order in surface species (epoxidation) and the dual site mechanism second order (combustion). Studying the effect of palladium as a promotor in silver catalysts, Cormack et al. [go] found that increasing amounts of Pd alloyed with silver drastically decreased the selectivity. No other partial oxidation products were found. 2.1.2 Acetaldehyde and acetic acid production The oxidation of ethylene t o acetaldehyde in the gas phase, carried out at rather low temperatures (100-2OO0C), is very similar to the Wacker liquid phase process. One of the main steps in this process is PdC12 + CZH, + H,O
=
CH3CHO + Pd + 2 HC1
and is followed by a copper-catalyzed oxidative regeneration Pd + 2 HC1+ f 0, = PdC12 + H,O The reaction is carried out at temperatures in the 125--130°C range. At the same low temperatures, Gerberich et al. [126] found that a palladium sponge catalyst converts ethylene in the gas phase into partial oxidation products with a selectivity of 30%. The product stream contains acetic acid, acetic anhydride and acetaldehyde. The analogy t o the Wacker process was even stronger for the results of Fujimoto et al. [119, 1201, who found that, at about lOO"C, palladium salts on active charcoal show a high selectivity t o acetaldehyde. The activity is markedly increased by steam showing a third-order dependence in PHz . It could very well be that capillary condenstion in the catalyst pores make it a real Wacker system if reoxidation of palladium is catalyzed by impurities in the charcoal. ?-Alumina and silica gel are not effective carriers. With charcoal, a selectivity of 98% t o acetaldehyde production is reached at 59.9% conversion with a temperature of 105°C and a ratio CzH4: O2 : H 2 0 = 5 : 2 : 15. The reaction time is long ( W / F = 2 5 g h mol-'). The activation energy is -18 kcal mol-'. Decrease of the reaction rate on raising the temperature can References p p . 253-262
134 be explained by the strong influence of steam, the concentration of which decreases with rising temperature. Gerberich e t al. [126] found an activation energy of 20 k 2 kcal mol-' for both complete and partial oxidation from initial rate data using a recirculation reactor, and 30 f 2 kcal mol-' from steady-state measurements. Apparently, the reaction products partly poison the catalyst. The reaction is first order in oxygen and first order in ethylene at low pressures, changing into negative first order in ethylene at higher pressures. This suggests that the catalyst is heavily covered by C2H4. It was not possible t o fit the rate data t o a Langmuir-Hinshelwood model. By alloying Pd with gold, it was found that the activity continuously decreased with increasingly concentrations of Au. The selectivity, however, went through a maximum of 50% near 80% Au, which is explained by the suppression of a dehydrogenating function of Pd on the olefin. The suppression is caused by filling of d-orbitals by gold. Other noble metals from the palladium group, e.g. Rh and Ru, only catalyze complete oxidation [ 72,731. Evnin et al. 11061 have found a true heterogeneous analogue of the Wacker system by combining palladium and vanadium oxide in a catalyst. At temperatures of 120--14O0C, 80% aldehyde and 12%acetic acid are formed a t a conversion of 20%. The concentration of Pd can be as low as 0.1 wt. %. The authors observe that the reducibility of V z 0 5 is increased by Pd and state that the fact that only Pd shows the particular effect is an indication that a complex Pd(C,E4) and not activated oxygen plays a role. The reoxidation of the catalyst is very fast. In accordance with a redox mechanism, TiOz, A1203 and SiOz are much less effective than
v
q / C
> 10
1
2 3 Residence time ( h )
2
1 ____.c_
3
Residence time (h)
Fig. 1 . Product yields from ethylene oxidation for Pd/V205 and pure V205. - - - - - -, 0.2%Pd/V205 ;,vzo5 ' Fig. 2 . Ethylene conversion for Pd/V205,pure V z 0 5 and Pd/A1203.
135 V 2 0 5 [243]. The kinetics are diffusion-controlled. Seoane e t al. [287] confirmed this result using 0.2 wt. 5% Pd on V z 0 5 at 230°C. Figures 1 and 2 show some typical data. Cant and Hall [72] report that the reaction rate with Pd on SiOz shows a first-order dependence on oxygen. The rate is strongly suppressed by olefin. 2.2 PROPENE OXIDATION
The oxidation of propene is at present the most extensively studied gas phase heterogeneous oxidation process. The selective production of acrolein over cuprous oxide has been known for a very long time. However, the discovery of bismuth molybdates as highly active and selective catalysts for the oxidation t o acrolein, and particularly the ammoxidation to acrylonitrile, has opened a new era in oxidation catalysis. Although “allylic oxidation”, yielding products like acrolein and acrylonitrile, is the most important and successful partial oxidation reaction, several other processes are of interest. Table 3 represents a summary of the nature of the various processes and the main partial oxidation products. This section concentrates primarily on the processes and products listed in the table. It excludes processes that yield predominantly partial oxidation products with less than three carbon atoms and cannot be regarded as very selective processes. Before discussing the various reactions in more detail in the following sections, a short characterization of the processes and a description of the types of catalyst involved will be given. The epoxidation of propene is analogous t o that of ethylene catalyzed by silver. However, the selectivity is much lower. Due t o the pronounced oxidation sensitivity of the allyl CH3-group, excessive combustion occurs as a side reaction. The heterogeneous process has no practical significance, therefore, as it has to compete with a highly selective liquid phase epoxidation process. Allylic oxidation constitutes the main group of oxidation processes
TABLE 3 Summary of propene oxidation processes .
_
_
_
_
_
_
_
~
~
Nature of the reaction .
Epoxidation (Amm)oxidation of the allyl CH3-group Oxyhydration Oxidative dimerisation
References p p . 2 5 3 - - 2 6 2
~
~
.
.
Main products ~
_
_
Propene oxide Acrolein, acrylic acid, acrylonitrile Acetone Hexadiene, benzene
~
136 which are of industrial importance for the synthesis of acrolein, acrylic acid and acrylonitrile. The commercial interest of the latter has markedly contributed t o the extraordinarily large number of metal oxide combinations that have been investigated and t o the complexity of the multicomponent catalysts applied in industry today. Some typical binary oxide combinations, which also form the basis of most multi-component catalysts, are Bi-Mo-0, U-Sb-0, Sn-Sb-0 and Fe-Sb-0. Characteristic features of the allylic oxidation reaction are the initiation by allylic-H abstraction, the formation of a symmetric allyl intermediate, and the role of the catalyst as supplier of oxygen according t o a redox mechanism. The oxidative dimerization has recently attracted attention, both from a fundamental viewpoint and as a means for synthesizing aromatics from lower olefins. The reaction is essentially a combination of allyl radicals, by which the oxidation is limited t o the abstraction of one hydrogen atom. Typically, the catalysts applied here do not contain MOO, or a similar component that promotes the selective incorporation of oxygen. The oxyhydration of propene t o acetone occurs at a much lower temperature than the allylic oxidation and demands, in principle, the presence of excess steam. The reaction is initiated by addition of a proton from the catalyst surface and the acetone formation involves oxygen originating from water.
2.2.1 Pro.Dene oxide production Attempts to produce propene oxide selectively by gas phase oxidation have been plentiful but not successful. One of the best results is reported by Zanderighi and Carra [ 3591, who investigated a number of tungstates in a pulse reactor at 250-350°C and with an oxygen/propene ratio of 3/2. Thallium tungstate appears to give propene oxide with a selectivity of almost 3096, besides considerable amounts of acetone and acrolein, at a temperature of 350°C. At lower temperatures, the last two products predominate. Tungstates were also studied in a batch reactor (autoclave) by Centola et al. [ 841. At 185°C with 10 atm propene and 12.5 atm oxygen, these authors also found thallium tungstate to be the most effective: a selectivity of 40% is reported at 37% conversion. Remarkably, even without catalysts, a considerable epoxide yield is observed under these conditions, The formation of propene oxide as a side product of the acrolein formation or dimerization reactions is reported by many authors. Daniel et al. [ 95,961 demonstrated that propene oxide is formed by surface-initiated homogeneous reactions which may involve peroxy radical intermediates. The epoxidation is increased by a large void fraction in the catalyst bed or a large postcatalytic volume. In view of these results, the findings of Centola et al. [ 841 are understandable, as the wall of the empty reactor may have been sufficiently active t o initiate the reaction.
137
2.2.2 Acrolein production The oxidation of propene t o acrolein has received much attention for several reasons. Firstly, the process is of industrial importance in itself, and it is also a suitable model reaction for the even more important, but a t the same time more complicated, ammoxidation. Secondly, propene oxidation is, in many aspects, representative of that of a class of olefins which possesses allylic methyl groups. Last, but not least, the allylic oxidation is a very successful example of selective catalysis, for which several effective metal oxide systems have been discovered. The subject has therefore attracted much interest from the fundamental point of view. Consequently, numerous metal oxide catalysts have been studied, ranging from single metal oxides t o complex multi-component mixtures, and accordingly the aim of the research has varied from purely fundamental aspects t o development and optimization of industrial catalysts. The flood of patent applications that started some years ago is still going on and today invariably concerns complex oxide compounds. Non-industrial research in the field of multi-component catalysts is still rather scarce at the moment, yet it is becoming more abundant as growing insight into the action of the simpler catalysts is gained. The oxidation process is carried out in the temperature range 300450°C, and generally studied at atmospheric pressure. Excess air is usually applied (with some exceptions) and substantial amounts of water vapour may be added t o the feed. High initial selectivities (>95%)are feasible, and, although some further oxidation (combustion) of the product is unavoidable, yields of 70-90% are reported in the patent literature. The main by-products are carbon oxides, in addition to minor amounts of acrylic acid, acetaldehyde and formaldehyde. Acrylic acid may be a main product depending on specific catalyst properties and reaction conditions, as described in more detail in Sect. 2.2.3. (a) Kinetics and mechanism in general
The allylic oxidation of propene is catalyzed by (compound) metal oxides, which essentially contain metal ions of variable valency. It is commonly accepted that a redox mechanism is operative in such a way that the catalyst acts as the oxidizer and that lattice oxygen is incorporated in the oxidation products. The assumptions have been proved for several catalysts by the analysis of cation valency changes and by experiments with labelled oxygen. The reaction between propene and the catalyst is, in general, rate-determining, as catalyst reoxidation is a relatively fast reaction. This implies that the degree of catalyst reduction under steady state reaction conditions is fairly low (i.e. less than 10% with respect to the total amount of oxygen that can be removed with propene). Thus the observed kinetics Rcfercnces p p . 253--262
138 for these catalysts are principally the kinetics of the reaction between propene and the oxidized catalyst. For many cases, the kinetics are adequately represented by a simple first-order model, viz. = 'PC3H6
An inhibiting effect of the reaction product (acrolein) is reported for several catalysts, and usually accounted for by an adsorption term in the denominator, leading to equations of the form
R
= -hPC3H6
1+ K P C 3 H 4 0 The first-order dependence with respect t o propene may be explained by the assumption that the initial reaction step (dissociative adsorption) is ratedetermining, while the inhibiting effect of acrolein is obviously connected with its strong adsorption on the active sites of certain catalysts. For some catalysts, the contrary situation occurs and reoxidation is the rate-determining step. A typical example is cuprous oxide. The observed rates are, in this case, dependent on the oxygen instead of the propene pressure. Kinetic redox models, as formulated by Mars and van Krevelen [ 2041, have not been considered in any recent work. Although the combined dependence on both propene and oxygen pressures does arise in certain investigations, the authors seem t o ignore redox mechanisms completely and correlate their data with Langmuir-Hinshelwood type models. A detailed treatment of the kinetics of groups of catalysts, and comparison between them is hardly possible due t o the widely different experimental conditions (e.g. catalyst preparation and pretreatment, reactor type, reaction conditions and experimental methods). Results of kinetic studies will be individually reported in the section on catalysts [Sect. 2.2.2.(d)].
( b ) Mechanism of acrolein formation The formation of acrolein comprises several steps. The first and ratedetermining step is generally assumed t o be abstraction of an allylic hydrogen atom. Evidence is provided from several sources, e.g. the deuterium effect of the reaction (Adams [ 2 ] ) and the analogy between oxidation and DzO exchange (Christie et al. [86]), and for various catalysts. The hydrogen abstracted is taken up by a surface oxygen anion t o form a hydroxyl group; the allyl radical is bonded t o the catalyst as a symmetrical complex. Hence, the first reaction step is represented by CH,=:CH-CH, +Me"++ 02-=(CHzLCH%H,),d,,
+Me("-')+ + O H -
Since it is well established that the allyl complex is neutral or (weakly)
139 positive, the reduction of a metal ion of the catalyst must take place, as indicated in the above equation. Metal ion reduction effected by propene adsorption has, in fact, been demonstrated for several metal oxides. The exact nature of the bonding between the allyl species and the catalysts is not known. The initial formation of a complex that is n-bonded t o a metal ion is often presumed, although n-0-complexes, or complexes which include oxygen anions, are proposed as well. The occurrence of a symmetrical intermediate at an early stage of the reaction with a variety of catalysts however, is unmistakably indicated by experiments with l 3 C- and l 4 C-labelled propene. The lifetime of a radical type of intermediate (if it exists) must necessarily be very short, as otherwise radical combination reactions might become important, yielding products like hexadiene and benzene. These products and free allyl radicals have indeed been detected in some studies (Dolejsek and Novakova [ 1021, Seiyama et al. [ 285,2861,and Weiss et al. [345]). It must be emphasized, however, that the catalysts concerned differed from the usual acrolein-forming catalysts, being hardly selective a t all. The selective action of acrolein-producing catalysts is very probably due to the ability to oxidize the initially formed complex rapidly to an allyl carbocation [ 3451. This assumption is the more plausible as selective catalysts are characterized by a strong electron affinity combined with a good electron conductivity. Abstraction of a second hydrogen atom from the initially formed allyl radical, as proposed by Voge and Adams 13431, is much less likely to be the second reaction step: this second hydrogen abstraction presumably requires more energy than the first. Hence, it excludes the first as the ratedetermining step. The exact mechanism of lattice oxygen incorporation and second hydrogen abstraction, and the precise sequence of elementary events is still a subject of speculation. Several authors assume that two distinct active sites are involved in the acrolein formation. The first, presumably a cation, participates in the formation of the initial allyl complex, while the second, which may contain a different cation and reactive oxygen anions, is the place where further hydrogen abstraction and oxygen incorporation take place. An attractive reaction scheme, which does not require transfer of a reaction intermediate from one site t o another, is proposed by Weiss et al. [ 3451. The authors assume that propene is adsorbed at the site containing the cation with the highest valency and that the metal oxygen double bond plays an essential role. The initial adsorption may be represented by
References p p . 2 5 3 - 2 6 2
140 Depending on the nature of M (which is not necessarily a metal), the reaction may proceed via a a-ally1 or Ir-ally1 complex, eventually leading t o an ally1 carbocation
L
A
-e-
A'
M-OH M-OH r-ally1 This cation is attached t o a lattice oxygen atom which activates the abstraction of the second hydrogen atom, followed by desorption of acrolein. The complex role of the catalyst is separately discussed in Sect. 3 , where a great many of the references concern propene oxidation studies. (c) By-products
The main by-products of acrolein formation are carbon monoxide and carbon dioxide, as well as minor amounts of acrylic acid and lower aldehydes and acids. Combustion takes place both consecutive and parallel t o the main reaction. Acrylic acid (in free or adsorbed form) is a possible intermediate in the acrolein combustion. Including this product, the following simplified scheme applies.
,CH2=:H*Ho
C,Hz
1 CO/CO,! The kinetics of the parallel formation of by-products is generally similar t o the kinetics of the main reaction (i.e. rate independent of the oxygen pressure and first order with respect t o propene), which presumes identical reaction steps. A study of the origin of acetaldehyde and formaldehyde carried out by Gorshkov et al. [144] is of interest in this connection. 14C-Labelled propene was oxidized over a bismuth molybdate catalyst a t 460" C; the results showed that acrolein, acetaldehyde and formaldehyde are formed via a symmetrical intermediate, presumably one and the same intermediate. The study, moreover, shows that acetaldehyde is exclusively formed from this intermediate, while formaldehyde may also be formed from the aldehyde group of acrolein and the methyl group of acetaldehyde.
141
Partly homogeneous reactions may also contribute to the formation of by-products, as demonstrated by the work of McCain et al. [214] and more recently by Daniel et al. 195,961. The latter showed that surfaceinitiated homogeneous reactions particularly occur when a large free volume is present after the catalyst bed. A typical product is propene oxide. A relatively large propene/oxygen ratio (e.g. 4/3) favours such reactions with bismuth molybdate at 430"C, while no homogeneous reactions are observed over Cu,O. Inactive as catalysts, but very active as initiators of homogeneous reactions, are Bi(OH)3 (360°C) and Sn(NO,)? (400°C). The authors suggest that the initiator is an allyl peroxide or allyl hydroperoxide species and propose the following reaction scheme for the formation of acrolein and propene oxide.
LO2
Allyl (hydro) peroxide Homogeneous reaction
According to the authors, the importance of the various reactions in this scheme depends o n the catalyst properties, particularly the possibility of reaction between the allyl radical (complex) and lattice oxygen and the possibility of desorption of an allyl peroxide species. ( d ) Catalysts
Historically, the first important acrolein catalyst was cuprous oxide. Since the discovery of bismuth molybdate as the first specimen of a group of superior catalysts, however, attention has been primarily focussed on binary and compound metal oxide mixtures in which copper plays no role. At present, the most promising binary metal oxides are Bi-Mo-0, It appears that these also form the basis of Sn-Sb-O and Fe-Sb-0. selective multi-component catalysts that are described in the literature or patent specifications.
(i) Bismuth molybdates. Catalysts based on bismuth molybdates are undoubtedly the most extensively studied for the (amm)oxidation of proReferences p p . 2 5 3 - - 2 6 2
142 pene and a number of related oxidation processes. The first successful development led t o a commercial catalyst with the composition Bi9PMo12O S 2 ,reinforced by SiOz (50 wt. 7%) for the use in a fluid bed reactor (Sohio acrylonitrile process). Bismuth molybdates have since then been investigated in forms varying from pure unsupported bismuth molybdates t o supported multi-component catalysts. Many studies have been devoted t o the clarification of the selective oxidation mechanism and the nature of the active sites. In the following, the known bismuth molybdate phases and their significance will first be briefly reviewed, followed by a discussion of the mechanism and kinetics. Active crystal phases. Detailed studies have been performed t o identify the distinct phases in the Bi-Mo-(P)-O system and t o determine their specific catalytic properties. It is well known that the individual oxides of bismuth and molybdenum are not useful catalysts themselves: B i 2 0 3 is moderately active but not selective while MOO, is selective but hardly active. Yet, the combination is very active and selective at a Bi/Mo ratio between 2/3 and 2/1. The following crystal phases are identified in this range (Schuit [ 2811).
(1) Bi203* 3 MOO,
= a-phase
(2) Bi,O, . 2 MOO, = 0-phase (3) Bi203. MOO,
= y-or
koechlinite phase
The structure of the a-phase is related t o scheelite (CaWO,) and can be written formally as Bi,,, ol,, MOO,, with ordered cation vacancies ( 3 ) [ 104,851. Molybdenum occurs in pairs of (distorted) tetrahedra which share edges. The y-phase has a structure of alternating bismuth and molybdenum layers. The molybdenum layer consists of corner-sharing Mood octahedra; the layer containing Bi3+ is similar to that encountered in BiOCl. The MOO, octahedra are deformed and possess oxygen atoms at three different distances [104]. The y-phase is unstable at high temperatures (>650°C).The 7'-phase which is then formed differs from the y-phase with respect to the molybdenum layer, which consists of separate Moo4 tetrahedra according t o the La2MoO4 structure [ 541. The structure of the 0-phase is not entirely clear. Moreover, this phase is said to be unstable and t o decompose into a and y. According to Erman et al. [105], it is related t o scheelite, but its unit cell is larger. Crystal phases that contain phosphorus in addition t o bismuth and molybdenum are not known. Phosphorus is probably present as bismuth phosphate. Remarkably, however, a fresh commercial catalyst seems to contain only a minor amount of the a-phase in spite of the Bi/Mo ratio which is close t o 2/3 (Schuit [Sl]). The phases (y' excluded) described above are all active and selective
143
catalysts for the propene oxidation. Several authors report that the highest activity is obtained at a Bi/Mo ratio of 1/1corresponding t o the 0-phase. As this phase is unstable, a correlation between activity and (bulk) structure seems doubtful. Arguments for both molybdenum tetrahedra and octahedra as the active species have been presented in the literature. A problem is the large difference in the conditions existing for the X-ray analysis and the condition of the steady state chemical reaction. Moreover, it is uncertain whether phases observed by X-ray analysis adequately represent the structure of the surface. There are quite a number of arguments in favour of the idea that the surface structure is, t o some extent, independent of the bulk structure. Studies concerning the effect of the method of preparation on the catalytic properties and comparison of catalysts prepared in different laboratories prove that large differences may occur a t the same Bi/Mo ratio and the same bulk structure [41,45,59,149,327], while, on the other hand, very similar properties may be displayed by catalysts with different bulk structures. Batist e t al. [42], for example, have shown that the bulk of a Bi9PMo,,052 catalyst initially consists mainly of the 0-phase, but gradually decomposes into 0 - and y-phases during its use in a reactor at 470"C, without a significant effect on the catalytic properties. Furthermore, it is noticeable that, although different activities are reported for distinct phases by various authors, no appreciable differences in the activation energy occur, indicating that it is not the quality but mainly the quantity of active sites that changes. Of particular interest is the study of Miura et al. [220] who investigated the existence of various types of oxygen by temperature-programmed reoxidation of partially reduced bismuth molybdates at 0-500°C. There was no relation between the recorded reoxidation peaks and the phases observed by X-ray analysis. However, the active and selective catalysts (Bi/Mo = 2/1,1/1 and 2/3) all showed a characteristic peak at 320"C, which also occurred in Moo3 at larger degrees of reduction (i.e. 4--10%). Addition of phosphorus t o the catalysts, P/(Mo + Bi) = 0.075, appeared t o intensify the 320°C peak; for the 1/1and 2/3 catalysts peaks other than the 320°C peak vanished with the addition of phosphorus; at the same time, the catalytic activity was increased, while the selectivity was invariably high (80%). The authors report a good reproducibility of the experiments and conclude that activity and selectivity are correlated with a specific type of oxygen, which is present irrespective of the bulk structure. Finally, Grzybowska e t al. [148] have studied structural changes in the surface layers of bismuth molybdate catalysts occurring on evacuation, reduction by hydrogen and interaction with propene/oxygen mixtures, by means of surface spectrography methods (XPS and UPS). They observed that, in a reducing atmosphere due t o either the action of high vacuum or the presence of reducing agents, rearrangements of the surface layers take place resulting in a lowering of the Bi/Mo ratio for the more Bi-rich catalysts and the formation of similar R e f e r e n c e s P P . 2 5 3-2 6 2
144 compositions on all three bismuth molybdates investigated. No rearrangements occurred by interaction with a reaction mixture (propene/oxygen/ nitrogen = 24/21/55) during 5 h at 440°C. It was found, however, that the surface became covered with a strongly bonded oxygenated hydrocarbon species. This species persisted at the surface even after outgassing at 430°C and could also be formed in situ in the spectrometer by adsorption of pr'opene at 500" C. Kinetics. The kinetics of the oxidation of propene over bismuth molybdate follow the general lines described above [Sect. 2.2.2(a)]. As acrylic acid, acetic acid and formaldehyde are minor by-products, a simple scheme well suited t o describe the acrolein production is
C3H40 C3H6 113) + T o , CO, The reactions in this scheme are first order with respect to the oxidized compound. Initial selectivities [k, /(k, + h , ) ] of 90% and more are possible at 400-500°C. The decrease in selectivity at higher conversions is mainly due to acrolein combustion (k,/k, = 0.2-0.3). The activation energy of acrolein formation is approximately equal for all bismuth molybdates (18-20 kcal mol-' ). The independence of the reaction rate on the oxygen partial pressure is confirmed by studies in which the reaction of the catalyst with propene and reoxidation of the catalyst are separately studied. Sancier et al. [ 2751 performed conductance measurements during reduction and reoxidation using a flow system at 327 and 387°C. The rate of change of the crystal voltage, dAV/dt, is a measure of the reaction rate and appears to satisfy the dependencies reduction:
dAV/dt
=
hrpCgH6
oxidation:
dAV/dt
=
kOp021/2
The calculated rate coefficients are listed in Table 4 and confirm the relative rapidity of the reoxidation for all compounds studied. The existence of a certain (albeit small) amount of reduction during steady state reaction is demonstrated; its level as expected, depends on the k , / k , and thepo,/pc,H6 ratios (see also ref. 251). Reduction and reoxidation rates were also measured by Brueckman et al. [61], who used a static circulation reactor at 220-460°C. The reduction with hydrogen or propene at 460°C proceeds t o Moo2 and Bio . The very fast reoxidation was studied at much lower temperatures. BiO is reoxidized first. The reduction process is rather complicated for the molybdenum-rich phases (Bi/Mo = 1/1and 2/3), which appear initially t o form a mixture of MOO, and the y-phase. Kinetic equations are presented by the authors, but do not seem relevant for catalysis in view of the too severe
145 TABLE 4 Rate coefficients for initial reduction and oxidation of catalysts a ~~
Catalyst
kr
( m V sec-’ tori--‘ 600K
x lo4)
660K
(mV sec-’ torr-ln)
krlko (torrln x
600K
600K
660K
0.35 1.7 0.16 0.06
4.2 5.0 1.1
k0 660K
lo3)
-
M003
Bi/Mo = 0.7 Bi/Mo = 6 Bi203 a
2.2 8.4 0.42 0.11
Estimated probable error Too small to estimate.
30 34 2.3 0.24 t
0.62 0.51 0.26 0.19
0.73 0.69 0.21 b
5%.
catalyst reduction. The same comment applies to some other catalyst reduction studies (Beres et al. [47]). A recent contribution with respect to the oxidation kinetics for a Biz03 * 2Mo03 catalyst is given by Cartlidge et al. [78,79], who used a well-stirred reactor. Contradictory t o the results of any other study, acrolein is reported t o accelerate its own formation, while,carbon oxides in turn accelerate the acrolein combustion. A check on these unusual effects by adding the products t o the feed is not reported. Misinterpretation of the data seems likely, e.g. by the fact that transfer limitations easily occur in this type of reactor. Mechanism. The mechanism outlined for the propene oxidation over metal oxides is, in general, fully applicable to bismuth molybdate. The occurrence of a symmetrical ally1 intermediate and the participation of lattice oxygen is well established (Hucknall [160], Voge and Adams [3431). With respect to the participation of lattice oxygen, some recent contributions concerning studies with labeled oxygen and experiments in the absence of gas phase oxygen must be mentioned. Results of investigations in a static circulation apparatus are reported by Gel’bshtein et al. [ 1231. An I6Oz atmosphere is replaced by a reaction mixture containing 1-2 torr propene and 3-6 torr ‘*OZ at 400°C, and circulated over a 1/1Bi/Mo catalyst. The results prove lattice oxygen participation and a large oxygen mobility in the catalyst. Carbon dioxide appears to be partially formed directly from the gase phase l80,presumably by an associative mechanism. Pulse experiments have been carried out by Sancier et al. [276], who tried to avoid the problem of non-stationary conditions by “seasoning” of the catalyst by I6O2 /propene pulses, followed by the “actual experiment” with l 8 0, /propene pulses. ESR measurements confirmed that the degree of catalyst reduction was indeed constant. The ‘ 8 0 / ’ 6 ratio 0 in the prodReferences PP. 253-262
146 ucts depended on the reaction conditions (temperature and O2/C3H6 ratio). This dependency proves that the kinetics of oxidation, reduction and diffusion through the catalyst determine the extent of gas phase l8O incorporation. Further, a good agreement was found with earlier work of Keulks [ 1741 and Wragg et al. [ 3521 performed with static systems. Experiments using a flow reactor under steady state reaction conditions were reported by Keulks and Krenzke [ 1751. A retarded breakthrough of l8O in the products was observed, after switching from I6O2 t o '*02in the feed which consisted of propene (9%), oxygen (10%)and helium (81%)at 1 atm. The 180/'60 similarly increased in both acrolein and C 0 2 , contrary to the results of Gel'bshtein et al. mentioned above (see also Sect. 3.2.2). The oxygen exchange between molybdates and water can be very rapid as appears from the work of Novakova and Jiru [238]. The authors suggest that this exchange might interfere with the I8O2 oxidation studies. However, as water is not actively participating in the oxidation process, its influence is probably of minor importance [ 3531. It is generally reported that propene oxidation can also occur in the absence of oyxgen. However, the activity rapidly decreases, even in the case of pulse experiments. This is in contrast t o the butene oxidation t o butadiene over the same catalyst, where the activity is remarkably constant. One might suppose that a special type of oxygen, which is readily consumed, is needed for the propene oxidation. A much more likely explanation is provided by Barannik et al. [38], who carried out pulse experiments with propene and acrolein over Bi203 3M003 at 350400°C. A gradual decrease in activity t o a final constant level occurs during the first five propene pulses, but the constant level is immediately attained by injection of an acrolein pulse. The authors conclude that, in absence of oxygen, the catalyst surface is saturated by strongly adsorbed acrolein. A strong adsorption of acrolein has indeed been observed by other authors (Forys and Grzybowska [117], Matsuura [209]). The former authors, moreover, report that the strong adsorption of acrolein is absent after preadsorption of water. This may also provide an explanation for the absence of a strong inhibition by acrolein in flow reactors, where water formed by the reaction is continuously present. As the main lines of the mechanism are well established, discussion in the literature at present is mainly focussed on the structure of the active sites and their method of participation. The hypothesis of a bifunctional mechanism involving ally1 radical formation and oxygen incorporation on distinct sites is advocated by Haber et al. [ 147,1521. This hypothesis is particularly based on experiments with M o o 3 , Bi203 and mechanical mixtures of these oxides, which are compared with bismuth molybdate catalysts. The reaction was carried out in cyclic operation (alternating feeds of oxygen and of propene diluted with nitrogen). The results are collected in Table 5. The authors con-
-
147 TABLE 5 Interaction of propene with mixtures of BizO, and Moo3 (C3H6 : NZ= 30 : 70 (vol.%); contact time = 2.5 sec.) Catalyst
Temp. ('C)
Conversion of C3H6
Selectivity (%)
1%)
Acrolein
1,5-Hexadiene
C02
~
MOO Bi203 9 Biz03 + 2 MOO, initial period after 1 h 4 Bi203 + 2 MOO, initial period after 1 h Bi20, + 2 MOO, initial period Bi2Mo2O9 a
480 480
0
10
75
18 a
480 480
10
70
30
480 480
10 0
30
10
60
480 44 0
8
79 90
Traces
20
0
8b 10
Other product: benzene. Other products: acetic acid, acrylic acid and propionic acid
clude from the surprising production of acrolein over the mechanical mixture that bismuth centres produce allyl radicals which may either combine to give hexadiene, o r further react to acrolein on molybdenum oxide. The significance of the experimental results may be questioned, however, in view of the instability of the mechanical mixtures, which the authors ascribe to migration of MOO, over the Bi203 surface and the formation of an inactive layer; solid state reactions between Moo3 and Bi203 leading to active bismuth molybdate are well known t o occur at such temperatures [246]. Important evidence in favour of this two-centre mechanism is provided by Gamid-Zade and Kisliev [ 1211, who carried out pulse experiments with allyl bromide as a model compound t o generate an allyl intermediate. Moo3 and BiAlly1 bromide and propene were oxidized over Bi-Mo-, S n 4 at 450°C. Very remarkably, the rate of oxidation of allyl bromide to acrolein over Moo3 approaches that over Bi-Mo-O, while for propene, the rates differ by three orders of magnitude. Over Bi-Sn--O (and other binary oxides containing Bi but no Mo), propene can be oxidized to diallyl and benzene. Similarly, no acrolein is formed in the oxidation of allyl bromide over Bi-Sn-P, although the rate of oxidation is higher. Bifunctionality of the catalyst is also assumed by Schuit [281] and is particularly based on the adsorption experiments carried out by Matsuura et al. [207-2121. Two sites are distinguished: (a) A-sites a t which activated and strong adsorption of butadiene, acrolein and ammonia takes place (20-30 kcal mol-' ) a t 25-200°C. The number of A-sites decreases linearly with increasing degree of reduction and each reduced site can References p p . 2 5 3 - 2 6 2
148 adsorb one molecule H 2 0 or a half molecule O 2 (reoxidation). These A-sites are supposed to be oxygen anions in between anion vacancies connected with bismuth ions and represented by “VBiOAVBi”. (b) B-sites which weakly adsorb olefins, acrolein, butadiene and ammonia (5-12 kcal mol-’). The sites are supposed to be anion vacancies connected to molybdenum in between oxygen anions, represented by “OBV M ~ O”,B viz .
@I = 02on top of M O ~ + @ = 0 2 -on top of Bi3’
B-site (@ 0 (@ A-site 0 @ @ B-site @
0
0
@
8
=Mo6‘ = Bi3’
The A- and B-sites are supposed to beintegrated in one reaction site, as depicted above. Based on the respective adsorption properties that are attributed to these sites, the mechanism proposed is (1)fast dissociative adsorption
p
F3HS
C3H6 + OBVM~OB + OBVMOOB (2) transfer of C3H5to VBi (rate-determining) 73HS
OB
V B ~ O A V+BVB~OAC~HS ~
(3) H-transfer and acrolein formation (fast) H VBiOAC3HS +- OB
I
--f
V B ~ V B ~+VOB B ~+ C3H40~
(4) H-migration to A-sites and water desorption
This mechanism implies that the ally1 complex is first bonded to Mo-0 and then to Bi, followed by the uptake of oxygen from the direct neighbourhood of the Bi-centre. These assumptions are in strong contrast with the ideas of Haber et al. [ 147,1521 and Gamid-Zade and Kisliev [ 1211 given above. Different types of oxygen are distinguished also by Mitchell and Trifiro [ 2191, who studied catalysts and model compounds by IR and UV reflectance spectroscopy. Activity appears to be correlated with MOO, species which contain three terminal (multi-bonded) oxygen atoms ( “Ot”)- Moreover, a M o - 0 bond was detected which is even weaker than Mo-0-Mo in Moo3 and assigned to Mo-0-Bi. Following the ideas of Schuit [ 2811,
149 above a Mo(O, ) ~species (-B-site) is ascribed the role of initial propene adsorption, while Mo-0-Bi (- A-site) contains the most loosely bonded oxygen and should serve for water desorption and reoxidation from the gas phase. A very attractive theory concerning the oxidation mechanism for scheelite-type molybdates containing bismuth is presented by Sleight and Linn [297] and is described below.
(ii) Multi-component catalysts containing Bi and M o . Defect scheelite systems were studied by Aykan e t al. [33] for catalysis of the oxidation and ammoxidation of propene at 400--500°C.The systems studied can be represented by A, - x & Moo4 in which A is a combination of metals including Bi and 4 denotes a cation vacancy (defect). The presence of defects appears essential for the catalytic activity. A typical example is presented Bi0.5+x@2xMo04. Some other complex phases and in Fig. 3 for Nao.5-3x their activity for the propene ammoxidation are listed in Table 13 (p. 169). The table shows that, in absence of defects, the activity is low. Moreover, it appears that the good catalytic properties are conserved when Mo is replaced by W , in contrast t o pure bismuth molybdate and tungstate
>
c 0
0
-
0
0 08
0.16
0.24
0.32
2 x In Na0,5-3x B10.5+x Moo4
Fig. 3 . Acrolein yield and propene conversion as a function of catalyst composition. References PP. 253-262
150 (the latter is not very active). The authors ascribe this fact t o the Mo-O and W - 0 coordination which is identical in the scheelites (tetrahedral), but may be different in bismuth tungstate (octahedral) and bismuth molybdate (tetrahedral or octahedral depending on the Bi/Mo ratio, but the authors assume that a tetrahedral coordination might prevail at the surface independent of the bulk structure). Finally, it is concluded that, besides defects, a (small) fraction of bismuth must be present for good activity and selectivity, as omission of bismuth or its substitution by Ce, Y or Ca yields rather inactive and non-selective catalysts. A very interesting hypothesis with respect t o the mechanism, particularly regarding scheelite-type catalysts, is given by Sleight and Linn [ 2971, Three essential elements are discerned at the catalyst surface: MOO, groups, cation vacancies and Bi3+ ions. The latter, due t o their surface position, are coordinated with only six oxygen anions of surrounding MOO, tetrahedra and possess a lone pair of electrons directed up out of the solid. Propene is dissociatively adsorbed on MOO, groups next t o a cation vacancy, forming an adsorbed allyl radical and a hydroxyl group stabilized by the cation vacancy, viz.
H o = oxygen anion c] = cation vacancy
C
H,C
8 ’A
CH, 0
0
Bi
0
Mo
0 0
0
0
Mo
0 0
Bi
This stabilization may also be interpreted in terms of oxygen anions, which, due, t o the vacancy, are initially double bonded t o Mo. One electron is transferred to the catalyst in this reaction step. To form acrolein, a second hydrogen atom is transferred (to form water) and an oxygen atom is bonded t o the allyl radical. In this (rather complex) process, another three electrons are transferred t o the catalysts and doubtless distributed over several Mo ions. Reoxidation takes place at the bismuth cations, where oxygen molecules are attracted by the free electron pair. The intermediate result is a surface bismuth with an oxygen coordination similar t o that in the bulk, viz.
The oxygen anions rapidly diffuse t o vacancies in the MOO, groups. The electron transfer certainly does not involve oxidation of Bi3+ t o higher oxidation states. On the contrary, the authors assume that the steady state situation will be something like leaving some electrons in the Bi6 p conduction band. This idea is the more attractive as it is well known
151 that bismuth molybdates are slightly reduced under steady state reaction conditions, while it is also known that the conductivity strongly increases on reduction (Sancier e t al. [275]). As this band is very close t o the Mo4d levels, a rapid transfer of electrons from reduced molybdenum ions to a (remote) Bi3' site is feasible. A number of Bi-Fe-Mo4 catalysts were investigated by Daniel and Keulks [94], using a flow reactor at 400°C. The catalysts were prepared by a slurry method from Bi2Mo3Ol2and Fe(OH), . At iron contents of 30-40 at.% with respect t o (Bi + Mo), a catalyst is obtained which is comparable with pure bismuth molybdate as to activity and selectivity, but remarkably resistant t o reduction and to heat treatment up t o 800°C. Regarding the kinetics, a first-order dependence with respect to oxygen is reported in contrast t o bismuth molybdate, but the experimental conditions are not given. The origin of the superior properties of the Bi-FeMo-0 catalysts is doubtlessly connected with the existence of a ternary phase. Several recent phase studies reveal the existence of a phase with (Bi + Fe)/Mo = 2 and a scheelite-like structure containing Moo4 and Fe04 tetrahedra. There is some discussion in the literature as t o the actual stoichiometry. According t o Lojacono e t al. [191] and Notermann e t al. [237], two compounds can form: Bi2Fe2Mo,Ol2 and- Bi3FeMo2012. Linn and Sleight [ 1881, however, argue that only the latter compound is formed, whereas the former consists of Bi3FeMozOI2mixed with known binary and single oxide phases.
(iii) Other Moo3 based catalysts. A large number of molybdates and compound oxides containing MOO, have been studied, but only a few of them approach the superior qualities of bismuth molybdates. Good selectivities can be obtained with Te-Mo-0, Sn-Mo-0 and probably with MnMo-0, but excessive combustion occurs on the major part of the studied combinations (e.g. MOO, with oxides of Cu, Fe, V, Zn, Co, Ni, etc.). A number of catalysts has been studied in more detail and will be briefly reviewed here. The system Te-Mo-0 has been investigated by Andrushkevich e t al. [ 271 and by Robin et al. [ 2671. A recirculation reactor was used by former (390°C; 1 atm; 3%C3H6;20% 0,; conversion 0-50%) while Robin e t al. used a pulse reactor (460-500°C; C3H6/O2 = 1/1.8; conversion 2%). Both studies reveal a strong promoting effect of TeO, . High selectivities (90-9576) are demonstrated over a large composition range (Te/Mo = 1/11-2/1). Andrushkevich e t al. [27] find amaximum activity a t 8 at. % Te, approximately coinciding with the maximum solubility of Te6+in Moo3 and with an observed maximum in the Mo5+ESR signal (Fig. 4). The Mo5+ signal intensity can be considered as a measure of the reducibility of the catalyst. Robin et al. [267] proved the existence of a Te,MOO, phase by X-ray and thermal analysis. They ascribe the catalytic activity t o this compound while Andrushkevich et al. propose the solid References p p . 253-262
152
0
b
c
X
.I
Y
iQ
N
E 3 3 u
-aJ
:
v
aJ
c
m
[r
1 ___le
Te Atomic ratio -
Mo
Fig. 4 . Correlation between catalyst activity and dation.
M05+ concentration
for propene oxi-
solution as the active phase. The latter also performed experiments which indicate a redox mechanism. Reoxidation of the catalyst is rapid and firstorder kinetics with respect t o propene are observed. The combination SnO2-MoO3 may also produce acrolein with reasonable selectivity. However, recent studies mainly concern acetone formation, which is favoured under appropriate conditions (excess of water and relatively low temperatures) (see Sect. 2.2.4). The properties of Mn-Mo-0 catalysts strongly depend on the Mn/Mo ratio. Machek and Tichy [193] report that at Mn/Mo > 1, the catalyst consists of MnMo0, and M n 2 0 3 , and is already very active a t 300°C. However, combustion is the only process. At Mn/Mo < 1, the catalyst consists of MnMoO, and M o o 3 , and may produce substantial amounts of partial oxidation products at 400-500°C. Machek and Tichy report, for example, that at 479°C and Mo/Mn = 7/1, a selectivity of 78% is obtained. However, formaldehyde is also formed and amounts t o 20-5096 of the total aldehyde yield. Viswanathan e t al. [342] have prepared a catalyst with a Mn/Mo ratio of 1/1which has catalytic qualities very similar to that of bismuth molybdates. At 400-450°C, a selectivity is 70% is obtained at
153 a 20% conversion level. Experiments in the absence of oxygen give evidence of a redox mechanism. Kinetics similar t o those with bismuth molybdates are observed. Mazzochia e t al. [213] report an acrylic acid selectivity of 30% a t 4 3 0 ° C and Mn/Mo = 1.01. Haber [151]studied solid state reactions occurring by reduction with hydrogen and noted a high mobility of Moo3 on the surface of Mn203. Iron molybdates, well known as selective methanol oxidation catalysts, are also active for the propene oxidation, but not particularly selective with respect t o acrolein. Acetone is the chief product at low temperature ( 200"C), whereas carbon oxides, besides some acrolein, predominate a t higher temperatures [ 182,2571. Firsova e t al. [ 112,1131 report that adsorption of propene on iron molybdate (Fe/Me = 1/2)a t 80--120°C causes cation reduction (Fe3++ Fez+) as revealed by y-resonance spectroscopy. Treatment with oxygen a t 4 0 0 ° C could not effect reoxidation (in contrast to similarly reduced tin molybdate). The authors assume that this phenomenon is related t o the low selectivity of iron molybdate. Cobalt molybdates have approximately the same activity as bismuth molybdates but the selectivity is much lower due t o combustion and acrylic acid formation as important side reactions. Vinogradova et al. [ 3411 report a kinetic study carried out with a flow-type gradientless reactor at 390°C Ssing excess water. The reaction orders with respect to propene and oxygen are 1 and 0. Experiments with propene/acrolein mixtures reveal complex kinetics and strong inhibition by acrolein. Grzybowska e t al. [ 1471 report equal activities for both CoMoOl modifications, and conclude that the rate-determining step must take place at C o sites, as Co has the same coordination in both modifications, whereas that of M o differs. Another conclusion of these authors concerns the low mobility of oxygen in these catalysts: apparently, a high mobility is not a necessary condition for a high activity. An iron-promoted cobalt molybdate catalyst (Fe,, 0 3 Coo., 7 M 0 0 4 ) was studied by Maksimov e t al. [195,196]with respect to the role of iron in the transfer of charge. Iron strongly enhances the catalytic activity and at the same time increases the conductivity by a factor of 100. Mossbauer spectroscopy reveals that 4% of the iron ions are present as Fe*+ "impurity". This fraction is doubled a t steady state reaction conditions, and indicates participation of iron in the charge transfer process. The interaction of propene and oxygen with sodium and potassium molybdates was studied by Burlamacchi et al. [ 661. ESR measurements reveal that Mo5+ is formed (at 380°C) although the catalytic activity is zero. The reoxidation appears t o be very difficult. The authors conclude that a cation like Bi3+or Fe3+is required to facilitate reoxidation. The activity of Moo3 supported on a high surface silica carrier was studied by Vaghi e t al. [331] using pulse and flow techniques at 400440°C. Oxidation activity and acrolein formation appear to be zero below 10 wt. % MOO,, but increase with the MOO, content above 10%.The References P P . 253-262
154 authors suppose that, below 10 wt. %, isolated tetrahedral M o - 0 species and silicomolybdic acid occur at the surface. Above 10 wt. 76 of MOO,, polymolybdates (octahedral coordination) and highly dispersed MOO, are formed, and the authors assume that these form the active sites for acrolein production; bulk MOO, is ruled out because of its low activity. Akimoto and Echigoya [ 161 studied Moo3 supported on SiOz , A l z 0 3 or TiO;, , modified by VA group or alkali metal oxides. These oxides were added to Moo3 by impregnation in amounts of 30 and 0.3 at. % (X/Mo), respectively. The activity of the modified oxides appears to increase in inverse proportion to the electronegativity of the modifying oxide. Thus the activity sequence is: Bi > Sb > As > P and Cs > Rb > K > Na > Li. The authors assume that the modifier affects the reactivity of the Mo=O double bond by decreasing the bond strength and giving the oxygen a more radical-like character. The addition of alkali metal must be confined to very small amounts to avoid the formation of alkali metal molybdates, in which case the level of activity is strongly decreased while the sequence is reversed.
(iu) Iron antimonates. The antimony-rich side of the Fe-Sb-0 system provides very selective and rather active catalysts for both the oxidation and ammoxidation of propene. Sramek and Tichy [ 3011 studied catalysts with Sb/Fe ratios ranging from 3/1 t o 9/1 in a flow reactor at 380-500°C The Sb/Fe ratio has only a small effect on the catalytic properties in this range. Some of the best results with respect to activity and selectivity are presented in Table 6. Regarding the kinetics, the authors report that the oxygen concentration and the presence or absence of water both hardly influence the rate (for the 3/1 catalyst). Gel'bshtein et al. [124] studied the same catalyst using flow and pulse techniques and reported a reaction order of 0.5 with respect to' propene, irrespective of the degree of reduction or the absence or presence of oxygen. As to the active components, little is known. Fattore et al. [lo91 ascribe the catalytic properties t o combination of FeSbO, and cu-Sb204 in a well dispersed mixture. Pure FeSbO, forms combustion products exclusively.
TABLE 6 Iron antimonates as catalysts for propene oxidation ~~~~
~~
~
Sb/Fe
Temp. ("C)
Conversion
Selectivity
311 31 1 311
385 425 496
12.1 19.1 31.2
92 84.4 69.2
155 (u) Tin antimonates. As with the Fe-Sb-0 system, very selective catalysts for acrolein and acrylonitrile formation are found over a large composition range (7-90 at. % Sb). Regarding the active phase, Godin et al. [ 1421 report that the catalyst consists of S b 2 0 4 and a saturated solution of Sb in SnOl, which contains 5 at. % Sb. The authors suppose that the solid solution is the active phase present at the catalyst surface. Christie et al. [86] proved that a redox mechanism is operative, by experiments with l80either in the gas or the catalyst phase. The authors confirmed the superior activity of the solid solution observed by Godin et al., and, moreover, observed a maximum C3H6-D20 exchange rate at an Sb/Sn ratio corresponding t o the saturated solid solution. The occurrence of this excharge reaction, which is confined to five H atoms and excludes the H atom of the central C atom of the propene molecule, is in agreement with the ideas of dissociative adsorption and formation of a symmetrical allyl complex as the initial step in the propene oxidation. Although the saturated solution demonstrates maximum catalytic activity, larger Sb contents are required for a high acrolein selectivity. As to the role of the two elements in the catalyst, Gamid-Zade and Kisliev [121] argue that Sn activates propene giving a reactive allyl complex, while the oxygen polyhedra of the anion-forming element of higher electronegativity (Sb) are the active centres for further oxidation t o acrolein. (A similar hypothesis is given for Bi and Mo in bismuth molybdate; see above.) Their arguments are based on a comparison of the results obtained with Bi,-Sb,., -Sn-0 and various binary oxides. Only binary oxides that contain Sb or M o are capable of producing acrolein (with 30-8096 selectivity), while Bi-Sn-0 and some other binary oxides containing Bi or Sn do not form acrolein but instead produce substantial amounts of diallyl and benzene (3-30%). The ternary oxide displays both properties depending on the Bi/Sb ratio. A kinetic study carried o u t with a well-stirred reactor was reported by Cartlidge et al. [78,79]. Temperatures below 420°C were used to avoid acrylic acid formation. At atmospheric pressure, the oxygen and propene concentrations were varied between 1 and 10, and 5 and 1576, respectively. Selectivities of 60-9096 and a maximum acrolein yield of 28% were reported at 400°C. The kinetic results were fitted t o a Langmuir-Hinshelwood type of rate equation
Although this model cannot correctly reflect a redox mechanism, it indicates that the reduction and reoxidation rates have the same order of magnitude, and hence both influence the kinetics. A commercial, iron-promoted ammoxidation catalyst (Fe/Sn/Sb = 0.25/1/4) was investigated by Crozat and Germain [93] using a flow reactor at 35O-48O0C, atmospheric pressure and a C3H6/02ratio of l/lO.The References p p . 253-262
156 initial selectivity is higher for oxidation than for ammoxidation, but the latter is better a t higher conversions due t o acrolein combustion. The kinetics are described by a simple parallel consecutive scheme. The ratio of reaction rates is almost independent of temperature. The initial selectivity is 9676, and propene oxidation and acrolein combustion have about the same apparent first-order rate coefficient. The overall activation energy is 20.4 ? 4.5 kcal mol-'. An overall reaction order of 1 is reported (approximately 0.5 with respect t o each of the reactants). The authors compared this catalyst with an unpromoted Sn-Sb-O catalyst and conclude that Fe has practically no effect on the oxidation reaction, although a substantial promoting effect on the ammoxidation is shown. Pulse experiments with a Sn/Sb = 2 / 1 catalyst in the absence of oxygen have been carried out by Barannik e t al. [38,39]. The activity rapidly decreases with increasing reduction, while the selectivity strongly increases. This is in contrast with bismuth molybdates, which demonstrate a similarly decreasing activity, but a constant (high) selectivity level.
(vi) Cuprous oxide. Cuprous oxide as an oxidation catalyst has been extensively studied in the past and amply reviewed (Hucknall [160], Margolis [203], Voge and Adams [343]). The active component is the cuprous oxide phase, and not cupric oxide, which only effects combustion. The selectivity is not very high (60-85% at'10-20% conversion). The conversion of propene is limited by the large excess of propene (e.g. C3H6/02 = 5) that is required to maintain the catalyst in the form of Cu20. The kinetics show a first-order dependence with respect t o oxygen and a zero-order for propene, which is different from the usual dependencies, partly because of the high propene/oxygen ratio. The mechanism is assumed t o follow the reduction-oxidation models and evidence has been provided for the occurrence of a symmetrical ally1 intermediate. There are only a few recent publications. Anshits et al. [29,30] have carried o u t adsorption studies with various C u - 0 phases and determined kinetics at low pressure in a static system. One of their conclusions is that the kinetics of partial and complete oxidation are very different. The mechanism of the latter is supposed to be of the associative type, contrary t o the redox mechanism of the partial oxidation. A kinetic study with a continuously stirred vessel (375-400°C, 1atm) was carried out by Lakshmanan and Rouleau [ 1851. In contrast t o the redox mechanism, a singlesite Langmuir-Hinshelwood model is proposed, for which the k values and activation energies are determined. The effect of methyl bromide, added as a modifier t o the feed, was studied by Holbrook and Wise [158]. The modifier appeared t o have a profound influence on the selectivity, even a t low concentrations (450 ppm CH,Br in propene). This resulted in a relatively high selectivity, which did not depend on the C3H6/02ratio. The action is explained as an effect on the Fermi level of the catalyst (see Sect. 3).
157
A CuO-MgO solid solution was investigated by Davydov and Budneva [97]. Propene adsorption complexes (n and Q) were detected a t room temperature and appeared t o react by heating t o 300°C.
(vii) Miscellaneous catalysts. A large number of Sb-containing binary oxides was studied by Sramek and Tichy [ 3011, i.e. combinations with the metals Fe, Ni, Cr, Sn, Pb, Cu, Ce, and Mn. At a 1/1 atomic ratio with Sb, the selectivities are generally low. Good results are obtained only with Fe-Sb-0 and Sn-Sb-0, which have already been mentioned. Zanderighi and Carra [359] investigated the tungstates of Cu, Mn, Pb, Bi, Fe and T1 in a pulse reactor (250--350"C, 02/C3H6 = 3/2). The tungstates of Cu (at 250"C), Bi, Pb and T1 are moderately selective and active. The latter is the most selective, producing acrolein, acetone and propene oxide. The main product at 250°C is acrolein, but a t 350°C propene oxide and acetone are the principal products. The epoxide formation indicates that peroxyradical species are produced by the catalyst, which may further react in the homogeneous gas phase. The reactivity data were correlated by the authors with the results of thermogravimetric analysis (reduction by hydrogen and propene). Zeolites of type X, containing various metal ions (Pd, Cu, Co, Zn, Ni, Mn, Cr, Fe) were investigated by Gentry e t al. [125]. Only very small amounts of partial oxidation products are found (acrolein, acetaldehyde, formaldehyde). Although V 2 0 5 catalysts are not very selective for the oxidation of propene, some studies were devoted to the investigation of the action of thk catalyst. Krylov [ 183,1841 studied V2OS catalysts supported on S O 2 , MgO, or A1203 by chemisorption of reactants and other techniques, and made assumptions about the types of oxygen and surface complexes involved from the results. ValdeliGvre e t al. [ 3331 investigated processes occurring in the surface layers of V z 0 5 and a 9/1 mixture of V z 0 4/ V 6 0 , , by a variety of techniques (adsorption, thermogravimetric analysis and analysis by IR, EPR and ESCA). The authors conclude that V5+is essential for oxidation activity. An oxidation-reduction mechanism is evidently operative. The selectivity of this catalyst is low. The fact that oxygen originates from different sites (V= 0 and V-0-V) is suggested to be the cause. Niwa and Murakami [ 235,2361 investigated various catalysts (Bi-Mo0, Bi-W-0, Sn-Sb-0, Sb-Mo-0, MOO, and Sn-P-0) with the periodic pulse technique. This method is distinguished from the conventional continuous flow reaction method by the alternate feeding of oxygen and propene. The reaction was carried out a t 386"C, and the products of the propene (P) and oxygen (0)pulse were separately analyzed. Acrolein is only formed during the P pulse, indicating reaction between propene and surface oxygen as the exclusive source of this aldehyde whereas carbon oxides are formed in both pulses. The data collected in Table 7 show that REfErenCes p p . 253-262
TABLE 7 Combustion products in pulse experiments for propene oxidation Catalyst
Production of COZ and CO -~
Bi-Mo (1/1) Bi-Mo/SiOZ Bi-W (1/1) Sn-P MOO, Sn-Sb (10/1) Sb-Mo (2/3)
Total amount (pmole per period per g cat.)
5% in 0 pulse
2.6 6.0 6.0 2.9 0.16 1.1 0.6
83 83 79 71 52 32 23
the most active catalysts (Bi-Mo-0 and Bi-W-) predominantly form the carbon oxides in the 0 pulse. (These catalysts are severely reduced by the P pulse.) Less active catalysts (Sn-Sb-0, Sb-Mo-0) maintain higher oxidation states and mainly form carbon oxides in the P pulse. MOO, and TABLE 8 Comparison between the periodic-pulse and continuous-flow reactions for propene oxidation Catalyst
Bi-Mo
(1/1)
Bi-Mo (1/1) Bi-W (1/1) Sb-Mo (2/3) Sn-Sb (10/1) Sn-Sb (4/1) Sn-P (10/1) Moo3
a
Technique
Pa fb P f P f P f P f P f P f P f
Flow
(mmol h - * )
coz
co
Acrolein
0.836 1.06 1.17 1.69 2.02 2.15 0.245 0.740 0.732 1.00 0.628 0.886 0.720 0.756 0.120 0.094
0.21 0.39 0.35 1.00 0.488 0.543 0.265 1.01 0.194 0.288 0.179 0.296 0.434 0.484 0.069 0.063
1.86 5.42 2.93 6.44 0.306 1.11 0.518 0.925 3.48 5.25 3.46 5.46 0.368 0.388 0.362 0.342
Pericjdic-pulse technique under t h e following conditions: period, 30 sec ( 1 5 0 , 15R); C3H6, 0.25 a t m ; 0 2 0.21 , atm. Continuous-flow technique a t the normalized pressure. 0.67 a t m of O2 in the 0 pulse.
159 Sn-P-O are intermediate. Comparisons with ordinary flow experiments (Table 8) reveals that much more acrolein is formed at normal flow conditions for Bi-Mo--O and Bi-W-0, while for the other catalysts the difference is small or similar for both acrolein formation and combustion. The system Te0,-SiOz was investigated by Castellan et al. [80] who found that mixtures containing more than 10% Te02 were catalysts at 440°C (air-propene ratio = 12). Selectivities of 30-7096 are reported. 2.2.3 Acrylic acid production
The one-stage conversion of propene to acrylic acid is much more difficult than the selective production of acrolein. The process is essentially a two-step process in which acrolein is the intermediate product. Further oxidation leads to acrylic acid. In fact, contrasting catalyst properties are required for these reaction steps. The acrylic acid production demands an acidic catalyst surface, while a basic, or only weakly acidic character is preferred for the selective acrolein formation. Therefore, enhanced combustion and by-product formation are unavoidable. It is doubtful whether a single-step process is at present competitive with the two-step process currently used in industry. In the latter, the oxidation of acrolein to acrylic acid is carried out with high selectivity over mixed-oxide catalysts based, for example, on Mo03-V205 or Moo3TeOz [160]. The catalysts that have been studied for the selective single-stage production of acrylic acid are all based on molybdates or modified MooJ.It appears that good yields can be achieved only with rather complex multicomponent catalysts. The quality of Moo3 combinations with iron, cobalt and manganese oxides has been investigated by Mazzocchia et al. [213]. Table 9 presents some results, obtained with a flow reactor at 300--430°C using a C3H6/02 ratio of 0.22-0.28. The best catalyst is MnMo04. The addition of small amounts of water (up to 2.5%) further increases the selectivity, but larger amounts cause rapid deactivation.
TABLE 9 Properties of molybdates in the conversion of propene to acrylic acid
X Fe-Mo-0 CO-Mo-0 Mn-Mo-0 Mn-Mo-0
a
a
X/Mo
Activity
Selectivity
112.3 111.02 111.01 116.03
Active Active Active Inactive
Small 15% 30%
~~
a
Contained some free MoO3.
References PP. 253-262
160 Cobalt molybdates are also the subject of the work of Alkhazov et al. [ 19,201, which includes extensive infrared spectrometry studies. The Mo/
Co ratio and preparation method were varied, while a standard feed consisting of 12% C3Ha, 15% Oz, 28% HzO and 45% Nz was supplied. The maximum yield of acrylic acid was found at Mo/Co = 2, i.e. with a mixture of equal amounts of CoMo04 and Moo3. This maximum was found to coincide with a maximum content of the P-CoMo04 modification, and the activity was ascribed to the terminal Mo=O double bonded oxygen present on the surface of this phase. The excess of Moo3 may also have a promoting effect due to its contribution to the acidity of the surface. More exotic combinations of Moo3 (20%) with AszO, (5--10%) and Nbz05 or TazO, (10%) were studied by Campbell et al. [71].A t 4OO0C, yields of about 50% can be obtained in the presence of water. Only very small amounts of acrolein are formed, and acetic acid is now the main byproduct. Unfortunately, the catalyst evaporates AszO3 during use. An interesting contribution with respect to the mechanism is given by Novakova et al. [ 2391. They studied the role of water in the formation of acrylic acid over a Mo-Te-W-Sn-0 catalyst. Water appears essential because, in its absence, acrolein is the only product. The use of HZ'*0 showed that one of the oxygen atoms incorporated in acrylic acid originates from water, while water does not participate in the formation of acrolein. The conclusion is that, in the oxidation of propene to acrylic acid, lattice oxygen is first introduced to form acrolein or an acrolein type of complex, followed by the introduction of an oxygen atom, or perhaps a hydroxyl group originating from water. This conclusion is in agreement with earlier work of Andrushkevich et al. (mentioned in ref. 239), who studied a similar type of catalyst (Mo-Te--Co--O). The latter authors, moreover, advance the idea of bifunctionality of these catalysts, by concluding that different active components of the catalyst are responsible for the axidation of propene to acrolein and of acrolein to acrylic acid. 2.2.4 Dimerization and aromatization Several single and binary oxides have a capacity to oxidize propene to dimerization products. The first compound formed is 1,5-hexadiene, which may undergo further dehydrogenation and cyclization leading to benzene. Many authors assume that the initial reaction step in the dimerization is identical with that in the acrolein production, namely hydrogen abstraction and formation of an allylic intermediate. Dimerization is then supposed t o occur because the ability to oxidize the ally1 radical to acrolein is absent. The best known dimerization catalysts are Biz03, bismuth salts and binary oxide mixtures containing Biz03. A very effective catalyst is BiZO3SnOz, in particular for the production of benzene.
161 Pure bismuth oxide has been investigated by several workers. It is generally established that high selectivities can be obtained in the absence of oxygen, up t o high degrees of reduction. According t o Swift et al. [311] and German et al. [136], hexadiene and benzene can be formed with selectivities of 54-76% and 30-4796, respectively, at a temperature of 475-500°C. Fattore et al. [lo81 measured the influence of temperature in the range 45O-60O0C, using a flow reactor and an oxygen-free reaction mixture. They report, for instance, that, at 450"C,the selectivities t o hexadiene and benzene are 60% and 2596, respectively, while at 550°C, the two selectivities are equal (45%), indicating that higher temperatures favour the dehydroaromatization reaction. Boersma [ 561 studied the kinetics with a differential flow reactor at 550°C. In the absence of oxygen, the reaction is first order with respect t o propene, with an activation energy of 20 kcal mol-' . It is concluded that the initial hydrogen abstraction is rate-controlling and, apparently, oxygen diffusion from the bulk t o the surface is fast. The kinetics of the formation of hexadiene are not affected by the presence of oxygen in a low concentration. However, carbon dioxide formation is strongly intensified, and predominates at higher oxygen concentrations. The author also carried out propene adsorption measurements on partially reduced catalysts and studied the relation between reaction rate and degree of reduction. The hexadiene production rate decreases t o zero as the reduction increases t o 10076,but initially the decrease is less than proportional. The propene adsorption capacity appears maximal on a partially reduced catalyst. The author hence concludes that the ally1 complex is adsorbed on an anion vacancy, while the abstracted hydrogen is taken up by an oxygen anion. Bismuth phosphates and various other bismuth salts (e.g. arsenate, basic sulfate, and titanate) are capable of producing benzene, as reported by Seiyama et al. [ 2831. A selectivity of 49% is reached with a combination 2Bi203. PzOs at 500°C. Sakamoto et al. [271] varied the Bi/P ratio and stated that a 2/1 ratio gives the maximum selectivity. Several other single oxides have been studied and compared with Biz0 3 .Fattore et al. [ 1081 report that S b 2 0 4 is very selective (75% t o hexadiene) but much less active than Bi2O3.The same applies t o pure SnO, . A high selectivity is obtained in the absence of oxygen, but reduction rapidly deactivates the catalyst [ 711. Other single oxides that remarkably demonstrate a dimerization or dehydroaromatization activity are ZnO, I n 2 0 3 and T1203 [ 286,3281. The Bi203-Sn02 combination was studied by Solymosi and Bozso [299] and by Seiyama et al. [284,285]. The former carried out pulse experiments in the absence of oxygen and report that even small amounts of SnOz added to Bi203 have a promoting effect and shift the product spectrum from hexadiene t o benzene. The best combination is a mechanical mixture of the two oxides in a 1/1ratio. With this catalyst, a selectivity of 80% (benzene) is reached at a 40% conversion level (at 500"C), R e f e r e n c r s PP. 2 5 3 - 2 6 2
162 which largely exceeds the capabilities of the individual oxides. Seiyama et al. studied the kinetics in a flow reactor at 500°C. At the low conversions applied (
163 at a relatively low temperature (150-300°C) and at a low percentage of MOO,, while the presence of water is a necessary condition. With higher temperatures and higher contents of MOO,, acrolein formation and combustion reactions predominate. Remarkably, the selectivity to acetone is lost when MOO, is exchanged for WO, or U 0 3 (Moro-oka et al. [ 2281). The essential role of molybdenum is obvious, but is active form in the catalyst is still uncertain. Several authors are of the opinion that molybdenum ions are dissolved to a small extent in the second metal oxide. Buiten [62], however, ascribes the activity to Moo3 islands, observed on the surface of his Sn02-Mo03 catalyst. Giordano et al. [141] report acetone formation over Moo3 supported on alumina in the presence of steam. With varying amounts of MOO, (10-30%), a correlation is observed between the catalytic activity and the presence of M o 5 + , as measured by ESR and reflectance spectra. The authors conclude that the active site is an acidic one, formed ~y the reversible adsorption of water on Mo5+surface cations. The authors suggest that any oxide combined with MOO, may yield a good catalyst, provided that Mo5+is stabilized and Bronstedt acid Mo sites are created. The characteristic reaction conditions and the special role of acidic sites in acetone formation accord very well with what is known about the mechanism. The first step in this process is the reversible uptake of a proton by the propene molecule, as evidenced by the D 2 0 experiments carried out with a SnOz-MoO3 catalyst by Buiten [63,64], viz. H+
CH3--CH=CH2 + CH3-kH--CH3 The proton originates from acidic OH-groups of the catalyst surface. The importance of acidic sites is stressed by several authors [11,141,312] of whom Takita et al. [ 3121 have demonstrated a linear relationship between activity and the concentration of acidic sites in the case of a Sn-Mo-O catalyst. Regarding the attachment of oxygen to the carbon formed, Moro-oka et al. proved by the use of H 2 I 8 0 together with 1 6 0 2 that the introduced oxygen mainly stems from water. (It is noted that acrolein, formed as a by-product, did not contain '*O.) On this basis, the mechanism proposed by Buiten is CH3--CH--CH3
It must be additionally assumed that the acidic OH-group is regenerated by water from the gas phase, while the remaining protons react on a different site with oxygen anions and desorb as water. A significant isotope effect for the hydrogen at the second carbon atom, observed by Buiten, may indicate that the second step in the above scheme is rate-controlling. A slightly different mechanism is proposed by Giordano et al. and References p p . 253-262
164 Moro-oka et al. who assume that 2-propanol (free or adsorbed) is a reaction intermediate, viz. CH3-CH=CH23 CH3-&H-CH3
OH+
(CH3--CH-CH3),d,.
0 +
I
OH CH3-C-CH3 II 0
+ H20
The very rapid conversion of 2-propanol t o acetone under these conditions can be consistent with such a mechanism, although it conflicts with the deuterium effect for the central H-atom reported by Buiten. The high velocity of the 2-propanol oxidation, on the other hand, hampers the finding of clear evidence in favour of one of the above mechanisms. The kinetics of the acetone production have not been studied in detail. Fractional orders with respect to both propene and oxygen are found by Moro-oka et al. for a Co/Mo = 9/1 catalyst, which agrees with the fact that the first reaction step is certainly not rate-controlling. It is commonly reported that oxygen increases the activity, but lowers the selectivity, which is consistent with the idea of Mo" active centres,
( b ) Paliadium-based catalysts The interest in palladium-based catalysts is due to the double bond oxyhydration capacity of palladium, unique among the noble metals, and well known from the Wacker process. Fuyimoto and Kunugi [ 1191 report that palladium salts on active charcoal are excellent catalysts for the oxidation of olefins, particularly ethylene but the higher olefins as well. A selectivity of 89% with respect to acetone beside 10%aldehyde production is obtained at a conversion level of 27%, using excess water and a very low temperature (105°C). Careful analysis of the charcoal does not indicate that metal oxide impurities are of importance. Other supports d o not provide stable catalysts. Cant and Hall [72,73], for example, used silica and found very low selectivities.
2.2.6 Arnmoxidation The ammoxidation of propene to acrylonitrile is of great industrial importance and accordingly the literature is abundant. The reaction is very similar t o the oxidation of propene t o acrylonitrile and carried out at the same conditions and over the same kind of catalysts. The famous bismuth phosphomolybdate catalyst developed by Sohio was the first of a series of highly effective mixed-oxide catalysts. The optimum yields are generally obtained at temperatures of 400-500" C. Initial selectivities over 95% and yields up to 80% are feasible. The superior selectivity of the ammoxida-
165 tion compared with the oxidation is mainly caused by the fact that acrylonitrile is much more stable with respect t o combustion than acrolein.
(a) Kinetics and mechanism The strong parallel with the acrolein formation initially suggested the idea that acrolein is a reaction intermediate in the ammoxidation, and can further react with ammonia and oxygen t o form acrylonitrile. Although the ammoxidation of acrolein is indeed a very rapid reaction, it is generally accepted today that a direct reaction path to acrylonitrile predominates. The differences between both theories are very small, however, when one assumes that the ammoxidation of acrolein and propene involves the same reaction intermediate. Thus the various kinetic schemes proposed in the literature can be derived from the general scheme below by omitting the reaction steps (3), (4)and/or ( 5 ) and variation of the ratio between (2) and (3).
-+
>
C3H40
(1)
C3H,
intermediate complex -- ( 5 )
$(4)
%T---- C ~ H ~ N Regarding the mechanism, it is generally accepted that, as with acrolein formation, a symmetrical allyl complex is formed (as already implied by the above scheme); for a number of catalysts, this has indeed been proved experimentally (for instance by Dozono et al. [103]). Little is known about the precise mechanism of nitrogen incorporation. It is well established, however, that nitrogen oxides do not play any role in the reaction, as demonstrated, for instance, by Wragg e t al. [ 3541. The reaction of propene with oxygen is much more rapid than the reaction with NO, while the selectivity of the reaction with NO is low, acrolein being the main product. Moreover, the oxidation of ammonia over selective ammoxidation catalysts (if it occurs under ammoxidation conditions) primarily produces Nz and no nitrogen oxides. It is very likely, therefore, that ammonia is dissociatively adsorbed (analogous t o water) on the catalyst in the form of NH; or NH2- fragments. The latter is very similar t o 02-and probably can replace 0'- surface anions of the catalyst, particularly at higher temperatures (400°C). Regarding the formation of by-products (acetonitrile, HCN, carbon oxides), little is known in detail. The first reaction step, formation of symmetrical allyl intermediate, is likely to be common for both main and side reactions. The work of Cathala and Germain (see below) indicates that these side reactions are complex processes that partially occur in the homogeneous gas phase. As to the kinetics, the reaction is first order with respect to propene and the rate is independent of the oxygen and ammonia concentrations References PP. 253-262
166 in most cases, indicating that the dissociative propene adsorption is ratedetermining. This is also confirmed by the activation energies, which are often equal for both oxidation and ammoxidation.
( b ) Catalysts As a consequence of the similarity between oxidation and ammoxidation, the same kind of binary and multi-component oxide mixtures are active and selective catalysts for both processes. Important, highly selective binary combinations are: Bi-Mo, U-Sb, Fe-Sb and Sn-Sb. The majority of the commercial catalysts are based on one of these combinations, but may contain minor amounts of several additional metallic and non-metallic oxide compounds. Because ammonia may replace active oxygen anions on the catalyst surface, catalysts exist which are selective for ammoxidation but not for oxidation. An example of this kind is the Bi-Sb--O catalyst studied by Barannik and Ven'yaminov [ 391. ( i ) Bismuth moly bdates. Results obtained with the original Sohio ammoxi(50 wt. 7%)) are reported by Callahan dation catalyst (Bi9PMo12052/Si02 et al. [70].With a reactant ratio C,H,/NH,/air of 1/1.1/13, a selectivity of 68% is obtained at a conversion of 96% (470°C). The main by-products are HCN, CH3CN and carbon oxides. The kinetics of acrylonitrile formation are described by a triangular scheme in which the reaction path via acrolein is of minor importance, viz.
C3H6
(1)
C3H3N 7
\
The individual reaction steps all appear t o be first order with respect t o propene (or acrolein), the rates being independent of the oxygen and ammonia partial pressures. Kinetic data are collected in Table 10. The kinetic studies of Cathala and Germian [81--831 with a Bi/Mo = 1 bismuth molybdate catalyst have a wider scope and include the effect of
TABLE 10 Kinetics of ammoxidation of propene at 430°C Reaction step 1+2 2
3
First-order rate coefficient (sec-' ) 0.2 <0.01 0.4
Activation energy (kcal mol-' ) 17-1 9 7
167 TABLE 11 Effect of steam o n the ammoxidation of propene (catalyst with Bi/Mo = 1 )
With water Without water
100 100
3.3 2.4
3.3 4.3
2.0 0
535 1680
18 40
0 0
37 18
steam and some side reactions. A flow reactor was used a t 460-475°C and a feed gas consisting of C3H6/NH3/H20/air= 1/1.5/0--1/10-9 at 1 atm. The maximum yield at 460°C was increased from 66.5 to 73.5% by the presence of steam. Numerical values of the (pseudo) first-order rate coefficients are summarized in Table 11according t o the reaction scheme adopted by the authors, viz.
Table 11 shows that water primarily inhibits the combustion of propene, and thus increases the (initial) selectivity. A comparison of rate coefficients of oxidation and ammoxidation is given in Table 12, and includes the separately studied (amm)oxidation of acrolein. A more complicated reaction scheme is proposed by the authors t o include the formation of the by-products acetonitrile, acetaldehyde and ethylene. However, appropriate rate coefficients cannot be given as the reactions appear to be partially homogeneous gas phase reactions, implying that factors like the reactor geometry are also involved. Regarding the oxidation mechanism, the authors assume that two hydrogen atoms are first abstracted from propene, followed by reaction with surface oxygen or NH species. TABLE 1 2 Kinetic comparison of ammoxidation and oxidation of propene and acrolein.
Ammoxidation Oxidation Ammoxidation Oxidation
C3H6 C3H6
C3H40 C3H40
References p p . 253-262
4.8
0.16 5.85
0.1 2.0
25.8 145
0.9 28
6.2 22 97
168 Wragg et al. [354] investigated two bismuth molybdate catalysts in a static reaction system: a Bi/Mo = 0.73/1 catalyst and a pure koechlinite phase catalyst (Bi/Mo = 2/1), prepared by Batist and Lankhuijzen [ 451. The first catalyst was equally active for oxidation and ammoxidation, and the activation energy was the same for both (25 kcal mol-'). The pure koechlinite catalyst, on the other hand, was twice as active (at 400°C) for ammoxidation as for oxidation, with activation energies of 9 and 20 kcal mol-' , respectively. Moreover, this catalyst showed an exceptionally high ammoxidation selectivity (95% instead of only 50% for the other catalyst). In the absence of oxygen, the remarkably low value of E A (=9 kcal mol-' ) was not observed, but instead a normal value of 21 kcal mol-' was found. The abnormally low activation energy is obviously due t o the occurrence of a different ratedetermining step, e.g. the reaction of the adsorbed allyl complex with an NH species instead of the dissociative adsorption of propene. The same koechlinite catalyst was recently investigated by Lankhuijzen et al. El86 and private communications] using a flow reactor. They observed that the ammoxidation at 400°C was practically zero order with respect to ammonia and propene, at an average feed composition (in mmol 1-') of C3H6 (1.5), NH3 (1.2), and O2 (2.6). A dependence of the oxygen partial pressure is presumably absent too, because of the well known very rapid reoxidation of partially reduced bismuth molybdates. The conversion space-time relation disagreed with first-order kinetics and is explained by strong adsorption (and hence inhibition) by the reaction products (acrylonitrile and acrolein). To explain the zero-order kinetics, the authors assume that the transport of oxygen from the reoxidation site to the reaction site, probably through the lattice, may be ratedetermining. At higher temperatures, e.g. 450"C, the kinetics change into normal first-order dependence with respect t o propene. Matsuura [ 2091 has proposed an ammoxidation mechanism similar t o that for the formation of acrolein [see Sect. 2.2.2(d)(i)]. Ammonia is assumed to adsorb dissociatively on B-sites, followed by reaction with A-site oxygen, finally leading t o the replacement of the A-site 0 2 -by a NHZ- species. The A-site is thus altered from a site which introduces oxygen into a site that introduces nitrogen in the allyl complex.
(ii) Other catalysts containing Bi and M o . Molybdates with the scheelite structure have been investigated by Sleight et al. [ 2981 for the oxidation [see Sect. 2.2.2(d)(iii)] and ammoxidation. There is clear evidence that, for a highly selective catalyst, the presence of defects (cation vacancies) and bismuth ions are essential. In the scheelites investigated, Bi was combined with several other metals the nature of which does not seem t o be very important. Selectivities up t o 99% (initial) and acrylonitrile yields up to 76% have been obtained at 459°C (Table 13). The mechanism of acrolein formation has already been mentioned above. Regarding the
169 TABLE 1 3 Complex phases in scheelites and propene ammoxidation Temperature, 45OoC; flow time, 4.0 sec; feed gas, 4.0% C3&/4.8% NH3/7.7% air, 43.5% NZ. Composition a
Conversion to acrylonitrile (%)
73 60 62 Pb0.45 5 C a 0 . 4 S 5 B i 0 . 0 6 ~ 0 . 0 3 M o 0 . 5 ~ 0 . 5 ~ 4 73 Pb0.44 CaO .4 4 BiO .O 8 4 0 .04 Mo04 67 Pb0.44Sr0.44Bi0 .0 840 .O4MO04 76 PbO. 8 8Bi0. 0 840.04M00, 5 WO. 5 0 4 70 Pbo.68Ago.i oBio.1
[email protected] 10 Pb0.834N~0.083~~0.083~004 64 Pb0.852Na0.042Bi0.08 540.02 IMo04 Pb0.42sNao.i ~ Y o . z ~ B ~ o . o s ~ o . o ~ s M ~ ~ ~50 33 Na0.38Y0.5Bi0.04~0.08Mo04 1.6 Pb0.5Bi0.5V0.5M00.504 41 Pb0.79Bi0.1
[email protected] 35 Bi0.9840.0Zv0.4 7AS0.47M00.0604 Pb0.9 I Bi0.06~0.03M00.SW0.504
[email protected]~004
a
4 represents cation vacancies.
action of ammonia, the authors assume that ammonia adsorbs at the defects and reacts with an oxygen anion to form water and a NHZ- group that replaces the oxygen anion. This NHZ- species reacts with the ally1 intermediate. A cerium-doped bismuth catalyst, Bi9CeMoIZOS2 /SiOz (50 wt. %), was studied by Giordano and Bart [140] using a flow reactor at various temperatures. Very high selectivities appear feasible with this catalyst, i.e. a 90% acrylonitrile selectivity, a t a 90% conversion level obtained at 480°C with a feed composition propene/ammonia/air of 1/1/10. More by-products (HCN, CH3CN, CO, COz) are formed a t lower temperatures (380420°C). The authors suggest that the essential parameter that governs the selectivity is the Mo6+/Mo5+ratio which is determined primarily by the propene/oxygen ratio. They assume that Mo6+sites effect the acrylonitrile formation, while Mo5+sites are involved in the formation of by-products.
(iii) Uranium an timonates. Uranium antimony oxide catalysts are well known commercial ammoxidation catalysts, and have been used in industrial acrylonitrile manufacture for several years. Due t o some disadvantages (e.g. instability under reducing conditions) catalysts based on Sb-Fe or Bi-Mo are preferred a t present. Two ternary compounds, USbOS and USb3OI0 occur in the U-Sb-O system, as was shown by Aykan and Sleight [ 341 usingvarious analysis techniques. ESR measurements revealed that, in both components, the formal References PP. 253--262
170
oxidation state of U is 5+;hence Sb must be in the 5+ state too. The relation between structure and activity has been studied by Grasselli and Suresh [146]. Both ternary phases are active, but USb3010displays the highest activity, and thus is the catalytically important phase. The formation of these phases during the catalyst preparation appears to be a very complex process. From experiments with deuterated propene, the conclusion was reached that the first hydrogen abstraction is ratedetermining, and that this step is followed by abstraction of the second hydrogen, before oxygen or nitrogen is incorporated. The authors further presume that the function of U is t o stabilize Sb in the 5+ state. The active site may be depicted by
The ally1 radical is thought t o adsorb on Sb, while the abstracted hydrogen forms a hydroxyl group connected t o U, which adsorbs as water when the second hydrogen atom is transferred. Ammonia is supposed t o react with the catalyst according t o
\A/
0
\ I / Sb + N H , 'l\0/l\
\ +
I ,NH,
U 'l'/I'
I
/
Sb
+H,O
Thus the introduction of NH is similar t o the introduction of 0 which takes place in the oxidation t o acrolein. Pulse experiments in the absence of oxygen proved that lattice oxygen is indeed used. However, the activity rapidly declines and the reduction should be confined to USb3OI0 USb309.945to avoid structural changes. +
(iu) Iron antirnonates. Iron antimonate-based catalysts are very selective, and are of industrial importance for acrylonitrile production. The iron antimony oxide system was investigated by Boreskov et al. (Hucknall [160], ref. 414), who reported (in 1969) that at the antimony-rich side of the system, a range of very selective and moderately active catalysts is formed. Remarkably, no recent publications have appeared. ( u ) Tin antirnonates. Tin antimony oxide catalysts are very selective
(amm)oxidation catalysts which are of industrial importance. Their structure and behaviour in the oxidation of propene to acrolein is described above [Sect. 2.2.2(d)(v)]. Recent contributions particularly concern pulse experiments, in which Sn-Sb-0 catalysts are compared with bismuth molybdate catalysts. Some essential differences have become evident. With regard to the selectivity of the ammoxidation, Sn--Sb--O catalysts are superior to bismuth molybdates. According to Germain and Perez
171 [135], an initial selectivity of 100% with respect t o both propene and ammonia is achieved in a pulse reactor a t 400°C for Sn/Sb = 1/4 and using a feed ratio C3H3/NH3/02/Heof 1/1.5/1.5/6. The high selectivity with respect t o ammonia (a striking difference from bismuth molybdates) is confirmed by Barannik and Ven’yaminov [ 391 who report that, remarkably, ammonia is not oxidized over a Sn/Sb = 2/1 catalyst at 350 and 400°C. These authors further observed that pre-adsorbed ammonia can participate in the nitrile formation, and increases the selectivity. Apparently, the most weakly bonded (i.e. the most active) oxygen ions are blocked or replaced by NH, species. Trifiro et al. [ 3251, studying a Sn/Sb = 5.66 catalyst at 400”C, report that (excess) ammonia has an inhibiting effect on the ammoxidation in contrast t o bismuth molybdates, where the rate is independent of the ammonia concentration. Other points of difference concern the oxygen dependence. A non-zero reaction order is observed when the OZ/C3H6 ratio is smaller than 1.2 indicating that the reoxidation rate is rather slow compared with that for Bi-Mo-0. A redox mechanism is certainly operative, but the oxidation capacity, demonstrated by pulse experiments in the absence of oxygen, is very small, indicating that this catalyst has much less the character of an oxygen “sink”. A commercial iron-promoted catalyst (Sn/Sb/Fe = 1/4/0.25) was studied by Germain e t al. [ 92,93,135,137]. Iron is reported t o improve the ammoxidation qualities of the catalyst although it has no effect on the oxidation [93]. The kinetics, determined in a flow reactor at 445°C and with a feed ratio C3H6/NH3/air= 1/1.2/10, are essentially similar for this catalyst and bismuth molybdate. The initial selectivity is 80% and the maximum yield is 65% (at 445°C). The initial selectivity markedly depends on the temperature (e.g. 91% at 415°C and 72% a t 507°C). The effect of water is hardly significant for this catalyst: the acrylonitrile formation is slightly inhibited, while some more acrolein is formed. Presumably, water and ammonia compete in the interaction with the catalyst, which is much less reactive with respect to ammonia than bismuth molybdate. The acrolein ammoxidation is very rapid (about six times the propene ammoxidation rate) and selective (86%).A comparison of the Sn-Sb-Fe-0 catalyst with bismuth molybdate is presented in Table 14.
(vi) Miscellaneous. A Bi-Sb-O catalyst (Bi/Sb = 65/35) was studied by Barannik and Ven’yaminov [39] with respect t o the influence of ammonia on the (amm)oxidation selectivity at 350 and 400°C. Pulse experiments proved that this catalyst has ammoxidation qualities comparable with Sn-Sb-0, although in the absence of ammonia the oxidation selectivity is very poor. Apparently ammonia completely blocks or replaces the active oxygen ion a t the catalyst surface. A Fe-Bi-P-0 catalyst (Fe/Bi/P = 44.5/44.5/11) was investigated by Japanese workers (Miyake et al. [221], Oka et al. [242]). Both propene References p p . 253-262
172 TABLE 14 Sn-Sb-Fe-0 catalyst for ammoxidation of propene compared with Bi-Mo-0 C3H6/NH3/air = 1/1.2/10; 460°C. Sn-Sb-Fe-0 Bi-Mo-0 Atomic ratio Initial selectivity to acrylonitrile (%) acrolein (%) Maximal conversion to acrylonitrile (%) acrolein (%) Activation energy (kcal mol-' ) acryloni trile acrolein Reaction velocity (mole h-' g-'
1/4/0.25
X
lo3)
80 2.5
73.5 5.7
66 1.5
67 1.6
17.5-20.8 35 3.5
17 38 6.6
and acrolein ammoxidation were studied in a flow reactor at 340-400 "C. Fractional orders with respect t o oxygen and ammonia are reported for both reactions. A number of Mo- or Sb-based mixed oxides was compared by Germain and Perez [135] using a pulse reactor. The main results with respect t o activities and selectivities are collected in Tables 15 and 16. The selectivity of the acrylonitrile formation with respect to ammonia is very low (
Bi-Mo 1/1 Fe-Mo 1/1 Ti-Mo 1/1 U-MO 3/2 Sb-Mo 2/1 Sn-Mo 2/1 V-Mo 6 / 1 Sb-Sn 4/1 Sb-Sn-Fe 4/1/0.25
Temperature ("C) at a conversion o f 2% m-*
8% m-2
330 410 348 415 360 490 265 400 400
350 472 378 48 2 428 540 314 455 440
173 TABLE 16 Selectivities t o nitriles €or ammoxidation of propene Catalyst Bi-Mo Fe-Mo Ti-Mo U-MO Sb-Mo Sn-Mo
V-MO Sb-Sn Sb-Sn-Fe
Temp. ("C)
Acrylonitrile
Acetonitrile
Total nitriles
(a)
(%)
(%)
330 430 410 470 348 468 415 470 360 440 490 590 265 390 400 500 400 500
57,5 87,5 26 45 8 35 26 26 19 55 29 18 10 35 100 68 100 83
37,5 6 43 50
95 93,5 69 45 58
60
86
50 3 26
69 58 55 22 573 65 100 68 100 83
4 47,s 30 0 0 0 0
A silica-supported Sn-V-P-0 catalyst (Sn/V/P = 1/9/3) was investigated by Onsan and Trimm [244].Working with a flow reactor at about 520"C, a maximum selectivity of 75% t o acrylonitrile was reached at a contact time of ca. 230 g sec 1-' and an oxygen/propene/ammonia ratio of 2/1/1.75. The authors assume that the six principal products (acrylonitrile, acetonitrile, HCN, CO, COz, N,) are formed by six parallel reactions and in the first instance apply power rate equations. A more detailed analysis reveals that a LangmuirHinshelwood type rate equation, surface reaction being rate-determining, properly describes the production of acrolein plus acrylonitrile from propene, viz. KRH[RH 1 K OL o 21 (1 4- KRH[RH]+ K0[021)* for which the parameters calculated are given in Table 17. Experiments were also done t o compare the propene ammoxidation R=k--
TABLE 17 Parameters for t h e ammoxidation of propene at 550°C _-_____-__
k = 1.3 i 0.2 X mol g-' sec-' KcH = 1.0 t 0.2 I mol-' K O = 8 3 t 21 mol-' --. References p p . 253-262
-__--
-
174 with the propene oxidation and with the acrolein ammoxidation. The conclusions are that propene ammoxidation and oxidation have about equal rates, while the acrolein ammoxidation is relatively slow. 2.3 BUTENES
The possibilities for selective oxidation of butenes present a complex picture because of the various ways of introducing oxygen into the molecule and the possibilities of dehydrogenation and isomerization. Finally, there are catalysts with a capacity for producing dimerization and aromatization. The conversion of isobutene to methacrolein is closely related t o the selective oxidation of propene to acrolein and demands similar catalysts. It has been verified that the same mechanism applies, involving a symmetrical allylic intermediate, viz.
8CH2
CH3-Cq5
CH2 Compared with propene, the oxidation of isobutene is more rapid but less selective, yet selectivities of over 75% appear feasible. Combustion is the main side reaction. One would expect that some considerable attention would be shown in the literature to isobutene oxidation as a route t o the industrially important methacrylic acid, but this is not the case. Nor is it with the production of methacrylonitrile, analogous t o the propene ammoxidation. Only in the patent literature is a high activity noticeable. With straight chain butenes, there is a larger number of possibilities. With mild oxidation catalysts, the reaction is directed t o dehydrogenation, e.g. for 1-butene CH,*H2--CH=CH2 -+
+
CH3-CH-CH-CH2
+
CH24H-CH-CH2
CH,=CH-CH=CH,
The stabilization by conjugated double bonds makes this sequence energetically attractive. It can be easily understood, however, that the activated butene will also have a tendency t o isomerize. Consequently, oxidation products are often found together with isomerization products. With the allylic intermediates, the introduction of oxygen is also possible and crotonaldehyde is one of the first oxygenated products t o be formed. This molecule is not very stable, certainly not under reaction conditions, and will be further oxidized t o crotonic acid or the relatively more stable furan. More severe oxidation will then result in the generation of maleic acid and its anhydride. A summary of the selective oxidation reactions is presented in the
175 scheme below. Besides these selective oxidation steps, side reactions occur, i.e. breakdown t o aldehydes and acids lower than C 4 , and combustion. The relative importance of these side reactions, which in fact apply to each component in the scheme, depends on the particular catalyst and reaction conditions employed. CH2=CH-CH**H3
[CH*-CH-CH*H3] /
diene formation / ;-H
+ O . -H
li!
CH3-CH=CH-CH0
CH2=CH-CH=CH2
(CHO-CH=CH-CHO) H HC=CH
HC-C HC-C
C-COOH
\ \\
0
Economically, the oxidation of butenes is of importance mainly for the production of butadiene and maleic acid anhydride. Dimerization and aromatization have been reported for isobutene and n-butenes, analogous t o propene, over catalysts like Bi203. Isobutene radicals dimerize more easily than n-butene radicals, which are less stable and rapidly form butadiene. Finally, oxyhydration reactions can also occur, analogous to the conversion of propene t o acetone. n-Butenes can yield methyl ethyl ketone. 2.3.1 Isobutene to methacrolein and methacrylonitrile
Research on the selective oxidation of isobutene is mainly focussed on bismuth molybdate-based catalysts. Under conditions which are optimal References p p . 253-262
176 TABLE 18 Kinetic parameters for isobutene oxidation
_ _ _ _ - - _ _ _ ~ - . - ~
Methacrolein
co
co2
10.87 62.05
15.11 196.48
15.06 230.97
~~
Activation energy (kcal mol-' ) Pre-exponential factor (mol g-' h-' atm-' )
for production of acrolein from propene, the best results are also arrived at in the production of methacrolein. Maissant e t al. [194]find a selectivity of 83% at 450°C (compared with 90% for acrolein) and an activity which is about ten times as large as in the case of propene oxidation (75 mol h-' g - ' ) if a Bi/Mo = 1/1catalyst is used. Ray and Chanda [261] studied bismuth molybdates (prepared by the method of Peacock [250,251]) in an integral flow reactor. At constant W / F = 8 g h mol-' and a feed ratio isobutene/oxygen = 1/6, a maximum selectivity of 75% was found at 400-450°C. As with propene, the reaction is first order with respect to isobutene and the rate is independent of the oxygen pressure. The reoxidation of the catalyst is very fast. Assuming that the kinetics can be described by three parallel first-order reactions for the production of methacrolein, carbon monoxide and carbon dioxide, rate coefficients were calculated (Table 18). Mann and KO [ 2021 likewise examined the selective oxidation of isobutene on bismuth molybdate. With an integral flow reactor, the highest selectivity was obtained at over 30% conversions for an oxygen/olefin ratio of 2/1 and a W / F = 2.5 g h mol-' (390°C). The data were correlated with a rather complicated Langmuir-Hinshelwood expression inconsistent with a redox mechanism. This was based on a rate-controlling step between adsorbed isobutene and adsorbed oxygen, and included an inhibiting effect of methacrolein by competitive adsorption with isobutene, viz. KSKHPH . KoPo (1+ KHPH + KMPM) 1+ KoPo where H = C4H,,M = C4H60and 0 = 0,; K , , KH, K M and K O are temperaturedependent, and can be expressed by the equations
R=
In In
KCqH8
=
5.18 + (17.79 X 103/T)
KC4H60
=
4.30 + (24.47 X 103/T)
=
0.75 + (3.64 X 1 0 3 / T )
=
7.44 - (7.32 x 103177)
In KO2 In K s
Muizebelt and Van Ooij [234] studied catalysts of general formula
TABLE 19 Influence of the composition of Fe,BibKo, ?CrCo6MoIoOy on its performance is isobutene oxidation catalyst in a flow reactor at 4OO0C Catalyst composition
a
b
0 1 1 1 1
1 2 1 0.5 0
Relative activity
0.2 0.5 1 1 0.07
Selectivity to methacrolein
Reaction order with respect to
(%)
Isobutene
85 75 83 85 45
0.7 -0.6 -0.7 -0 1.0
Bi/Mo ratio at the surface ( X 1 0 ) (ESCA)
Surface area (m2 g 4 )
3.8 0.8 0.5
4 .O 6.5 3.5 3.4 8.9
Oxygen
-0.4 1.1 1.4 1 0.1
178 Fe,Bi,KCCrCo,Mo,,O,. It appears that both bismuth and iron considerably influence the activity as well as the oxidation kinetics, as shown by Table 19. Bismuth is indispensible for both activity and selectivity of the catalyst. The addition of iron increases the activity but at the same time changes the kinetics: in the absence of iron, the kinetics approach the usual dependencies, i.e. first order in olefin and zero order in oxygen, while, in the presence of iron, the process is approximately first order in oxygen and zero o r a negative order in isobutene. The oxygen dependence signifies that reoxidation has become a relatively slow step in the reduction-oxidation cycle. ESCA analysis has accordingly revealed that a higher degree of reduction was present in the working catalyst. It should be noted that a similar effect of iron is reported by Daniel and Keulks [94] for Fe-Bi-Mo-0 catalysts compared with pure bismuth molybdates. In that case, the difference was ascribed t o the formation of a ternary phase, which was recently identified as Bi3FeMo,012 [188]. To explain the observed negative order with respect to isobutene, Muizebelt and Van Ooij suggest that isobutene may strongly adsorb on the reduced sites of the catalyst (presumably bismuth sites) and thus hinder the reoxidation. A rather complicated, modified redox model is proposed to describe the reaction kinetics. Morita e t al. [222] compared bismuth molybdate (1/1)with U-Sb oxides (1: 2) at 400°Cin a continuous flow system. The methacrolein selectivity for U-Sb is significantly higher than in the case of Bi-Mo (see Table 20). These values increase slightly with increasing conversion of isobutene. Isobutene itself retards the oxidation. In contrast t o the propene oxidation, addition of steam accelerates the reaction up t o a factor 4 with U-Sb and t o a smaller degree with Bi-Mo. With the first catalyst, the activation energy decreases from 27 t o 18 kcal mol-' (0.23 atm steam). U-Sb seems to be less stable than Bi-Mo, but steam has a beneficial effect here t o o (Table 20). TABLE 20 Effects of added steam o n isobutene oxidation a t 400°C a Partial pressure of steam added ( a t m )
Cat a1yst
0
0.05
0.10
0.23
0.30
0.40
~
U-Sb
Xb S C
Bi-Mo
Xb S C
0.249 0.702 0.190 0.630
0.309 0.822
0.368 0.861 0.209 0.792
a The time factor was 7 ( g cat h mol-I).
Isobutene conversion. Methacrolein selectivity.
0.440 0.915 0.223 0.816
0.132 0.917 0.216 0.875
179
Zhiznevskii et al. chose another group of catalysts for the production of methacrolein, based on the solid solution systems Mo-Te oxides [110,360] and Sb-Te oxides [361] which were investigated in a pulse reactor at 340°C. Selectivities of 70--90% are reported. The Mo-Te catalyst (1 : 4) has a unique dependence of the activity upon the degree of reduction of the catalyst surface. Initial reduction t o 6% increases activity, but further reduction lowered this again. The selectivity changes in a similar manner. X-Ray investigations led the authors t o the conclusion that some Te-Mo-P-phase, which is formed with reduction but decomposed if the reduction is taken too far, is responsible. Combinations of Sb205and TeO, can likewise be selective catalysts. The relatively low activity at 400°C is much improved by the addition of molybdenum (60 mol.%). It can be concluded that bismuth molybdates and uranium antimonates give the highest yields of methacrolein coupled with high activities. F'romotors have a strong effect on the kinetics. The results are much better than with the older Cu,O catalyst, although Mann et al. [200,201]mention that Cu,O can give good yields when halogen compounds are added to the feed as modifiers. The ammoxidation of isobutene has not received much attention. The only contribution in this field is by Onsan and Trimm [2.44]for a rather unusual catalyst, a mixtyre of the oxides of Sn, V and P (ratio 1/9/3) supported on silica. At 520 C, a maximum selectivity to methacrylonitrile + methacrolein of 80% was reached with a Sn-V-P oxide catalyst (ratio 1/9/3), an isobutene/ammonia/oxygen ratio of 1/1.2/2.5 and a contact time of 120 g sec 1-' . The kinetics are very similar t o those for the propene ammoxidation. Again, the data are initially analysed by means of (parallel) power rate equations, for which the parameters were calculated, while a more detailed analysis proves that a Langmuir-Hinshelwood model with surface reaction as the rate-controlling step provides the best fit with regard to the two main products. At 520°C, the equation which applies for the production of methacrolein plus methacrylonitrile is
R=
3.7 [ RH] (1 + 240[RH] + 5.7[0,])*
It is interesting t o note that the reactivity of an allylic hydrogen atom in isobutene is about 2.5 times the reactivity of such an atom in propene. Onsan and Trimm note that the same ratio is found for oxidation by bismuth molybdates and also for hydrogen abstraction in an iso-octane solution by methyl radicals.
2.3.2 Normal butenes t o butadiene, furan and maleic anhydride The dehydrogenation of butenes t o butadiene is the most selective butene oxidation process and has received considerable attention in the References PP. 253-262
180 literature. Two types of catalyst are effective for this process: (a) mild catalysts, i.e. the same catalysts that oxidize propene to acrolein and isobutene to methacrolein and (b) iron oxide catalysts and iron-containing spinels, which, apart from selective dehydrogenation, only catalyze combustion reactions. Deeper oxidation products like furan and, particularly, maleic acid anhydride, can be produced by catalysts that have a stronger oxidative power than the above type-(a) catalysts, but, at the same time, have retained the capacity to transfer oxygen selectively t o the organic molecule (a capacity which is absent in the type-(b) catalysts). Besides, a more acidic character of the catalyst surface is probably required t o produce an acidic product like maleic anhydride effectively. The most interesting catalysts of this group are V z 0 5-based catalysts and certain molybdates and Moo,-based catalysts. There is no strict distinction between mild catalysts that produce mainly butadiene and the stronger oxidizing catalysts that can yield maleic anhydride. The reaction conditions are also an important factor. The contributions are classified, therefore, according t o the catalysts studied, instead of a division with respect t o the main product that is formed.
(a) Moly bdates ( i ) Bismuth molybdates. Bismuth molybdates are very active catalysts for the dehydrogenation of n-butenes at 430-500°C. Very high selectivities (90-95%) can be obtained, even at high conversion levels (>80%).It is generally found that, in this temperature region, the reaction is first order in butene and zero order in oxygen. Steam may have a positive effect on the activity according t o Komarovskii et al. [ 1791. A doubling of the reaction rate was observed by adding up t o 20%steam t o the oxidation of a mixture of n-butenes in a flow reactor over a Bi/Mo catalyst of unknown composition at 420-480°C. The same authors [179] also studied the influence of the oxygen concentration, which was found t o have no effect on the kinetics at O,/butene > 0.4. Furthermore, a rather complex set of kinetic equations was derived to describe side reactions (isomerization, and formation of carbonyls, acids and furan). Butadiene has a strong inhibiting effect on the reaction at lower temperatures (320-400°C), correlated with a strong increase of the apparent activation energy. Batist et al. [43,172] studied Bi/Mo = 2/3 and 2/1 catalysts using various reactor types. An activation energy of 10-11 kcal mol-' is found by pulse experiments, irrespective of the temperature, and assigned t o the uninhibited reaction. However, a value of about 36 kcal mol-' is observed in the range 30O-37O0C, both for experiments in flow and recirculation reactors and for pulse experiments in which butadiene is added to the pulses. At higher temperatures, the lower value of about 11 kcal mol-' is again found.
181 The oxidation reactivities of the three butenes are different. A ratio of 1/0.63/0.5ri is found by Pulse experiments using a Bi/Mo = 2 / 3 catalyst at 385°C and similar results are found with a Bi2Mo0, catalyst [ 1721. The difference in rate proves that, for bismuth molybdates at this temperature, the isomerization is slower than the oxidation. At much lower temperatures ( 15O-25O0C), the situation is reversed, isomerization being the prevailing reaction, as shown by Batist e t al. [40] for a koechlinite catalyst. Remarkably, the isomerization is also strongly inhibited by butadiene, indicating that the isomerization and oxidation sites are probably the same. The activation energy of the isomerization reaction was found t o increase from the normal value of 2-4 t o 17-19 kcal mol-' in presence of butadiene. The selectivity of the diene formation as a function of temperature was investigated by the same authors [ 1721 for various catalysts by means of pulse experiments. The selectivity appears t o decrease from above 90% to about 50% with increasing temperature (300-420°C) and catalysts with Bi/Mo = 2/1 and 1/1,while a rather constant level of 70-7576 is observed for the Bi/Mo = 2/3 catalyst at 350-420°C. These rather low selectivities may have been caused by the relatively high O2/butene ratio (3.3) and the applied pulse technique. It is well known that the butene oxidation rate is practically the same in the presence as (initially) in the absence of oxygen, which confirms the validity of a redox model. In contrast t o the propene oxidation, the activity does not seem t o decline very rapidly with increasing reduction as demonstrated, for instance, by Batist e t al. [43]. This is evidence that much more than one surface layer of oxygen can be consumed and, moreover, implies that diffusion of oxygen through the lattice is a fast process. The latter is also confirmed, for instance, by the reduction-oxidation experiments of Beres e t al. [ 471. With respect t o the mechanism, the adsorption measurements of Matsuura and Schuit [207-2091 are of interest. The assumption of Aand B-sites is reported in some detail for the propene oxidation in Sect. 2.2.2(d)(i). As for the oxidation of propene, the abstraction of (both) H-atoms is assumed to occur on the molybdenum layers by initially forming HOB groups, followed by H-transfer from OB t o OA and desorption of water ( H 2 0 A ) . Abd. El-Salaam e t al. [ 11 have studied the catalytic activity of various bismuth molybdates for the oxidative dehydrogenation of 2,5-dihydrofuran t o furan. A close correlation with the oxidation of butene to butadiene is expected and was indeed observed. Bismuth molybdenum oxide combinations that additionally contain some phosphorus were investigated by Ai and Suzuki [ 8,121. The bismuth content was varied at a constant ratio P/Mo = 0.2. The effect of this variation on the overall activity for the oxidation of cis-2-butene as a function of the temperature is shown in Fig. 5. The Bi/Mo ratio appears t o have a RL-butene/Rcis-2-butene/Rtrans-2-butene
References p p . 2 5 3 - 2 6 2
182
C
e
L
al C
0
u
-
I
I
I
I
4 00
450
500
550
Temperature ( " C )
Fig. 5 . Dependence of butene conversion on temperature and composition of catalyst (Bi/Mo ratio).
strong influence on the activity, with a maximum at Bi/Mo = 0.1. This maximum coincides with a maximum in the selectivity with respect to maleic anhydride and with a maximum of the butene isomerization activity. The selectivity (with regard to maleic anhydride) amounts to 33% at a 80% conversion level of butene, while a selectivity of 63%at 90% conversion is found when butadiene is used as the starting material. The rate of oxidation of butene over the Bi/Mo = 0.1 catalyst can be described by
in which the order of magnitude of h is about 1sec-' at 460°C.The authors correlate the catalytic qualities of the catalysts with their acid-base properties. A maximum of surface acidity is found at Bi/Mo = 0.1, while the basicity is very low at low bismuth content, but sharply increases at
I
183 Bi/Mo = 1. The possible significance of acid-base properties in oxidation catalysis is more generally discussed in Sect. 3. Villa et al. [340] have shown that the bismuth tungstates are comparable with bismuth molybdates with respect t o dehydrogenation catalysis, although activities and selectivities are somewhat lower. Although the phase structures are different, interesting catalysts are formed in a similar composition range Bi/W = 2 / 3 t o 2/1. (Note that, in case of propene (amm)oxidation, tungstates are definitely inferior to molybdates.)
(ii) Multi-component catalysts containing Bi and M o . Scheelite-type catalysts containing defects (non-ordered cation vacancies) were studied by Sleight and Linn [297]. It appears that, in general, catalysts which contain both bismuth cations and defects are the most selective. For all catalysts studied, the activity strongly increases when the fraction of cation vacancy increases from 0 t o 0.1. Most catalysts are practically inactive when defects are absent. The selectivity is also increased by the presence of defects. Examples of very selective catalysts, studied in a differential reactor are presented in Table 21. Two catalysts which d o not contain bismuth are also included in the table. Although these catalysts display a somewhat lower activity, the difference is less pronounced than in the case of the propene oxidation, where bismuth 'is indispensible for a good selectivity [see Sect. 2.2.2(d)(ii)]. The kinetics appear t o depend on the composition of the catalyst and the temperature. A close correlation is generally observed for the rate dependence on oxygen and butene, inhibition by butadiene and the apparent activation energy. Some results are collected in Table 22 and show that, in the case of a non-zeroorder reaction with respect t o oxygen, a high activation energy occurs (butadiene inhibition) while in the case of first-order kinetics with respect to butene, a low activation energy is found (absence of product inhibition). Iron bismuth molybdenum oxides have been studied by Linn and Sleight [188] and Batist et al. [44].The latter studied catalysts conTABLE 2 1 Scheelite catalysts Catalyst
a
~~
~
Pb0.8 2 BiO. 1 2 40.0 6 M 0 0 4
[email protected]
Na0.455Bi0.S I
[email protected] Bi0.9@0. I V0.7M% .30 4
[email protected] [email protected]
a
4 represent
defects.
References p p . 2 5 3 - 2 6 2
Temp. ("C)
Selectivity
325-425 430 430 430 430 430
98 98 99 99 96 75
(%)
~
184 TABLE 22 Kinetics with scheelites -
___~__.-._
Catalyst
a
Temperature range ("C)
Reaction order with respect to Oxygen
BizMoOG Biz(M004)3
[email protected]
Pb0.88Ce0.0
[email protected]'004
[email protected] 1 0 4
a
3 2 5-390 390-450 325-4 25 3 50-390 39 0-4 50 350-390 4 10-4 50
0 0 Fractional
0.39 0.16
Activation energy (kcal mol-I)
Butene <1 1 0
0 1
35-3 8 12-14 35 36 16 35 11
@ represents defects. Reaction orders determined for
[email protected] 1Mo0.03V0.9704 at 370 and 450°C.
sisting of BiZ(MoO4), and Fez(Mo04), in various ratios and discovered the presence of a ternary phase over a wide range of composition (Fe/Bi = 4/1to 1/5). Linn and Sleight studied the system in detail and found that there probably exists only one ternary phase, having a scheelite structure represented by Bi3(Fe04)(Mo04)2[see also discussion in Sect. 2.2.2(d)(ii)]. This pure phase was studied for the oxidation of 1-butene. Both this pure Bi3FeMo2Ol2and the iron-promoted bismuth molybdates of Batist et al. have catalytic properties that are very similar t o those of pure BizMo06. The oxidation is first order with respect to 1-butene, the rate being independent of the oxygen concentration. The activation energy strongly depends on the temperature. Values of 50 and 18 kcal mol-' are found at 325-390 and 400-450°C, respectively, indicating butadiene inhibition in the low-temperature region. The oxidation rates are essentially equal for the iron-containing catalysts and BizMoo6 , while selectivities above 95% are found for all catalysts. Only the ratio of isomerization t o oxidation rate is slightly different. The great similarity between Bi2Mo06 and Bi3FeMoz0,, is remarkable as the structures are very different (e.g. octahedral M-0 coordination in the former and tetrahedral in the latter). One might suppose that the same phases are present at the surface. However, ESCA measurements have not revealed essential differences between surface and bulk compositions for either of the catalysts. Complex multi-component catalysts were studied by Wolfs et al., [ a l l , 351J . The catalysts had the general composition Me:'Me:"BiMo I Z O ~ ~ . The small fraction of Bi-cations present in these catalysts appears sufficient t o achieve an active and very selective catalyst. The authors suggested that the catalytic qualities are connected with the presence of an outer shell of BizMoO, around a nucleus of other molybdates. However,
185 more recent ESCA measurements, e.g. those of Van Ooij and Muizebelt [ 2451, indicate that the differences in composition between surface and bulk cannot be all that large.
(iii) Other moly bdates and molybdenum oxide-based catalysts. Iron molybdates are not very selective catalysts for the oxidation of butene. Butadiene and maleic anhydride are the main partial oxidation products, but combustion predominates. The best results (i.e. selectivities up t o 30%) are obtained at low oxygen partial pressures or in pulse experiments which allow partial reduction of the catalyst. According t o Trifiro et al. [ 3171 P-FeMo04 is the preferred catalytic phase and the formation of Fe,(Mo04)3 must be avoided. Sazonova et al. [279] compare Fe-Mo0 with Bi-Mo-0 and report that these catalysts are simiIar with respect to the oxygen bond energy (80 kcal mol-' for both at 4OO0C), high oxygen mobility and the occurrence of a redox mechanism. The low selectivity of Fe-Mo-O is ascribed t o the strong interaction between butene and the catalyst, probably causing combustion, while reversible adsorption is found in case of Bi-Mo-O. Dente et al. [ l o o ] report a kinetic study, which includes the formation of formaldehyde. Butene, as we11 as partial oxidation products, were fed to a flow reactor at 320-380°C. The results were interpreted by means of power rate equations, and parameters were calculated for each reaction step. With regard to the butadiene formation, reaction orders of 0.5 with respect to both butene and oxygen are reported, and an activation energy of 18 kcal mol-' is found. Some other molybdates, i.e., MnMo04, cr-CoMoO, and CdMoO, have also been studied by Trifiro e t al. [ 317,3201. These molybdates are compared with /3-FeMo04 with respect t o their catalytic properties. Using a stirred tank reactor at 310--420"C, maleic anhydride selectivities of 2040% can be obtained at low oxygen partial pressures. High butadiene selectivities (>70%) were found with a pulse reactor in the absence of air. The adsorption properties of MnMo04 towards oxygen, l-butene and furan were studied by Veleva and Trifiro [337] using a gas chromatographic method. The authors conclude that apart from lattice oxygen, three irreversibly adsorbed oxygen species are present at the surface. They suggest that lattice oxygen is active in the formation of butadiene from butene, while the different adsorbed species are assigned roles in the oxidation of furan t o maleic anhydride and in combustion reactions. Molybdenum oxide activated by various supporting materials (MgO, S O 2 , A1203,TiOz) was investigated by Akimoto and Echigoya (13-15, 17) for the production of maleic anhydride from butadiene. The most interesting combination studied was MOO, -TiOz (1 : 3) modified with 1.5 at. % Sb. A selectivity of 45% is reached with this catalyst at 380°C, using a flow reactor and a feed consisting of 1.5 mol. 5% butadiene in air. Furan occurs as a co-product with a selectivity of about 10%at short conReferelices p p . 2 5 3 - 2 6 2
186 tact time. The authors conclude that maleic anhydride is in part formed via furan as an intermediate. The studies further concern the interaction between butadiene and the various catalysts, related to the reaction mechanism. ( b ) Antimonates
Several antimony-based oxide combinations are very good catalysts for the dehydrogenation of butene t o butadiene. Attention has particularly been paid to the three binary oxide combinations that are well known for their propene ammoxidation qualities: Sn-Sb-0, Fe-Sb-0 and U-Sb-0. There are some essential differences between antimonates and bismuth molybdates, although it is generally assumed that a redox mechanism is operative for both. First, the amount of lattice oxygen that is involved in the oxidation process is very small. There is no rapid exchange between bulk and surface oxygen, or, in other words, only a small amount of oxygen is mobile. Secondly, the reoxidation rate is much smaller for antimonates than for molybdates. It appears that the rate of reaction between butene and the catalyst is of the same order of magnitude as the rate of oxidation. Consequently, the kinetics depend on both butene and oxygen concentrations. Finally, inhibition by butadiene is not reported by investigators of these catalysts. ( i ) Tin antimonates. Active catalysts are found over a wide range of com-
positions. There are no specific binary crystal phases in this system. However, antimony slightly dissolves in S n 0 2 , with a maximum of 5-7 at. % Sb for the saturated solid solution [142]. Thus t i n a n t i m o n y oxide catalysts, in general, consist of a mixture of S b 2 0 4 or S b 2 0 j and the saturated solid solution. This solid solution displays catalytic properties very superior t o that of pure SnOz and is probably the catalytically important phase. This has been shown for the oxidation of butene t o butadiene by Sala and Trifiro [273] using a flow reactor at 370-440°C. Adding S b 2 0 4 t o S n 0 2 resulted in an increase of the selectivity from 40 t o about 93%, with a maximum at 5% Sb (in contrast t o the oxidation of propene, where more Sb is required t o obtain the maximum selectivity, although a similar activity maximum was found at 5% Sb). The authors also observed that the oxidation activity was coupled with an isomerizing capacity. However, addition of water destroyed this capacity without affecting the oxidation power. The kinetics are unanimously reported t o depend on both the butene and oxygen partial pressures. Reaction orders close to 0.5 for both reactants are found by several authors, and for various catalysts (Sn/Sb = 1/4, 1/3, 2/1) at about 450°C [36,278,329]. Sazonova et al. [278] proved that the reaction proceeds on a partially reduced surface through a redox mechanism. They used a Sn/Sb = 2 / 1 catalyst at about 450°C with flow
187 TABLE 23 LangmuirHinshelwood parameters for butene oxidation
Isomer -
_
_
1-Butene trans -2-Butene cis-2-Butene
_
Temp. ("C) _
k (mol g-' sec-')
~
~
KB
KO
(1 mol-')
(1 mol-')
~ _____ ~ 60.2 2.42 x 10-4 7 . 7 0 x 10-5 56.0 9.35 x 10-5 55.6
438 450 450
i t
i
2.0 2.0 2.0
2.0 7.4 3.0
i
*
0.3 1.0 1.0
-
and pulse reactors. The oxidation and reduction rates (individually determined in the pulse reactor) are equal mol O2 m-' sec-') and in good agreement with the reaction rate in the flow reactor. Several authors have studied the kinetics in detail and have proposed kinetic models. According t o Trimm and Gabbay [329], the kinetics of the reaction over a Sn/Sb = 1/4 catalyst is best described by a LangniuirHinshelwood type of model, which reflects that hydrogen abstraction from butene, involving dissociatively adsorbed oxygen, is rate-controlling; viz.
R
= _
hKB[C4H81 (KO[02 _ ~
[I + KB[C4H81
1
+ (KO[021)"212
The parameters for this expression are given in Table 23. Equal activation energies of about 1 7 kcal mol-I are found for the three butenes. The authors further report that, besides combustion products, furan is the major by-product (in yields of 2-776 depending on the conditions). Minor products (<0.5%) are acrolein, n-butyraldehyde and acetaldehyde. A rather complex network of isomerization, butadiene formation and a number of side reactions was analyzed and, based on simple power rate equations, over 100 kinetic parameters (rate coefficients, activation energies and reaction orders) were estimated. Trifiro and Pasquon [324] conclude that their kinetic results for the same catalyst can be equally well described by Langmuir, redox and power rate models. They used three types of reactor: differential flow, recirculating flow and pulse reactors. The data from the two flow reactors were individually treated and led t o six kinetic equations (Table 24). Pulse experiments by the same authors [ 3241 demonstrated a great influence of the oxygen partial pressure in the pulse on the conversion, indicating that only a very small amount of lattice oxygen is available. Further, it was shown by thermogravimetry with different Sn/Sb ratios (9, 5.66, 4 and 0.25) that the reducibility rises with increasing amounts of Sb,05 . This runs parallel t o an increase in selectivity (up to a certain value): the most easily reducible catalysts are the most selective. The reverse is true for mixed oxides based on M o o 3 . The authors feel that the difference between molybdates and tin antimonates is large enough t o suggest that yet another mechanism applies. Instead of a redox mechanism they R c f e r e n c c s p p . 2 5 3 - 262
TABLE 24 Kinetic models for tin antimonate catalysis of butene oxidation Mechanism
Type of reactor
Langmuir
Differential flow
Error (%)
Rate equation
12
9.33
Differential re-circulation Redox
Power
Differential flow
R=
Differential re-circulation
R=
DiI'l'erential flow
R
Differential re-circulation
pC4H8pod:
4.6 x
=
l o v 3 exp(400/T) Pgi + 0.3 exp(4400/T) P C , H ~ pC4Hp,pod:
2
X
12
8.6
exp(4000/733) + 0.3 e x p ( 4 4 0 0 / 7 3 3 )Pc,
1 6 P o d , 5 ~ , Pexp(--5000/T) ~~
10 7.4
189 suggest a role for “activated adsorbed oxygen”. An inhibiting effect of steam is reported by Bakshi et al. [36] using a Sn/Sb = 1/3 catalyst in a flow reactor at 350-470°C. A very complex rate equation is proposed but no constants are specified. Secondary reaction kinetics are reported in a subsequent paper by the same authors [ 361. Finally, Ven’yaminov et al. [ 3381 compare a Sn/Sb = 4 / 1 catalyst with bismuth molybdate (213) at 450°C and note that the Sn-Sb catalyst is less hampered by a drop in selectivity at high degrees of conversion. On the other hand, lower yields are found in the oxidation of 2-butene (compared with I-butene). The superiority of the bismuth molybdate catalyst in this respect is probably connected with its greater isomerizing capacity.
(ii) Iron antimonates. Iron antimonates are very effective catalysts for the dehydrogenation of butenes, and have properties very similar to those of tin antimonates, although the phases are different. A binary component FeSb04 occurs as a separate phase, which, depending on the Fe/Sb ratio, is present in addition to S b 2 0 4 or F e 2 0 3 . A Fe/Sb = 1 / 6 catalyst was investigated by Bakshi et al. [36] in a differential flow reactor, and compared with a Sn/Sb catalyst. With both, approximately half-order kinetics are found with respect t o both the olefin and oxygen. At product partial p r e s s u r e s ~ ~and , ~p~H Z obelow 0.02 atm the rate of butadiene formation is described by R
=
k p ~ ~ ~ 8 p ~ : s
The rate coefficient, k , is equal t o 2.7 X mol m-’ h-’ atm-’ at 370°C,with an activation energy of 10-11 kcal mol-’ . Shchukin e t al. [ 288-2901 confirmed that iron antimonates consist of a mixture of FeSb04 with either Sb204 or a - F e , 0 3 , and report that a maximum selectivity of 90% occurs between Fe/Sb = 0.06 and 1.7 while the activity hardly changes in this region. As with propene oxidation, a high selectivity thus requires an excess of Sb over Fe. Ammosov and Sazonov [ 21,22,24] demonstrated that for iron antimonates the initial selectivities are lower than for bismuth molybdates due to a higher rate of the parallel combustion reaction. It is proved that both selective oxidation and combustion occur by a redox mechanism. In another publication [23], the same authors report the kinetics of the butene and butadiene combustion reactions. A very low selectivity (10%) is found by Shchukin et al. [288] for a fully oxidized catalyst in a pulse reactor at 425°C. However, in the absence of oxygen, the selectivity increases t o 90% upon reduction of the catalyst t o an extent that corresponds t o 40% of a monolayer.
(iii) Other antimonates. Uranium antimonates have received some attention because this type of catalyst was of industrial importance for some time for the ammoxidation of propene. The U-Sb-0 system and its References pp. 253-262
190 active phases are described in Sect. 2.2.6(b)(iii). The catalyst is also a very active and selective butene oxidation catalyst. Simone et al. [ 2941 report that isomerization reaction do not occur with U - S b - 0 catalysts. They further note that even a slight reduction irreversibly destroys the activity, and conclude that, as with other antimonates, only surface oxygen anions are available for the oxidation reaction. Other combinations with antimony oxide were tested by Sala and Trifiro [272] by adding Co, Ni, Mn, Zn, c e , Cd and U oxides. These binary oxides are compared with tin antimonate. The selectivities d o not differ much, although the activities do. The explanation given by the authors is that the active site is always S b 5 + = 0 ,which is reduced. The second metal is assumed t o adsorb oxygen and effect a rapid reoxidation of Sb4+.
( c ) I r o n oxide-based catalysts and iron-containing spinels Iron-based oxide mixtures, e.g. ferrites, and supported iron oxides can be very effective catalysts for the dehydrogenation of butene t o butadiene, as also appears from the patent literature [ 1601. Pure iron oxide in the form of a-Fe,O3 was investigated by Sazonova et al. [280]. Pulse experiments revealed that high selectivities can be obtained under conditions that cause a partial reduction of F e 3 0 4 . In that case, a redox mechanism is operative. The required conditions are a rather low temperature (325°C) and a low oxygen concentration. Binary and ternary oxide compounds with a spinel structure have been studied by several authors. Butadiene can be produced with very high selectivities over these compounds. Rennard et al. [ 263,2641 have studied a variety of oxidation reactions over this type of catalyst, particularly MgCrFeO, and ZnCrFeO, , and conclude that only the dehydrogenation t o a conjugated product is selectively catalyzed. In the absence of such a possibility, only combustion takes place. The latter is also the conclusion of Zanderighi et al. 13581 from a study of the propene oxidation over spinels. A magnesium ferrite catalyst, MgFe,04, was studied by Gibson and Hightower [138,139]. A good selectivity is reported (89% at 20% conversion), but the catalyst suffers rapid deactivation and loss of selectivity, which is attributed to a too extensive reduction of Fe3+ to Fe2+.The inverse spinels CuFe204 and CoFe,04 were studied by Cares and Hightower [ 76,771. The kinetics are rather complex and the process appears t o involve both lattice and adsorbed oxygen. Particularly in the case of CoFe204 the selectivity is strongly influenced by the oxygen t o hydrocarbon ratio, while this factor is not important for CuFe20, . Both catalysts are inactive for the isomerization of butenes. The substitution of Cr for a part of the iron in the lattice of the binary spinels MgFe204 and ZnFe204 markedly increases their efficiency as dehydrogenation catalysts and strongly improves their stability. Selectivi-
191 ties over 90% can be achieved at conversion levels of about 60% for both MgCrFeO, and ZnCrFeO, at 360°C. This is confirmed by Sterrett and McIlvried [ 3061. A correlation between the catalytic qualities and the reducibility of this type of catalyst is suggested by Massoth and Scarpiello [ 2051. They performed reduction experiments both with hydrogen and with butene. Reduction may destroy the lattice, and the best catalysts appear to be those that are only superficially reduced. The effect of the introduction of Cr in the ferrites, mentioned above, is shown to be essentially due to an increase of the stability against bulk reduction. With respect t o the mechanism, the authors cited here all agree with the scheme that was initially proposed by Rennard and Kehl [ 2641, viz. C4H, + C4HB
t
Evidence for the Fe2+/Fe3+redox cycle was provided later by ESR measurements [ 2051, while recent experiments with deuterium-labelled butene indicate that C-H cleavage is involved in the rate-controlling step [ 1381. In agreement with the views of Schuit [ 2811, chemisorption of the olefin on an anion vacancy is assumed, but 0 - is postulated as the active oxygen species rather than 0’-. An argument in favour of 0- is that otherwise much more, and rather complicated, elementary reaction steps are required to account for the transfer of charge. ( d ) Vanadium pentoxide-based catalysts
Catalysts based on V z 0 5 are particularly effective for the production of maleic acid anhydride. Pure V 2 0 5 ,as well as several modified vanadium pentoxide catalysts, have been investigated. Pure V z 0 5 was investigated by Ai [9]using a flow reactor at 350°C. A maximum butadiene yield of 46% is reported, while furan and maleic acid anhydride can be produced (from butadiene and furan) with maximum selectivities of 7 2 and 6076, respectively. The depth of oxidation can be controlled by the oxygen pressure and the contact time. Isomerization reactions d o not occur. Crotonaldehyde is formed as a by-product. The References P P . 253-262
192 TABLE 25 Oxidation rate coefficients (sec-') at 350°C Reactant
14
k'(V205) h"( v2 0.j-P, 0 5 ) hII/kI
0.4 1/35
18 0.6 1/35
18 0.6 1/30
product spectrum qualitatively confirms the reaction scheme on p . 175. Vanadium pentoxide catalysts modified with P 2 0 5 or SnO, were studied also by Ai and Suzuki [5,9,10]. PzOs strongly reduces the activity as shown in Table 25 for a ratio of P/V = 1.6. At the same time, the formation of maleic anhydride is suppressed, and maximum selectivities for the
-
0
0.2
0.4
0.6
0.8
0
V Atomic ratio ___ V + Sn
Fig. 6 . Selectivity of butene oxidation t o diene and anhydride as a function o f catalyst composition (V-Sn-0).
193 TABLE 26 Rate of oxidation over V 2 0 5 - M o 0 3 (20%M o o 3 ) at 350°C and a reactant/air ratio of 0.6-1.0/100 Compound
Apparent first-order rate coefficient (sec-' )
Butene Butadiene Furan
20 35 35
~
~~
formation of butadiene (from butene) and furan (from butadiene) increase t o 88 and 9396, respectively. The effect of the Sn/V ratio for Sn0,-modified catalysts is demonstrated in Fig. 6 with respect to the selectivity of butadiene formation and to butadiene conversion into maleic anhydride. The highest activity is obtained below a Sn/V ratio of 0.1. The variation of the selectivity with the Sn/V ratio is interpreted by the author as an effect of the degree of catalyst acidity. Increasing amounts of V 2 0 5 make the catalyst more acidic, which, according t o the author, parallels a tendency to deeper oxidation. Adsorption measurements on P205-modified V 2 0 5 were carried out by Rozhkova et al. [ 2681 under precatalysis conditions (150°C). Catalysts modified with Moo3 were studied by Delgrange and Blanchard [ 991 and by Ai [ 71. As for the oxidation of benzene, the best activity and selectivity with respect t o partial oxidation products is obtained at 20-30% Moo3. Delgrange and Blanchard note that, in the absence of oxygen, such a composition also exhibits a maximum isomerization activity (at 405" C). Ai presented some quantitative information with respect to oxidation rates and selectivities at 350°C. The data are collected in the Tables 26 and 27. A P 2 0 5-promoted V205-Mo03 catalyst was investiTABLE 27 Maximum selectivities depending on the M o o 3 content of the vzo5-MO03 catalysts at 350'C
Moo3 < 20%
M o o 3 > 50%
(incl. pure V2O5)
Ouerall Selectivities w i t h respect to maleic anhydride C4Hs + C4H203 C4H6 C 4 H ~ 0 3
20 43
24 53
Selectivities of individual reaction steps C4Hs + C4H6 C4H6 C4H40 C4H40 C4H203
46 72 60
46 72 73
+
+
References PP. 2 5 3 - 2 6 2
194 gated by Slavinskaya et al. [ 2961. Roughly the same results are reported, i.e. the effect of P z 0 5 is not very significant (see also ref. 292). Finally, it may be noted that vanadium-based catalysts can oxidize butenes t o acetic acid. Kaneko e t al. [ 168,1691 report a selectivity of 50% at a relatively high conversion of 67%. V z 0 5 is not active but combinations with Sn, W, Ti are, especially at temperatures in the range 200300°C. In accordance with the hypothesis of Ai about the effect of the catalyst acidity, a deep oxidation is favoured by the highest acidity. Hence Sn-V-0 catalysts with 50% V, which display the maximum acidity per unit of weight, are the most suitable.
2.3.3Dimerization and aromatization o f iso- and n-butenes The dimerization of butenes is very similar to that of propene. Bismuth oxide and several other oxides have a capacity t o catalyze this reaction. The combination Bi-Sn-0 (Bi/Sn = l / l ) , investigated by Seiyama et al. [2841,appears to be particularly active. The dimerization and aromatization of olefins occurs in consecutive reaction steps. Isobutene, for example, reacts as follows
CH, Accordingly, the maximum production of 2,5-dimethyl-l,5-hexadiene from isobutene over a Bi-Sn-0 catalyst is reached at low conversions (10%) and short contact times (0.1-4.2 sec) as appears from the work of Seiyama et al. [284]. Isobutene leads to higher dimer yields than n-butenes, which preferentially form butadiene. The reaction orders for various butenes were determined by the same authors using a flow reactor at 500°C and applying a constant oxygen pressure (Table 28). The kinetic data fit a rate expression according to a Langmuir-Hinshelwood mechanism, based on a bimolecular surface reaction as the rate-controlling step. TABLE 28 Reaction order in olefin for oxidation of butene isomers Reactant
Deh ydro-dimers
COZ
1.5 2.0 1.9 1.9
1.1 1.1
Butadiene
~
Isobutene Butene-1 cis-Butene-2 trans-Butene-2
. _ _ _ _ ~ _ . _ _ _
1.1 0.8
1.2 1.1 0.8
195 Remarkably, however, the logarithm of the rate constant varies linearly with the dissociation energy of the allplic C-H bond, which indicates that the rupture of the C-H bond is included in the ratedetermining reaction step. Mixed olefin feeds (propene and butene) were also used. It appears that codimerization can occur yielding C, -dimers. Bismuth phosphate has been investigated as a catalyst for aromatization of the four different butene isomers at 550°C. An optimal catalyst has an atomic ratio Bi/P = 2 (Sakamoto e t al. [271]).Isobutene is converted a t short contact times (7= 0.3 sec) t o dimers and t o aromatics, with a selectivity of 29% each. n-Butenes give much lower yields. Parera and Trimm [248] converted isobutene on In2.03 mainly to dimethylhexadiene (S = 36%). The yield of aromatics is very low.
2.3.4 Ox y hy dration Analogous t o the production of acetone from propene, Sn0,-MOO, catalysts may convert 1-butene t o methyl ethyl ketone. Tan et al. [313] find 85.3% selectivity at 4.8% conversion (135°C).At a higher temperature, mainly the dimer is formed. 2-Butene gives a somewhat lower selectivity. Co,04-Mo0, also produces this ketone at 240°C but with more acetic acid as a co-product. The isobutene structure does not allow oxyhydration. The hydration step, however, does occur with a high selectivity over Sn02-Mo0, at 105°C and yields tert-butanol. The occurrence of this reaction supports the assumption of the alcohol as an intermediate in the oxyhydration reaction of unbranched olefins (see propene oxyhydration, p. 162). 2.4 HIGHER OLEFINS
Since the review of Voge and Adams [ 3431,not much kinetic research has been carried out on the selective oxidation of olefins with more than four carbon atoms. This is unexpected because isopentenes can be selectively oxidized t o isoprene, which is an important material in the production of thermoplastic rubbers. Isopentenes are available from hydrocarbonconversion operations. Maissant et al. [194] used a Bi203: MOO, = 1 : 2 catalyst at 450°C t o convert 1-methyl-2-butene and 2-methyl-2-butene into isoprene. A selectivity of 37% was reached with an activity about 10 times as large as the dehydrogenation of butene with the same catalyst under comparable conditions. Trimethyl-2,2,3-butene-l gave 7 5% of tert-butylacrolein. Investigating a bismuth molybdate catalyst with Bi : Mo = 1 oxidizing pentenes to isoprene, Watanabe and Echigoya [ 3441 found that isomers of pentenes were less reactive than those of the corresponding butenes in flow experiments, but the reverse was true in pulse experiments. Heat of adsorption measurements make it clear that the active sites are not uniReferences p p . 253-262
196 form. This leads to different kinetic data from flow and pulse experiments. A Langmuir-Hinshelwood model was found t o apply. V2OSshifts the product spectrum t o maleic acid anhydride. Higher selectivities were obtained by Rennard et al. [263] by using MgCrFe04 and ZnCrFeO,, catalysts. At temperatures above 350" C 2-methyl-2-butene and 2,3-dimethyl-2-butene are converted a t a level of about 30%t o diolefins with a selectivity of 90%. Molecules with conjugated bonds are favoured. Pichler e t al. [255] report the conversion of monomethylbutenes to isoprene on a catalyst containing 9.8 wt. % V 2 0 5 , 19.4 wt. % Moo3 and 0.8% P2OS on a T i 0 2 carrier. At a temperature of 335"C, the selectivity is maximal (30%)at a residence time of 0.2 sec. With Sn-Mo oxides, Tan e t al. [343] oxidized 1-pentene at 158" and 185°C t o give an oxyhydration. The main product was methyl propyl ketone with a selectivity of 6 5 and 53%, respectively, and related conversion levels of 8 and 22%. The olefin isomerized rapidly. 2.5 AROMATIC HYDROCARBONS
The oxidation of aromatic hydrocarbons originating from coal is one of the first organic gas phase oxidation processes carried out on an industrial scale. The development of these processes was initiated by the discovery that the V 2 0 5 catalyst used for the oxidation of sulphur dioxide was also applicable to the partial oxidation of benzene to maleic anhydride and naphthalene t o phthalic anhydride. Remarkably, V2O 5-based catalysts are still used in these processes today as they appear superior to any other type of catalyst. There are a number of analogies between the oxidation of aromatic hydrocarbons and olefins. Two classes of aromatic oxidations are t o be distinguished. (a) The mild side-chain oxidation of alkyl aromatics and in particular methyl aromatics, which yields aldehydes and acids similar t o the oxidation of propene and is carried out over similar catalysts. (b) The more severe oxidation of aromatic hydrocarbons over V 2 0 sbased catalysts, which yields acid anhydrides as the stable final products. These anhydrides may be formed either by oxidative decomposition of the aromatic nucleus, o r by combined oxidation of two ortho-positioned methyl groups. An example of the latter is the oxidation of o-xylene t o phthalic anhydride which is similar t o the oxidation of butene-2 t o maleic anhydride over the same type of catalyst. Ammoxidation can be successfully applied t o methyl aromatics (e-g. toluene and xylene) as it can t o propene. However, the subject has not received much attention in the literature, mainly due t o the fact that there are no important applications for aromatic nitriles at present. Catalysts based on V 2 0 5 resemble the binary oxide catalysts such as
197
Bi-Moand U-Sb-0 in many aspects, The participation of oxygen from the V 2 0 5 lattice is proved by the selective oxidation capacity of the catalyst in the absence of gas phase oxygen, while ESR investigations have shown the reduction of V 5 +t o V4+ and demonstrated a correlation between the activity and the V4+ content of a working catalyst. These facts strongly support the assumption of a redox mechanism. In contrast to the aforementioned binary oxides, VzOs has a stronger oxidation power and is able t o attack hydrogen attached t o the aromatic nucleus. Sometimes attention is drawn t o the importance of a layer structure in the catalyst or t o geometric factors (e.g. Sachtler [270]). Unexpectedly, however, very effective vanadium-based catalysts exist which operate in the molten state, indicating that a fixed structure is not important. The catalytic activity of molten oxide phases seems to occur exclusively in the oxidation of aromatic hydrocarbons over V 2 0 5-based catalysts, such systems have not been reported for the selective oxidation of olefins. With respect t o the kinetics of aromatic oxidations, (extended) redox models are suitable, and often provide an adequate fit of the data. Not all authors agree on this point, and Langrnuir-Hinshelwood models are proposed as well, particularly t o describe inhibition effects. It may be noted once more that extended redox models also account for certain inhibition effects, for mixtures of components that are oxidized with different velocities. The steady state degree of reduction (surface coverage with oxygen) is mainly determined by the component that reacts the fastest. This component therefore inhibits the reaction of a slower one, which, on its own, would be in contact with surface richer in oxygen (see also the introduction t o Sect. 2).
2.5.1 Benzene The gas phase oxidation of benzene with air is an important process for the production of maleic anhydride, which is the major partial oxidation product, besides products of complete oxidation. By-products are benzoquinone, in particular at low conversion, and fumaric acid which is formed at the high conversion levels used in industrial installations [ 281. Only traces of phenol and other by-products are formed. The important catalysts are based on V 2 0 5 and a maximum yield of 6 0 4 0 % maleic anhydride is obtained at 350-500°C. The oxidation of benzene to maleic anhydride is generally described by the simplified reaction scheme (Scheme 1, p. 198). Complete oxidation products (CO, C O P )are mainly formed from benzene and not by combustion of maleic anhydride itself. Therefore, the parallel character of the reaction scheme predominates, which implies that a high initial selectivity enables high yields t o be obtained. References P P . 253--262
198
0
Scheme 1.
The above scheme can be expanded to the more complicated scheme below t o include the formation of benzoquinone, viz.
0
0 Scheme 2.
As indicated, in this scheme maleic anhydride is formed both directly and from benzoquinone. The latter route, however, contributes only 2--10% to the maleic anhydride formation. To describe the oxidation kinetics, a reduction-oxidation model is generally accepted. Mars and van Krevelen [204] were the first t o apply this model to the benzene oxidation. The overall oxidation rate is expressed by
The suitability of this model, at least for a non-too-wide conversion range, has been confirmed by several authors. However, a correct description of the maleic anhydride production obviously demands splitting of k B into individual rate coefficients for maleic anhydride formation and for combustion, while in fact a separate equation should be added for maleic anhydride combustion. Such multi-step redox models have not been reported in the literature. Sets of first-order rate equations, however, are widely used,
199 assuming zero order with respect t o oxygen at sufficient low ratiop,,,,,,,/ Poxygen.
The fact that the kinetic parameters are slightly dependent on conversion is probably due t o inhibition by the product (maleic anhydride), as shown by the circulation flow experiments of Lyubarskii e t al. [192], and represented by the negative order with respect to maleic anhydride in their power-rate equations. Remarkably, a very low temperature dependence of selectivity is generally reported, indicating that the activation energies of the various reaction steps hardly differ (the order of magnitude is 20 kcal mol-I). This might be explained by the same or similar rate-determining steps. Scarcely anything is known about the reaction mechanism. For example, it is not known whether H atom abstraction and formation of a phenyl complex, or oxygen introduction into a p-diradical complex (associative adsorption of benzene), or both initiate the oxidation reaction. (a) Catalysts
The most interesting catalysts for the oxidation of benzene are VzOsbased, and particularly supported V 2 0 5-MOO,. Up t o the present, the V,O,-based catalysts have proved t o be superior t o any other metal oxide or metal oxide mixture. Catalysts based on V,Oj have been extensively studied in the past [ 1011, including the effect of promotors. Some more recent contributions will be mentioned here. A V205-K2SOj catalyst was used in a kinetic study by Jaswal et al. [165] with a differential flow reactor at 350-400°C and varying benzene/ oxygen ratio. The overall benzene oxidation rate was adequately described by a simple redox model. Kinetic parameters are given in Table 29. The initial selectivity is not reported by the authors, but a value of 5 0 4 0 % can be derived from the stoichiometric number n, assuming COz and maleic anhydride as the main products. Unsupported V2O5-Moo3 catalysts of variable composition were studied by Germain and Peuch [127] using a differential reactor at 370440°C. A maximum in activity and selectivity is found at 25 mol. 5% TABLE 29 Kinetic parameters for benzene oxidation
k 400°C ( I ( g cat.)-' sec-')
______
___ Benzene reaction Catalyst reoxidation __._
~
Re f c wncvs p p . 2 5 3 --2 6 2
-
-~
-~
5 x 105 28 x 105 -
EA (kcal mol-')
n
22.5 25
5-6 5-6
(mol oxygen/ mol benzene)
~
- _______________
200
Moo3 (V/Mo = 6), correlating with the maximum solubility of MOO, in V205which is about 30%.With a triangular reaction scheme (Scheme 1,p. 198) and first-order rate equations (a large excess of air was used) the ratio of rate coefficients h , : h2 : h 3 = 1: 0.43 : 0.14 is found which corresponds to an initial selectivity of 70% and a maximum yield of 55% with respect t o maleic anhydride. The above ratio is almost independent of temperature and an activation energy of 18.5 kcal mol-’ is reported. The kinetics are not influenced by C 0 2 and H 2 0 which were added t o the feed in some experiments. The effect of some additives t o the catalysts was also investigated. Alkali metal oxides decreased the maleic acid selectivity while phosphoric acid had the opposite effect. These results are very likely connected to the desorbability of the acidic reaction product. Unsupported V 2 0 5 - M o 0 3 catalysts were also studied by Bielanski and Ingolt [49,50], using an integral flow reactor at 350°C. Several hours of pre-treatment with the reaction mixture were needed t o obtain a stable and selective catalyst (the catalyst was prepared from a melt instead of by the usual co-precipitation method). In this period, partial reduction and formation of V4+ (ESR analysis) take place. The activity increases with the actual V4+ concentration in the working catalyst (Fig. 7 ) until a 100%
I
1
-
0
I
2
1
4
I
6
I
8
Concentration ~ ~ + t r n og i- e ’ )
Fig. 7 . Rate of benzene oxidation as a function of V4+ ( 3 5 O O C ) .
I
10
201 reduction t o V 4 + is achieved. Further reduction and formation of V3+ again, decreases the activity. The selectivity with respect t o maleic anhydride (measured at 50% benzene conversion) initially increases with the V4+ concentration, but a mole V4+(g cat.-’). Catamaximum of about 60% is reached at 8 X lysts with 20-30 mol. 5% MOO, give the best results with respect to both activity and selectivity. Vaidyanathan and Doraiswamy [332] studied the kinetics for a silicasupported V205-Mo03 catalyst. A flow reactor was used at 300--400°C and excess air was present (air/benzene 2 50). First-order rate coefficients were determined according t o Scheme 1. Almost independent of temperature and composition (V/Mo = 1and V/Mo = 2) the ratio h , : k 2 : h , = 1 : 0.3 : 2 was found indicating an initial selectivity 77% with respect to maleic anhydride. A remarkable shift of the activation energy was found from 20 kcal mol-’ below 350°C to about 2 kcal mol-’ between 350 and 400°C. This low value cannot be explained by diffusion limitations. Low activation energies (2-6 kcal mol-’) are also reported by Ahmad et al. [3,4] for silica-supported V205-M003 and V 2 0 5 - C r 2 0 3 catalysts, studied with a fluid bed reactor a t 350-550°C. Rather low initial selectivities (<40%) are reported in this work. The possible inhibition of the oxidation reaction by the maleic acid produced is not considered in the above two studies, but may have been important. The activity and selectivity of 19 oxides at 400--450°C were investigated by Germain and Laugier [133]. The activities are compared with those for the oxidation of toluene in Fig. 8, and show a linear relationship for the major part of the oxides, the toluene oxidation being approximately twice as fast as the benzene oxidation. The only selective catalysts, i.e. those that produce substantial amounts of benzoquinone and maleic anhydride from benzene, and benzaldehyde and benzoic acid from toluene are the oxides of V, Mo and W. Remarkably, these oxides clearly deviate from the average correlation in Fig. 8 and show a much higher toluene/benzene activity ratio (about l O / l ) . The order of activity, maximum yield of maleic aldehyde and initial selectivity with respect t o benzoquinone is the same for these oxides: V > Mo > W. The authors explain the formation of the products by a “rake” type of mechanistic scheme. B(enzene)
Ph( enol)
Q( uinone)
A(nhydride)
COX
a2 I b 2
B-Ph’20
kl
Q-
A k2
.c k3
The scheme involves a number of unknown surface intermediates, while the appearance of product molecules in the gas phase depends on the relative rates of adsorption, desorption, formation and conversion. The calcuReferences p p . 253-262
202
aJ C N aJ C
aJ
n L
0 CL, c. m L
rs,
0 -
:
I 1
co
-1
-2
NbO
I
-
I
I
I
log r a t e of toluene oxidation
Fig. 8 . Correlation between oxidation rates (log scales) of benzene and toluene (400'C).
lated relative values for the oxides of V, Mo and W are collected in Table 30. The same authors [132] studied a number of MOO,-containing binary oxide mixtures. The results are summarized in Table 31 and show that V205-M00, and TiOz-Mo03 are the most selective, the first one being very superior in activity. system was studied in more detail by DCchelette and The Ti-Mo-0 Germain [98]. The promoting effect of TiOz is apparent for all catalysts, but the best results were obtained with Ti02-rich mixtures. Maximum initial selectivities of 63-7776 and maximum yields of 30-35% were obtained at 450°C using excess air (75/1). First-order kinetics with respect t o benzene and an activation energy of 18 kcal mol-' were observed under these conditions. No binary phases could be detected by X-ray analysis, and the fact that the best catalyst resulted from a mechanical mixture of rutile and MOO,
TABLE 30
Kinetic parameters f o r benzene oxidation Catalyst
400°C
wo3
MOO3
v2 0 5
450°C
400°C
450°C
V2 0 5 - M 0 0 3 V/Mo = 6
400°C
45C"C 420°C
b2h2 b3/k3 k2X2Ikoho a k3h3/kOXO a2lkoho a3/k0X0
0.12 1.25 171 14.4 20.5 18
0.19 2.50 79 7.75 15 19.4
0.008 1.74 2300 4.4 18.6 7.65
0.03 2.12 930 15.1 29 32
0 0.72
0.052 1.9
0 0.82
29
22.6
21
18.5
TABLE 3 1 Activity and selectivity of different catalysts for benzene oxidation a t 400°C Catalyst Ti-Mo Activity (mol h-' m-2) Initial selectivity (%) Quinone Maleic anhydride Maximum yield (%) Quinone Maleic anhvdride
8.5
V-Mo 73
1 47
8 55
0.04 30
0.3 30
Fe-Mo
U-Mo
1.5
4.7
0 30 0 0.12
3 36
0.1 6.2
Sn-Mo 25.6
Sb-Mo 2.6
1 32
12 51
0.04 23.5
0.4 20
Bi-Mo 0.13 0 12 0 0.13
Mo 0.38 0.8 42 0.04 8
LQ
0
204 led t o the hypothesis that a surface phase is formed of either bidimensional molybdenum oxide or a Mo-Ti-0 mixture formed by surface reaction. This idea is in agreement with results of Akimoto and Echigoya [16], who studied supported-Moo3 catalysts with low MOO, content in the butadiene oxidation t o maleic anhydride. 2.5.2 Toluene and xylene
The gas phase oxidation of toluene and xylene is not of practical importance except for the conversion of o-xylene into phthalic anhydride, which will be separately discussed in Sect. 2.5.3. The main products of partial oxidation are aromatic aldehydes and acids formed by oxidation of the reactive methyl group. However, the yields that can be obtained are rather poor, in contrast to the catalytic liquid phase oxidation, which is much more selective. The poor yields are due partly to further oxidation (combustion) of the primary products, and partly t o direct oxidation of the aromatic nucleus (also mainly combustion). Activity and selectivity patterns for the oxidation of toluene over a variety of single oxides and MOO, -containing binary oxides were investigated by Germain and Laugier (129-131). A flow reactor was used at 400-500°C with a large excess of air (75/1). Yields above 30% with respect t o the sum of partial oxidation products were not reached for any of the catalysts studied. The main partial oxidation product was benzaldehyde. Initial selectivities above 50% were obtained with V 2 0 s , Moo3, VM o - 0 , Fe-Mo-0 and U-Mo-0. Benzoic acid was only of importance for V20,-based catalysts. Maleic anhydride was the main product at high conversion levels. Maximum yields above 10% were obtained with V z 0 5 and the molybdates of V, Ti and Sb. Remarkably, SnO, almost exclusively TABLE 32 Oxidation of toluene Catalyst
v2 0 5
M003 V-Mo-0 Fe-Mo-0 Sb-Mo-0 Bi-Mo-0 U-Mo-0
Temp. ("C)
400 450 400 450 400 400 400 400 400
Initial selectivity (%) with respect to benzaldehyde
Maximum yield ( % ) Benzaldehyde
Benzoic acid
Maleic anhydride
48 51 42 64 54 63 31 40 58
9.1 13.5 11 18 4.1 13.1 5.8 9 14.3
8.8 9.5 2.9
8.8 13.0
8.2 1.0 3.4 0.2
0.8
7.3 9.6 12.5 0.5 15.6 0.3 1.8
205 produced benzene as the partial oxidation product; a maximum yield of 10% was reported. A summary of initial selectivities and maximum yields is given in Table 32 for the most interesting catalysts. Reddy and Doraiswamy [262] used 15% Mo0,-5% W 0 3 as a catalyst at 500°C and found a selectivity of 75%with respect t o benzaldehyde by keeping the degree of conversion at a low level (5%). A formal reaction scheme, given by Germain and Laugier to describe the formation of all observed partial oxidation products, is shown below. This scheme comprises three reaction paths corresponding to side chain oxidation, oxidative coupling and direct oxidation of the nucleus as initiating reactions. In reality, an even more complex reaction network is t o be expected.
CHO
COOH
:$...;1:;.;&&&+ 0
0
T
r
1
0
0
The kinetics of the toluene oxidation over Bi,MoO, and a commercial Bi-Mo-P-0 ammoxidation catalyst were investigated by Van der Wiele and Van den Berg [348]. A flow reactor was used at 45O-55O0C, 1-3 atm and varying feed rate, toluene and oxygen partial pressures. Benzaldehyde formation and combustion reactions are the main process; the parallel-consecutive scheme which applies is Benzaldehyde
Toluene References p p . 2 5 3 - 2 6 2
(2)
>CO, co,
206 The kinetics are adequately described by an extended redox model, in which the rate of each of the three reactions is described by
Ri
=
kiPA8
in which A is the reacting component in reaction i and 8 is the fractional coverage with oxygen, expressed by
8=
~ _ _ _ _ _ _ _ _IZreoxid _
2
kreoxid.P02 + ( k l k 2 ) Ptoluene ( n - h3Pbenzaldehyde where n is the number of oxygen molecules required t o convert toluene into carbon oxides. The ratio of rate coefficients is practically equal for = 1/1/6/ both catalysts, and is given approximately by k , /k2/h3/kreoxld 12, corresponding t o an initial selectivity of 50%and a maximum yield of about 10%. This ratio is almost independent of the temperature. An apparent overall activation energy of 16 kcal mol-' is found. The same commercial Bi-Mo-P-0 catalyst was used in a study by Van der Wiele [347], which included the oxidation of the xylenes at 400500°C. In contrast t o the oxidation of toluene, dealkylation cannot be neglected. Table 33 presents an example of the product distribution at a 71-74% conversion level for both the xylenes and toluene. Remarkably, substantial amounts of the dialdehyde are only formed from p-xylene, while an enhanced benzene production is found in the case of o-xylene. The reaction schemes shown on p. 208 are proposed; the combustion reactions, applicable t o each component in the scheme are left out for simplicity. The overall oxidation rates of o-xylene and p-xylene are identical over the entire temperature range studied. At 500"C, the rate is three times higher than that with toluene, while an apparent overall activation energy of 33 kcal mol-' is found. For rn-xylene, a close resemblance t o toluene is observed: an oxidation rate approximately twice as high, in agreement with the presence of two methyl groups, and about the same activation energy (20 kcal mol-'). The influence of substituents on the catalytic oxidation of toluene was investigated by Trimm and Irshad [ 3301. Toluene, chlorotoluenes and xylenes were oxidized over a Moo3 catalyst at 350-500°C. Partial oxidation products are aldehydes, acids and phthalic anhydride (in the case of o-xylene). Unexpectedly, both xylenes and chlorotoluenes are oxidized faster than toluene. The authors conclude that apparently the electromeric effect of the chlorosubstituent is more important than its inductive (-1) effect. The activation energies of the xylenes and chlorotoluenes all fall in the same range (17-18 kcal mol-'), while a much higher value is reported for toluene (27 kcal mol-'). The kinetics of the p-xylene oxidation over tin vanadate was studied by Mathur and Viswanath [206]. A differential reactor was used at 320380°C. p-Tolualdehyde, maleic anhydride and p-toluic acid are the main
N 01 cc,
I
N 0)
N
TABLE 33 Example of product distribution in xylene oxidation Reaction conditions
p-Xylene o-Xylene m-Xylene Toluene
Temp. ("C)
Press. (atm)
Flow (set)
463 463 485 550
1 1 1 1
0.73 0.75 0.75 3.0
Xylene, toluene in feed (mol.%)
Total conversion (70)
2.3 1.9 2.0 2.2
74 73 71 72
Aromatic products (% of feed) Benzene
Toluene
Benzaldehyde
Tolualdehyde
Terephthalaldehyde
Total
2.8 5.0 2.1 3.4
4.1 3.6 4.0
1.1 0.5 0.8 10.5
11.3 12.3 15.0
4.9
24.2 20.9 21.9 13.9
N
0
cx,
209 partial oxidation products, besides traces of terephthalic acid. The overall oxidation rate is adequately described by a simple redox model. Remarkably, the activation energy appeared very dependent on temperature, ranging from very low values at 320°C t o 35-50 kcal mol-' at 420°C. Demethylation with conservation of the aromatic nucleus is normally an unimportant reaction compared with aldehyde formation or combustion. However, under reducing conditions some catalysts specifically promote dealkylation. The conversion of toluene to benzene was studied by Steenhof de Jong et al. (302-304), by both pulse and flow experiments. In the absence of oxygen, selectivities of up t o 70%are obtained with bismuth uranate catalysts at 400-500°C. The best catalyst is Bi2U06.The catalyst is reduced in the oxidation process, as gas phase oxygen is absent. The reduction proceeds to metallic bismuth and UOz. The activity decreases during reduction, but is completely restored by reoxidation with air. Therefore, a regenerative mode of operation is proposed for practical application of this process. Dealkylation may also become the predominant reaction with bismuth molybdate. This was shown by van der Wiele [347] for the progressive reduction of the catalyst with toluene pulses. While initially the product spectrum is similar t o that obtained in presence of air, a shift to benzene formation occurs at increased reduction. ( a ) Mechanism
The mechanism of the toluene and xylene oxidation bears a close resemblance t o the oxidation of propene. Abstraction of a H-atom from the reactive methyl group and formation of a complex between the resulting radical and the catalyst is the first and probably the ratedetermining step for both. However, the effect of the mesomeric stabilization of this radical complex is different. While a symmetrical ally1 structure is formed from propene, an asymmetrical situation occurs for toluene and xylene, which is illustrated below for the case of toluene, viz.
Thus oxidation of the nucleus may be activated and may form a possible explanation of the relative low initial selectivity generally observed [ 345, 3471. As to side chain oxidation, introduction of oxygen and dissociation of a second hydrogen atom yields the aldehyde as the first desorbable aromatic product. On V,O,-based catalysts the aldehyde is easily converted into the acid. The transition of an adsorbed aldehyde into a symReferences p p . 253--262
210 metrical carboxylate complex has been clearly demonstrated by IR spectroscopy (Sachtler et al. ref. 1 2 from [270]). The strong bonding of a complex of this type is probably the principal reason that benzoic acid is hardly formed by most of the catalysts. Instead, decarboxylation and combustion take place. Demethylation as the mechanism of dealkylation is also suggested by Steenhof de Jong [304] for the benzene formation from toluene over bismuth uranate catalysts, mentioned above. Another mechanism for benzene formation and parallel combustion is proposed by Germain and Laugier [ 1291. They suggest that toluene is 7r-adsorbed on a surface cation via the nucleus, and then looses two benzylic H-atoms t o form an cu,a(o)-x-adsorbed carbenoid complex, viz.
t
V
Rotation of the nucleus (x-bond rupture) is necessary t o introduce surface oxygen into the methyl group. Therefore a weak x-bond favours aldehyde formation, while a strong bond causes C% bond rupture and benzene formation or oxidation of the nucleus. Not much is known concerning the mechanism of the oxidation of the nucleus. Complete oxidation is the main reaction while minor amounts of maleic anhydride are formed over some catalysts, in particular those based on V 2 0 5 . Blanchard and Vanhove [ 521 demonstrated with 14C labelling that, for o-xylene oxidation over V 2 0 5 ,anhydride is exclusively formed from nuclear carbon atoms. This may be generalized t o other methyl benzenes.
2.5.3 Ortho-xylene Phthalic anhydride is the most important product in the oxidation of s-xylene, which has become competitive with naphthalene as a feedstock for the industrial production of this component. This process is carried out at 350-400°C and the industrial catalysts consist of doped V 2 0 5 or V205-Ti02 mixtures, pure or supported. Maximum yields of 70-75 mol. % (95-105 wt. %) are reported. Carbon oxides are the main by-products, besides minor amounts of tolualdehyde and maleic anhydride. Tolualdehyde is the main product at low conversion and an essential intermediate in the phthalic anhydride formation, while maleic anhydride is mainly formed as a side-product directly from o-xlyene. A suitable simplified reaction scheme is presented below. A variety of more complicated networks is reported in the literature, covering a
211 larger spectrum of intermediates and by-products (e.g. toluic acid, phthalide, C,-compounds), but will not be discussed in detail here, as they do not essentially modify the kinetic description of the main reactions. 0 II
\
I
, ,/'
0
0 II
iI
c,
0
--- C'
II 0
An interesting contribution t o determining the detailed mechanism is given by Blanchard and Vanhove [ 521, who oxidized labelled o-xylene (methyl-I4C)and phthalic anhydride (-l4CO--) over VzOs at 420-470°C. Examples of the product distribution and I4C content are given in Table 34 for conversions of 30 and 96.5%. Remarkably, maleic anhydride is only formed from carbon atoms of the nucleus. This is also confirmed by the oxidation of labelled phthalic anhydride which yields inactive maleic anhydride with a selectivity of lo%, besides CO and COz. This low selectivity also implies that phthalic anhydride hardly contributes to the maleic acid formation. The kinetics of the o-xylene oxidation appear to be rather complex due to the fact that several reaction steps seem more or less inhibited by reaction products. Lyubarski et al. [ 1921 ptudied the oxidation of tolualdehyde, phthalic anhydride and maleic anhydride, separately and in mixtures, using a recirculation reactor with a high temperature VzOs catalyst at 400-460" C. o-Tolualdehyde (and o-xylene) were found to stabilize phthalic anhydride, even with a 1 O : l excess of the latter. Phthalic anhydride, in turn, appeared to inhibit the maleic anhydride oxidation. Consequently, the tolualdehyde oxidation seemingly follows a parallel scheme
fl?
Tolualdehyde
Phthalic anhydride Maleic anhydride
for which the ratio of reaction rates is approximately given by R 1 : R 2 : R3 = 1 : 0.07 0.3. An integral approach to the kinetic analysis, including statistical methReferences p p . 253-262
TABLE 34 14C activities and maximum yields for the oxidation of o-xylene -_ _ ~ _ _ _ _ . _ _ _ _
Conversion 30%; temp. 420°C; ~_ air - 100; W I F = 1 5 o-xylene Composition (%)
o-Xylene o-Tolualdehyde Maleic anhydride Citraconic aldehyde Dimethylmaleic anhydride Phthalic anhydride Carbon monoxide Carbon dioxide
70 12
0.2 Traces Traces 7 2.6 7.3
Activity (mCi mo1-l) 6.3 6.3 0.15 3 16.3 6.3 -0.75 20.72
Conversion 96.5%; temp. 470°C; air_ __ - 100, W I F = 48 o-xylene Composition (%)
3.5 1.7 1.5 0.3 0.1 43.8 13.2 35.8
Activity (mCi mo1-I) 3.8 3.8 =O
Not measurable 3.8
Not measurable
213 ods, is adopted by Boag et al. [55]. Silica-supported V 2 0 5 was used as the catalyst in a recirculation reactor at 310-330” C with 8-7576 conversion of o-xylene. A satisfactory description of the kinetics was obtained with an “extended steady state adsorption model”, which is identical t o a redox model. The rate of the reaction step in which compound i is converted into j is given by
Rij = hG[i]% where 8, the fractional coverage of the surface with oxygen, is given by
in which ha is the oxygen adsorption (or catalyst reoxidation) rate coefficient, nij is the oxygen consumption (stoichiometric coefficient) for reaction i + j and C represents a summation over all reaction steps. The reaction scheme used by the authors is slightly different from the simplified scheme presented above, and includes the formation of phthalide, ViZ.
The kinetic parameters are presented in Table 35. According t o these data, the maximum yield of phthalic anhydride (at 100% xylene conversion) is about 45%. Pant and Chanda [247] report a kinetic study with a sintered Vz05 catalyst carried out with a spinning basket reactor. The following selectivities are observed at 450-517” C, independent of conversion (2040%) and temperature: 80% phthalic anhydride, 10% maleic anhydride, 10% carbon oxides. The authors propose a parallel reaction scheme. However, these data, and the absence of tolualdehyde at this conversion level prove that the reactor did not function as a well-stirred reactor, but suffered considerable bypassing (mass transfer limitation with respect t o References PP. 253-26-3
214 TABLE 35 Estimates of the kinetic parameters at 320°C for o-xylene oxidation ha k1.2 k1,3 k1.4
k1.5 k2.3 k2,4 k2.5
a Including
0.874 x 0.405 X 0.209 X 0.832 X 0.134 X 0.146 X 0.151 X 0.972 X
10-4
a
+ 0.200 x 10-5 I0.363 x 0.226 X 0.105 x * 0.140 X z 0.350 X t 0.160 X * 0.180 X t
lo-'
?-
95% confidence limits.
the basket). Nevertheless, the high ratio of phthalic anhydride t o carbon oxides points t o an excellent selectivity for this pure V 2 0 5catalyst. Calderbank [69] confirms this. While the above studies all concern pure or supported V 2 0 5catalysts, considerable attention has been gwen t o other catalysts, in particular the binary mixture V2OS-TiO2 which is a favourite among industrial catalysts. The VfTi ratio of these catalysts has a marked influence on activity and selectivity. Vanhove and Blanchard [ 3341 studied a series of V20s-Ti02 catalysts, prepared by coprecipitation, using a differential flow reactor at 450" C. The effect of the composition on the activity and the selectivity of the o-xylene oxidation (Fig. 9) is similar t o those mentioned above for butene (Fig. 6), with a remarkable sharp selectivity minimum at 35%T i 0 2 . However, X-ray analysis does not reveal specific binary phases in this system. Slightly different results were obtained by Grabowski et al. [ 1451, mainly due t o the method of catalyst preparation. (Solid T i 0 2 , ammonium vanadate and some water were mixed, dried and calcined.) The authors report a selectivity maximum at low V z 0 5 content (2-lo%), observed with pulse experiments a t 350°C. An ESR study by Yabrov et al. [355] revealed that, a t least at low V 2 0 s content (0.05-5 wt. %), vanadium forms a solid solution of V4+and V3+ in T i 0 2 . The samples investigated were sealed in the reactor after steady state operation of the o-xylene oxidation at 350°C. The V4+solid solution, which is considered the active phase, is not formed by the catalyst pretreatment at high temperature, but requires the interaction of the reaction mixture as was shown by the analysis of fresh catalysts. Solid state reactions between V 2 0 5 and TiOz were also studied by Cole et al. ~391. The influence of alkaline promotors (Na, K , Rb, Cs) on the catalytic properties of V205-Ti02 ( 5 at. 5% V z 0 5 )has been investigated by Boreskov et al. [ 571. The promotor addition (1: 8 with respect t o V) increases the initial selectivity while it considerably decreases the activity. The effect increases with increasing atomic number of the alkali metal. The rate of oxidation of partial oxidation products (tolualdehyde, phthalic anhydride), however, is hardly affected. The authors conclude that dif-
215
-
mole % T i 0 ,
Fig. 9. Selectivity and activity as a function of composition of V205-Ti02 for o-xylene oxidation.
catalysts
ferent sites are responsible for the oxidation of o-xylene and partial oxidation products, although the relative increase of the phthalic anhydride combustion rate may also be explained by decreased desorbability due to decreased surface acidity of the catalyst. The net effect of the promotor addition on the phthalic acid yield is not reported, but would be expected to be negative. Kinetic and mechanistic investigations on the o-xylene oxidation over Vz05-TiOz catalysts were carried out by Vanhove and Blanchard [335, 3361 using a flow reactor at 450°C. Possible intermediates like o-methylbenzyl alcohol, o-xylene-a,&'-diol,toluic acid and phthalaldehyde were studied by comparing their oxidation product distribution with that of toluene. Moreover, a competitive oxidation of o-methylbenzyl alcohol and 14C-labelledo-xylene was carried out. The compounds investigated are all very rapidly oxidized, compared with o-xylene, and essentially yield the same products. It is concluded, therefore, that these compounds, or their adsorbed forms may very well be intermediates in the oxidation of o-xylene t o phthalic anhydride. The ratio in which the partial oxidation products are formed appears t o depend on the nature of the oxidized compounds, i.e. o-methylbenzyl alcohol yields relatively more phthalide, whereas o-xylenediol produces detectable amounts of phthalan. This References p p . 2 5 3 - 2 6 2
216
dependency implies that there exist parallel reaction paths. The following (simplified) scheme is proposed by the authors.
UCH3
(4)-
pJC%
C6
Regarding the kinetics, the oxidation of o-xylene and o-tolualdehyde were compared for catalysts with different V/Ti ratios (Table 36). The ratio between partial and complete oxidation (X for o-xylene and Y for o-tolualdehyde) are influenced similarly, indicating that a change in the catalyst structure influences all the reaction steps. The oxidation of o-tolualdehyde in mixtures with o-xylene revealed that o-tolualdehyde reduces the o-xylene oxidation rate by a factor of about 2. The authors conclude that a redox model is inadequate and that hydrocarbon adsorption cannot be ratedetermining. Adsorption of various products should be included, and equations of the Langmuir-Hinshelwood type are proposed. It should be noted that the observed inhibition is not necessarily caused by adsorption competition, but may also stem from different TABLE 36 Comparison between selectivities for o-xylene and o-tolualdehyde oxidation o n V205-Ti02 catalyst a t 450°C Catalyst V-Ti: mol.% Ti02 12.5 o-Xylene oxidation selectivity o-Tolualdehyde oxidation selectivity X = kllkS Y = (k2 + k3)/k6
x/y
0.74 0.90 2.85 9 0.32
35 0.62 0.85 1.63 5.65 0.29
96 0.83 0.94 4.9 15.7 0.31
217 reactivities of the reactant molecules (“reaction competition with respect to oxygen”). The performance of a number of single oxides of transition metals was studied by Skorbilina et al. [295] using a differential reactor. As usual, o-tolualdehyde, phthalic anhydride and carbon oxides are the main reaction products. The initial selectivity with respect t o partial oxidation products decreases in the order Co > Ti > V > Mo > Ni > Mn > Fe > CU from 71% to 33%. The relatively high initial selectivities demonstrated by the “deep oxidation” catalysts (e.g. Co, Ni, Mn) indicates that the primary activation is probably the same for all these catalysts, while the differences that actually determine the character of the catalyst are connected with the stability of intermediates and products.
2.5.4 Naphthalene The gas phase oxidation of naphthalene t o phthalic anhydride over V,O,-based catalysts is one of the oldest successful partial oxidation processes and is still of industrial importance today. Common commercial catalysts are modified silica-supported V-K-S-0 catalysts and catalysts similar to those used for benzene or o-xylene oxidation. Maximum phthalic anhydride yields of 80-85 mol. 5% (92-98 wt. 76) at 350--400°C are reported. By-products are naphthoquinone (2-5%), maleic anhydride (25%) and carbon oxides. Naphthalene oxidation is very similar t o benzene oxidation except for the much greater importance of naphthoquinone, compared with benzoquinone, as a reaction intermediate. Roughly equal amounts of phthalic anhydride and naphthoquinone are initially formed from naphthalene. A suitable simplified reaction scheme is
\
0 The kinetics of the naphthalene oxidation obviously depend on the properties of the catalyst used, but some general statements can be made for the majority of V,O,-based catalysts. Refcrcnces p p . 253-262
218 (i) The rates of reactions (1)+3) are of the same order of magnitude, while small fractions of both naphthalene and naphthoquinone are converted into CO, CO, and maleic anhydride. The relatively high stability of the anhydrides has been pointed out already in the case of the o-xylene oxidation, and implies that the phthalic anhydride decomposition is almost negligible until complete conversion of naphthalene is achieved. (ii) Regarding the form of the rate equations, the overall oxidation rate appears t o depend on the partial pressures of naphthalene, or oxygen or both, and to be best described by a redox model. Individual reaction steps have been amply investigated [ 1011, but disappointingly, no integral kinetic analysis, based on an extended redox model, has been reported. The initial selectivities, as well as the integral product distribution, are hardly dependent on temperature, which implies that the activation energy has approximately the same value (25-30 kcal mol-I) for all reaction steps involved. Extremely few new contributions have appeared in the literature. The participation of the lattice oxygen of a pure V 2 0 5catalyst was studied by pulse experiments (Andreikov e t al. [26]). Although the catalyst is capable of oxidizing naphthalene in the absence of gas phase oxygen, the latter was indispensable for achieving a good conversion and selectivity, and this was apparently related t o the strong adsorption of phthalic anhydride on the partially reduced catalyst. Butt and Kenney [68] have demonstrated the catalytic activity of a V205/K2S04melt. A naphthaleneair mixture was fed over the surface (15 cm') of a carefully stirred liquid consisting of 39% v205 and 61% K2S04 (m.p., 433°C) at 440-470°C. The same experiments were carried out at lower temperatures with the solid catalysts. The activity rises with increasing temperature up t o 380"C, then falls steeply (45°C below the melting point). At the same time, the selectivity falls t o values below 20% for both phthalic anhydride and naphthoquinone. Above the melting point, the activity increases again. A rough analysis of the kinetics indicate the validity of a redox model for both temperature regions, although the kinetic parameters differ. The ability of the melt to participate in oxidation-reduction processes was demonstrated. The melt appeared to release oxygen when the atmosphere of air was replaced by nitrogen, corresponding to the conversion of V204,,2 to V204.85.
2.5.5 Anthracene Anthraquinone is the primary product of the oxidation of anthracene over V,O,-based catalysts. The reaction is very selective and high yields of anthraquinone are possible due t o its relatively high stability. An iron vanadate catalyst is used in the industrial process and yields of 80-90 mol. % are reported at 320-370" C. Phthalic anhydride, maleic anhydride and carbon oxides are the by-products.
219 TABLE 37 Activation energies and pre-exponential factors for t h e first-order rate coefficients for anthracene oxidation
Activation energy (kcal mol-I) Pre-exponential factor (mole g-' min-' mm Hg-')
Reduction step
Reoxidation step
18.01
19.20
2.465
X
lo3
4.0119
The kinetics of this reaction have not been extensively studied. Redox kinetics are suggested by Mars and van Krevelen for V2OSand the same kinetics are recently reported by Subramanian and Murthy (307-309) for V205-K2S04 and CoMoO, catalysts, both supported on silica. The oxidation was carried o u t in a flow reactor a t 270--360"C, with negligible formation of by-products. Activation energies and pre-exponential factors for the cobalt molybdate catalyst are collected in Table 37, while the results for the V2OS-KZSO4 catalyst demonstrate a remarkable change in activation energy at 330°C. Above this temperature, the activation energies are more than twice the original values. Power rate equations are proposed by Andreikov and Rosyanova [25] for V20,-K2S04/Si02 at 330-370°C. These d o not seem very appropriate, as the coefficients depend on the temperature and the oxygen partial pressure. The negative order ( - 0 . 2 4 ) with respect to anthraquinone suggests a rather strong inhibiting effect of this product on its formation.
2.5.6 Other aromatic hydrocarbons Some aromatic hydrocarbons have not been mentioned in the previous sections and will be briefly discussed here. Aromatic hydrocarbons which d o not have side chains in general form p-quinones and acid anhydrides. Benzene, naphthalene and anthracene have been dealt with above. In the case of phenanthrene, no p-quinone is formed as the adjacent C-H groups of the central nucleus are the most reactive. Phthalic anhydride is the main partial oxidation product, in addition to minor products such as 9,lO-phenanthraquinone. Andreikov and TABLE 3 8 Relative oxidation rates of some aromatic hydrocarbons -~
_ -
~
Benzene Toluene Naphthalene Anthracene Phenanthrene
References p p . 253--262
1 24 1700 300-400 17,000
~~~
220 Rusyanova [ 251 describe the oxidation of phenanthrene t o various products (using the V,05 catalyst above) by power rate equations according to a parallel reaction scheme I
0
CH-C
It 11
0 \ 0 /
CH-C 0
1
Q
co
Product inhibition is reported for reactions (1)and (2) in this scheme. Of interest are the relative overall oxidation rates for some aromatic hydrocarbons reported by the authors (Table 38). Aromatic hydrocarbons which have methyl side chains mainly behave like toluene and form aldehydes, while combustion is stimulated and selective oxidation of the nucleus is repressed. The oxidation of methylnaphthalene, for example, exhibits a low selectivity with respect t o phthalic anhydride formation, combustion and maleic acid formation being the dominating reactions. Durene is a special case because it resembles 0-xycatalyst at 420°C is reported lene. The oxidation of durene over a V-W-0 t o produce pyromellitic dianhydride, phthalic and maleic anhydride, although combustion dominates (Geiman et al. [ 1221 ). 1,2,4-Trimethylbenzene yields dimethylbenzene and trimellitic acid if oxidized on a SnV-0 catalyst. Kinetic data have been measured by Balsubramanian and Viswanath [ 371. Aromatic hydrocarbons with ethyl and longer side chains are easily attacked at the side chain, which is either completely oxidized or reduced t o one C atom and converted into the aldehyde. In the case of ethylben-
221 zene, quite a number of investigations have been carried out to develop catalysts that would direct the oxidation t o styrene, which is analogous to the oxidative dehydrogenation of butene t o butadiene. A selectivity of 80% is reported by Cortes and Seoane [91] for Ni-W-0 catalysts with an atomic W/Ni ratio between 2 and 4. A very high selectivity (>95%) was found by Joseph et al. [166], using a cobalt molybdate catalyst in a flow reactor at 500-600°C. However, a low oxygenlethylbenzene ratio (below 0.5) is necessary t o achieve this high selectivity and coke formation problems are to be expected. Industrial alumina-supported cobalt molybdate catalysts were studied by Russo e t al. [269]. Selectivities of 6C-70% were obtained at a conversion level of 20--30%. Aluminasupported MOO, appears t o have the same qualities, while unsupported Moo3 exclusively produces benzaldehyde (beside carbon oxides). Alumina itself also has some activity and may be used as the catalyst. Lisovskii et al. 11891 suggest the addition of alkali metal oxides t o alumiria in order to reduce the surface acidity, and thus t o prevent poisoning of the catalyst by condensation and cracking products. Although some progress has been made, the oxidative dehydrogenation is far from competitive with the highly selective non-oxidative dehydrogenation process used in industry today.
2.5.7 Ammoxidation Methyl side chains of aromatic hydrocarbons can be selectively ammoxidized to nitrile groups. The process is very similar t o the ammoxidation of propene and the same catalysts are found to be effective. Identical mechanisms have been proposed, and will not be discussed here. The selectivity of the ammoxidation of molecules like toluene and xylene is much higher than that of the oxidation of these compounds to aldehydes. The selectivity difference is more pronounced here than in case of propene. The initial selectivities of the propene oxidation and ammoxidation are practically the same, and the selectivity difference is mainly due to the high stability of acrylonitrile compared with acrolein. For aromatic (amm)oxidation, however, the initial selectivities also differ. Apparently, ammonia interacts with the catalyst in such a way that the activity for oxidation of the aromatic nucleus is reduced. A few contributions with respect t o the ammoxidation of aromatic hydrocarbons that have appeared in the literature concern toluene and xylene.
( a ) Toluene Simon and Germain [293] investigated a number of molybdates at 450°C with a molar feed of toluene/ammonia/air = 1 : 5 : 50. The main results are presented in Table 39, in which selectivities to benzonitrile and Refewncc.s P P . 253-262
222 TABLE 39 Selectivities and activities with molybdenum-based catalysts for toluene ammoxidation Selectivity
Conversion
(%)
("/.I
Activity (mmol h-' m-2)
85 83 82 61 68 87 85
80 87 95 87 85 93 94
0.09 0.33 0.62 0.14 0.78 0.69 1.77
~~
Bi-Mo Sb-Mo Sn-Mo U-MO Fe-Mo Ti-Mo V-Mo
~~
~
activities are given. The authors conclude that U-Mo and Fe-Mo catalysts, which are the most selective in toluene oxidation, are the least selective in ammoxidation. Because the overall rates of oxidation and ammoxidation are equal, the rate-determining step occurs before formation of the C6HsCH: complex. It can very well be that imine is an intermediate, viz.
The selectivity t o nitrile is higher than the comparable selectivity t o benzaldehyde. This is probably due to the greater stability of the nitrile or a difference in desorption velocity of the imine compared with benzaldehyde. Nitriles are only weakly adsorbed. A combination of VzOs and SnO, (weight ratio 70 : 30) is a reasonable catalyst at 300-360°C in giving about 50% yield, as has been shown by Lodaya e t al. [ 1901.The yield was measured a t a 5-8% level of conversion and is hardly dependent on temperature in the given region. The optimal NHJtoluene ratio is 6.
TABLE 40 Selectivities and activities in the formation of nitriles from p-xylene _.
Catalyst
Temperature ("C)
Selectivities a (%) p-Toluylnitrile
Terephthalonitrile
_ _ _ _ _ _ _ _ ~ _ Sn-Mo Ti-Mo V-Mo V
460 415 430 430 -
a
__.-
64(52) 92156) 77(64) 78(66) -
Activity (mmol h-' m-2)
_ 30(78) 6(87 ) 19(88) 19(80) _ ~ _ _ _
The respective conversion levels are given in parentheses.
~ 1 0.5 0.6 12 -__
223 TABLE 4 1 Kinetic parameters for ammoxidation of p-xylene over V 2 0 5
kl
k2 12 3
k4
Rate coefficient EA (sec-' ) (kcal mol-') ___ ___.. ___________ 1.3 x 104 19.4 2.4 x 104 13.9 3.8 x 10-3 2.2 28.6 15.2
-_
( b ) Xylene
Simon and Germain [293] tested some Moo3-based catalysts and compared these with Vz05. Ammoxidation with a reactant ratio hydrocarbon/ NH3/air = 1 : 10 : 100 gives the results for p-xylene shown in Table 40. V 2 0 5 was also used by Novella e t al. as a catalyst in the ammoxidation of p-xylene [241]. These researchers carried out experiments at 390,400,410 and 420°C, with varying feed ratios (p-xylene/NH,/air = 1/3--5/60-80). They proposed the kinetic scheme
CN
\
CH,
CN
CN
The kinetic parameters are given in Table 41. m-Xylene can also be ammoxidized as was shown by Rizaev et al. [ 2651, who used a recirculation reactor with a V-Mo catalyst (6%V 2 0 5 , 2% Moo3 on Al,03). The kinetic scheme is
According t o the authors, the kinetics are zero order, provided that sufficient oxygen and ammonia are present. Referetices p p . 253-262
224 2.6 METHANOL
Selective oxidation of methanol is the industrial route t o formaldehyde. In practice, two types of process are used, differing with respect t o the catalyst and process conditions. Silver is a very active catalyst at 600700°C and requires a high methanol/oxygen ratio for a good selectivity, while iron molybdate catalysts are already active at 350°C and may be used with low methanol/oxygen ratios.
2.6.1 The silver process The silver process is the older one and is still used in many formaldehyde manufacturing plants today. Yields of about 90 mol. 5% are reported, and combustion to carbon dioxide and water is the main side reaction. As significant amounts of hydrogen are formed, it has long been assumed that formaldehyde is essentially formed by dehydrogenation of methanol, accelerated by the combustion of a large part of the liberated hydrogen. Recently, however, several authors explain the kinetics on the basis of direct interaction of methanol with oxygen. The reaction is carried out over a silver gauze or low surface supported catalyst at 600-7OO0C, indicating a very fast chemical reaction. This implies that determination of the intrinsic reaction rate in laboratory reactors is complicated by the interference of heat and mass transfer limitations. To avoid this problem, studies have been made at much lower temperatures, which in turn run the risk of being non-representative. A Langmuir-Hinshelwood type of model is suggested by Robb and Harriott [266] who studied the reaction at 420°C. They find that the intrinsic kinetics can be represented by
The equation reflects dual site reversible adsorption. Methanol and formaldehyde compete for sites, while oxygen is dissociatively adsorbed on different sites. At a not-too-low oxygen pressure (>0.01 atm) the coverage of the oxygen sites is complete and the equation reduces t o KCH30HPCH30H
R-k +
KCH30HPCH30H
+
Kprod.Pprod.
An Eley-Rideal model with dissociative adsorption of oxygen is proposed by Bhattacharyya et al. [48]. Because the oxygen adsorption is assumed t o be irreversible, the model is identical with a redox model and
225 The experiments were carried out a t a very low temperature (264-290°C) resulting in an unusually low conversion. The relevance of the calculated kinetic parameters is therefore doubtful. Recently, Hodges and Roselaar [157] used gold and platinum as catamixture was stoichiometrically lysts. At 400" C, a rich methanol-xygen converted to formaldehyde and water with a residence time of 270 psec. Larger partial pressures of oxygen and higher temperatures raised the degree of combustion t o carbon dioxide. With platinum, the maximum yield of formaldehyde was reached a t 210°C. The authors assumed that methanol was dissociatively chemisorbed and reacted with adsorbed oxygen atoms. 2.6.2 Iron molybdate and other metal oxide catalysts The use of iron molybdate in industrial plants started about 1960. Yields of about 90% are reported for this process, applying either excess air or excess methanol and recirculation of the latter. Carbon dioxide is the chief by-product. Kinetic investigations have appeared in the literature since 1965. A redox mechanism is generally accepted [254], and has been confirmed by pulse experiments which demonstrated the equal activity of the catalyst in the presence and absence of oxygen. The results of Pernicone [254] and Liberti et al. [187] seem t o indicate that the rate-determining step is either hydrogen abstraction from methanol or desorption of formaldehyde. The structure of the iron molybdate catalyst in relation to its oxidation properties has been studied by several authors. It is stated by Pernicone [254] that there is an excess of Mo6+ and 0'- ions in the Fe, ( h I 0 0 ~lattice )~ giving rise t o an enlargement of the unit cell in one direction. Two iron ions can be replaced by two molybdenum ions as the insertion of three 02-ions in the lattice is possible. The activity of such a structure is higher than with Moo3 and Fe2(Mo04)3in pure forms, although MOO, is very selective. Carbucicchio and Trifiro [75] have shown, however, that the specific activity is the same, when the different surface areas of the pure and the iron-deficient molybdate are accounted for. The selectivity to formaldehyde is also practically the same. Another property of the irondefective molybdate is the presence of Mo= 0 double bonds on the surface. The hydrogen-abstracting capacity of the catalyst is closely related t o Mo6+ contained in the Mo=O as is shown in Sect. 3. There the role of iron is also discussed. It is, however, interesting t o note here that pure iron oxides accelerate combustion and that a W03-Fe2 (WO,), catalyst is practically inactive [254]. Replacement of iron by chromium is possible but leads t o a lower activity [ 2531. Baussart et al. [ 461 prepared stoichiometric NiMo04 which showed selective behaviour towards formaldehyde in a pulsed column below 375°C. References p p . 253-262
226 TABLE 42 Vanadium-chromium catalysts in methanol oxidation Cr (at.%)
Surface area (m2 g-l)
ko
(x
E (kcal mol-')
Order in O 2
5 8 9 14 14 19 20 15 6 28
1.4 0.81 0.59 0.14 0.18 0.001 0.0019 0.0001 0.0007 0.14
44-46 44 45 44 44 39 38 33 35 34
0.5
0.65 0.4
There is considerable evidence that surface acidity influences the catalytic activity of iron molybdate [254]. It was found by studying the adsorption of ammonia using infrared spectroscopy that, under reaction conditions, the acidity is due t o Lewis sites. The conclusion is that surface acidity is a necessary, but not a sufficient, property. Another group of binary oxides has been tested by Koval and Boreskov [ 1801. These authors studied 10 different compositions of VzOs--Cr2O3, starting with pure VzOs and adding increasing amounts of C r 2 0 3 .The rate data are given in Table 42 for a temperature of 300°C. The methanol concentration in the feed was 3.6-3.7 vol. 7%. Activities and selectivities are shown in Fig. 10. l80exchange rate measurements in the range 400-
-
Composition (mole % )
Fig. 1 0 . Selectivity and conversion in methanol oxidation on V z 0 5 - C r 2 0 3catalysts as a function o f composition at 30OoC.
227
D
m
1
I
I
20
A t o m i c ratio
I
I
40 (O/d-
I
I
60
v +Me
I 80
I
I 100
(Me=Fe.Co,Nl)
Fig. 11. Relation between selectivity (conversion) and catalyst composition for methanol oxidation at 31OoC.
500°C indicate that the oxygen in V-containing compositions, but not in C r 2 0 3 , is all exchangable. The selectivity decreases with decreasing strength of the oxygen bond on the surface, while the activity increases. Malinski et al. [199] combined vanadium pentoxide with oxides of nickel, iron and cobalt and reported that these mixed oxides have a much higher selectivity than the pure oxides. The results obtained a t 310°C are shown in Fig. 11, the methanol concentration in the gas phase being 44% and the C H 3 0 H / 0 , mole ratio 2.2. Selectivity and activity are given as a function of the atomic ratio V/Me. The ratio V/Me 2 1 gives the highest conversion to formaldehyde. A V-Ni catalyst gives the best results and does not show any activity for the side reaction which produces some hydrogen with other catalysts. The authors suggest that the latter group of oxides contain active oxygen centres which are not regenerated at a sufficient rate. Aldehyde molecules then get an opportunity t o decompose on the catalyst surface with simultaneous hydrogen evolution. 2.7 AMMONIA
The oxidation of ammonia can produce nitric oxide, nitrous oxide and nitrogen according to the stoichiometries References p p . 2 5 3 - 2 6 2
228
2 NH3 + 2; 0, = 2 N O + 3 H,O 2 NH3 + 2
0 2
= NZO
2 NH3 + If 0, = N,
+ 3 H,O + 3 HzO
The production of N O is of great industrial importance for the manufacture of nitric acid. The other two reactions do not have practical applications.
2.7.1 The production of NO The industrial process is carried out with platinum gauze as the catalyst at 750-900°C. Selectivities of 95-97% are reported for this extremely fast chemical reaction. The main by-product is N,, and only traces of N,O are formed. The kinetics were reviewed by Dixon and Longfield [ l o l l , since when the subject has not received much attention.
2.7.2 The formation of N, and N z O The conversion of ammonia t o N2 and N z O is catalyzed by metal oxides. Depending on the type of catalyst, N, or NzO may be the main product. The situation is analogous to the oxidation of hydrocarbons in that mild oxidation catalysts (e.g. MOO,, V 2 0 5 ) favour formation of nitrogen, while the more severe oxidation catalysts (e.g. Co304, MOO,) produce the largest amounts of NzO. Ilchenko et al. [ 161-1631 compared the oxides of Mn, Co, Cu, Fe and V, and found that MnO, gives a selectivity t o N z O of 42% at 155°C and p N H 3 = 0.1 atm at contact times, T , of 1.5-4 sec. Co304produces less N z O and more nitrogen at 143°C (selectivity = 18% at p N H 3 = 0.2 atm, T = 5-15 sec). At these low temperatures, the selectivity to N 2 0 was not very sensitive to variations in T , suggesting that the products are formed by parallel reactions, viz.
2 NH3 + oxygen
,N, + water - N 2 0 + water
If the temperatures are raised, catalytic N z O decomposition is observed, viz . N z O = N2 + 0, In principle, nitrogen can also be formed by the catalyzed reaction with ammonia.
3 NzO + 2 NH3 = 4 N, + 3 H,O The general rate equation for the oxidation of ammonia t o nitrogen is of the redox type (161-163).
TABLE 43 Rate coefficients (molecule cm-’ sec-’ atm-’) and activation energies (kcal mol-’) of ammonia to nitrogen Catalyst
Temp. ( ’ C )
kl x
k, x
El
EZ
MnOz CO304 CUO
145 143 240 250 290
0.66 0.26 5.13 0.38 6.50
1.21 0.61 16.13 0.55 2.57
30
17
20 16 23
20 21 20
Fe203
VZOS ~
in which n is a stoichiometric coefficient. Table 4 3 presents values of h l and k 2 and activation energies. The catalysts show a steadily increasing selectivity with increasing surface coverage of oxygen. It is clear that the formation of nitrogen is a “milder” oxidation than the one leading t o nitrous oxide. The major role of oxygen coverage has been confirmed by experiments in the absence of oxygen in which rate data have been determined for the reduction of the metal oxides with ammonia. Selectivities for the formation of N2 increase in the sequence MnOz < Co304< F e 2 0 3< CuO < MOO, < VzOs. The same pattern has been found in the mild oxidation of hydrocarbons and methanol. Ilchenko et al. [ 161-1631 relate the difference in selectivity t o the metal-oxygen bond strength; this is considered in Sect. 3. Holbrook and Wise [159] worked with crystalline Cu,O as a catalyst in the oxidation of ammonia at about 300°C. In this case, there is also a strong correlation between the amount of excess oxygen and selectivity. When the catalyst surface changes its defect structure from copper-rich t o oxygen-rich, the nitrogen concentration goes through a maximum. The rate of disappearance of NH3 is independent of the ammonia concentration and is first order in oxygen, comparable with the kinetics of acrolein production on Cu,O. The catalytic properties of Cu,O are controlled by the electronic properties (see Sect. 3). Another copper catalyst, prepared by treating a NaY zeolite with copper nitrate, for ammonia oxidation (160--185°C) has been studied by Williamson et al. [349]. The reaction is first order in NH3 and zero order in oxygen. The mechanism here is based on a Cu(II)(NH,):’ complex formed in the large cavities of the zeolite. The ratedetermining step is the reduction of Cu(I1) by ammonia. Wise [ 3501 investigated the parallel between ammoxidation and oxidation of ammonia over bismuth molybdates. It was shown that the rate of conversion t o nitrogen is first order in NH3 and independent of oxygen concentration, analogous to the selective oxidation of propene. Under conditions in which propene combusts, NH3 is converted t o nitrogen oxides. References p p . 253-262
230 Bismuth molybdate and other binary compositions (Fe-Mo, Sn-Sb and others) were tested by Germain and Perez [128) using a pulsed reactor. The authors demonstrate that a qualitative analogy may exist between ammonia and propene oxidation but if activities are compared, different sequences of catalytic efficiency arise. It must be noted, however, that these conclusions are based only on pulse experiments. These can be quite different from results in flow reactors, depending mainly on the nature of the steady state. From the different contributions, it may be concluded that, in the oxidation of ammonia, the same type of redox mechanism is operative for metal oxides as in the selective oxidation of hydrocarbons. As a consequence, the hydrogen atoms will be abstracted successively from the NH3 molecule by a stepwise mechanism. 2.8 SULPHUR DIOXIDE
Although there is only one oxidation reaction possible with sulphur dioxide and hence a selectivity problem does not exist, recent results from kinetic research are included in this chapter, since there is a close analogy with other oxidations, especially on V205-based catalysts. The oxidation of sulphur dioxide to trioxide is one of the oldest heterogeneous catalytic processes. The classic catalyst based on VzOs has therefore been the subject of numerous investigations which are amply reviewed by Weychert and Urbaneck [ 3461. These authors conclude that none of the 34 rate equations reported is applicable over a wide range of process conditions. Generally, these equations have the form of a power expression, in which the reverse reaction is taken into account within the limits imposed by chemical equilibrium, viz.
Also, Langmuir-Hinshelwood models have been proposed as well as models based on a redox mechanism. Recently, Happel et al. [154] using data from Kadlec et al. 1167,2171 conclude that a model based on the dissociative adsorption of oxygen, which is ratedetermining, fits the experimental results best, viz.
With h = A 1 exp(-E1/RT) and K = A z exp(E,/RT), the values of the parameters for the temperature region 380-480°C are
-
~
Parameter
Value
A1 El
7.34 X 1014 mol h-' atm-' (g cal)-' 4 . 7 1 X l o 4 cal mol-' 1.22 X mol h-' atm-' (g cal)-' 2.72 X l o 4 cal mol-' 8.20
_ -.
A2 E2
n
~~
__-
The industrial catalyst consists of a mixture of V z 0 5and KZSzO7supported on silica. Under technical reaction conditions (>440°C),this mixture forms a viscous molten phase on the surface of the porous silica structure. Apparently a redox model can also be applied to such a system [154]. Putanov et al. [ 2591 investigated K-V-S-0 catalysts carried by S O z . By different techniques, it was noted that compounds such as KV4010.4, K 2 V 5 0 1 3and K3V5014 are present. It was demonstrated that SiOz as a support plays an active role in transformations of the catalytic layer. In the binary system K2S04-VZ05, compounds with even higher K/V ratios were confirmed. Kato et al. [170] also drew attention t o the importance of the vanadium-potassium ratio. Working in the region 500-6OO0C, they found a simpler rate equation This is obviously valid for initial conditions only. The same comment applies t o the work of Herce e t al. [ 1551, who also d o not account for the effect of chemical equilibrium.
3. Role of the catalyst 3.1 THERMODYNAMIC CONSIDERATIONS
Thermodynamically, the oxidation of hydrocarbons t o carbon dioxide and water is preferred t o any partial oxidation reaction. The possibility of forming partial oxidation products is thus entirely dependent on the kinetics of the oxidation process. The oxidation of hydrocarbons, is in general, a stepwise process. One way to confine the depth of oxidation, therefore, is t o apply a low oxygen t o hydrocarbon ratio and a short reaction time. However, to avoid a multitude of products with different oxidation depths, the use of a catalyst is obviously required. In that case, the above two factors (oxygen deficient conditions and short reaction time) may loose their importance. Basically, the role of the catalyst can be twofold. (a) Activation of the hydrocarbon molecule by chemisorption in a specific way. The attack of oxygen may thus be selectively directed t o a particular site on the hydrocarbon molecule. References p p . 2 5 3 - 2 6 2
232 TABLE 44 Free energy for the transition of a higher to a lower oxide (kcal per mole of liberated oxygen), calculated from ref. 362 _ _ ~ ~ _ _ ~ ..___
Temperature (“C)
Ago Biz03 CUO C U 2 0
Fez03 Fe304 M003 SbzO, Sb2O4
SnOz Ti03
+Ag BiO
+
+cuzo +cu +Fe304 +FeO +Moo2 +Sbz04 +Sb203 +SnO +Ti203
uo3
+u,ox
u3ox
’U02 +V?04 +V203 +W02
v205
V204 W03
350
400
-3 6 I2 37 59 72 113 64 10 51 109 152 14 52 31 66 103
-4 4 65 30 54 62 105 51 3 44 102 145 9 46 30 61 98
~ _ _ _ _
(b) Reducing or “tempering” the activity of oxygen. The amount of energy liberated by the formation of C-0 or H-O bonds by reactions between hydrocarbons and molecular oxygen is roughly 100 kcal per mole of oxygen. This energy is so large that bonds within the hydrocarbon molecule can be broken and fragments result, which can be easily further oxidized. The catalyst can effect the distribution of the energy of oxidation over two partial reactions, i.e. the reaction between molecular oxygen and the catalyst and the reaction between the “loaded” or “oxidized” catalyst with the hydrocarbon. In the case of metal oxide catalysts, the degree of “tempering” can be derived from the free energy of the transition from a higher t o a lower oxide (Table 44). A value close t o zero means that the oxide has almost the same oxidation potential as molecular oxygen, while values of 100 kcal and more signify that the reactivity towards hydrocarbons is practically zero. The intermediate region, therefore, is of interest for catalysis by metal oxides. Quantitatively, the meaning of the figures in the table is very limited, however, because on a catalyst surface the situation is different from the bulk and the strength of the oxygen bonds is not uniform. It is also influenced by edges, corners and defects of lattices. With inorganic compounds, there can also be a selectivity problem, as illustrated by the oxidation of ammonia t o nitrogen. Deep oxidation leads t o nitrogen oxides. With sulphur dioxide, no selectivity problem rises.
233 In the following section, the metal-oxygen more detail. 3.2. METAL-OXYGEN
bonds will be treated in
BOND STRENGTH
Especially in those cases where 0’- is the active form of oxygen and the catalyst operates according to a redox mechanism, it is reasonable t o assume that the metal-oxygen bond plays an important role. It would be expected that the rate of oxidation should be inversely correlated with the bond strength, provided that the reduction of the catalyst by the hydrocarbon molecule is the rate-controlling step. Exceptions to such a correlation can easily occur, however, because of the heterogeneity of the surface. Indeed, it is found that the bond strength often depends on the degree of coverage. Another factor is the special geometry at the active site of the catalyst. Finally, it may be remarked that a concerted mechanism can occur in which the M e 4 bond strengths are only relevant in close connection with the complex to be oxidized. The most important properties used as a measure of the bonding strength are the heat of formation of the metal oxides, the heat of oxygen desorption, the reducibility of the metal oxide and the activation energy for isotope exchange between I8O2 in the gas phase and oxygen in the catalyst.
3.2.1 Heat of formation of metal oxides, AHf
AHf can be calculated, in principle, from thermochemical data. It is then necessary t o take into account the variable valency of most metals and t o fix the different oxidation states which occur during stationary or non-stationary reaction conditions. Some difficulties with this method are thy scarcity of data for mixed oxides, the difference in conditions between those on the surface of the catalyst and those in the bulk and the inaccuracy of a number of data obtained by measuring differences in AH. Attempts to correlate the activity with AHf have not been very successful. A fairly good inverse correlation was found by Moro-oka e t al. [223, 2241 but it concerns complete oxidation to carbon oxides. Some patterns of activity for various selective oxidation reactions, related t o LW,, are described by Germain [ 1341. With respect to the selectivity, the situation is even more complex. Only a rough classification into three groups can be made. The first one consists of metals which bind the oxygen loosely, e.g. noble metals, and generally promote complete oxidation. A second one has strongly bound oxygen but adsorbs oxygen loosely, which also favours combustion (e.g. Co, Mn, Ni, Cr). A third group is characterized by moderately bound oxygen, often coupled with variable oxidation states of the metal oxide. This group, in particular, effects a selective oxidation. Refere1ici.s P P . 2 5 3 - 2 6 2
234 3.2.2 Heat of oxygen desorption Measurements with a vacuum system of equilibrium oxygen partial pressures as a function of temperature indicate desorption energies. There is some difficulty in choosing a representative state of comparison. Generally, investigators evacuate a t increased temperature for a long time. Of interest is the new flash technique applied by Halpern and Germain [153]. This technique reveals that mobile oxygen generally occurs in discrete binding states. The authors compared V 2 0 5 and CuO with other catalysts mainly concerning total oxidation. Figueras et al. [ l l l ]emphasize the importance of the entropy of the oxygen bond, which can be considered as a measure of the surface mobility of oxygen. Unfortunately, their assumed positive correlation between entropy and selectivity is only based on two V 2 0 5 catalysts which differ with respect t o the carrier (Si02 and A1203). Portefaix e t al. [256] measured the equilibrium oxygen pressure as a function of temperature for iron molybdates. The authors demonstrated that, in the case of the system Fe-Mo-0, the bonding energy of oxygen increases with increasing degree of reduction if the composition is rich in iron. In the case of an iron-deficient combination, the bonding energy decreases with increasing degree of reduction. Only in the case of Fe,O, . 3Mo0, does the bonding energy remain constant. 3.2.3
isotope exchange
The activation energy for isotope exchange between "02in the gas phase and oxygen in the catalyst is a measure of the metal-oxygen bond strength. With selective catalysts, the exchange between gas phase and catalyst oxygen (hetero-exchange) is about as fast as the exchange between gas molecules via the catalyst (homo-exchange), implying that in both reactions the same oxygen species is involved, i.e. 02-.With nonselective catalysts, however, the homo-exchange rate may be considerably faster, and apparently involves a more loosely bonded, adsorbed form of oxygen. These principles are illustrated, for example, by Haber and Grzybowska [152], as shown in Table 45, in which a number of oxides are ordered according t o their homo-exchange activity. Indeed, the most selective catalyst is found at the t o p of the table, while a catalyst like Co304only effects combustion. Attempts t o correlate the exchange rate for selective catalysts with the activity for hydrocarbon oxidation have not been very successful, mainly due t o the fact that the oxidation activity of such catalysts is much greater than the exchange activity. The difference is often so large that the reactions must be studied in different temperature regions. The origin of this difference is obvious: liberation of oxygen from the catalyst is facilitated by the presence of a reducing agent (i.e. the hydrocarbon molecule),
235 TABLE 45 Activity of catalysts towards IRO2exchange Catalyst
Temperature
Rate (g O2 m-2 h-* 1
(“C)
Bi/Mo = 2 : 1 Bi/Mo = 1 : 1 Co/Mo = 1 : 1.7 MOO Fe/Mo = 1 : 1 Fe2°3 c o30 4
250-500 474-500 599-634 5 80-60 1 508-552 350-450 125-250 -__~____
None None 1.8 x 10-4 9 x 10-4 10-3 4 x 10-1 12.7
which may form an intermediate complex (transition state) involving the oxygen t o be transferred. The observed differences in activation energy between oxidation and oxygen exchange are considerable. Successful correlations may be found, however, within binary oxide systems, i.e. by comparing catalysts with different ratios of the same oxides. Blanchard e t al. [51,53], for example, studied the V2O5-MoO3 and V,Os -Ti02 systems and found a striking correspondence between the activation energy of isotopic exchange and the hydrocarbon oxidation selectivity, both as a function of the V/Mo and V/Ti ratios. Interesting reviews on the subject of isotopic oxygen exchange are those of Novakova [ 2401 and Parravano [ 2491. 3.3 OXYGEN TRANSFER
In heterogeneous catalytic oxidation, the reaction is always between a molecule t o be oxidized (in adsorbed form or not) and oxygen which is attached to, or is part of, the surface. A number of different oxygen species is possible, ranging from free oxygen molecules t o oxygen anions. The species in between can be represented by the scheme @2
0 (ads-)\
O,(gas) L_ O,(ads.) 0; (ads.)
@
2 0-(ads.) _7 2 0’- t--,2 0 2 surface
bulk
From left t o right, the oxidizing power will decrease. Which of the different oxygen species are active on a catalyst is determined by (a) the rates of the different steps in the scheme; these depend on temperature, oxygen pressure, state of the surface and type of catalyst; (b) the reactivity of the oxygen species with respect to the molecule to be oxidized. This depends on the oxygen bonding energy, the adsorbed References p p . 2 5 3 - 2 6 2
complex, the underlying geometry and the temperature. It is generally assumed that, in the group of transition metal oxides, the intermediates between 0, (gas) and 0’- (surface) do not come into the picture. 0’- (surface) is considered t o be the reactive oxygen species. In the case of a high oxygen anion mobility, surface 02-may rapidly exchange with bulk oxygen. Consequently, a large fraction or even all of the catalyst oxygen may appear t o participate in the oxidation reaction. Another consequence of a high mobility is that the sites at which oxygen reacts with the component t o be oxidized, and those at which oxygen is taken up, may be quite remote. Apart from the intrinsic properties of the catalyst lattice, the reaction conditions can also influence the oxygen mobility in the catalyst: as the transport of oxygen through the solid can be regarded as the diffusion of holes (anion vacancies), the number of these is an essential factor. This number obviously depends on the degree of reduction and thus on the rates of the reactions between the catalyst and oxygen, and between the catalyst and the compound t o be oxidized. These rates in turn depend on the reaction conditions. Several recent contributions concerning the participation of lattice oxygen in selective oxidation processes have appeared and fully agree with the above concepts. They will be discussed in more detail below. In a second group of metal oxides, which are not easily reduced, the oxygen is strongly bound and the catalyst is generally in a fully oxidized state. Thus 02-is not reactive, but an adsorbed form of oxygen, much more weakly bound, is active. This leads only to combustion. Quite a number of these metals are non-transition metal oxides. A third group contains those metal catalysts which d o not form specific crystal phases in an oxidized state. The common types of oxygen on the surface are then O2 (adsorbed) and 0 (adsorbed) which generally do not lead t o selective oxidation. One of the exceptions is silver, which very probably catalyses the selective oxidation of ethylene by providing 0; on the surface. However, an active role of surface oxides, which may be formed particularly by the action of promotors, is not excluded.
3.3.1 Participation of lattice oxygen The participation of lattice oxygen is inherent to the redox mechanism, which is operative in many of the oxidation processes that are catalyzed by metal oxides. Reviewing the processess described in Sect. 2, participation of lattice oxygen appears t o be the case for the majority of them, namely for the allylic (amm)oxidation of olefins, for the (amm)oxidation of aromatic hydrocarbons and for the oxidation of methanol, ammonia and sulphur dioxide. Two types of experiment are commonly used t o give evidence of participation of lattice oxygen: (a) experiments in the absence of gas phase oxygen and (b) experiments with labelled oxygen.
237 ( a ) Experiments in the absence of gas phase oxygen The activity of an oxide catalyst in the absence of gas phase oxygen provides direct evidence that lattice oxygen can perform the selective oxidation process, although it does not exclude the possibility that, in the presence of gas phase oxygen, other forms of oxygen also participate in some stage of the reaction. Pulse experiments are the most suitable for this purpose, because rapid catalyst reduction is then avoided. As pulse experiments have been amply reviewed in Sect. 2, only the conclusions will be discussed here. The activity of oxide catalysts in general declines as reduction proceeds. Characteristic of the processes that involve lattice oxygen is that the initial activity (i.e. that measured by the first pulse) approaches that in the presence of oxygen, while the selectivity is either identical in the presence or absence of oxygen, o r higher in the latter case, because side reactions due to adsorbed oxygen are excluded. The rate a t which the activity falls during reduction is dependent on both the nature of the catalyst and on the process studied. After a certain initial activity decrease, often a lower, but rather constant, activity level is reached. Different explanations are given for the fact that a part of the initial activity may be rapidly lost. Several authors suggest heterogeneity of the catalyst surface and conclude that more loosely bonded oxygen is consumed first. Another possible cause is the effect of the increasing of anion vacancies and reduced cations on the electronic properties of the solid, which in turn may affect the oxygen reactivity and the adsorption capacity for the reactant molecule. Finally, the irreversible adsorption of reaction products may be of importance. Barannik e t al. [ 38,391, for example, have shown that this is the predominating factor in the fall in activity during the pulse reduction of bismuth molybdates by propene. The occurrence of an almost constant, albeit rather low, activity level, which is reached after a number of pulses, signifies that a certain quasiequilibrium concentration of active sites is mzintained by transport of bulk oxygen anions t o the surface. Such a mobility of oxygen is particularly observed for bismuth molybdates and some related catalysts (see below). Typical examples of catalysts which completely loose their activity a t a low degree of reduction are the antimonates; this is primarily caused by the absence of anion mobility.
( b ) Experiments with labelled oxygen Most experiments concern the application of labelled gas phase oxygen in reaction mixtures, while only in a few studies has labelling of the solid phase been used. Catalysts that have received particular attention are the bismuth molybdates and the antimonates of U, Fe and Sn, all very selective catalysts for the oxidation of propene t o acrolein and similar allylic oxidations. References p p . 2 5 3 - 2 6 2
238
(i) Bismuth moly bdates. Bismuth molybdates have been extensively studied, mainly by using propene/”O, mixtures. The experiments have been performed in static systems [ 174,252,3521 or static recirculation systems [51,123] at rather low pressures, but also in a pulse reactor [276] and, very recently, in a flow reactor under atmospheric conditions [ 1751. It is clearly shown in all these studies that lattice oxygen is consumed in the selective oxidation, while the gas phase oxygen that reoxidises the catalyst diffuses into the solid. At temperatures of 400°C and higher, the mobility of oxygen anions appears t o be very large, and the oxygen introduced at the surface appears t o equilibrate with essentially all oxygen anions present in the lattice. The effect of temperature and oxygen partial pressure was studied by Sancier e t al. [276], They showed by pulse experiments that, at lower temperatures and higher oxygen t o hydrocarbon ratios, a certain amount of “short-circuiting” occurs between the catalyst reoxidation process and the transfer of oxygen from the cataiyst t o the reactant, as shown by a partial break-through of labelled oxygen in the reaction products. This short-circuiting is obviously caused by the decrease in ratio between the rate of diffusion into the lattice and the rate of reaction at the surface. Temperature primarily influences the diffusion rate, which has the highest activation energy, while the oxygen partial pressure may influence both: a higher pressure implies a higher oxidation state of the catalyst, i.e. it decreases the number of anion vacancies and thus the diffusion rate, while at the same time the reaction rate at the surface is increased. There is some uncertainty with respect t o the participation of lattice oxygen in the formation of carbon oxides parallel t o acrolein. Some authors report that an enhanced amount of l80is found in the COz produced, while others d o not observe any diEference between acrolein and carbon dioxide, with respect t o the l80/l6Oratio. The equal ratio in both products, however, may also be caused by the exchange of oxygen between C 0 2 and the catalyst. Gel’bshtein e t al. [ 1231 report that, in a static recirculation system, the amount of CI8O2formed is maximal in the beginning, and then decreases due t o exchange with the catalyst; Sancier et al. [276] find that, in pulse experiments, larger amounts of acrolein-”O are formed in the presence of C1802,and calculate an activation energy of only 4 kcal mol-‘ for the exchange reaction. On the other hand, carbon oxides are also formed in the absence of gas phase oxygen, while it is further known that, under the usual process conditions, the kinetics of acrolein formation and parallel combustion are the same, and both involve an allylic intermediate. One must conclude, therefore, that the initial reaction steps are very likely identical and involve lattice oxygen, but that, in the combustion of the ally1 intermediate, probably both lattice oxygen and adsorbed forms of oxygen can participate. Selective labelling of the catalyst is applied in an interesting study by Otsubo et al. [ 2461. Starting with labelled and unlabelled oxides of bis-
239 muth and molybdenum, y-Bi2’*03 . Moo3 and y-Biz03 - Mo1803 were prepared by solid state reaction between the oxides. Reduction by hydrogen was studied in a circulation reactor at 400°C. Initially significant differences occur between the H 2 ’ * 0 content of the produced water and the average “0 content of the catalyst, indicating that isotopic scrambling did not occur before the reduction took place. The results prove that the oxygen attached t o bismuth reacts with the hydrogen, while reoxidation proves that oxygen is introduced at the molybdenum sites. This with lS02 implies that oxygen transfer from molybdenum t o bismuth is a part of the redox cycle. The authors report that the same is indicated by experiments with propene. Details of this promising work have not been published a t the time of writing. ( i i ) Antimonates. The antimonates of tin, iron and uranium have been studied by using propene/l8O2 mixtures in static (circulation) systems [86,123,252]. As with bismuth molybdates, it has been shown that lattice oxygen is the only source of oxygen in the selective oxidation, while both lattice oxygen and adsorbed oxygen may be involved in the carbon dioxide formation. Compared with bismuth molybdate, however, a rapid break through of l 8 0 is observed, which proves that the exchange capacity of the antimonates is very small and is, in fact, restricted t o one or two surface layers, at least at the usual reaction temperatures (300-400” C). Apparently the mobility of anions in the antimonates is small, which also implies that reoxidation must take place practically on the reaction site, in contrast t o bismuth molybdates where reaction and reoxidation sites may be quite remote. This difference in mobility, therefore, may be one of the reasons why the kinetics of the selective propene oxidation differ for bismuth molybdates and antimonates.
(iii) Other catalysts. Vanadium pentoxide-based catalysts ( Vz05-Mo03 and V205-Ti02) have been studied by Blanchard and Louguet [ 511, using a butene/’*O, mixture in a static circulation apparatus. Labelled oxygen is immediately observed in the oxidation products, indicating that the mobility of oxygen is low. The authors d o not believe that adsorbed oxygen is involved, but assume short circuiting via a partially reduced catalyst surface that cannot receive oxygen anions from the bulk. 3.3.2 Role of Me=O type of oxygen Several workers correlate the catalytic activity of metal oxides with the presence and the nature of double bond type oxygen at the surface. This type of oxygen is coordinated with one cation and can be regarded as “terminal oxygen” in contrast to a-bonded oxygen (or “bridging oxygen”) that is coordinated with two cations (Me-O-Me). Oxygen anions in different coordination states can be detected by IR spectroscopy, while Refrrcriccs p p . 253-262
240 reflectance spectroscopy is particularly suitable for an investigation of the surface of a catalyst. One of the early studies in this field was that of Sachtler [270] concerning V 2 0 , in the oxidation of aromatic hydrocarbons. It was shown that the hydrocarbon interacts with V=O, which is abundantly present in the V 2 0 s structure. Much work concerning molybdates and the allylic oxidation of olefins was carried out by Trifiro et al. [ 219,318,3191. A strong correlation between the activity and the presence of oxygen double bonded t o molybdenum was observed particularly for bismuth molybdates. It was concluded, therefore, that Mo=O is the most reactive with respect t o the olefin molecule. Mitchell and Trifiro [219] specify the nature of the active sites more precisely as Mo(O,)~,i.e. a molybdenum ion with three terminal oxygen anions. The Mo(OJ3 configuration can be expected t o occur only at corners, edges and defects in the lattice. A remarkable parallel, therefore, exists with the observation of Sleight et al. [ 33,2971 that the activity of scheelite-type molybdates is strongly correlated with the presence of defects (cation vacancies) and the conclusion that olefins are adsorbed on M o - 0 polyhedra next t o these vacancies. With regard t o iron molybdates, the correlation is less clear. Trifiro and Pasquon [318]defend the view that Mo=O is also of importance for iron molybdates and state that pure F ~ , ( M o O ~is) ~inactive because of the absence of such oxygen species. However, Carbucicchio and Trifiro [75] have recently reported that no differences in selectivity and specific activity exist between iron-deficient molybdate and the pure compound, although Mo=O oxygen is only detected in the former. Double-bonded oxygen as the active oxygen species is also observed for tin antiinonate (Sb=O) by Sala and Trifiro [ 2721 and for Bi,WO, (W=O) by Villa et al. [340]. The fact that only Bi2W06is an active and selective propene (amm)oxidation catalyst, while other compositions ( B i 6 W 0 I 2 , Bi2W209,Bi2W30,,) primarily cause combustion, is ascribed t o the acidic W-0-W configuration, which is only absent in Bi,WO,. The W+-W sites are presumably responsible both for isomerization and combustion reactions. Several suggestions have been made with respect t o the particular acitivity of double-bonded oxygen. Trifiro [ 3191 assumes a parallel between gas phase oxidation and the oxidation of olefins in solution (at a low temperature), which is also catalysed by Me= 0-containing oxide compounds (e.g. Os04, R u 0 4 , SeO,), implying that similar complexes may be formed. The olefin oxidation mechanism proposed by Weiss et al. [ 3451, and presented in Sect. 2.2.2(b), is, in fact, based on this parallel. Trifiro even extends the parallel t o the positive effect of steam on the acrolein selectivity in the gas phase oxidation of propene, which might be analogous t o the solvolysis effect in solutions. Kazanskii [ 1711 suggests that the mechanism of selective oxidation possibly involves “electronic excitation” of a double bonded oxygen anion t o
241
an anion radical ( 0 - )induced , by the interaction of the olefin molecule with the metal cation, as presented in the scheme
CH 2 7CH-C H 3
Mt+
___L
= 0 2 -
CH2=CH-CH CH2-CH-CH2 + M(?l-l)+ -0M("-l)+ - O H -
The transition of Mo=O oxygen t o 0 - radicals is also assumed by Akimot0 and Echigoya [ 13,151, who investigated Moo3-based catalysts by ESR and IR spectroscopy. They distinguish Mo6'=0 from Mo5'=0 and state that the oxygen in the latter has the strongest radical character, in agreement with the observation that the maleic anhydride production from butene increases in parallel with the Mo" content of the catalyst. They further state that modifiers can influence the reactivity of Mo=O oxygen. The reactivity of this oxygen decreases as the electronegativity of the modifier increases, according t o the sequence Bi203> Sb203 > As203 > P 2 0 , . Correspondingly, the reducibility and activity for the propene oxidation is highest for the Mo03-Bi203 combination. Although this hypothesis has interesting aspects, an oversimplification of the nature of the binary compounds is obviously present. Finally, it must be noted that the assignment of catalytic activity t o Me=O type of oxygen does not imply that other types of oxygen are not involved. Assuming that the olefin molecule is indeed attacked by double bonded oxygen, o-bond oxygen may take part in other steps of the oxidation process. The alternation of double bonded and bridging oxygen in the reduction-oxidation mechanism is a possibility suggested by Akimoto and Echigoya [ 161 for Mo03-based catalysts, and represented by
reduction
-Mo6'=0
O=M06+-
by olefin
-MoS++-Mo5+-
reoxidation b y oxygen
In bismuth molybdates, the bridging Mo-0-Bi oxygen is the most weakly bound oxygen and therefore supposed t o participate in the transfer of oxygen [ 219,2811. 3.3.3 Significance of 0; and 0 - radicals From the large amount of work with labelled oxygen, it is clear that active oxygen in a selective reaction is a type of lattice oxygen. Neverthelrss, a number of publications, mainly of Russian origin, which investigate the presence of 0; and 0-radicals, are interesting. This is often done at rather low temperatures at which catalysis does not occur (pre-catalysis). Several authors state that such energetic radicals give rise t o combustion. Kazanskii 11711 observes 0; and 0 - radicals with ESR when oxygen is adsorbed on Mo03/Si02, V 2 0 5 and MgO. At -196"C, the presence of References p p . 2 5 3 - 2 6 2
only 0; is established, at room temperature 0 - and O;, at 160°C mainly 0 - and above 400°C the oxygen is present as 0 2 - The . oxygen is dissociated more easily if electrons are supplied. Transition metal oxides require small energies for electron transfer. With these n-type semi-conductors, the number of conducting electrons is so large that the reaction will be faster than the reaction of 0 - with the ccmpound to of 0- to 02be oxidized. In that case, a selective oxidation will take place instead of combustion. Yoshida et al. [356] show the presence of 0 - and 0 ; on V 2 0 5 supported by S O 2 . With strong reduction, it is mainly 0-.It was proved that reaction of 0; with propene gives rise t o aldehydes (propionaldehyde, acrolein and formaldehyde) at temperatures below 150°C. Yoshida et al. [357] confirm this and find that oxygen at room temperature is mainly adsorbed as molecular oxygen, only 10% is the sum of 0 - and 0;. 0 - is the oxygen species reactive towards carbon monoxide, 0; is not. Further spectroscopic research has been carried out by Krylov [183, 1841 on adsorption of oxygen on MOO,, WO,, V 2 0 5 and CuO, supported on A1203, MgO and BeO. In all cases, 0; radicals were formed. Extra stabilization occurs when the catalyst is reduced with hydrogen. On a number of active oxides, the 0; intensity is increased drastically by simultaneous adsorption of propene. It is suggested that 0 ; is attached t o the carrier cation. The electron transfer with simultaneous adsorption is then supposed t o be
0;
H+----C3Hs
ag-&-ho5+
Combustion sets in with a high coverage of the transition metal oxide, diminishing the number of Mg-O-Mo bonds. It is also increased by a high concentration of “mobile oxygen” (0; or split-off singlet oxygen). Burlamacchi e t al. [65,66] also used ESR techniques and found that CdMo04 differs from Bi2Mo0, in adsorbing molecular oxygen. The oxygen is activated through the formation of radicals on the surface and leads to deeper oxidation than bismuth molybdate. One has to be careful with the statement that combustion is exclusively caused by adsorbed radicals like 0 - (and 0;). Many oxides can oxidize hydrocarbons mainly t o carbon oxides in the absence of gas phase oxygen. Finally, it can be noted that Me=O type oxygen may be regarded as an intermediate between 02-anions and 0 - radicals, thus providing some relation between the respective theories. 3.4 ASPECTS OF CHARGE TRANSFER
The catalyst plays an important role in transporting electrons from the molecule t o be oxidized t o the reacting oxygen. It can be expected that
243 the capacity of the catalyst t o furnish electrons, to take these up and to transport them internally will influence the catalytic properties. Many authors accordingly have studied the electrical properties of catalysts; these are mainly semi-conductors. However, the correlation of the electrical properties of the bulk phase with the catalytic properties of the essentially heterogeneous catalyst surface is a classical difficulty. This may be one of the reasons why no general correlation between these properties is found when a variety of different metal oxide catalysts is compared. A close relationship is often shown, on the other hand, when a particular catalyst is modified or doped with minor amounts of an additional metal oxide. It is very likely that the correlation is successful in this case, because the nature of surface sites is not essentially changed. Studies have also been carried o u t which are more specifically aimed at charge transfer on an atomic scale and deal with the atomic situation within the lattice. This is especially so in the case of binary oxides. Many authors assume that, in these systems, both types of cation participate in electron transfer. The reactivity of the binary oxides is then explained by the hypothesis that the cation on the active site obtains an electron supply from the second type. 3.4.1 Bulk electrical properties
Holbrook and Wise [158,159]pay special attention to the electrical conductivity of copper oxide which catalyzes the selective oxidation and ammoxidation of propene. It was ascertained that selectivity is promoted by an oxygendeficient Cu,O in the case of propene conversion, as well as in the oxidation of ammonia t o nitrogen. The selectivity is lowered by increasing the oxygen-opper ratio and, with an oxygen-rich CuzO and CuO, complete combustion leading t o the formation of N 2 0 is the main reaction. There is a large change in the slope of a plot of conductivity as a function of oxygen pressure which coincides with a rapid selectivity change. The authors conclude that a charged oxygen species is responsible for this behaviour. The surface coverage with this species depends on the relative difference between the surface state energy level of oxygen and the Fermi level. The value of the activation energy for the NH3 oxidation ( < l o kcal mol-') is of the same order as the temperature coefficient of the conductivity. Sala and Trifiro [274] give evidence that dissolving antimony in SnO, increases and stabilizes the number of free electrons. Morrison [232,233] finds that the free energy of electrons in the bulk phase (Fermi energy) is about the same for different selective and active catalysts. He notes that this value is very near (or just above) the electron exchange level of oxygen and hence makes reduction of oxygen possible. References pp. 253-262
244 3.4.2 Charge transfer o n an atomic scale It is generally accepted that valency transitions of cations are connected with the redox mechanism. It is obvious therefore, that activity and selectivity demand that t h e cation in the active site has the right oxidation state before the hydrocarbon is adsorbed, and that it is effectively reoxidized afterwards. Accordingly, correlations are often found between activity, selectivity and the concentration of cations in specific oxidation states, e.g. V4+ in V z 0 5 . The improvement of selective catalytic qualities of metal oxides by addition of modifiers, or by combination in mixed oxides may hence be explained by stabilization of the essential cation in the proper oxidation state. In some cases, stabilization of a (partially) reduced cation appears to yield the most effective catalyst; however, more often it is the higher oxidation state that should be maintained, and accordingly the role attributed t o a second cation often concerns facilitating the reoxidation, for instance, by direct electron transfer between the cations or, in general, by increasing electron conductivity. Techniques that enable the observation of specific valencies of cations include E.S.R., y-resonance spectroscopy and ESCA, and have been considerably improved in the recent years. Bismuth molybdate (Bi/Mo = 0.7), MOO, and BiO, were investigated by Sancier et al. [275] who carried out ESR spectroscopy and conductivity measurements simultaneously. Reduction, reoxidation and steady state conditions were examined a t 325-380°C using propene and air. The kinetics of initial reduction and oxidation were treated in Sect. 2.2.2. As a measure of the conductivity, AV, the change in crystal voltage, was taken. Figure 1 2 demonstrates the relation between the degree of reduction with AV, on the one hand, and with the ESR signal strength for Mo5+ on the other. The highest degree of reduction in the steady state was observed for the bismuth molybdate sample. The reduction level depended o n the C3H6/02 molar ratio as illustrated in Fig. 13. It is remarkable that the signal strength of MoS+levels off in bismuth molybdate *. The authors explain this by supposing that “reduction by propene results in oxygen vacancies that form an impurity band and cause the conductivity t o increase. However, the concentration of Mo5+will reach a constant value because ionization of the vacancy levels t o the conduction band (to form Mo5+species) is limited by the pinned Fermi energy”. Various molybdates (Bi-Mo-O, Al-Mo-, Sn-Mo-, Fe-Mo-0) and a Vz0,-K2S04 catalyst were investigated by Maksimovskaya [ 1971 by ESR measurements. The Mo5+and V4+signals appeared t o increase by reduction with butene and t o disappear by successive reoxidation. The correlation between the ESR signal and the degree of reduction is good for
* Similar effects were reported earlier by Peacock et al. [250,251].
245
/
Bi /M0=0.7
.
7! BI/MO
=
o7
BI/Mo = 6 40 BipOq
20
Time ( s e c )
Fig. 1 2 . Dependence of ESR signal strength and change of crystal voltage o n degree of reduction of bismuth molyhdate catalysts during propene oxidation. , av;
-.
- . _ - -AI.
the vanadate catalyst, but less clear for the molybdates, in agreement with the results of Sancier et al. [ 2751. Iron moly bdates were investigated by several authors. I t is generally observed that iron is reduced first (Fe3+ Fez+),while deeper reduction is required to reduce the molybdenum ions as well. Both cations occur in partially reduced states during the reaction with butene. Pernicone [ 2541 concludes from his ESR work that under stationary reaction conditions the iron ions stay in the reduced state and that the redox process only involves Mo6+ and Mo". However, Trifiro and Pasquon [318]and Matsuura and Schuit [207] are of the opinion that reoxidation initially may lead to Fe3+which in turn (rapidly) oxidizes the Mo5+ ions at the hydrocarbon reaction sites of the catalyst. However, direct evidence is not provided. --f
246
-> E v
a,
m m
5 0
>
-
m +. Ul Ir L
U L
0 a, Is)
C
m
.c U
>‘ Q
t
-
C,H6/02
mole r a t i o
Fig. 13. Steady state value of crystal voltage as a function of C 3 H 6 / 0 2 mole ratio at constant flow (1 1 min-1 g cat-’) and constant P o 2 (0.12 atm).
An iron-modified Co-Mo--O catalyst was studied by Maksimov and Margolis [ 196,2031, using y-resonance spectroscopy. Replacing 3%of the cobalt in CoMo04 by Fe strongly increases the catalytic activity, while a hundredfold increase of the conductivity results. These effects are attributed to the occurrence of Fe”, the concentration of which doubles reversibly during reaction with propene air mixtures at 310-330°C. Strangely, the doubling is not observed with propene alone. The amount of M o 5 +on the surface of Mo-Ti-0 and Mo-Te-0 catalysts has been assessed with ESR techniques by Akimoto and Echigoya [13,15,17] and Andrushkevich et al. [27]. These workers find a strong correlation between the maximum intensity of the Mo5+signal with maximum activity in the oxidation of propene t o acrolein (at 8 at. % Te) and conversion of butadiene t o maleic anhydride (75 at. % Ti). Antimonates were also the subject of ESR investigations, and it was combinations only shown by Suzdalev et al. [310] that in Sn-Sb-0 reduction of Sb5+ takes place. The initial value of the Sb5+/Sb3+ratio These last two is 1 for S b z 0 4 , 2.3 for Sn-Sb-0 and 4.3 for Fe-Sb-. values are strongly reduced by chemisorption of acrolein. The Sb3+spectrum shows a change after formation of a complex with the adsorbate.
247 Margolis [203] confirms such results for antimonates and reports the existence of a surface compound containing Sb3+--O-C. Aykan and Sleight [34] examined the system U-Sb-0 in air up t o 1000°C by different techniques (e.g. ESR) and found the ternary components USb05 and USb3OI0. Since U S b 0 3 is paramagnetic, the formal oxidation state of U must be 5+, hence Sb must also be in the 5+ state. The authors conclude that USb3OI0 also contains pentavalent uranium. Finally, it may be noted that, although variable valency in binary oxides is important, it is not a sufficient requirement, as can be concluded from the fact that even in the systems Bi-Mo-0 and U-Sb-4 not every crystal phase is active and selective [294]. What matters is the configuration of the ions at the active site. Apparently, the character of the typical Me-0 bands is a function of the situation of oxygen in the lattice. 3.5 NATURE OF T H E ACTIVE SITES
It is generally accepted today that the oxidation activity of catalysts is not merely due t o the presence of a particular metal ion, but to the ensemble of metal and oxygen ions that forms the active site. The reactive properties of individual sites, where the interaction with molecules t o be oxidized takes place, and the determination of their geometry is the greatest challenge in catalytic research. With regard t o the geometry, a classical difficulty is the fact that the surface structure may differ considerably from the bulk. Only if surface and bulk structure are closely related may it be expected that specific crystal phases are responsible for active and selective oxidation. Otherwise these properties cannot be attributed t o a specific lattice structure. Although a discussion of the nature of active sites should, in fact, include all aspects of catalysis, attention will be focussed here on two aspects which receive considerable attention in the literature: the acidity or electron affinity of surface sites and the possible participation of different sites in one oxidation process, i.e. the bifunctional action of a catalyst. Some remarks must be made about the role of oxygen coordination. Several authors have remarked that the coordination in catalytic oxides is of major importance. Mitchell and Trifiro (e.g. ref. 219) concluded that a bismuth molybdate catalyst is most active if the amount of tetrahedrally coordinated molybdenum is large in comparison with octahedrally coordinated molybdenum. However, V , 0 5 and SbzO, are structures with specific octahedral coordination [ 1421 and often the coordination is changed by reduction of the catalyst or by the support [203]. In a - and 0-cobalt molybdates the coordination differs, but the catalytic behaviour is really the same. The low temperature Bi2Mo06(y phase) has an octahedral coordination but is an effective catalyst. It can be concluded from these and other investigations that the oxygen Rcfrrences P P . 253-262
248 coordination in the bulk is not a principal factor. It may very well be, however, that the type of coordination at the surface is important. Unfortunately, hardly any data are available. It may be expected that extension of electron spectroscopic techniques will throw light on this problem.
3.5.1 Acid -b use p ropert ies The interaction between selective metal oxides and molecules t o be oxidized is, of course, based on electron-accepting and electron-donating properties, respectively. In this way, Mo6+, V5+,etc. act as electron acceptors and molecules with 7r-bonds as donors. Ai et al. [5-121 have drawn attention t o the fact that this can also be described by acid-base properties. An electron donor molecule like butene is a basic entity interacting with acidic sites on the catalyst. Hence it follows that activity and selectivity depend on the relative acidity and basicity. MOO,, for example, is an acidic oxide, while Bi203is a basic oxide. Different compositions Bi: Mo have different acidities. The rate of oxidation depends on the number of acid sites (=acidity) and the acid strength, viz.
R
a
acidity X f (acid strength)
The same applies t o the rate of isomerization. The Ai and Suzuki [ 5,9] investigated the combination V,O,-P,O,. acidity was measured indirectly by the activity for dehydration of isopropanol and was shown t o decrease with increasing P 2 0 5 content. The activity for the oxidation of butene-1 and butadiene t o maleic acid anhydride decreased accordingly. It was shown that the adsorption equilibrium constant of the olefin on the catalyst also decreased in the same way. Ai [6,10,11] also reports work on Sn0,-based catalysts, i.e. Sn0,MOO,, Sn02-P20, and Sn0,-V,O,. SnO,, as such, does not have an acidic character but MOO, and V,Os change this effectively (more than P,O,). At 30-60 at. 3' 6 Mo, the acidity is highest and activity for isomerization and selective oxidation are a maximum. With tin vanadates, the selectivity for the formation of butadiene goes through a maximum at an atomic ratio Sn/V = 9. Below this ratio, the acidity is greater, leading t o more maleic acid anhydride in the reaction products. Butadiene will adsorb more with increasing acidity and will have a greater opportunity to be oxidized. The resulting acid anhydride will desorb relatively easily from an acid catalyst. A basic catalyst will result in more combustion products. Combinations of Bi203 and MOO,, promoted by P,O, at a constant P/Mo ratio (0.2) were studied over a full composition range by Ai and Ikawa [6]. Acidity (and basicity) were measured directly by adsorption of compounds like ammonia, pyridine and acetic acid. The effect of the Bi/Mo ratio on the acidity (Fig. 14) parallels the effect on the overall butene oxidation activity [presented in Fig. 5, Sect. 2.3.2(a)(i)].
249
,-. 0 c
-
X
0
-
E v
c
o_
+-
Cl L
0, D m
mK 0
E E Q
t
--
Fig. 1 4 . Acidity of 0.2).
I
I
0.2
0.4
I
0.6
I
0.8
, 3
BI
Atomic ratio ___ B I + Mo
Bi203-Mo03-P205
as a function of bismuth content (P/Mo =
With respect to the reaction products, the catalysts can be classified into three groups. The first group is very acidic in nature (Bi/Mo = 0 - 0 . 3 ) and converts olefins t o acidic products (e.g. butene t o maleic anhydride), the second group has medium acidity (Bi/Mo = 0.5-3) and provides the optimal conditions for the dehydrogenation of butene t o butadiene, while the third group (Bi/Mo > 3), which has a basic character, only forms combustion products. Pernicone et al. [ 253,2541 bring forward some evidence that surface acidity also plays a role with iron molybdate catalysts. Hammett indicators adsorbed over the molybdate assume the acid colour. Pyridine poisons the oxidation of methanol t o formaldehyde. A correlation is reported between acidity and activity [253]. The authors agree with Ai that the acid sites are connected with Mo6+ions. Ai finally notes that, with regard t o bismuth molybdates, such acid References p p . 253-262
250 sites can very well be equivalent t o the B-centers of Matsuura (p. 240) and the M o ( O ~ sites ) ~ of Trifiro (p. 240). Basic sites are then probably the oxidizing sites, equivalent t o Matsuura's A-centers. 3.5.2 Bifunctionality
Bifunctionality means that sites with different functions are present on the surface of a catalyst. In this general sense, two types of bifunctionality in hydrocarbon oxidation catalysis can be discerned.
(a) Bifunctionality connected with the redox mechanism Strong indications are present for some mixed oxide catalysts that the interaction with the molecule to be oxidized and the oxygen that reoxidizes the catalyst take place on different sites and involve different cations. These two sites may together form one ensemble that performs the complete reaction. However, they may also be actually separated and quite remote, provided that the transport of anions and the conduction of charge between such sites is sufficiently large. Bismuth molybdate-based catalysts are well known examples for which these conditions apply. Unfortunately, there is no agreement as to which function must be connected with either cation. For scheelite-type bismuth molybdates, Linn and Sleight [188] have advanced the theory that Bi cations with their free electron doublets at the surface are the favoured centres for reoxidation, while propene oxidation takes place at the Mo-0 tetrahedra. Schuit [ 2811 also assumes that, for bismuth molybdates, oxygen is introduced at the bismuth sites, but his mechanism is more complicated, as both Bi and Mo interact with the hydrocarbon substrate. Direct evidence to the contrary, i.e. oxygen introduction in the molybdenum layers of bismuth molybdate, has been provided by Otsubo et al. [ 2461. They proved that catalyst reduction by hydrogen and reoxidation with "02yields Bi203 MoI8O3, while hydrogen primarily consumes bismuth oxygen. For several other binary oxide catalysts, this type of bifunctionality is indicated by the fact that both cations are partially reduced under reaction conditions, as observed by ESCA and y resonance techniques. An example of a catalyst investigated is FeMo04 [ 751.
-
( b ) Bifunctionality related to different reaction steps Several authors have suggested that the allylic oxidation of olefins t o aldehydes requires a bifunctional catalyst. The two functions then concern the formation of an allylic radical and the coupling of such a radical with lattice oxygen. This idea is primarily based on the fact that several single oxides (e.g. Bi203, S n 0 2 , T1203)catalyze the formation of ally1
radicals, but lack the capability t o transfer oxygen; hence allyl dimers are formed. Molybdenum oxide, on the other hand, appears t o have a capacity t o oxidize allyl radicals t o acrolein, a capacity which largely exceeds that for the oxidation of propene [121]. Unfortunately, no other oxides have been investigated with respect t o their specific reactivity towards allyl radicals. The possible bifunctionality of bismuth molybdates is amply discussed in Sect. 2.2.2(d)(i). The fact that active and selective catalysts in general comprise two or more oxide components is certainly not a sufficient argument t o assume bifunctionality ; the combination of oxides may also cause modification of sites or formation of one type of new sites which combine the specific properties required for a sequence of reaction steps. Such properties may concern the geometry, the type of oxygen bonding, oxygen and charge mobility in the solid, acidity, etc. It can be concluded that the occurrence of a two-centre mechanism is not easily distinguished from that of a mechanism involving multi-function reaction sites, the more so as the separation of catalytic sites and of reaction steps is a practical difficulty. An example of real bifunctionality appears t o be the case of acrylic acid formation, because two reaction steps which can be individually studied, are involved, i.e. the formation of acrolein, in which lattice oxygen is incorporated, and the aldehyde t o acid conversion, which involves water as the oxygen source. The most effective catalysts are multi-component catalysts, which very likely possess different sites, probably on different catalyst phases (see Sect. 2.3.3). 3.6. ADSORPTION AND REACTION COMPLEXES ON T H E CATALYTIC SURFACE
In the foregoing discussion, the emphasis has been mainly on the properties of the catalyst, but it is evident that these must be regarded in close connection with the nature of the adsorbed hydrocarbons. Important information about this interaction can be gained from structure analysis of adsorption and reaction complexes, as well as adsorption measurements, Infrared spectra of propene and isobutene on different catalysts were measured by Gorokhovatskii [ 1431. Copper oxide, which converts olefins to butadiene and aldehydes, shows adsorption complexes different from structures on a V z 0 5 - P 2 0 5 catalyst which produces maleic acid anhydride. Differences also exist between selective oxidation catalysts and total oxidation catalysts. The latter show carbonate and formate bands, in contrast to selective oxides for which 7r-allylic species are indicated. A difficulty in this type of work is that only a few data are available under catalytic conditions; most of them refer t o a pre-catalysis situation. Therefore it is not certain that complexes observed are relevant for the catalytic action. References p p . 2 53-262
252
Sachtler [270] notes that the n-ally1 complex can be attached t o a metal ion or to an oxygen anion but doubts that a n-ally1 metal complex can be stable at the high temperatures normally used. He draws attention t o the fact that, in the case of aromatic oxidations, benzoates, maleinates, etc. are observed spectroscopically, indicating that a carbon-metal bond is not formed. 0 / Trifiro et al. [322], however, did not find (R-C=O)- groups in an investigation of Sn02-V205 catalyst by infrared spectroscopy. The spectra reveal the presence of MOO, on the surface. If propene is adsorbed, the Mo=O band of the oxide is influenced. The Mo=O band disappears when acrolein is adsorbed (at room temperature). Desorption at 225°C restores this peak. Electron spectroscopic studies were carried out by Haber et al. [150] on CoMo04, MOO, and Mooz at -200 t o +5OO0C. It was demonstrated that, during interaction with acrolein, a change in the spectra can be observed, indicating the change of a vinylic carbon atom t o a paraffinic one. Simultaneously, carbonyl peaks change into carboxyl. The conclusion is 0 //
that acrolein becomes bonded by a C
group. With desorption, decarboxy-
\
0 lation occurs leaving the hydrocarbon part at the surface. Studies with propionic acid showed that reduction of the surface favours decarboxylation of the acid molecules which does not occur on oxidized CoMo04. Haber suggests that the relation between active sites of different types is associated with the nature of the active complex (ref. in Butt [67]). The final products can be classified in the following manner. Active center
Active complex
Product
Cations, Bi3', Co2+, Sn4+ 0 2 -in polyhedra of Mo, W, Sb, Nb OH- of basic character OH- of acidic character
n-allylic complex
Dienes
o-bonded allylic species Carboxylate type of complex Carbonium ion
Unsaturated aldehydes and ketones Unsaturated acids
_ ___.
Saturated ketones
~ ~ _ _ _ _ _ _ _
Trifiro and Carra [ 3231 used the amount of c i s t r u n s isomerization and double bond shift as a method of investigating the type of intermediate. It is concluded that three groups of oxides exist. The first is a group of which Bi-Mo-O, Bi2W06 and Fe-Te-Mo-O are typical. These catalysts give only isomerization a t temperatures at which selective oxidation also occurs, probably via the same intermediate (allylic). A second group gives isomerization at much lower temperatures (Sn-Sb-0, Fe-Sb-0, for
253 example) and it is suggested that acidic Lewis centres are the cause of double bond isomerization. A third group (Co, Mn, Fe molybdates) carries the oxidation further and probably contains Bronstedt acid centers, which would act via carbocations according t o the authors. Finally, a classification of catalysts by Matsuura [212] may be mentioned, in which the relation of adsorption entropy t o heat of adsorption of butene-1 appears, surprisingly, t o be linear. The conclusion can be drawn that moderate heats of adsorption (about 40-50 kcal mol-’) characterize suitable catalysts. Only here is the right combination of surface mobility and adsorption intensity found. Apparently, the oxygen is then “tempered” sufficiently t o make a selective oxidation possible. Otherwise, the oxides are non-active (e.g. low heat of adsorption in FeP04 and low mobility) or active but non-selective because of high mobility coupled to a large heat of adsorption (e.g. Fe,04).
References 1 Abd. El-Salaam, E. Echigoya and M. Akimoto, Z. Phys. Chem. (Frankfurt am Main), 96 (1975) 311. 2 C.R. Adams, Proc. 3rd Int. Congr. Catalysis, Amsterdam, 1964, p. 240. 3 S.I. Ahmad, S.H. Ibrahim and N.R. Kuloor, Indian J. Technol., 8 (1970) 131. 4 S.I. Ahmad, S.H. Ibrahim and N.R. Kuloor, Indian J. Technol., 9 (1971) 251. 5 M. Ai and S. Suzuki, Bull. Chem. SOC.Jpn., 47 (1974) 3074. 6 M. Ai and T . Ikawa, J. Catal., 40 (1975) 203. 7 M. Ai, Bull. Chem. SOC.Jpn., 44 (1971) 761. 8 M. Ai and S. Suzuki, Bull. Jpn. Pet. Inst., 1 6 (1974) 118. 9 M. Ai, Bull. Chem. SOC.Jpn., 4 3 (1970) 3490. 10 M. Ai, J. Catal., 40 (1975) 318. 11 M. Ai, J. Catal., 40 (1975) 327. 12 M. Ai and S. Suzuki, J. Catal., 26 (1972) 202. 1 3 M. Akimoto and E. Echigoya, J. Catal., 31 (1973) 278. 14 M. Akimoto and E. Echigoya, Bull. Chem. SOC.Jpn., 48 (1975) 3518. 15 M. Akimoto and E. Echigoya, Bull. Jpn. Pet. Inst., 1 6 (1974) 8. 16 M. Akimoto and E. Echigoya, J . Catal., 35 (1974) 278. 1 7 M. Akimoto and E. Echigoya, J . Catal., 29 (1973) 191. 18 F. Alfani and 5.3. Carberry, Chim. Ind. (Milan), 52 (1970) 1192. 19 T.G. Alkhazov. K.Y. Adzhamov and N.K. Allakhverdova, Kinet. Katal., 1 5 (1974) 1492. 20 T.G. Alkhazov and K.Y. Adzhamov, Kinet Katal., 1 5 (1974) 201. 21 A.D. Ammosov and L.A. Sazonov, Kinet. Katal., 16 (1975) 1241. 22 A.D. Ammosov and L.A. Sazonov, React. Kinet. Catal. Lett., 3 (1975) 255. 23 A.D. Ammosov, L.A. Sazonov and S.A. Venyaminov, Kinet. Katal., 1 2 (1971) 608. 24 A.D. Ammosov and L.A. Sazonov, React. Kinet, Catal. Lett., 2 (1975) 403. 25 E.I. Andreikov and N.D. Rusyanova, Kinet. Katal., 10 (1969) 722. 26 E.I. Andreikov, Y.A. Sveshnikova and N.D. Rusyanova, Kinet. Katal., 1 6 (1975) 919. 27 T.V. Andrushkevich, G.K. Boreskov, L.L. Kuznetosova, L.M. Plyasova, Y.N. Tyurin and Y.M. Shchekochikhin, Kinet. Katal., 1 5 (1974) 424.
28 Anonymous, Hydrocarbon Process, Petrochemical Handbook Issue, 54 (1975) 97. 29 A.G. Anshits, V.D. Sokolovskii, G.K. Boreskov, Y.D. Pankrat’ev, V.F. Malakhov and A.M. Kolchin, Kinet. Katal., 1 4 (1973) 1209. 30 A.G. Anshits, V.D. Sokolovskii and G.K. Boreskov, Kinet. Katal., 15 (1974) 730. 31 A.G. Anshits, V.D. Sokolovskii and G.K. Boreskov, Kinet. Katal., 1 5 (1974) 1033. 32 A. Ayame, H. Kano, T . Kanazuka and H. Baba, Bull. Jpn. Pet. Inst., 1 5 (1973) 150. 33 K. Aykan, D. Halvorson, A.W. Sleight and D.B. Rogers, J. Catal., 35 (1974) 401. 34 K. Aykan and A.W. Sleight, d. Amer. Ceram. SOC.,53 (1970) 427. 35 Y.M. Bakshi, R.N. Gur’yanova and A.I. Gel’bshtein, Kinet Katal., 16 (1975) 443. 36 Y.M. Bakshi, N.K. Danilova, R.N. Gur’yanova and A.I. Gel’bshtein, Neftekhimiya, 11 (1971) 810, 817. 37 S . Balasubramanian and D.S. Viswanath, Ind. Eng. Chem. Fundam., 14 (1975) 158. 38 G.B. Barannik, S.A. Ven’yaminov, G.K. Boreskov and H. Wolf, React. Kinet. Catal. Lett., 2 (1975) 191. 39 G.B. Barannik and S.A. Yen’yaminov, React. Kinet. Catal. Lett., 4 (1976) 437. 40 Ph.A. Batist, P.C.M. van der Heijden and G.C.A. Schuit, J. Catal., 22 (1971) 411. 41 Ph.A. Batist, J.F.H. Bouwens and G.C.A. Schuit, J. Catal., 25 (1972) 1. 42 Ph.A. Batist, J.F.H. Bouwens and I . Matsuura, J. Catal., 32 (1974) 362. 43 Ph.A. Batist, H.J. Prette and G.C.A. Schuit, J. Catal., 1 5 (1969) 267. 44 Ph.A. Batist, L.G.M. van de Moesdijk, I. Matsuura and G.C.A. Schuit, J . Catal., 20 (1971) 40. 45 Ph.A. Batist and S.P. Lankhuijzen, J. Catal., 28 (1973) 496. 46 H. Baussart, R . Delobel, G. Francois and J.M Leroy, C.R. Acad. Sci. Ser. C, 281 (1975) 375. 47 J. Beres, K. Brueckman, J. Haber and J . Janas, Bull. Acad. Pol. Sci., Ser. Sci. Chim., 20 (1972) 813. 48 S.K. Bhattacharyya, N.K. Nag and N.D. Ganguly, J. Catal., 2 3 (1971) 158. 49 A. Bielanski and A. Inglot, Bull. Acad. Pol. Sci., Ser. Sci. Chim., 22 (1974) 785. 50 A. Bielanski and A. Inglot, Bull. Acad. Pol. Sci., Ser. Sci. Chim., 22 (1974) 773. 51 M. Blanchard and G. Louguet, Kinet. Katal., 1 4 (1973) 30. 52 M. Blanchard and D. Vanhove, Bull. SOC.Chim. Fr., (1971) 4134. 5 3 M. Blanchard, G. Louguet, G.K. Boreskov, V.F. Muzykantov and G.I. Panov, Bull. S O C . Chim. Fr., (1971) 814. 54 G. Blasse, J . Inorg. Nucl. Chem., 28 (1966) 1124. 55 I.F. Boag. D.W. Bacon and J. Downie, J. Catal., 38 (1975) 375. 56 M.A.M. Boersma, Thesis, Eindhoven University of Technology, 1977. 57 G.K. Boreskov, A.A. Ivanov, O.M. Iluyich and V.G. Ponomareva, React. Kinet. Catal. Lett., 3 (1975) 1. 58 G.K. Boreskov, Kinet. Katal., 14 (1973) 7. 59 P. Boutry, R. Montarnal and J . Wrzyszcz, J. Catal., 1 3 (1969) 75. 60 A.M. Bradshaw, A. Engelhardt and D. Menzel, Ber Bunsenges. Phys. Chem., 76 (1972) 500. 61 K. Brueckman, J . Haber and J. Janas, Bull. Acad. Pol. Sci. Ser. Sci. Chim., 21 (1973) 763. 62 J . Buiten, J . Catal., 10 (1968) 188. 63 J. Buiten, J. Catal., 1 3 (1969) 373. 64 J. Buiten, J . Catal., 27 (1972) 232. 65 L. Burlamacchi, G. Martini and F . Trifiro, J. Catal., 33 (1974) 1.
66 L. Burlamacchi, G. Martini, F. Trifiro and G. Caputo, J. Chem. SOC., Faraday Trans. I, 7 1 (1975) 209. 67 J.B. Butt, AIChE J., 22 (1976) 1. 68 P.V. Butt and C.N. Kenney, Proc. 6th Int. Congr. Catal., The Chemical Society, London, 1977, p . 779. 69 P.H. Calderbank, Adv. Chem. Ser., 133 (1974) 646. 70 J.L. Callahan, R.K. Grasselli, E.C. Milberger and H.A. Strecker, Ind. Eng. Chem. Prod. Res. Dev., 9 (1970) 134. 7 1 W.E. Campbell, E.L. McDaniel, W.H. Reece, J.E. Williams and H.S. Young, Ind. Eng. Chem. Prod. Res. Dev., 9 (1970) 525. 72 N.W. Cant and W.K. Hall, J. Catal., 1 6 (1970) 220. 73 N.W. Cant and W.K. Hall, J . Catal., 22 (1971) 310. 7 4 J.J. Carberry, G.C. Kuczinski and E. Martinez, J. Catal., 26 (1972) 247. 75 M. Carbucicchio and F. Trifiro, J. Catal., 4 5 (1976) 77. 76 W.R. Cares and J.W. Hightower, J. Catal., 23 (1971) 193. 7 7 W.R. Cares and J.W. Hightower, J. Catal., 39 (1975) 36. 7 8 J. Cartlidge, L. McGrath and S.H. Wilson, Trans. Inst. Chem. Eng., 52 (1974) 222. 79 J . Cartlidge, L. McGrath and S.H. Wilson, Trans. Inst. Chem. Eng., 5 3 (1975)117. 80 A. Castellan, A. Vaghi, J.C.J. Bart and N. Giordano, J. Catal., 39 (1975) 213. 81 M. Cathala and J.E. Germain, Bull. SOC.Chim. Fr., (1970) 4114. 82 M. Cathala and J.E. Germain, Bull. SOC.Chim. Fr., (1971) 2174. 8 3 M. Cathala and J.E. Germain, Bull. SOC.Chim. Fr., (1971) 2167. 84 P. Centola, C. Mazzocchia, G. Terzaghi and I. Pasquon, Chim. Ind. (Milan), 54 (1972) 859. 85 M. Cesari, G. Perego, A. Zazzetta, G. Manari and B. Notari, J. Inorg. Nucl. Chem., 33 (1971) 3595. 86 J.R. Christie, D. Taylor and C.C. McCain, J. Chem. SOC., Faraday Trans. I , 72 (1976) 334. 87 R.B. Clarkson and A.C. Cirillo, J. Vac. Sci. Technol., 9 (1972) 1073. 8 8 R.B. Clarkson and A.C. Cirillo, J. Catal., 33 (1974) 392. 89 D.J. Cole, C.F. Cullis and D.J. Hucknall, J. Chem. SOC.,Faraday Trans. I, 72 (1976) 2185. 90 D. Cormack, D.H. Thomas and R.L. Moss, J. Catal., 32 (1974) 492. 9 1 A. Cortes and J.L. Seoane, J. Catal., 34 (1974) 7. 92 M. Crozat and J.E. Germain, Bull. SOC.Chim. Fr., (1972) 3526. 9 3 M. Crozat and J.E. Germain, Bull. SOC.Chirn. Fr., (1973) 2498. 94 C. Daniel and G.W. Keulks, J. Catal., 29 (1973) 475. 95 C. Daniel and G.W. Keulks, J. Catal., 24 (1972) 529. 96 C. Daniel, J.R. Monnier and G.W. Keulks, J. Catal., 3 1 (1973) 360. 97 A.A. Davydov and A.A. Budneva, Kinet. Katal., 15 (1974) 1557. 98 J. Dechelette and J.E. Germain, Bull. SOC.Chim. Fr., (1976) 874. 99 J.C. Delgrange and M. Blanchard, Bull. SOC.Chirn. Fr., (1971) 1093. 100 M. Dente, E. Ranzi, D.D. Quiroga and G. Biardi, Chim. Ind. (Milan), 55 (1973) 563. 101 J.K. Dixon and J.E. Longfield, in P.H. Emmett (Ed.), Catalysis, Vol. VII, Reinhold, New York, 1960, p. 183. 102 Z. Dolejsek and J. Novakova, J. Catal., 37 (1975) 540. 103 T. Dozono, D.W. Thomas and H. Wise, J. Chem. SOC., Faraday Trans. I, (1973) 620. 104 A.F. van den Elzen and G.D. Rieck, Acta Crystollogr. Sect. B, 29 (1973) 2433, 2436. LO5 L.Y. Erman, E.L. Galperin, I.K. Kolchin, G.F. Dobrzhanskii and K.S. Chernyshev, Zh. Neorg. Khim., 9 (1964) 2174.
106 107 108 109 110
A.B. Evnin, J.A. Rabo and P.H. Kasai, J . Catal., 30 (1973) 109. G. Fagherazzi and N. Pernicone, J. Catal., 1 5 6 (1970) 321. V. Fattore, Z.A. Fuhrman, G. Manara and B. Notari, J. Catal., 37 (1975) 215. V. Fattore, Z.A. Fuhrman, G. Manara and B. Notari, J . Catal., 37 (1975) 223. E.V. Fedevich, V.M. Zhiznevskii, M.V. Nikipanchuk, L.F. Yakubovskaya and I.M. Golub, Kinet. Katal., 1 5 (1974) 1288. 111 F. Figueras, M. Forissier, L. de Mourgues, J.L. Portefaiz and M. Rosa-Brussin, Kinet. Katal., 1 4 (1973) 25. 112 A.A. Kirsova, N.N. Khovanskaya, A.D. Tsyganov, I.P. Susdalev and L.Y. Margolis, Kinet. Katal., 1 2 (1971) 792. 113 A.A. Kirsova, I.P. Suzdalev and L.Y. Margolis, Kinet. Katal., 14 (1973) 1469. 114 E.L. Force and A.T. Bell, J . Catal., 38 (1975) 440. 115 E.L. Force and A.T. Bell, J . Catal., 44 (1976) 175. 116 E.L. Force and A.T. Bell, J. Catal., 40 (1975) 356. 117 J. Forys and B. Grzybowska, Bull. Acad. Pol. Sci. Ser. Sci. Chim., 23 (1975) 269. 118 P. Forzatti, A. Klimasara, G.C. Kuczynski and J.J. Carberry, J . Catal., 29 (1973) 169. 119 K. Fujimoto and T. Kunugi, Proc. 5th Int. Congr. Catal. Miami Beach, 1972, North-Holland, Amsterdam, 1973, Vol. 1 , p. 445. 120 K. Fujimoto, H. Takeda and T. Kunugi, Ind. Eng. Chem. Prod. Res. Dev., 1 3 (1974) 237. 121 E.G. Gamid-Zade, A.R. Kuliyev, E.A. Mamedov, R.G. Rizayev and V.D. Sokolovski, React. Kinet. Catal. Lett., 3 (1975) 191. 122 1.1. Geiman, D.R. Kreile and V.A. Slavinskaya, Kinet. Katal., 1 6 (1975) 957. 123 A.I. Gel’bshtein, Y.A. Mishchenko, P.P. Nechiporuk and N.D. Gol’dshtein, Dokl. Akad. Nauk SSSR, 217 (1974) 377. 124 A.I. Gel’bshtein, S.S. Stroeva, A.N. Parfenov and T.V. Zinoveva, Kinet. Katal., 17 (1976) 412 (Engl. Transl. p. 357). 1 2 5 S.J. Gentry, R. Rudham and M.K. Sanders, J. Catal., 35 (1974) 376. 126 H.R. Gerberich, N.W. Cant and W.K. Hall, J. Catal., 1 6 (1970) 204. 127 J.E. Germain and J.C. Peuch, Bull. SOC.Chim. Fr., (1969) 1844. 128 J.E. Germain and R. Perez, Bull. SOC.Chim. Fr., (1975) 739. 129 J.E. Germain and R. Laugier, Bull. SOC.Chim. Fr., (1971) 541. 130 J.E. Germain and R. Laugier, Bull. SOC.Chim. Fr., (1971) 650. 131 J.E. Germain and R. Laugier, C.R. Acad. Sci., Ser. C, 276 (1973) 1349. 132 J.E. Germain and R. Laugier, C.R. Acad. Sci., Ser. C, 276 (1973) 371. 1 3 3 J.E. Germain and R. Laugier, Bull. SOC.Chim. Fr., (1972) 2910. 134 J.E. Germain, Intra-Sci. Chem. Rep., 6 (1972) 101. 1 3 5 J.E. Germain and R. Perez, Bull. SOC.Chim. Fr., (1975) 1216. 1 3 6 K. German, B. Grzybowska and J. Haber, Bull. Acad. Pol. Sci., Ser. Sci. Chim., 21 (1973) 319. 137 E. Ghenassia and J.E. Germain, Bull. SOC.Chim. Fr., (1975) 731. 138 M.A. Gibson and J.W. Hightower, J. Catal., 4 1 (1976) 420. 139 M.A. Gibson and J.W. Hightower, J. Catal., 4 1 (1976) 431. 140 N. Giordano and J.C.J. Bart, J. Roy. Neth. Chem. SOC.,94 (1975) 28. 141 N. Giordano, A. Vaghi, J.C.J. Bart and A. Castellan, J. Catal., 38 (1975) 11. 142 G.W. Godin, C.C. McCain and E.A. Porter, Proc. 4 t h Int. Congr. Catal., Moscow, 1968, Akademiai Kiad6, Budapest, 1971, p. 271. 1 4 3 Y.B. Gorokhovatskii, Kinet. Katal., 1 4 (1973) 83. 144 A.P. Gorshkov, G.V. Isagulyants, Y.I. Derventesev, L.Y. Margolis and I.K. Kolchin, Dokl. Akad. Nauk SSSR, 1 8 6 (1969) 827. 145 R. Grabowski, B. Grzybowska, J. Haber and J. Sloczynski, React. Kinet. Catal. Lett., 2 (1975) 81.
146 R.K. Grasselli and D.D. Suresh, J. Catal., 25 (1972) 273. 1 4 7 B. Grzybowska, K. German, J . Haber, T. Kowalski and A. Mazurkiewicz, Proc. Symp. Mechan. Hydrocarbon React. Hungary, 1973, p. 265. 148 B. Grzybowska, J . Haber, W. Marczewski and L. Ungier, J . Catal., 42 (1976) 327. 149 B. Grzybowska, J . Haber and J . Komorek, J . Catal., 25 (1972) 25. 150 J . Haber, W. Marczewski, J. Stoch and L. Ungier, Proc, 6 t h Int. Congr. Catal., The Chemical Society, London, 1977, p. 827. 151 J. Haber, in P.C.H. Mitchell (Ed.), Proceedings of t h e Conference o n the Chemical Uses of Molybdenum, Climax Molybdenum Co. Ltd., London, (1973, p. 146. 152 J . Haber and B. Grzybowska, J. Catal., 28 (1973) 489. 153 B. Halpern and J.E. Germain, J. Catal., 37 (1975) 44. 154 J . Happel, R.E. Lief and R. Mezaki, J . Catal., 31 (1973) 118. 155 J.L. Herce, J . Besombes-Vailhe and R. Bugarel, Kinet. Katal., 14 (1973) 90. 156 W. Herzog, Ber. Bunsenges. Phys. Chem., 74 (1970) 216. 157 C.N. Hodges and L.C. Roselaar, J . Appl. Chem. Biotechnol., 25 (1975) 609. 158 L.L. Holbrook and H. Wise, J. Catal., 20 (1971) 367. 159 L.L. Holbrook and H. Wise, J. Catal., 27 (1972) 322. 160 D.J. Hucknall, Selective Oxidation of Hydrocarbons, Academic Press, New York, 1974. 161 N.I. Ilchenko and G.I. Golodets, J. Catal., 39 (1975) 57, 73. 162 N.I. Ilchenko and G.I. Golodets, Kinet. Katal., 1 5 (1974) 394. 163 N.I. Ilchenko, I.M. Avilova and G.I. Golodets, Kinet. Katal., 16 (1975) 679. 164 Y.V. Ionov, L.K. Zhidkova, N.V. Kulkova and M.I. Temkin, Kinet. Katal., 14 (1973) 642. 165 I.S. Jaswal, R.F. Mann, J.A. Juusola and J. Downie, Can. J. Chem. Eng., 47 (1969) 284. 166 A. Joseph, R.L. Mednick, L.M. Shorr, S.W. Weller and P. Rona, Isr. J. Chem., 12 (1974) 739. 167 B. Kadlec, J. Michalek and A. Simecek, Chem. Eng. Sci., 25 (2970) 319. 168 K. Kaneko, T. Hoshino and S. Wada, Bull. Jpn. Petrol. Inst., 1 6 (1974) 24. 169 K. Kaneko, T. Koyama and S. Wada, Bull. Jpn. Petrol. Inst., 1 6 (1974) 17. 170 A. Kato, K. Tomoda, I. Mochida and T. Seiyama, Bull. Chem. SOC. Jpn., 45 (1972) 690. 1 7 1 V.B. Kazanskii, Kinet. Katal., 1 4 (1973) 95. 172 K. Keizer, Ph.A. Batist and G.C.A. Schuit, J. Catal., 15 (1969) 256. 173 R.E. Kenson and M. Lapkin, J . Phys. Chem., 74 (1970) 1493. 174 G.W. Keulks, J . Catal., 19 (1970) 232. 175 G.W. Keulks and L.D. Krenzke, Proc. 6 t h Int. Congr. Catal., The Chemical Society, London, 1977, p. 806. 176 P.A.Kilty, N.C. Rol and W.M.H. Sachtler, Proc. 5th Int. Congr. Catal. Miami Beach, 1972, North-Holland, Amsterdam, 1973, p. 929. 1 7 7 P.A. Kilty and W.M.H. Sachtler, Catal. Rev., 1 0 (1974) 1. 178 P.D. Klugherz and P. Harriott, AIChE J., 17 (1971) 856. 179 N.A. Komarovskii, A.L. Tsailingol’d, M.E. Basner, F.S. Philipenko and M.G. Slin’ko, Kinet. Katal., 1 4 (1973) 1483. 180 L.M. Koval and G.K. Boreskov, Kinet. Katal., 1 5 (1974) 1193. 181 R. Krabetz, Chem. Ing. Tech., 46 (1974) 1029. 182 G. Kremenic and P. Perez Torio, Acta Cient. Venez., 34 (1973) 204. 183 O.V. Krylov, Kinet. Katal., 14 (1973) 35. 184 O.V. Krylov, Proc. 5th Int. Congr. Catal., Miami Beach, 1972, North-Holland, Amsterdam, 1973, p. 263. 185 R. Lakshmanan and D. Rouleau, Can. J. Chem. Eng., 47 (1969) 45. 186 S.P. Lankhuijzen, P.M. Florack and H.S. van der Baan, J . Catal., 42 (1976) 20. 187 G. Liberti, N . Pernicone and S. Soattini, J. Catal., 27 (1972) 52.
188 W.J. Linn and A.W. Sleight. J. Catal., 4 1 (1976) 134. 189 A.E. Lisovskii, T.G. Alkhazov, A.M. Dadasheva and S.A. Feizullaeva, Kinet. Katal., 1 6 (1975) 456. 190 K.D. Lodaya, A.H. Tijoriwala and M.R. Chivate, Indian J. Technol., 14 (1976) 36. 1 9 1 M. Lojacono, T. Notermann and G.W. Keulks, J. Catal., 40 (1975) 19. 192 A.G. Lyubarskii, A.G. Gorelik, V.P. Petoyan, E.V. Lyapin and V.S. Beskov, Kinet. Katal., 1 4 (1973) 410. 193 J. Machek and J. Tichy, Coll. Czech. Chem. Commun., 40 (1975) 3774. 194 J.M. Maissant, D. Vanhove and M. Blanchard, Bull. SOC.Chim. Fr., (1973) 2787. 195 Y.V. Maksimov, I.P. Suzdalev, V.I. Gol’danskii, O.V. Krylov, L.Y. Margolis and A.E. Nechitailo, Dokl. Akad. Nauk SSSR, 211 (1975) 880. 196 Y.V. Maksimov and L.Y. Margolis, Chem. Phys. Lett., 34 (1975) 172. 197 R.I. Maksimovskaya, S.A. Ven’yaminov and G.K. Boreskov, Dokl. Akad. Nauk SSSR, 211 (1973) 1389. 198 V.V. Malakhov and F.G. Abdikova, Kinet. Katal., 1 3 (1972) 196. 199 R. Malinski, M. Akimoto and E. Echigoya, J. Catal., 44 (1976) 101. 200 R.S. Mann and K.C. Yao, Ind. Eng. Chem. Prod. Res. Dev., 10 (1971) 25. 201 R.S. Mann, K.C. Yao and M.K. Desi, J . Appl. Chem. Biotechnol., 22 (1972) 915. 202 R.S. Mann and D.W. KO, J. Catal., 30 (1973) 276. 203 L.Y. Margolis, Catal. Rev., 8 (1973) 241. 204 P. Mars and D.W. van Krevelen, Chem. Eng. Sci., 3 (1954) 4 1 (spec. suppl.). 205 F.E. Massoth and D.A. Scarpiello, J. Catal., 2 1 (1971) 294. 206 B.C. Mathur and D.S. Viswanath, J. Catal., 32 (1974) 1. 207 I. Matsuura and G.C.A. Schuit, J. Catal., 20 (1971) 19. 208 I. Matsuura and G.C.A. Schuit, J. Catal., 25 (1972) 314. 209 I. Matsuura, J. Catal., 3 3 (1974) 420. 210 I. Matsuura, J. Catal., 35 (1974) 452. 211 I. Matsuura and M. Wolfs, J. Catal., 37 (1975) 174. 21 2 I. Matsuura, Proc. 6 t h Int. Congr. Catal., The Chemical Society, London, 1977, p. 819. 213 C. Mazzocchia, P. Centola, R. Del Rosso, G. Terzaghi and I. Pasquon, Chim. Ind. (Milan), 55 (1973) 687. 214 C.C. McCain, G. Gough and G.W. Godin, Nature (London), 198 (1963) 989. 215 E.A. Melnik, L.P. Shapovalova. I.Y. Tyuryaev and V.M. Belousov, Kinet. Katal., 1 4 (1973) 381. 216 P.L. Metcalf and P. Harriott, Ind. Eng. Chem. Proc. Res. Dev., 11 (1972) 478. 217 R. Mezaki and B. Kadlec, J. Catal., 25 (1972) 454. 218 J. Mikami, S. Satok and H. Kobayashi, J. Catal., 1 8 (1970) 265. 219 P.C.H. Mitchell and F. Trifiro, J. Chem. SOC.A, (1970) 3183. 2 20 H. Miura, Y. Morikawa and T. Shirasaki, J. Catal., 39 (1975) 22. 221 K. Miyake, H. Oka, H. Sakamoto, Y. Harano and T , Imoto, J. Chem. SOC.Jpn., (1972) 284. 222 Y. Morita, M. Miyazaki and E . Kikuchi, Bull. Jpn. Pet. Inst., 17 (1975) 71. 223 Y. Moro-oka and A. Ozaki, J. Catal., 5 (1966) 116. 224 Y. Moro-oka, Y. Morikawa and A. Ozaki, J. Catal., 7 (1967) 23. 225 Y. Moro-oka, S. Tan and A. Ozaki, J. Catal., 1 2 (1968) 291. 226 Y. Moro-oka, S. Tan, Y. Takita and A.Ozaki, Bull. Chem. SOC. Jpn., 4 1 (1968) 2820. 227 Y. Moro-oka, S. Tan and A. Ozaki, J. Catal., 17 (1970) 125. 2 28 Y. Moro-oka, Y. Takita and A. Ozaki, J. Catal., 23 (1971) 183. 2 29 Y. Moro-oka, Y. Takita and A. Ozaki, Bull. Chem. SOC.Jpn., 44 (1971) 293. 230 Y. Moro-oka, Y. Takita and A. Ozaki, J. Catal., 27 (1972) 177.
231 Y. Moroitka and Y. Takita, Proc. 5th Int. Congr. Catal., Miami Beach, 1972, North-Holland, Amsterdam, 1973, p. 1025. 232 S.R. Morrison, J. Catal., 20 (1971) 110. 233 S.R. Morrison, J. Catal., 34 (1974) 462. 234 W.J. Muizebelt and W.J. van Ooij, J. Catal., in press. 235 M. Niwa and Y. Murakami, J. Catal., 26 (1972) 359. 236 M. Niwa and Y. Murakami, J. Catal., 27 (1972) 26. 237 T. Notermann, G.W. Keulks, A. Skliarov, Y. Maximov, L.Y. Margolis and 0 . V . Krylov, J . Catal., 39 (1975) 286. 238 J. Novakova and P. Jiru, J. Catal., 27 (1972) 155. 239 J. Novakova, Z. Dolejsek and K. Habersberger, React. Kinet. Catal. Lett., 4 (1976) 389. 240 J. Novakova, Catal. Rev., 4 (1970) 77. 241 E.C. Novella, F.C. Rubio, J.L. Sotelo and R.A. Pedreira, An. Quim., 69 (1973) 1187. 242 H. Oka, K. Miyake, Y. Harano and T. Imoto, J. Appl. Chem. Biotechnol., 25 (1975) 663. 243 A. Omar, G. Djega-Mariadassou, F. Bozon-Verduraz and G. Pannetier, Bull. SOC. Chim. Fr., (1974) 2740. 244 I.Z. Onsan and D.L. Trimm, J. Catal., 38 (1875) 257. 245 W.J. van Ooij and W.J. Muizebelt, J. Catal., in press. 246 T. Otsubo, H. Miura, Y. Morikawa and T. Shirasaki, J. Catal., 36 (1975) 240. 247 G.S. Pant and M. Chanda, Can. J. Chem. Eng., 54 (1976) 305.. 248 N.S. Parera and D.L. Trimm, J . Catal., 30 (1973) 485. 249 G. Parravano, Catal. Rev., 4 (1970) 53. 250 J.M. Peacock, M.J. Sharp, A.J. Parker, P.G. Ashmore and J.A. Hockey, J:Catal., 1 5 (1969) 379. 251 J.M. Peacock, A.J. Parker, P.G. Ashmore and J.A. Hockey, J. Catal., 1 5 (1969) 387. 252 P. Pendleton and D. Taylor, J. Chem. SOC.,Faraday Trans. I , 72 (1976) 1114. 253 N. Pernicone, G. Liberti and L. Ersini, Roc. 4th Int. Congr. Catal., MOSCOW, 1968, Akademiai Kiad6, Budapest, 1971, p. 287. 254 N. Pernicone, J. Less-Common Met., 36 (1974) 289. 255 H. Pichler, H. Schulz and A.Z. El Deen, Erdoel Kohle Erdgas Petrochem., 29 (1976) 418. 256 J.L. Portefaix, M. Forissier, C. Pralus, F . Figueras and P. d e Montgolfier, C.R. Acad. Sci., Ser. C, 282 (1976) 523. 257 C. Pralus, F. Figueras and L. d e Mourgues, C.R. Acad. Sci., Ser. C, 27 (1970) 20. 258 G. Prauser, H. Kiemer and K. Dialer, Ber. Bunsenges. Phys. Chem., 8 0 (1976) 49. 259 P. Putanov, D. Smiljanec, B. Djukanovic, N. Jovanovic and R . Herak, Proc. 5th Int. Congr. Catal., Miami Beach, 1972, North-Holland, Amsterdam, 1973, P1061. 260 Y.I. Pyatnitskii and G.I. Golodets, React. Kinet. Catal. Lett., 2 (1975) 143. 261 S.K. Ray and M. Chanda, Ind. Eng. Chem. Process Des. Dev., 1 5 (1976) 234. 262 K.A. Reddy and L.K. Doraiswamy, Chem. Eng. Sci., 24 (1969) 1415. 263 R.J. Rennard, R.A. Innes and H.E. Swift, J. Catal., 30 (1973) 128. 264 R.J. Rennard and W.L. Kehl, J. Catal., 21 (1971) 282. 265 R.G. Rizaev, V.E. Sheinin and S.D. Mekhtiev, Kinet. Katal., 1 3 (1972) 1496. 266 D.A. Robb and P. Harriott, J. Catal., 35 (1974) 176. 267 J.Y. Robin, Y. Arnaud, J. Guidot and J.E. Germain, C.R. Acad. Sci. Ser. C, 280 (1975) 921. 268 E.V. Rozhkova, S.V. Gerei and Y.B. Gorokhovatskii, Kinet. Katal., 1 5 (1974) 694. 269 G. Russo, S. Crescitelli and P. Ciambelli, Chim. Ind. (Milan), 55 (1973) 634.
270 271 272 273 274 275 276 277 278 279 280 281 282 283 284 285 286 287 288 289 290 291 292 293 294 295 296 297 298 299 300 301 302 303 304 305 306
W.M.H. Sachtler, Catal. Rev., 4 ( 1 9 7 0 ) 27. T. Sakamoto, N. Egashira and T . Seiyama, J. Catal., 16 ( 1 9 7 0 ) 407. F. Sala and F. Trifiro, J . Catal., 4 1 ( 1 9 7 6 ) 1. F. Sala and F. Trifiro, Z. Phys. Chem. (Frankfurt a m Main), 9 5 ( 1 9 7 5 ) 279. F. Sala and F. Trifiro, Z. Catal., 3 4 ( 1 9 7 4 ) 68. K.M. Sancier, A. Aoshima and H. Wise, J. Catal., 3 4 ( 1 9 7 4 ) 257. K.M. Sancier, P.R. Wentrcek and H. Wise, J. Catal., 39 ( 1 9 7 5 ) 1 4 1 . N. Sato and M. Seo, J. Catal., 2 4 (1972) 224. N.N. Sazonova, S.A. Ven’yaminov and G.K. Boreskov, Kinet Katal., 15 ( 1 9 7 4 ) 419. N.N. Sazonova, S.A. Ven’yaminov and G.K. Boreskov, Kinet. Katal., 1 6 ( 1 9 7 5 ) 1518. N.N. Sazonova, S.A. Ven’yaminov and G.K. Boreskov, Kinet Katal., 11 ( 1 9 7 3 ) 1169. G.C.A. Schuit, J . Less-Common Met., 3 6 ( 1 9 7 4 ) 329. I.W. Schulz and J. Scheve, Proc. Symp. Mechanism Hydrocarbon React. Hungary, 1 9 7 3 , p. 283. T. Seiyama, M. Egashira, T. Sakamoto and I. Aso, J. Catal., 24 ( 1 9 7 1 ) 76. T. Seiyama, T. Uda, I. Mochida and M. Egashira, J . Catal., 3 4 ( 1 9 7 1 ) 29. T. Seiyama, M. Egashira and T. Sakamoto, J. Catal., 2 4 ( 1 9 7 1 ) 7 6 . T. Seiyama, N. Yamazoe and M. Egashira, Proc. 5 t h Int. Congr. Catal., Miami Beach, 1 9 7 2 , North-Holland, Amsterdam, 1 9 7 3 , p. 997. J.L. Seoane, P. Boutry and R. Montarnal, C.R. Acad. Sci., Ser. C., 278 ( 1 9 7 4 ) 565. V.P. Shchukin, S.A. Ven’yaminov and G.K. Boreskov, Kinet. Katal., 1 2 ( 1 9 7 1 ) 621. V.P. Shchukin, G.K. Boreskov, S.A. Ven’yaminov and D.V. Tarasova, Kinet. Katal., 11 ( 1 9 7 0 ) 153. V.P. Shchukin, S.A. Ven’yaminov and G.K. Boreskov, Kinet. Katal., 11 ( 1 9 7 0 ) 1236. N. Shimizu, K. Shimokoshi and I. Yasumori, Bull. Chem. SOC.Jpn., 4 6 ( 1 9 7 3 ) 2929. Y.M. Shulga, I.N. Ivleva, M.V. Shimanskaya, L.Y. Margolis and Y.G. Borodko, Russ. J . Phys. Chem., 4 9 ( 1 9 7 5 ) 2976. G. Simon and J.E. Germain, Bull. Soc. Chim. Fr., ( 1 9 7 5 ) 2617. Th.G.J. Simons, P.N. Houtman and G.C.A. Schuit, J. Catal., 2 3 ( 1 9 7 1 ) 1 . T.G. Skorbilina, Y.I. Pyatnitskii and G.I. Golodets, Kinet. Katal., 1 6 ( 1 9 7 5 ) 1207. V.A. Slavinskaya, D.R. Kreile, D.Y. Eglite and 1.1. Geiman, Kinet. Icatal., 13 ( 1 9 7 2 ) 1 4 8 2 (Engl. trans]. p. 1 3 2 1 ) . A.W. Sleight and W.J. Lonn, Ann. N.Y. Acad. Sci., 272 ( 1 9 7 6 ) 22. A.W. Sleight, K. Aykan and D.B. Rogers, J. Solid State Chem., 13 ( 1 9 7 5 ) 231. F. Solymosi and F. Rozso, Proc. 6th Int. Congr. Catal., The Chemical Society, London, 1 9 7 7 , p. 365. H.T. Spath and K.D. Handel, Adv. Chem. Ser., 1 3 3 ( 1 9 7 4 ) 395. B. Sramek and J. Tichy, Collect. Czech. Chem. Commun., 4 0 ( 1 9 7 5 ) 3500. J.G. Steenhof de Jong, C.H.E. Guffens and H.S. van der Baan, J. Catal., 3 1 (1973) 149. J.G. Steenhof de Jong, C.H.E. Guffens and H.S. van der Baan, J . Catal., 26 (1972) 401. J.G. Steenhof de Jong, Thesis, Eindhoven University of Technology, 1 9 7 2 . G. Stefani, Chim. Ind. (Milan), 5 4 (1972) 984. J.S. Sterrett and H.G. McIlvried, Ind. Eng. Chem. Process Des. Dev., 1 3 ( 1 9 7 1 ) 54.
307 P. Subramanian and M.S. Murthy, Chem. Eng. Sci., 29 (1974) 25. 308 P. Subramanian and M.S. Murthy, Ind. Eng. Chem. Porcess Des. Dev., 1 3 (1974) 112. 309 P. Subramanian and M.S. Murthy, Ind. Eng. Chem. Process Des. Dev., 11 (1972) 242. 310 I.P. Suzdalev, A.A. Firsova, A.U. Aleksandrov, L.Y. Margolis and D.A. Raltrunas, Dokl. Akad. Nauk SSSR, 204 (1972) 408. 311 H.E. Swift, J.E. Bozik and J.A. Ondrey, J. Catal., 2 1 (1971) 212. 312 Y. Takita, A. Ozaki and Y. Moro-oka, J. Catal., 27 (1972) 185. 313 S. Tan, Y. Moro-oka and A. Ozaki, J . Catal., 17 (1970) 132. 314 S. Tanaka and T. Yamashina, J. Catal., 40 (1975) 140. 315 M.I. Temkin and N.V. Kulkova, Kinet. Katal., 16 (1975) 1211. 316 J . Tichy and B. Sramek, Collect. Czech. Chem. Commun., 39 (1974) 2796 317 F. Trifiro, G. Caputo and P.L. Villa, J. Less-Common Met., 36 (1974) 305. 318 F. Trifiro and I. Pasquon, Chim. Ind. (Milan), 5 3 (1971) 577. 319 F. Trifiro, Chim. Ind. (Milan), 5 6 (1974) 835. 320 F. Trifiro, G. Caputo and P. Forzatti, Ind. Eng. Chem. Prod. Res. Dev., 1 4 (1975) 22. 321 F. Trifiro, Chim. Ind. (Milan), 56 (1974) 541. 322 F. Trifiro, L. Kubelkowa and I. Pasquon, J . Catal., 1 9 (1970) 121. 323 F. Trifiro and S. Carra, React. Kinet. Catal. Lett., 2 (1975) 411. 324 F. Trifiro and I. Pasquon, Chim. Ind. (Milan), 52 (1970) 228. 325 F. Trifiro, and C. Lambri and I. Pasquon, Chim. Ind. (Milan), 5 3 (1971) 339. 326 F. Trifiro, Ann. Chim. (Rome), 65 (1975) 113. 327 F. Trifiro, H. Hoser and R.D. Scarle, J . Catal., 25 (1972) 12. 328 D.L. Trimm and L.A. Doerr, J. Catal., 23 (1971) 49. 329 D.L. Trimm and D.S. Gabbay, Trans. Faraday Soc., 67 (1971) 2782. 330 D.L. Trimm and M. Irshad, J. Catal., 18 (1970) 142. 331 A. Vaghi, A. Castellan, J.C.J. Bart, N. Giordano and V. Ragaini, J. Catal., 42 (1976) 381. 332 K. Vaidyanathan and L.K. Doraiswamy, Chem. Eng. Sci., 23 (1968) 537. 333 M. ValdeliGvre, H. Baussart, R. Delobel and J.M. Leroy, Bull. SOC. Chim. Fr., (1975) 2467. 334 D. Vanhove and M. Blanchard, Bull. SOC.Chim. Fr., (1971) 3291. 335 D. Vanhove and M. Blanchard, J . Chim. Phys., Phys. Chim. Biol., 7 3 (1976) 51. 336 D. Vanhove and M. Blanchard, J. Catal., 36 (1975) 6. 337 S. Veleva and F. Trifiro, React. Kinet. Catal. Lett., 4 (1976) 19. 338 S.A. Ven’yaminov, N.N. Sazonova and N.I. Alferova, Kinet. Katal., 1 5 (1974) 1500. 339 A. Verma and S. Kaliaguine, J . Catal., 30 (1973) 430. 340 P.L. Villa, G. Caputo, F. Sala and F. Trifiro, J. Catal., 3 1 (1973) 200. 341 O.M. Vinogradova, G.F. Vytnov, I.V. Luiksaar and O.V. Al’tshuler, Kinet. Katal., 1 6 (1975) 671. 342 B. Viswanathan, V . Ramalingam and M.V.C. Sastri, Indian J. Chem., 1 2 (1974) 205. 343 H.H. Voge and C.R. Adams, Adv. Catal., 1 7 (1967) 151. 344 T. Watanabe and E . Echigoya, Bull. J p n . Pet. Inst., 1 4 (1972) 100. 345 F. Weiss, J. Marion, J. Metzger and J.M. Cognion, Kinet. Katal., 1 4 (1973) 45. 346 S. Weychert and A. Urbaneck, Int. Chem. Eng., 9 (1969) 396. 347 K. van der Wiele, Thesis, Delft University of Technology, 1976. 348 K. van der Wiele and P.J. van den Berg, J. Catal., 39 (1975) 437. 349 W.B. Williamson, D.R. Flentge and J.H. Lonsford, J. Catal., 37 (1975) 258. 350 H. Wise, F’roc. Symp. Mechanism Hydrocarbon React., Hungary, 1973, p. 297. 351 M.W.J. Wolfs and Ph.A. Batist, J . Catal., 32 (1974) 25.
352 353 354 355 356 357 358 359 360 361 362
R.D. Wragg, P.G. Ashmore and J.A. Hockey, J. Catal., 22 (1971) 49. R.D. Wragg, P.G. Ashmore and J.A. Hockey, J. Catal., 28 (1973) 337. R.D. Wragg, P.G. Ashmore and J.A. Hockey, J. Catal., 3 1 (1973) 293. A.A. Yabrov, G.K. Boreskov, A.A. Ivanov and O.M. Il’inich, Kinet. Katal., 1 6 (1975) 1534. S. Yoshida, T. Matsuzaki, T. Kashiwazaki, K. Mori and K. Tarama, Bull. Chem. Soc. Jpn., 47 (1974) 1564. S. Yoshida, T. Matsuzaki, S. Ishida and K. Tarama, Proc. 5th Int. Congr. Catal., Miami Beach, 1972, North-Holland, Amsterdam, 1973, p. 1049. L. Zanderighi, M.P. Faedda and S. Carra, J. Catal., 35 (1974) 427. L. Zanderighi and S. Carra, Chim. Ind. (Milan), 56 (1974) 815. V.M. Zhiznevskii, E.V. Fedevich, D.K. ToIopko, M.V. Nikipanchuk, L.F. Yakubovskaya, I.M. Golub and V.Y. Shipailo, Kinet. Katal., 1 5 (1974) 381. V.M. Zhiznevskii, E.V. Fedevich, D.K. Tolopko and I.M. Sulima, Kinet. Katal., 1 2 (1971) 431. Handbook of Chemistry and Physics, The Chemical Rubber Co. Press, Cleveland, Ohio, 49th edn., 1968.
263 Chapter 3
Heterogeneous Eliminations, Additions and Substitutions L. BERANEK and M. KRAUS
1. General features 1.1CORRESPONDENCE BETWEEN HOMOGENEOUS AND HETEROGENEOUS
REACTIONS
Elimination, addition and substitution reactions over solid catalysts are treated together in this chapter on the basis of some common features of their mechanisms and the acid-base nature of the catalysts. They behave in such an analogous way t o liquid phase reactions, both catalysed and uncatalysed, that electron shifts solely in pairs (heterolytic) have never been seriously doubted and free radical-like (homolytic) mechanisms have been considered only by few authors. The discovery of parallelism between acid-base reactions in solution and over solids helped t o advance the understanding of reaction mechanisms in this branch of heterogeneous catalysis much more than, for example, in catalysis over metals. The theory of organic reactions has been developed mostly with the help of experimental material concerning substitution and elimination in the liquid phase and the accumulated knowledge and proven research methods were utilised in interpretation of transformations over catalysts with acidic and basic properties. The first step in this approach was the recognition [l--31 that the cracking reactions of hydrocarbons over strongly acidic silica-alumina catalysts have patterns similar t o the reactions in the liquid phase catalysed by strong mineral Brplnsted or Lewis acids for which the carbonium ion mechanism has been suggested [4]. It took some time t o adopt a similar view of other heterogeneous elimination and substitution reactions. Most efficient experimental tools have been found in stereochemical studies, correlation of structure effects on rates and measurement of deuterium kinetic isotope effects. The usual kinetic studies were not of much help due t o the complex nature of catalytic reactions and relatively large experimental error. The progress has been made possible also by the studies of surface acid-base properties of the solids and their meaning for catalysis (for a detailed treatment see ref.
5). The analogy between homogeneous and heterogeneous eliminations and substitutions has been pursued further. Joint action of an acidic and a basic site, suggested quite early for the heterogeneous dehydration of alcohols [ 6 ] ,has been gradually accepted as a general mode of operation
264 in acid-base catalysis over solids (e.g. refs. 7-9). No basic difference is now seen between the action of a surface acid-base double-centre and heterolysis of the bonds in an organic molecule caused by an attack of a base (or an acid) assisted by the solvent acting as a conjugated acid (or base) [ 91. Also, the nomenclatures for homogeneous elimination and substitution mechanisms have been adopted for heterogeneous reactions with only a slightly modified meaning. Of course, steric requirements are more restrictive in surface processes than in solution because the surface sites are immobile. On solid acid-base catalysts, beside elimination, addition and substitution, some other reactions also proceed. Of these, especially skeletal isomerisation of hydrocarbons and double bond shift should be mentioned. The latter can influence the product composition in olefin-forming eliminations and thus distort the information on orientation being sought. 1 . 2 NATURE OF THE CATALYSTS
From the point of view of chemical composition, the solid acid-base catalysts are oxides (like alumina, silica, thoria, magnesia), mixed oxides (like silicaalumina, silica-magnesia), crystalline aluminosilicates (zeolites), metal salts and ion exchange resins. The last type differs from the others in the character of reactant transport into the catalyst grain. With organic ion exchangers, which may or may not possess pores in the dry state, the important component of the reactant penetration into the grain is the diffusion through the more or less swollen macromolecular mass; then, in a favourable case, almost all acidic and basic functional groups may serve as active centres. With inorganic solid catalysts, the reactants reach the internal surface of a porous catalyst grain by means of diffusion through the pores; the bulk of the solid is not utilised for catalysis. Therefore, for understanding the ways in which a catalyst influences a reactant, the surface chemistry of the inorganic solids is important. In spite of much effort, the nature of the active sites on acid-base inorganic catalysts is still not completely understood. However, the work on this problem has shown how complicated the surface structure may be and that several types of active centres may be simultaneously present on the surface; the question is then which type plays the major role in a particular reaction. Also, the catalytic activity may be influenced t o a large extent by impurities present in the feed (catalytic poisons) or by-products of the reaction. The last point is often not taken into account and it will be discussed specially in Sect. 1.2.6. First, the models of surface sites on the most important and best-studied catalysts will be described.
1.2.1 Silica The surface of silica (for detailed description of results see refs. 5, 11 and 1 2 ) contains a variable amount of hydroxyl groups and adsorbed wa-
265 ter molecules. Even after heating t o 900°C in vacuum, it retains some OH groups (e.g. ref. 10). The absolute number of hydroxyl groups may differ from sample t o sample according t o the methods of preparation resulting in different participation of various crystal planes on the surface. Extensive research by means of IR spectroscopy and chemical reactions has shown (e.g. refs. 10-18) that two types of surface hydroxyl groups are present: single or free (A) and paired (B). Their relative proportions on partially dehydroxylated surfaces were estimated t o be 1 : 1 [14],1 : 2 [ 181, 1 : 9 [ 171. The nature of the paired hydroxyl groups is still a matter of discussion: vicinal ( B l ) and geminal (B2) structures are both possible. With vicincl OH groups, interhydroxyl hydrogen bonding is assumed H H I I 0 0 \ I
/St
/o\
0 0 0 I\ I \ I Si Si Si Si Si
0 0 0 I \ I\ Si Si Si B2
A
H I 0 I Si \ 0 \
H
H I 0 \
/si
0
I Si
B1
The surface of silica is highly reactive and hydroxyl groups exchange hydrogen for deuterium with DzO [ 14-16] but not with Dz. They can be replaced by C1 from Clz or CC14 [16] and they react with silanes and aluminium chloride [ 15,191. Surface alcoholates are formed when silica is contacted with primary or secondary alcohols [20] either by the reaction with hydroxyl groups \ \ -Si--OH + CH30H = S i - O - C H 3 + H,O l I
or by the rupture of a surface siloxane bond [21]
R f f r rc nccs u p . 38 5-3 98
266
A number of other substances react with surface hydroxyl groups forming surface compounds I221. However, for catalysis, the hydrogen bonding seems to be more important. With alcohols, the hydrogen bonds are formed in such a way that surface hydroxyl groups act as donors of hydrogen [ 231, viz. Si-O-H...O-R
I
H 1.2.2 Alumina
Aluminium oxide exists in many crystalline modifications, usually designated by Greek letters, some with hexagonal and some with cubic lattices (cf. refs. 11 and 24). The best known and mostly used forms are a- and y-alumina but practical catalysts are seldom pure crystailographic specimens. This makes the surface chemistry of aluminas rather complicated. Moreover, the catalytic activity of alumina depends very much on impurities. Small amounts of sodium (0.08-0.65%) poison the active centres for isomerisation but do not affect dehydration of alcohols [ 101. On the other hand, traces of sulphates and silica may increase the number of strong acidic sites and change the activity pattern. The surface hydroxyl groups and adsorbed water are important factors determining the surface properties of alumina (e.g. refs. 5, 11and 24). At present, we have at our disposal a model of the (100) plane which is probably exposed on the surface of spinel-type y-alumina (cf. ref. 24). The model is due to Peri [25] and is based on his detailed investigations by IR spectroscopy [ 261, by gravimetry [ 261, by ammonia adsorption [ 27 1 and by Monte Car10 modelling of surface dehydration [ 251. According t o Peri [ 2 5 ] , the (100) plane of alumina, fully hydrated at low temperatures, exposes a square lattice of OH groups [Fig. l ( a ) ] . If the dehydration were ideal, a regular surface of equally spaced 0 2 -ions would be formed [Fig. l(b)]. However, the splitting off of water molecules is a random process and, consequently, only two-thirds of the original OH groups can be removed without disturbing the original order. Further dehydration is possible only at the expense of some disorder. Ultimately, only isolated hydroxyl groups, which have no partner in the neighbourhood for the formation of water, remain on the surface. Five different types of these isolated hydroxyl groups can be distinguished according t o the number of neighbouring oxygen atoms in the surface layer (Fig. 2); their frequency depends on the degree of dehydration. The hydroxyl groups act as Br@nsted acidic sites and the exposed aluminium atoms in the second layer [Fig. l ( b ) ] as Lewis acidic sites. Rehydration of the surface changes the Lewis into Brqhsted sites. The Peri model of alumina also demonstrates that basic sites of various strength, consisting of oxygen atoms in various arrangements (isolated
267
0000 0000 000
(A)
Fig. 1. Ideal surface (100) plane of alumina after Peri [ 25 1. (A) T o p layer viewed perpendicualrly to the plane; (B) section through the three t o p layers. (a) Fully hydrated surface. ( b ) dehydroxylated surface. Open circles denote oxygen, filled circles hydroxyl, small black points aluminium,
0
8 00 (d)
0
00
o@o (e)
Fig. 2. Schematic representation of five different arrangements of oxygen atoms around the surface hydroxyl groups (filled circles) o n t h e (100) plane of alumina after Peri [25]. References p p . 385-398
268 atoms in the upper layer or two or three oxygen atoms on adjacent sites), are available on the surface. Experimental evidence for the presence of basic sites comes from adsorption of BF3 [ 281, titration with benzoic acid [29] and poisoning of the dehydration of alcohols over alumina by tetracyanoethylene [ 81 and by acetic acid [ 301. Different types of hydroxyl groups and oxygen atoms have different properties and the surface is therefore non-homogeneous. This heterogeneity manifests itself not only in the varying acid and base strengths of the sites but, and this might be more important for catalysis, in the frequency of suitably spaced pairs of acidic and basic sites. Strong evidence from mechanistic studies shows that such pairs are a prerequisite for the concerted elimination mechanism which predominates over alumina. The surface of alumina is highly reactive, not only t o water, ammonia or acetic acid, but also t o a number of other substances. Surface alcoholates are products of the interaction with alcohols [31] and carboxylate surface structures are formed from a fraction of adsorbed alcohol molecules [ 321. The action of hydrofluoric acid [ 33-35], as well as impregnation by BF3 [ 31,341, increases the acidity of alumina. 1.2.3 A lu m inosilica tes Aluminosilicates are the active components of amorphous silicaalumina catalysts and of crystalline, well-defined compounds, called zeolites. Amorphous silica-alumina catalysts and similar mixed oxide preparations have been developed for cracking (see Sect. 2.5) and quite early [36,37] their high acid strength, comparable with that of sulphuric acid, was connected with their catalytic activity. Methods for the determination of the distribution of the acid sites according t o their strength have been found, e.g. by titration with t-butylamine in a non-aqueous medium using adsorbed Hammett indicators for the Ho scale [ 381. The chemistry of silica-alumina catalysts has been reviewed several times (e.g. refs. 39-41) and the nature of acidic active sites has been discussed in numerous papers, very often from the point of view of whether Lewis or BrQlnsted sites are responsible for catalytic activity. The experimental methods for their separate determination are not very conclusive and in the actual catalytic process one type of centre may be converted t o another by the action of reagents, products or impurities. The experiments with various substances added t o the feed indicate (see the following sections dealing with individual reactions) that different types of reaction require sites of different strength. The great variety of Lewis and Brgnsted sites which may exist on the surface of silicaalumina has been demonstrated by Peri [42] on the basis of a simplified model of the reaction of AlCl, with a silica surface and subsequent hydrolysis. Peri has constructed eight different surface sites by combining possible groups on the surface of silica with possible aluminium ion structures; more arrange-
269 ments can probably be thought of. Some of the Peri sites are
0 Al’ I \ 0 0
‘si’
‘ A1 I \
0 0 \ I Si
do‘ lAf
I \ 0 0 I di si \ I 0
0 I
0 I si s i \ I 0
The original view, that in the reaction of silica with aluminium hydroxides a strong aluminosilicic acid, which possesses a dissociable proton (e.g. ref. 2), is formed has not been proved. H-aluminosilicates are unstable and spontaneously convert t o aluminium aluminosilicates [ 191. Crystalline aluminosilicates (zeolites, molecular sieves) catalyse a number of organic reactions [43] and the striking difference between them and amorphous silicaalumina is that they are active for cracking even in the form of Na’ of CaZ+salts [ 44,451 ; these cations are poisons for silicaalumina. However, metal salts of zeolites exhibit strong acidity [ 51. This acidity is of both the Lewis and Brq5nsted type and strong Lewis sites are converted t o Brgnsted by water [46]. The catalytic activity of zeolites depends on the nature of the cation but it seems (cf. refs. 47 and 48) that the active centres are not metal cations or hydroxyl groups attached t o such ions. As with amorphous silica-alumina, the active centres in zeolites are probably situated on the aluminosilicate surface. The function of metal cations is not clear; they might stabilise the structure and influence the degree of hydroxylation and hydration of the surface which are important factors for catalysis. In the section dealing with alumina and silica, the necessity of basic sites on the surface, which cooperate with acidic sites, has been stressed. Also, for both amorphous silica-alumina and zeolites, the simultaneous presence of acidic and basic sites has been proved and it has been suggested that OH groups act as amphiprotic centres according t o the nature of the adsorbed species [ 491. 1.2.4 Metal salts
Solid metal sulphates and phosphates also exhibit acid-base properties; their acid strength is lower than that of silicaalumina but they are stronger acids than some oxide catalysts [ 51. Correlation of activity with electronegativity of cations has been obtained for several reactions [ 9, 50,511.
270
1.2.5 Ion exchange resins Organic ion exchangers are macromolecular substances containing chemically welldefined acidic or basic functional groups, The macromolecular skeleton may be formed by polycondensation or, more frequently, by copolymerisation. The use of basic (anion) exchangers as catalysts (e.g. in aldol condensation) is rather rare; the main representatives of acidic sites in cation exchangers are sulphonic (-SO,H), phosphonic (-PO(OH),) and carboxylic (-COOH) groups. In the kinetic studies reported in this chapter, sulphonated styrenedivinylbenzene copolymers were used almost exclusively. They may be of two types: (i) non-porous (standard) ion exchangers whose grains do not possess internal porosity in the sense usual in catalysis, and (ii) porous (macroreticular) ion exchangers with artificially developed porous structure (pores of about 10-20 nm prevailing) and a large inner surface area. Ion exchangers can be used as catalysts both for liquid (standard ion exchangers are preferred) and vapour phase (macroreticular ion exchangers are more convenient) reactions. The main factors determining the catalytic activity of ion exchangers are: (i) the acid strength of the functional groups (sulphonated resins are much more active than the others), (ii) the concentration of functional groups in the protonated form (ion exchangers fully neutralised with cations are catalytically inactive) and (iii) the degree of crosslinking of the copolymer, i.e. the content of divinylbenznee (DVB). There is no doubt that the functional groups in ion exchangers are responsible for the catalytic activity. Although they are chemically defined, it is not clear in which form they participate in the catalytic reaction, since a certain amount of water is always present in the resin which cannot be removed easily. It has been proven by IR spectroscopy [52541 that in polystyrenesulphonic acid several hydrated forms of the --S03H groups may occur (mono-, di-, tri- and tetra-hydrates) and, in consequence, the mobility of the proton of a -S03H group and also the catalytic activity may change. Lower hydration states are also possible through hydrogen bonding of one water molecule to two or more sulphonic acid groups. Even fully dehydrated sulphonic groups may be hydrogen bridged, for example in the form
P-H***o o=\\
-S=O
ii 0.e.H-O
-
\\
The lack of information about relative activities of different forms and the unknown dependence of their relative concentrations on catalyst pretreatment and reaction conditions, and the influence of reactants, products (water) and solvents, introduce uncertainty into the interpretation of kinetic measurements.
271
It seems probable, in view of the idea presented in Sect. 1.1,that, in elimination, addition and substitution reactions over ion exchangers, also, two types of catalytic sites are involved, viz. acidic (protons of the functional groups) and basic, which are likely t o be represented by oxygen atoms of the functional groups. A typical property of ion exchange resins which distinguish them from inorganic catalysts is swelling; this is the more important factor the lower is the degree of crosslinking of the copolymer. Due t o swelling, a considerable amount of reactants, products and solvents can be retained (absorbed) by the resin and the functional groups inside the polymer mass may also be utilised for catalysis. Thus, the accessibility of the catalytically active groups can be facilitated, not only by a artificial porous structure (which increases only the number of the groups on the surface of the polymer mass), but also by swelling. In this situation, the rate of reactant transport (diffusion), not only through the pores (if their are present), but also through the more or less swollen polymer mass, may become important. If the rate of diffusion through the polymer is much larger than that of the chemical reaction, then, in the extreme case, all functional groups may be utilised for catalysing the reaction. In the opposite case, when the diffusion through the polymer mass is much slower than the chemical reaction, only the surface groups will act as catalytic sites. This latter was observed with highly crosslinked ion exchangers and large reactant molecules and the term “sieve effect” was used t o describe it. 1.2.6 The working surface
The surface structures outlined in the preceding sections have been determined under conditions very far from those of an actual catalytic reaction. At partial pressures of reactants used in flow reactors and in the steady state, the catalyst surface is very probably almost covered by starting substances and products. This is indicated by the type of kinetics found for various reactions (see following section); very often zero-order expressions or Langmuir-Hinshelwood type rate equations with high values of adsorption coefficients have been found. Some products of the catalytic reactions are of special interest in this connection. Water formed in elimination, esterification or condensation reactions is present in sufficient quantities to change almost all Lewis sites into Br@nsted sites. Much more fundamental changes can be caused by hydrogen halides produced in the decomposition of alkyl halides on oxides; it is well known that the catalytic activity of alumina can be enhanced by the action of hydrochloric or hydrofluoric acids. It is evident that the study of free surfaces and of surfaces covered only partially by various substances at temperatures much lower than those needed for a catalytic reaction to proceed can give only indirect inforReferences P P . 385-398
272 mation about possible states on working surfaces. Better evidence is obtained by observing the influence of substances added t o the feed which can interact with some surface sites a t reaction conditions. For example, in this way the importance of basic sites has been confirmed. Linear correlations of effects of reactant structure on rate and adsorptivity are also helpful and especially the interpretation of their slopes may yield valuable information (e.g. refs. 55 and 56). The transient-response technique, in which the changes in product composition after an abrupt stop or start of the feed flow are observed, is also promising. 1.3 TYPE OF KINETICS
The complex nature of heterogeneous catalytic reactions, which consist of a sequence of at least three steps (adsorption, surface reaction and desorption), the possible intervention of transport processes and the uncertainty about the actual state of the surface makes every attempt t o obtain a complete formal kinetic description without simplifying assumptions futile. In this situation, some authors prefer fully empirical equations of the type
r
=
kpi&
...
(1)
which bear no connection to the mechanism. With the exception of zero-, first- and second-order expressions, the interpretation of the constants h, a, b, ..., cannot be used as a basis for the elucidation of the laws governing catalytic reactions. However, simple kinetic models, especially of the Langmuir-Hinshelwood type, can serve with advantage for correlation of experimental data in spite of simplifying assumptions which are necessary for their derivation. Experience shows that heterogeneous acid-base catalysis is the very field where they fit best. Their most frequent general form
where Ki denotes the adsorption coefficient of the substance i, a, b, ..., = 1 o r $ and n = 1, 2, 3, .,., is well suited to the estimation of the competition of all substances present in the system for active centres. However, because the same equation may be obtained on the basis of various different assumptions (cf. ref. 57), its form cannot be used as a proof of a certain mechanism. Of the assumptions accepted for the generation of LangmuirHinshelwood type and related equations, the most controversial seems t o be that the surface is homogeneous. It has been shown in the preceding section that inorganic oxide catalysts and even ion exchangers contain a number of differing acidic and basic sites, i.e. they possess an inherent heterogeneity. The question is how this “static” non-homogeneity manifests itself
273 under the dynamic conditions of a catalytic reaction. Some sites may be ineffectual for steric reasons when they d o not find a basic (or acidic) partner site within a suitable distance. Out of the residual spectrum of sites differing in strength, some are probably too weak t o be able t o initiate bond reorganisation in adsorbed molecules. Other sites can bind the reactants or products too strongly and thus be blocked out. Working sites come, therefore, from a band which is narrower then the original one estimated on the basis of adsorption measurements (including determination of the number of acidic sites by titration with a base etc.). The position and width of this working band must depend on the chemical nature of the reagent (e.g. cracking of alkanes requires other sites than dehydration of alcohols) and on the form of the distribution curve of sites according t o their strength. Some experimental results are available which show the influence of surface heterogeneity on the kinetics and the contribution of sites of different strength t o the over-all rate. The surface of acidic catalysts has been divided into several fractions by acidimetric [ 581 or thermochemical [ 591 titrations and on the basis of group analysis [59]and partial poisoning [58,60] the contribution of these fractions has been calculated. It has been found that the over-all rate of dehydration is determined by the performance of a single narrow fraction, the contribution of the others being almost negligible. Another approach t o this problem involved modelling of acidic catalysts with different sites by mixing ion exchangers containing functional groups of different acidity [ 611. For dehydration; the over-all activity was again given by the activity of the strongest (-S03H) group. For re-esterification, the contribution of weaker centres (-PO(OH),) could not be neglected but the over-all kinetics could still be correlated by a single Langmuir-Hinshelwood rate equation. Summarising, it seems that the surface heterogeneity is not such a serious problem for the formal kinetic description of acid-base catalysis on solids as would be expected from the studies of the surface by non-kinetic methods. Moreover, the rate equations for non-homogeneous surfaces, developed by the Russian school (Temkin, Roginskii and Kiperman, see ref. 62) are similar t o eqn. (2); the term 1 is not present and n can have any value greater than 0 (cf. also ref. 63). Only their further drastic simplification leads to equations of type (1). The next problem of the LangmuirHinshelwood kinetics, the validity of the ratedetermining step approximation, has not been rigourously examined. However, as has been shown (e.g. refs. 57 and 63), the mathematical forms of the rate equations for the LangmuirHinshelwood model and for the steady-state models are very similar and sometimes indistinguishable. This makes the meaning of the constants in the denominators of the rate equations somewhat doubtful; in the Langmuir-Hinshelwood model, they stand for adsorption equilibrium constants and in the steady-state models, for rate coefficients or products and quotients of several rate coefficients. References P P . 385-398
274 The problem discussed in Sect. 1.2.6, i.e. what composition the working surface has, also has its kinetic counterpart. If the number of active sites of a certain type depends on the partial pressures of some reaction components, then the question arises whether rate equations of type (2) are sufficient for the description of such changes. All these facts and unsolved problems require that the rate equations of type (2) be taken as semi-empirical expressions. They may be directly utilised for engineering purposes with higher certainty than eqn. (l), but they reflect the actual react.ion mechanism only in general features. However, the constants are a good source of values for comparison of reactivities and adsorptivities of related reactants on the same catalyst. Such interpretations of experimental data are usually quite meaningful as is confirmed by successful correlations of the constants with other independent quantities. 2. Elimination reactions In organic chemistry, elimination processes are those decompositions of molecules whereby two fragments are split off and the multiplicity of the bonds between two carbon atoms or a carbon atom and a hetero atom is increased. Such a broad definition also embraces the dehydrogenation of hydrocarbons and alcohols which is dealt with in Chap. 2. Here we shall restrict our review t o the olefin-forming eliminations of the t Y Pe I I I I --(+?-C,=Cc =
X
+ HX
H
Although some observations (e.g. ref. 7) indicate that the process need not (Y, 0-(or 1 , 2-) elimination, practically all experimental results have been interpreted on the assumption that 1,3- and 1,4-eliminations d o not participate significantly. The substituents X may have very different structures but heterogeneous catalytic eliminations with X = halogen, OH, alkoxyl, NR2 (R = H or alkyl), SH, OCOCH3 and alkyl or aryl only have been described. The individual reactions are usually named according to the compound HX which is the product, i.e. dehydrohalogenation, dehydration etc. but some exceptions exist (e.g. cracking). The reverse reactions are additions t o the C--C multiple bonds which will be dealt with in Sect. 3 of this chapter. Homogeneous olefin-forming eliminations have been studied extensively, especially in the liquid phase and comprehensive treatments of the subject are available [ 64,651. The rules governing the course of homogeneous eliminations and their mechanisms are well established and the interpretation of the results obtained with heterogeneous catalytic sys-
to be always an
275 tems can obtain useful assistance from these. In this connection, a recent review on catalytic eliminations is especially valuable [9]. 2.1 COMMON FEATURES O F HETEROGENEOUS CATALYTIC ELIMINATIONS
2.1.1 Mechanism
In discussing the mechanism of eliminations over solids, the nomenclature which has been developed for homogeneous reactions will be used. Therefore the basic mechanisms of olefin formation have first to be outlined and their meaning in heterogeneous catalysis defined. The E2 mechanism is so called because the process is bimolecular and in solution consists of an attack by a base on the P-hydrogen atom with synchronous splitting of the substituent X in the form of an anion. In heterogeneous catalysis, the most important feature is the timing of the fission of the two bonds C,-X and CB-H: in the E2 or E2-like mechanism, these bonds are broken simultaneously. Because this can be achieved only by the action of two different centres, a basic one and an acidic one with both present on the sudace, the kinetic distinction of the mechanism loses its original sense under these circumstances. The E l mechanism has, as the ratedetermining step in solution, the ionisation of the reactant forming a carbonium ion which then decomposes rapidly. For heterogeneous catalytic reactions, the important features are the occurrence of the reaction in two steps and the presence on the solid surface of carbonium ions or species resembling them closely. Again, the kinetic characterisation by way of an unimolecular process is of little value. Even the relative rates of the two steps may be reversed on solid catalysts. A cooperation of an acidic and a basic site is also assumed, the reaction being initiated by the action of the acidic site on the group
X. The ElcB mechanism is a two-step process beginning with the abstraction of a proton from the P-position by a base to give a carbanion. The second step is the loss of the group X as an anion. In heterogeneous catalysis, the corresponding mechanism consist of the primary action of a basic site assisted later by an acidic site which temporarily accomodates the group X-. It is evident that the simple model of heterogeneous catalytic eliminations assumes the same adsorption complex for all mechanisms, written schematically as
-c-c-
I I X. H.
. . 00 The only distinction between various mechanisms is the timing of the References p p . 385-398
276 fission of the bonds C,-X and C,-H. Usually, a continuous spectrum of mechanisms is assumed in which E l , E2 and ElcB are processes with a clearly defined character. This idea, which has been slowly developed during the last decade, has been discussed in detail in a recent review [ 91. It is certainly compatible with views on the nature of elimination catalysts. These solids are typically oxides o r metal salts which have positively and negatively charged atoms on their surface. In the array of electron-donating and electron-accepting centres, pairs of required acidic and basic sites with suitable spacings can be found. Because the strength of the sites is different on individual catalysts according t o their structure, the catalysts can be put into a sequence, from those where the basic character predominates through those where basic and acidic properties are in balance to those with prevailing acidic nature. It is d e a r that a catalyst wifl transform a reactant by means of the mechanism which corresponds t o the predominating acid o r base strength of the sites. It is well known from homogeneous reactions that the mechanism depends also on the strength of the C,-X and C,-H bonds and this applies also t o heterogeneous catalysis. The double influence on the “choice” of mechanism, i.e. of the nature of the catalyst and of the reagent, has been graphically represented by Mochida et al. [66] (Fig. 3). They have
-
-
Cd-H Bond strength Fig. 3 . Schematic representation of the influence of reactant structure, of catalyst nature and of temperature on the elimination mechanism. Numbers in parentheses denote the rate-determining steps on Scheme 1 .
Cp- H
277
*
xI
I
H .
A h <=CX-H' I I A B
. . X. H. . .
-
1
x.
:
H' I
A B
*=C-
X- H' I 1 A B
*=CX- H' I 1 A B
Scheme 1. A = acidic site, B = basic site.
inserted into the mechanism scale two additional fixed points, E2cA and E2cB, and attributed to every mechanism a single ratedetermining step. Scheme 1 and Fig. 3 are modified versions of their original representations; as will be shown later in sections dealing with individual elimination reactions, their hypothesis can accomodate various experimental facts very well. This somewhat simplified picture of possible transitions from one mechanism to another can be expanded and supplemented by a finer differentiation of the factors influencing bond strength and catalyst acidbase properties. Such structural parameters are the number and nature of substituents on C, and C, and the nature of the group X. The action of a catalyst depends on its cation charge and radius, on anion basicity and on lattice and surface arrangement (for some details see ref. 67). A temperature increase usually shifts the mechanism in the direction of E l . 2.1.2 Orientation
Most heterogeneous catalytic eliminations proceed according to the Saytzeff rule, i.e. the most stable olefin is formed. This is in agreement with the prevailing situation in E2 and E l homogeneous eliminations and the reasons might be the same (cf. refs. and 64 and 65). The transition state is probably considerably in the direction of the double bond formation, and hyperconjugation with the groups on C, and C, helps t o stabilise the species with more alkyl substituents on the double bond. The notable exception is the dehydration of alcohols on thorium oxide which is governed by the Hofmann rule. However, in this case the ElcB mechanism has been shown to occur [ S S ] . References P P . 385--398
278 More complex is the problem of syn and anti eliminations (cis and trans in the older nomenclature), i.e. whether the substituents on C, and C, leave the parent compound from synperiplanar or from antiperiplanar positions, viz. H
XH
X
SYn anti The mode of the elimination can be recognised by the composition of the products if the olefin formed has such substituents on the double bond that cis and trans stereoisomers may be distinguished. The question is of interest with respect to the concerted E2 mechanism, because in the pure E l and ElcB processes, the intermediate carbonium ion or carbanion, respectively, usually have enough time to rotate around the C,-C, bond, equilibrate and give the same cidtrans ratio from different conformers. The factors influencing the syn/anti elimination ratio were extensively studied with various catalysts, using pairs of compounds with different steric arrangements like meso and df forms of 2,3-dihalobutanes, threo and ery thro forms of 2deutero-3-X-butanes or cyclic stereoisomers. The syn/anti ratio depended on the nature of the catalyst and was low where a more or less concerted mechanism could be assumed (i.e. on not too strongly acidic catalysts) but, in general, a preference for anti-elimination was observed (cf. refs. 9 and 69): Whereas the syn-elimination proceeding on the surface of a solid catalyst can be easily visualised, there were considerable difficulties in explaining the anti-elimination where the leaving components of HX are on the opposite sides of the C,--C, bond. This problem is not encountered in liquid phase eliminations since the leaving groups can be temporarily bound by other reactant molecules or by the solvent. The first explanation of the anti mode of heterogeneous elimination was that it takes place in narrow pores or crevices, the fragments X and H being bound to opposite walls of the openings [ 71. However, this seems improbable since pores or crevices with a suitable distance between the walls must be rare and therefore the reaction rates would be rather low. Noller and Kladnig [ 91 considered the possibility of hydrogen tunelling through the electron cloud of the other atoms of the reactant; however, this hypothesis lacks support for which it would be necessary to do quantum chemical calculations. The proton can be taken and transported t o the surface by means of another molecule of the reactant [70] or by the product HX [9]. The feasibility of such assistance has been confirmed by quantum chemical cal-
279 culations for the dehydration of 2-propanol [ 701. Knozinger e t al. [71] have suggested a model of anti-elimination on the surface of solids which has been considered further by Noller and Kladnig [9] and by SedlaEek [72]. It explains, without further assumptions, the anti-elimination over heterogeneous catalysts as a natural reaction course. On most oxides, water or its components (H' and OH-) are firmly bound to the surface under the conditions used for elimination. It can be shown by suitable atomic models or drawings that a molecule with X and H in antiperiplanar positions can easily find a pair of sites, an acidic site (formed, for example, by a surface hydroxyl group) and a basic site (formed by a oxygen atom), to which it fits well without notable deformation of bond angles and interatomic distances, and thus gives rise t o the anti-eliminaaxis must not lie parallel or tion. The only condition is that the C,-C, perpendicular to the surface plane. Figure 4 shows the model of such adsorption of isopropanol on an alumina surface. Also, for syn-elimination, suitably spaced acidic and basic sites may be found on the surface of alumina [ 721. The question then arises, what is the reason for the preference for anti-elimination when conditions exist for both modes? This has been explained on the basis of quantum chemical calculations [73]. An
Fig. 4. Adsorption complex of 2-propanol o n a partially dehydroxylated (100) plane of alumina after SedlaEek [72]. Small open circles denote H, medium circles with thin hatching 0, medium circles with dense hatching C, large open circles Al. F o r consistency, covalent atomic radii have been used f o r 0 in alumina although its structure is partly ionic ; therefore aluminium atoms appear unusually large. The diagram shows the possibility of a two-point interaction of a n alcohol with a surface hydroxyl group and a surface oxygen without any distortion of t h e molecule. References p p . 385-398
280
attack by a positively charged species on X o r by a negatively charged one on H changes the distribution of electrons in the parent molecule and its bond lengths. The C,--CB bond is strengthened and the C,-X and CD-H bonds are weakened. The rotation of substituents around the C a - C o axis brings about two minima of total energy, a smaller one for the synperiplanar conformation of X and €3 (rotation angle 0") and a larger one for the antiperiplana; conformation (rotation angle 180").The unti-elimination is therefore energetically more favoured than the syn mode and the syn/ anti ratio must depend mostly on the difference in the energy of the two conformations of the particular substance. Of course, the relative frequency of suitable pairs of sites for anti and syn eliminations must also play a certain role, as the variation of syn/unti ratios from catalyst t o catalyst show.
2.1.3 Kinetics For elimination reactions of the general form
A=R+S the simplest Langmuir-Hinshelwood type rate equation, assuming a surface reaction on a single centre as the ratedetermining step
has been found suitable in a number of cases (see sections on individual elimination reactions). However, more often it has been applied in simplified forms. The first one is valid for the case where the adsorption of reactant A is very strong and, in consequence, the surface is almost entirely covered by it all the time, i.e. KAPA >> 1 + K R p R + &pS. This gives the zero-order rate equation
r=k (4) The second simplified form corresponds t o the case where A, R and S are all weakly adsorbed, i.e. 1>> KApA + KRpR + K s p s . Then eqn. ( 3 ) is reduced t o a first-order rate expression
r
= kKAp, = k'p,
(5)
The complications begin when one of the products (the HX molecules particularly) can influence the reaction rate, not only by adsorption on active sites blocking a fraction of them but by forming new active centres of a different nature. Then the parameter k is no longer a constant, but changes with the composition of the reaction mixture. This possibility has received only limited attention until now, but could explain some unusual empirical rate equations which have been found for the dehydration of alcohols on oxide catalysts [8,69].As has been outlined in
281 Sect. 1.2, the surface of an oxide can change its structure quite easily. For example, cracking catalysts exchange oxygen with water almost instantaneously at higher temperatures [74,75]; the interaction of water with such catalysts must therefore be a complex process, involving not only an adsorption in molecular form. In the literature on elimination reactions, it is found that mechanistic conclusions are quite frequently made on the basis of the values of the activation energy. This is a dubious practice, especially when the data have been obtained by the gas chromatographic (pulse-flow) technique, i.e. when there is a non-stationary state on the catalyst surface, and on the basis of supposed first-order kinetics. True activation energies are obtained when the reaction order is zero and probably also when the rate coefficient, k, and adsorption coefficient, K A , have been separated by treatment of rate data by means of eqn. (3). In the case of the first-order rate equation, the apparent activation energy, calculated from k ’ values [eqn. (5)] by means of the Arrhenius equation, is the difference between the true activation energy and the adsorption enthalpy of the reactant A
(6) Therefore there is little emphasis in this chapter on the values of the activation energy. = Etrut- - AHA
Eapp.
2.2 DEHYDRATION
2.2.1 Types of dehydration reactions In general, dehydration means loss of water molecules from chemical substances, irrespective of their structure. Even if all cases where water is bonded in hydrate form are excluded, a number of reactions remain which also include formation of nitriles from amides, lactones from hydroxy acids etc. However, the present treatment will concentrate on the heterogeneous catalytic decomposition of alcohols in the vapour phase, which can be either olefin-forming or ether-forming reactions, and on the related dehydration of ethers t o olefins. The dehydration of alcohols on solid catalysts is one of the first catalytic reactions discovered and has been studied intensively for many decades. During years of experimental work, a general parallel-consecutive reaction scheme (Scheme 2 below) has been developed by gradual addition of /ether
2 alcohols
[H*O\ olefin + alcohol + H,O
\2 olefins + 2 HPY2@0 Scheme 2. Refereiices p p . 385-398
further steps t o the original parallel scheme [ 76-84]. The relative participation of the various steps depends on a number of factors. Prominent is the structure of the starting alcohol. Alcohols which have no @hydrogens, like methanol or benzyl alcohol, yield only ethers. The tendency t o give an olefin increases with the substitution on C, ; with secondary alcohols, the olefin/ether ratio is much higher than with primary and from tertiary alcohols only olefins are formed. Also substituents on C, may influence the ratio. The next important factor is the temperature, on increasing which more olefin is formed than ether. Although all steps are reversible, the thermodynamic calculations for ethanol have shown that, over 2OO0C, the equilibria are shifted t o the production of ethylene [ 831. Other factors influencing the olefin/ether ratio are the partial pressure of the starting alcohol because olefin-forming and ether-forming reactions obey different kinetics (see Sect. 2.2.3) and the nature of the catalyst. The dehydration of alcohols over solids has been the subject of several excellent reviews which summarise most of the vast literature [ 7,69,85871. Therefore in this chapter, reference will be made only to the papers which are most significant, those that are newer or which have not obtained adequate attention in preceding reviews.
2.2.2 Catalysts Hundreds of substances of many types have been tested as dehydration catalysts and found active. Lists can be found in the literature [69,76,85] and we need t o name here only such catalysts which show high activity and selectivity. The latter parameter is more important because a number of solids, especially oxides, can catalyse both the dehydration and the dehydrogenation of alcohols. The formation of aldehydes or ketones is then a parallel reaction t o the dehydration, and the ratio of the rates depends on the nature of the catalyst. Only few oxides are clean dehydration or dehydrogeneration catalysts, but the selectivity may be shifted t o some extent in either direction by the method of catalyst preparation. The important groups of dehydration catalysts are oxides, aluminosilicates (both amorphous and zeolitic), metal salts and cation exchange resins. Most work on mechanisms has been done with alumina.
2.2.3 Experimental kinetic results
( a ) Formal rate equations The complicated reaction scheme for the dehydration of alcohols (Scheme 2) makes kinetic analysis rather difficult. However, initial reaction rates have been measured, without special problems, for secondary
and tertiary alcohols, and even for primary alcohols. Low conversions are also desirable because water is adsorbed more strongly on alumina than are the alcohols [ 881 and modifies its surface. Initial reaction rates can be obtained separately for olefin and ether formation. Complete kinetic descriptions of a somewhat simplified Scheme 2 have been attempted several times [ 82,831 for the reaction of ethanol on the basis of data from integral o r differential flow reactors. Tables 1 and 2 summarise published results on the kinetics of alcohol and ether dehydration. The data are organised according t o the rate equation found or assumed t o be the best one. Table l shows that, for olefin TABLE 1 Rate equations for t h e catalytic dehydration of alcohols t o olefins Alcohol Catalyst Temperature (“C)
Ref.
r=k 1-Propanol, 2-propanol, 1-butanol, 2-butanol, tert-butanol Ethanol, 1-propanol, 2-propanol, 1-butanol, 2-butanol, 2-methyl-l-propanol, tert-butanol 4-Heptanol, 2-methyl-3-hexano1, 2,4-dimethyl-3-pentanol,cyclopentanol, cyclohexanol, cycloheptanol, cyclooctanol, cisand trans-2-methylcyclohexanol, cis- and trans-4-tert-butylcyclohexanol 2-Methy1-3-butano1, 2-methyl3-hexanol, 2,4-dimethyl-3-pentanol, 2-methyl-4-ethyl-3-hexano1 1-Phenylethanols (p-CH3, p-F, H, rn-OCH3, rn-F) Tert-butanol 3-Deutero-2-butanol,cisand trans-2-methylcyclohexanol Ethanol, 1-propanol, 2-propanol, 1-butanol, 2-butanol, 2-methyl-l-propanol, tert-butanol, 1-pentanol, 2-pentanol, 2-methyl-l-butanol, 3-methyl2-butanol, 2,2-dimethyl-l-propanol, 2-methyl-2-butanol Refererices p p . 385-398
Bauxite
100-250
89,90
A1203
300-4 00
88, 91-93
A1203
180-210
94
A1203 + NaOH Ti02 Si02 Zr02 A1203 + NaOH
300
55
220
95
50 2 10-350
8 96
28 2-39 5
56
Ti02
SiO? Zr02 A12 03-Si02 A12 0 3
hydroxyapatite O3
hydroxyapatite
TABLE 1 (continued) Alcohol
r
=
Catalyst
Temperature ("C)
Ref.
A12 O3
350,400 380,400 2 5 0-3 00 346-465 274-3 1 4 425 121-177 232-343 204-260 204-371
79 a 98,99 100 101 82 98 102 102 103 104
200 300-360 170-195
105 106 107
108 109 110 111 112
113
kKpgA/(l+ CKipi) i
Ethanol
Tert-butanol 1-Hexanol, 1-heptanol 2-Butanol Ethanol, 1-propanol, l - b u tan01 1-Phenylethanols 2-Phenylethanol 1-Propanol
AI203-Si02 A1203-Si02 Alz03-Si02 Alz03-Si02 Al203-Si02 A12 0 3
ZnO BP04
r = k K o A / (1 + C K i p i ) 2 i
Tert-butanol
Ion exchanger Ion exchanger
346-4 6 5 2 50-350 210-240 90-110 95
Tert-butanol
A1203
130-180
Ethanol CyclohexanoI 2-Propanol
a
A12 0 A120
3 3
A12 O3
Without 1 in the denominator of the rate equation.
pw denotes partial pressure of water. In its absence, the equation is reduced t o the zero-order expression.
formation, the zero-order rate equation and its parent (see Sect. 2.1) single-site Langmuir-Hinshelwood type expression predominate. The few cases where the Langmuir-Hinshelwood rate equation with the second power in the denominator has been found are rather puzzling; only for the reaction on an ion exchanger has some explanation been suggested [ 1121. The smaller number of results in Table 2 concerning ether formation are less consistent. The rate equations with the square-root of the partial pressure of the alcohol are either empirical expressions without an underlying kinetic model or are based on a very complicated assumption about the mechanism.
285 TABLE 2 Rate equations for the catalytic dehydration of alcohols to ethers Alcohol
Methanol Ethanol 2-Propanol
Catalyst
Temperature (“C)
Ref.
Al2 O3
2 50-3 00
100
Ion exchanger A12 O3 Ion exchanger
Ion exchanger
119 274-314 80-120 210-240 90-110
131 82 113 110 111
Ion exchanger
111-150
139
A1203
152-279 150-195 130-240
115 116 117
AIZ O3
r = kKApA/[l + ( K f l ~ ) ” ~ ] ~ Methanol
r = k p i 5 / @ i 5+ bpw) a Methanol Ethanol 1-Propanol, 2-methyl-1-propanol, cyclohexanol, phenylmethanol
r = k p y / ( l + ap:’
+
Methanol a
0 3 0 3
bpw) a Alz03--Si02
160
8
p w denotes partial pressure of water.
For the reaction of diethylether giving either ethylene and ethanol or ethylene and water, the validity of the Langmuir-Hinshelwood type rate equations has again been confirmed [ 821. Different approaches to the kinetics of alcohol dehydration were attempted by two groups of authors [ 118,1191. In one case, it has been assumed that the active surface of alumina is formed either by free hydroxyl groups or by surface alkoxyl groups. The rate equation was then derived on the basis of the steady-state assumption; a good fit t o the experimental data was obtained [1118]. The second model was based on the fact that water influences the adsorption of an alcohol and diminishes the available surface. The surface concentrations of tert-butanol and water were taken from independent adsorption measurements and put into the first-order rate equation; a good description of integral conversion data was achieved [ 1191. References p p . 385-398
( b ) Structure and reactivity
Kinetic deuterium isotope effects have been measured with alcohols deuterated in various positions and have yielded information about ratedetermining steps. As expected, deuterium placed on C, does not influence the rate; this has been shown by comparing the dehydration of 2-butanol and 2deutero-2-butanol on some phosphate [ 1201 and sulphate catalysts [121]. In contrast, deuterium atoms on C, change the rate in comparison with non-deuterated compounds [ 96,122,1231. The value of the kinetic isotope effect depends on the catalyst [123] and on the temperature [122], as both factors affect the mechanism (cf. Sect. 2.1.1). Also, deuterium in the hydroxyl groups can influence the rate to various extents according t o the nature of the catalyst. Table 3 shows how the kinetic isotope effects in the dehydration of differently deuterated 2-propanols change with the catalyst, indicating a change of mechanism. The secondary isotope effect was not observed in catalytic dehydration [96, 120,1211. The dehydration rate depends very strongly on substitution on C , . Large differences in reactivity of primary, secondary and tertiary alcohols over solid catalysts were reported as early as in 1931 by Dohse [go]. Also, substituents on C, affect the rate. Both influences can be quantitatively described by the Hammett and Taft relationships; the published correlations are summarised in Table 4. Of special interest is the extensive set of alcohols of the type R'R2R3COH [56], which includes primary, secondary and tertiary alcohols and gives a single Taft correlation with an excellent fit. The values of p and p * which can give information about the mechanism and catalyst nature will be discussed in the following sections. The structure of the reactant also influences the orientation, i.e. the ratio of 1-t o 2-alkenes in the dehydration of 2-alkanols and the ratio of cis t o trans alkenes. Table 4 shows that these ratios can also be correlated by the Taft equation. For the cis/trans ratio, a better fit was obtained with steric E , constants of substituents than with polar constants [ 1271.
TABLE 3 Kinetic isotope effects of deuterium ( h H / h D ) in the dehydration of 2-propanol o n various catalysts at 300°C [ 1 2 3 ] Catalyst
~ ( C H .JCHOH . CH3) h(CH3 . CHOD . CH3)
h(CH3 . CHOH . CH3) k ( C D 3 . CHOH . CD3)
A1203 + NaOH Zr02
1 .o 1.1 1.2 1.5
1.5 1.4 1.2 1.0
Ti02
SiOz
287 The steric requirements of the surface during the formation of the adsorption complex or transition state also manifest themselves in the dehydration of rigid alcohols with fixed conformations, e.g. of cyclic alcohols. Cis- and trans-2- and 4-alkylcyclohexanoIs differ markedly in their rate of dehydration on alumina (see Table 5). Most significant are the data on 4-tert-butylcyclohexanols where the bulky tert-butyl group is in an equat6rial position, and thus the differences in the reactivity of the cis and trans isomers indicate the differences in the reactivity of axial and equatorial hydroxyls. The high reactivity of cis-2-tert-butylcyclohexanol is caused most probably by steric acceleration of the elimination, which is, however, absent in the case of 2,2-dimethylcyclohexanol. Table 5 also gives data on the effect of ring size on the rate which are consistent with observations on other reactions of cyclic compounds. This influence can be explained by the change of the strain in the ring (Istrain) in consequence of the change from sp3 t o sp2 hybridisation [ 1291. The above results give a clear picture of structure-reactivity relationships at a chosen temperature. The published information about the influence of temperature on these relationships is less consistent. In cases where the kinetics were clearly established (usually being zero order), the activation energy for individual members of a homologous series of compounds (e.g. 1-alkanols) has been found t o be structure-independent within the experimental error [92,94,104,124,128], On the other hand, the slightly different activation energies for a series of primary alcohols have been correlated successfully, together with the considerably different activation energies for secondary and tertiary alcohols, by means of the Taft equation [ 561. A similar decrease of the activation energy with substitution on C, was reported very early [90,92]. The activation energy of dehydration is further influenced by steric effects as indicated by the data in Table 5.
(c) Direct ion of elimination Over most catalysts, with the notable exception of thoria [130], the thermodynamically more stable olefins are formed (Saytzeff rule) as primary products when two elimination directions are possible. This is in agreement with the results concerning other elimination reactions, both in the liquid phase (cf. refs. 6 4 and 65) and over solid catalysts. The striking difference in the action of thoria has been explained on the basis of a different mechanism [ 681. The frequently observed preference for anti-elimination over syn-elimination on alumina (for a summary of earlier results see ref. 7 , later especially ref. 96) has been a cause of much controversy. However, as has been explained in Sect. 2.1.2, it is a natural reaction course for concerted elimination, provided that suitably spaced acidic and basic sites are available on the surface. Catalysts which operate by means of the El-like mechanism References p p . 385-398
f\3
00 00
TABLE 4 Linear free energy relationships for the dehydration of alcohols over solid catalysts Substi tuen t s (X o r R )
Alcohols
X . C,H4
. CHOH . CH3b
H, p-CH3, p t - C 4 H 9 , m-F, p-F, m-OCH3 H, p-CH3, rn-F, p-F, m-OCH3
Catalyst
A1203
+ NaOH
Zr02
Ti02 SiOz A12
O3
A1203 + NaOH ZrOz TiOz Si02 A12 0 3
ROH
A1203-Si02 + NaOH
Temp. ("C)
Correlation of a
porp*
Ref.
200
log k-0'
-2.6
105
220 220 220 220 380
log k-u' log k-' log k-0' log k-<+
-2.2 -2.1 -2.2 -2.4 -15.8
95 96 95 95 124,125
300 300 300 300 220 250
log k-0' log k-0' log k-' log k-' log k-' log r-u'
1.2 0.3 -0.8 -2.8 -2.0 -13.3
55 55 66 55 56 126
log F
0'
Non-stoichiometric hydroxyapatite
R . CH2 . CHOH . CH3b
= RZ = H, R3 = CH,; = R 2 = H , R 3 = C2 H5 ; = RZ = H, R3 = l - C 3 H 7 ; = RZ = H, R3 = 1C4H9 ; = R2 = H, R3 = 2-C3H7; = R2 = H, R3 = [-C4 H 9 ; = H, RZ R3 = CH 3 ; = H, R2 = CH3, R 3 = C 2 H S ; = H, R2 = CH3, R3 = 1-CjH7; R' = H, R2 = CH3, R3 = 2 - c ~ H 7 ; R'=HR2=R3=C 2 H5 ; R' = Rz = R3 = CH 3 ,. R' = R2 = CH,, R3 = C2H5 CH,, C2HS, l-C3H7, 2-C3H7,1-C4Hg
ROH
CH,, C2HS, 1 G H 7 , C6HSCH2
A1203
R'R~R~COH
tJ
P
a
R' R' R' R' R' R' R' R' R'
1
230
log h-u*
282 350
-5.1
56
-4.5 -3.9
Hydroxyapatite
395
log h-u*
A1203
210 210 160
log S21-u* log Sct--E, log t-u*
-2.3
2.2 0.3 2.9
127 128
k , rate coefficient; r , reaction rate; S 2 1 , ratio of rates of 2-alkene and 1-alkene formation; S,,, ratio of rates of cis- and trans-2-alkene formation; u*, Taft substituent constant; u+, Hammett substituent constant for mesomeric interaction between reaction centre and aromatic ring; E,, Taft steric constant. Olefin formation. Ether formation. Value of t h e 6 parameter.
290 TABLE 5 Effect of ring size and substitution on dehydration of cyclic alcohols at 200°C [94] ____
Alcohol
kRl
____ Cyclopentanol Cyclohexanol Cy cloheptanol Cyclooctanol 2,2-Dimethylcyclohexanol trans-2-Meth ylcyclohexanol cis-2-Meth ylcyclohexanol trans-4-Methylcyclohexanol cis-2-Tert-bu tylcyclohexanol trans-4-Tert-butylcyclohexanol cis-4-Tert-butylcyclohexanol
-
Activation energy (kJ mol-')
150 145 145
1.9 1.0 2.3 8.6 0.9 1.8 12.1 1.9 40 2.1 13.6
140 180 85 160 80 160 115
give cis and trans products in comparable quantities as experiments with threo- and erythro-3deutero-2-butanols have revealed [68,96,132].
2.2.4 Mechanism
( a ) Surface complex A number of contradictory views have been published concerning the structure of adsorbed alcohols and the nature of adsorption sites (for review see ref. 69). Experimental evidence from IR investigations has shown that, on alumina, alcohols form several surface complexes of very different chemical natures (e.g. refs. 31, 32, 117, 133-137): (i) alcohol molecules weakly bonded t o the surface, very probably by hydrogen bonds (I) (such complexes are sometimes denoted as physically sorbed alcohols); (ii) surface alkoxides (alcoholates) (11); (iii) surface carboxylates (111). Less certain is the existence of species with partial double bonds or of ketone-like species. The formation of the various surface complexes is dependent on the structure of the alcohol. For examples, weakly bonded species (I) have been found with all alcohols, alkoxides (11) mostly with primary alcohols, sometimes also with secondary alcohols, but have never been reported for tertiary alcohols.
sp
T0
H I 0
y
0 0
0 A1 A1
A1
IH
Al
A1
I
R I 0
R I C
(-%a
0 0 0 A1 A1
291 The question now is which of these complexes are intermediates in the dehydration of alcohols. There is general agreement that the carboxylate structure (111) is a product of a side reaction and has therefore no connection with dehydration. Species (Ia) and (Ib) could be natural precursors of the surface alkoxides (11) but such a transformation has not been observed t o be the main reaction. Hydrogen-bonded species either desorb without any change in the alcohol structure (or configuration [135]) or dehydrate directly t o an olefin [ 1371. However, the alkoxides (11) are very probably intermediates in the formation of ethers as will be shown later. Also, on an ion exchanger with sulphonic groups, the adsorption complex of methanol is hydrogen bonded [ 1381. ( b ) Mechanism of the surface reaction
For the olefin-forming reaction, two alternative paths have been considered with minor variations concerning the number and nature of surface metal and oxygen atoms which take part in the elementary steps. (i) The alkoxide mechanism assumes complex (11) as the intermediate; it was suggested by Sabatier [76] and advocated mostly by Topchieva et al. [81]. \I
A1
A1
A1 A1
A1
A1
(ii) The cyclic mechanism assumes cooperation of an acidic and a basic site; it was suggested by Eucken and Wicke [6] and supported by Pines and Manassen [ 71 and numerous other authors.
A1 A1 References p p . 385-398
A1 A1
Al A1
292 All recent results suggest the second mechanism. The arguments for its validity may be summarised as: (1)high stereospecificity of elimination on a number of catalysts; (2) existence of both basic and acidic sites on dehydration catalysts; (3) the possibility of treating all elimination reactions in a common way from the point of view of mechanism (cf. Sect. 2.1). Recently, a new argument has been added on the basis of quantum chemical calculations [70]which has shown that the attack of an acid on the hydroxyl group activates the hydrogenes on C, for elimination whereas the loss of hydrogen from the hydroxyl or its substitution by a metal ion (corresponding t o the formation of the surface alkoxide 11) activates the C,-H bond for dehydrogenation. The cyclic mechanism is probably seldom a fully concerted (E2) process, and the different timing of individual electron shifts results in a transition towards the E l or ElcB mechanisms (cf. Sect. 2.1.1). The “choice” of the mechanism depends on the reactant structure as well as on the catalyst nature. As an indicator of the mechanism, either the degree of stereoselectivity (see refs. 68, 121, 132 and 141) or the value of the reaction parameter of a linear free energy relationship, e.g. p or p * constants of the Hammett and Taft equations (cf. ref. 551, may be used. Pines and Manassen [ 71 suggested that tertiary alcohols are dehydrated by the E l mechanism involving the formation of more or less free carbonium ions, secondary alcohols by a mechanism lying somewhere between E l and E2 (i.e. synchronous with a ionic contribution) and primary alcohols by a concerted E2 mechanism. However, the large kinetic isotope effect for the dehydration of fully deuterated tert-butanol on alumina [122] indicates that, even in this case, some synchrony must exist. Alumina strongly prefers the concerted process and with other catalysts the situation may differ. In some cases, the effect of reactant structure may outweigh the influence of catalyst nature. This is seen by comparison with the dehydration of aliphatic secondary alcohols and substituted 2-phenylethanols on four different oxide catalysts (Table 4). With aliphatic alcohols, the slope of the Taft correlation depended on the nature of the catalyst (A1203 + NaOH 1.2, ZrOz 0.3, TiO2-O.8, Si02-2.8 [55]) whereas for 2-phenylethanols, the slope of the corresponding Hammett correlation had practically the same value (from -2.1 t o -2.4) for all catalysts of this series [95]. The resonance stabilisation of an intermediate with a positive charge on C, clearly predominates over other influences. In contrast to olefin-forming dehydrations, the transformation of alcohols t o ethers very probably includes surface alkoxides as intermediates. It is assumed that one molecule of the alcohol forms the alkoxide which is then attacked by the second alcohol molecule either from the gas phase or from a weakly adsorbed state. Again, cooperation of acidic and basic sites seems to be necessary [ 116,142,1431. The important step of ether forma-
293 tion is a nucleophilic substitution, viz. \I/ C \I/ I c ... 0 I 3 I
O
H
The arguments for the suggested mechanism are: (1) Similar products are obtained by the decomposition of metal alkoxides containing no 0-hydrogens and by the reaction of corresponding alcohols on alumina at lower temperatures [ 1421. (2) Acetic acid and pyridine are poisons for the formation of ethers [ 1431. (3) The different degrees of water inhibition on the ether and olefin formation from ethanol on alumina, and the agreement of etherlethylene selectivity ratios found experimentally with those calculated by the Monte Car10 simulation of the hydrated surface of alumina [ 1441. (4) Correlation between the rate of ether formation from ethanol and the surface concentration of ethoxide species determined by IR spectroscopy [ 1361. (5)The positive value of the Taft reaction parameter for the formation of ether in contrast t o negative values for the olefin formation on the same catalyst (Table 4).
(c) Influence of catalyst nature The necessity of cooperation between surface acidic and basic sites for splitting off the elements of water from alcohol molecules was intuitively suggested quite early [ 6 ] and used as a working hypothesis by an increasing number of authors. However, it took some time t o recognise which type of acidity and basicity is suitable for dehydration. The experiments with reversible poisoning of alumina by small amounts of bases like ammonia, pyridine or piperidine revealed [ 8,137,142,145, 1461 relatively small decreases of dehydration activity, in contrast t o isomerisation activity which was fully supressed. It was concluded that the dehydration requires only moderately strong acidic sites on which weak bases are not adsorbed, and that, therefore, Lewis-type sites d o not play an important role with alumina. However, pyridine stops the dehydration of tert-butanol on silica-alumina [ 81. Later, poisoning experiments with acetic acid [ 1431 and tetracyanoethylene [ 81 have shown the importance of basic sites for ether formation, but, surprisingly, the formation of olefins was unaffected. The picture has been clarified by surface acidity and basicity distribution measurements for several catalysts using thermometric titrations with References p p . 385-398
294
1-butylamine and trichloracetic acid [ 59,601. A fine balance between acidic and basic strengths of the working sites is necessary: ethylene formation requires the cooperation of moderately strong acidic sites with weak basic sites, whereas for diethylether formation, moderately strong basic sites play an important role. Moreover, the poisoning by bases increases the basicity; the disappearance of moderately basic sites manifests itself in a decrease of ether formation. These relations seem to be valid for the dehydration of primary alcohols, but secondary and tertiary alcohols may need other combinations of acidic and basic sites. It has been observed that the dehydration of tertbutanol was more sensitive t o the presence of strongly acidic sites than the reaction of methanol, but both processes required basic sites [ 81. All this is in accordance with the dynamic mode1 of elimination mechanisms presented in Sect. 2.1, which allows transition from E l t o E2 or further t o ElcB according t o the structure of the reactant and the nature of the catalyst. The relation between the acid strength of the catalysts and the mechanism has also been demonstrated by correlations [55,123] of the reaction parameter, p * , of the Taft equation for the dehydration of secondary alcohols on A1,03 + NaOH, ZrO,, TiOz and SiO, (see Table 4) with the sensitivity to pyridine poisoning, the heat of adsorption of water and diethylether and the kinetic isotope deuterium effects (Table 3) on the same catalysts (Fig. 5). The parameter p* reflects the mechanism being
-AHads Log 7 log a Fig. 5. Correlation of the Taft reaction parameter for the dehydration of secondary
alcohols (see Table 4) on four different oxide catalysts with the heat of adsorption,
AH,,,, of water and diethylether, with the sensitivity of the rate to pyridine poisoning 7,[55] and with the value of the deuterium kinetic isotope effect [ 1 2 3 ] for the same catalysts.
295 related t o the free energy change on the interaction between the reactant's reaction centre and the catalyst's active centre. The kinetic isotope effect also depends directly on the mechanism and the other quantities show the acid strength of the active centres. It is believed that, in this series of catalysts, the transition toward the E l mechanism is observed in the direction from alkalised alumina t o silica. The question remains open whether the addition of alkali t o alumina shifts the mechanism fully t o ElcB. Another diagnostic method for determining the effect of the nature of the catalyst on the mechanism is the observation of the stereoselectivity of elimination [68,121,132]. It has been found, using the reaction of threo- and ery thro-3deuter0-2-butano1, that, in a series of salt catalysts, only the phosphates Ca3(P04)*,CaHP04, Ba3(P04)*and A1P04 preferred the E2 mechanism while on other metal phosphates and all carbonates, the E l mechanism predominated [ 1321. Another fine distinction among salt catalysts was obtained by following the activity and olefinlether selectivity of metal sulphates in the dehydration of ethanol and 1-propanol. A linear correlation between the electronegativity of the metal ion and the activity has been found, but the selectivity gave a curve with a minimum [ 511. 2.3 DEAMINATION
2.3.1 Types of deamination reactions The transformation of amines over acidic catalysts is more complicated than that of other compounds which can undergo elimination. Like alcohols, amines form either an olefin or a higher substituted amine (the formal counterpart of an ether). With trivalent nitrogen, the reaction scheme includes more compounds than divalent oxygen allows, viz. RNH,
J
.
olefin + NH3
-NH3 +NH3
. R2NH
J
olefin + RNHz
1
etc.
-NH3 +NH3
R3N
J.
olefin + R2NH
1
etc.
Scheme 3 .
At temperatures above 250" C, the olefin-forming reactions are irreversible but the transformations in the first line of Scheme 3 are reversible. Thus, starting with an arbitrary amine, all other derivatives are obtained by these reactions, called disproportionations, transalkylations or dismutations (the nomenclature is also inconsistent in that the analoguous formation of ethers from alcohols is named dehydration). Similarly, like secondary and tertiary alcohols, amines with the alkyl groups branched on a-carbon atoms (i.e. containing the grouping R'R'CHReferences P P . 385-398
296 N- and R’RZR3C-N-) tend more to olefin formation than t o disproportionate. N-Alkylidenalkylamines were found in the reaction products of the transformation of l-butylamine [ 147,1481, cyclohexylamine and isopropylamine [149] on alumina, and were probably formed by the dehydrogenation of the primary amine to an imine, followed by its condensation with a second molecule of the amine [ 1481, rather than by the dehydrogenation of the dialkylamine [ 1471. The N-alkylidenalkylamines R=N-R decompose to an olefin and an imine; a cyclic process has been postulated [ 1481 which explains the increases in reactivity of amines with secondary alkyls. Also butyronitrile has been detected in appreciable amounts in the reaction products of l-butylamine on alumina at 500°C [ 1471. Much less information is available on the deamination and related reactions over solid catalysts than on some other elimination reactions but it suffices for comprehension of the general features.
2.3.2 Catalysts Only fews solids have been used as catalysts for deamination and disproportionation reactions. Among them, alumina has been studied most frequently, and some attention has also been paid to silicaalumina and t o molecular sieves [ 1491. The activity of alumina for the disproportionation of weakly basic aniline t o diphenylamine can be enhanced by impregnation with HC1 [ 1491 or H3B03 [ 1501.
2.3.3 Experimental kinetic results ( a ) Formal rate equations In most cases, the rate equation, eqn. (3) (p. 280), was suitable t o describe the deamination of alkylamines t o alkenes on alumina. It has been applied t o the decomposition of mono-, di- and trialkylamines [ 1511 and of cyclohexylamine [ 1481. In some cases, the simpler zero-order rate expression (cf. Sect. 2.1.3) has been observed [ 149,1531. Equation (3) can also describe the disproportionation of amines such as aniline [ 151,1541 and diethylamine [ 1531. However, the related expression
~KAPA (6) (1+ K A P A + K B P B + KCPC)’ where A denotes the starting amine, B the amine formed by the reaction and C ammonia, also suits the data on disproportionation quite well [ 152,1531. In two cases, more elaborate kinetic models were necessary in order to
r=
297 obtain a good fit of the experimental data. For the deamination of triethylamine and diisopropylamine on alumina, the superposition of two parallel processes [ 1553, one of zero order and the second described by eqn. (3), has been suggested; the final expression had the form
To explain the retardation by the very weakly basic aniline of the deamination and disproportionation of diethylamine and diisopropylamine on alumina, the acidic and basic sites were separately balanced in the derivation of the rate equation and the following expression was obtained [156], viz.
where the subscript P denotes aniline and a, b the acidic and basic sites, respectively. The complete system of deamination and disproportionation reactions has been treated with success by means of eqns. (3) and (6) on the basis of integral data with the exclusion of the time variable, i.e. with relative concentrations of reactants and products [ 1521. ( b ) Structure and reactivity
Two extensive sets of data on the reactivity of various amines in deamination and disproprotionation on alumina give an insight into the influence of structure on rate [ 149,1521. However, the picture is complicated by different effects of the number and nature of the alkyl substituents on the reaction rate coefficient and on the adsorption coefficient. Adsorption studies have revealed [157-1601 that the amount of adsorbed amine depends mostly on the size of the molecule. A linear correlation was obtained between the reciprocal cross-section of the molecule and the adsorptivity [ 1571. However, practically no differences have been observed for the series of primary amines RNHz with the straight chain alkyl groups; this was explained by a perpendicular orientation of the adsorbed amine molecule with respect t o the surface [160]. On the contrary, the adsorptivity decreased with branching of the alkyl group [160] and with the number of alkyl groups on the nitrogen [157--1591; steric hinderance seems t o be the obvious cause. Similarly, large differences in adsorptivity of pyridine and 2,6-dimethylpyridine on alumina have been reported [ 1581. Also, the adsorption coefficients, Ki, determined by means of eqn. (3) from kinetic data [ 1521 show the same trend. The rate coefficient of deamination increases with the inductive effect of the alkyl group as has been demonstrated by the published kinetic data [152] and their correlation by the Taft equation [125]. This was later References p p . 385-398
298 confirmed by experiments which were not complicated by disproportionation and in which the size of the amine molecule was the same or almost the same in the whole series. The deamination of N-ethyl-N-propyl-Nbutylamine showed an increase in the reactivity of the alkyl groups in the order ethyl < propyl < butyl [ 1491. The same trend was observed within the series of alkyldimethylamines [ 1491. It has also been found that the reactivity of the isopropyl group is influenced by other (less reactive) alkyl groups on nitrogen in the order methyl < ethyl < propyl [ 1491. No data are available on the direction of elimination in the deamination on acidic catalysts.
2.3.4 Mechanism The adsorption of ammonia and amines has been studied many times as a method of estimation of the acidity of solid surfaces. Some of the results are pertinent t o the mechanism of amine transformation on these catalysts. Depending on the structure of the catalyst surface, several types of adsorbed species have been observed. Nitrogen bases can be adsorbed on both Lewis and Br@nstedacid sites by means of their free electron pair. Other modes are realised by hydrogen bonding, the hydrogen originating either from a surface OH group or from NH groups of primary and secondary amines. Under suitable conditions, dissociative adsorption of NH3 has been observed by IR spectroscopy; surface NH; ions are formed [7]. Ammonium ions NH'4 have also been found [ 1611. Methylamine [ 1621, aniline [ 1631 and pyridines [ 1361 are adsorbed by means of their nitrogen free electron pairs on Lewis acid sites. Several types of adsorbed species can exist simultaneously on the surface. Measurements of the heat of adsorption of ammonia on silicaalumina catalysts indicated two types of adsorbed forms, one with a low energy of adsorption of about 30 kJ mol-', the second form having double that value [ 1641. Thermal desorption and gas chromatographic studies of adsorption of amines on alumina led t o the suggestion that three types of adsorption have t o be distinguished: below the coverage of 0.1 mmol g-l, strong irreversible adsorption on A13+ sites, in the range 0.1-0.3 mmol g-', reversible adsorption on OH groups and over 0.5 mmol g-' adsorption o n 02-ions also [159]. The reaction mechanism of amine deamination and disproportionation has been put forward by analogy with other eliminations, namely dehydration and dehydrochlorination [ 149,1551, its characteristic feature being the cooperation of acidic and basic sites. In the deamination, 0-hydrogen and the NR, group ( R is hydrogen or alkyl) are eliminated by an E2-like mechanism on alumina, but by an El-like mechanism on silicaalumina. The nature of the acidic sites is not clear, protons from surface hydroxyls o r aluminium ions are possible candidates. However, it seems
299 quite probable that strong A13+sites are poisoned at a very early stage by bases present in the system and consequently do not contribute much t o the reaction (cf. ref. 159). Saturation of the dehydrated alumina surface with water, creating new hydroxyl groups, increased the rate of deamination [ 1491. All this is in favour of the surface hydroxyl groups as acidic sites. The basic sites are most probably the surface 0'- ions. Then the surface complex can be written as
Quantum chemical calculations for some alkylamines, RNH', and their protonised forms supported this model [ 1651. As in the case of alcohols [139], the action of an acid on the nitrogen atom activates the @hydrogens and the weakening of the C,-H bond is most pronounced when one hydrogen atom is in the antiperiplanar position with respect t o nitrogen. Thus, the anti (or trans) elimination is again suggested as the most favoured mode of elimination. The disproportionation of amines is visualised as a nucleophilic substitution at the 0-carbon atom of an amine adsorbed on an acidic site [149,153]. The attacking species is an amine adsorbed by a hydrogen bond onto a basic site
This is again a direct analogy with ether formation from alcohols (see Sect. 2.2.4). The acidic sites might be the A13+ions because rehydration of the alumina surface does not enhance the rate, in contrast t o deamination [ 1491. Very little is known about the behaviour of different catalysts; only a few comparisons of alumina and silica-alumina have been made. On Alz03-Si02, the disproportionation of diethylamine is more rapid by one order of magnitude than its deamination; on A1203, the rates are comparable [ 1491. The activity of alumina for aniline disproportionation is higher than of silica-alumina [150]. The steric demands of the alumina surface are higher than those of silica-alumina as the comparison of the chemisorption of pyridine and 2,6dimethylpyridine has shown [ 1581. Rc9ferences p p . 385-398
The complex interplay of basic and acidic sites in the deamination and disproportionationn of amines is the probable cause of the “stop-effect” which has been observed in the reaction of triethylamine on alumina [155] and, more recently, of other amines [149]. When the steady state on the catalyst surface in a flow reactor is rapidly changed by substituting the amine feed for a nitrogen stream, a rapid temporary increase in olefin production is observed. This phenomenon has been explained as the result of the increased availability of basic centres which were previously blocked by adsorbed molecules [ 149,1551. 2.4 DEHYDROHALOGENATION
2.4.1 Types of d e hy d rohalogena t ion reactions The 0-elimination of hydrogen halides HX from organic halogen compounds yields olefins o r acetylenes, depending on the structure of the starting substance and the number of HX molecules which have split off, viz. X-A-C-H I
I I =C=C + H X I I
(A)
The reaction is reversible and, in general, higher temperatures and lower partial pressures favour the decomposition by shifting the equilibrium t o the right-hand side. The position of the equilibrium is influenced by the groups attached to C, and C,; for example, the heating of 2,2,2-trichloro1,l-bis(pchloropheny1)ethane with FeC13 t o 115--120°C results in a rapid and quantitative conversion t o 2,2dichloro-l,l-bis(p-chlorophenyl)ethylene [ 1661, whereas for high conversions of 1,2dichloroethane t o vinyl chloride, temperatures over 400°C are zeccssary. The equilibrium of reaction (B) is less favourable for the formation of an alkyne and, in order t o achieve equal conversion, much higher temperatures would be required than are necessary for the olefin-forming elimination (A). Therefore little attention has been paid to this type of reaction and this section will be devoted solely to type (A) dehydrohalogenation.
2.4.2 Catalysts A large variety of catalysts, both homogeneous and heterogeneous, has been found active for dehydrohalogenation. The catalysts include a number of Br4nsted and Lewis acids (liquid or soluble, as well as solid), metal oxides, active carbon, carbides, nitrides and some metals. However, in the latter case, the actual catalysts are most probably surface metal halides
301 TABLE 6 Comparison of the activity of different catalysts for the dehydrochlorination of 1chlorobutane [ 1 7 4 ] Catalyst
Temperature ("C) of incipient decomposition
of appreciable decomposition
~.
Thoria Zirconia Calcium chloride Calcium phosphate Alumina Barium chloride Bone charcoal, HCI washed Active carbon
205 205 224 240 24 5 24 5 255 280
220 220 245 260 260 210 275 305
formed from the metal and the reactant [ 1671. Even glass shows some catalytic effect [ 1681 making purely thermal rate measurements impossibly in glass vessels; again, surface metal halides might be the cause of the activity. From the point of view of the present review, the most important catalysts are metal salts and metal oxides which have been used for dehydrohalogenation of simpler organo halides in the gas phase and studied with respect t o kinetics and mechanism. The metal salt group is represented mainly by A1F3, FeF, and MgF, [ 169,1701 and chlorides, sulphates, carbonates and phosphates of the Group Ia, I1 and 111 metals which were employed by many authors but systematically studied by the Noller school (for review see ref. 67) and by the group of Mochida and Yoneda. Occasionally, alkali and alkali earth metal bromides were used. Oxides of the Group Ia, 11, I11 and IV elements are active catalysts and comparative descriptions of their properties has been published by several authors (e.g. refs. 67, 1 7 1 and 172). However, there is some difficulty in finding out the relative activities since most papers have dealt with selectivity problems or with comparisons of the activity in closely related catalyst groups (e.g. in a series of barium salts with different anions [173]). Some information about the activity of quite different types of catalyst is brought together in Table 6.
2.4.3 Experimental kinetic results ( a ) Formal rate equations Only a few kinetic studies on dehydrochlorination and dehydrobromination have been published. They are summarised in Table 7 and the general impression is that the more complicated rate equations have resulted References p p . 385-398
TABLE 7 Kinetics of catalytic dehydrohalogenation Catalyst
Technique
Temperature Rate equation range (“C)
Activation energy (kJ mol-’)
Ref.
Flow Flow Flow Flow Flow Flow Flow Flow
520-600. 368-420 355-404 38 7-42 1 354-4 13 404-438,
100 54 46 96 63 120 88 117
175 175 175 175 175 175 176 176
Chloroethane A12 0 3
SiOz CaClz Na2S04 CaS04 SrC12 Alz 0 3 MgS04
r=
kKAPA (1 + C K i p i ) i
1 -Chloropropane
MgS04
177
Flow
1,2-Dichloroethane Active Flow carbon
r
=
kp2’
1,1,2-Trichloroethane Pulse-flow NiS04 1 -Chlorobutane Zeolite Flow 2-Chlorobutane Zeolite Flow
or
120-140
178
179 48
180
38
180
a4
181
63
181
Stop-flow Stop-flow Stop-flow Stop-flow Stop-flow
79 125 120 84 79
181 182 182 182 182
2-Brorno-2-rnethylpropane Al203-LiCl Stop-flow
130
182
1 -Brornobutane A1203-KBr
Stop-flow
2-Brornobutane A1203-KBr Stop-flow
r = kpA
1-Brorno-2-methylpropane
Al203-KBr A1203-LiCl A1203-KCl A1203-NaCl A1203--C~Cl
from more careful work with continuous flow reactors in which the conditions were varied in a broader range. The chromatographic (pulseflow) method usually requires the postulation of first-order kinetics as a basis for the treatment of data. However, Langmuir-Hinshelwood types
303 TABLE 8 Deuterium kinetic isotope effects in the dehydrobromination of the 1,2-dibromoethanes CH2Br. CH2Br and CDzBr. CDzBr [ 1 7 2 J
--_ SrO NaOH-SO2 CaO K2S04-Si02
2.15 1.83 1.65 1.65
KOH-Si02 NaOH-SiO? NiS04-Si02 Al203-Bz03
1.59
1.31 1.23 1.05
M a O3 Si02-AI2O3 NiSO4
1.02 1.00 1.00 1.00
rate equations suggested by some authors can easily transform t o the firstorder expression if the adsorption of reactants is small (see Sect. 2.1.3). f b ) Structure and reactivity
The rate of catalytic dehydrohalogenation is influenced by the structure of the reactants, but the extent of this effect varies from one catalyst to another with change of mechanism, i.e. with the timing of the fission of the C,-X and C,-H bonds. This is best seen from the published data on the deuterium kinetic isotope effect in Table 8. Their significance for the elucidation of the mechanism will be dealt with in Sect. 2.4.4 and here we can simply state that the value of the isotope effect depends on the nature of the catalyst. However, with a different reactant and within a series of related catalysts, k H / k D values independent of the catalyst were obtained (Table 9) [ 1831. A clearer picture emerges from studies of substituent effects on the rate with a single catalyst. A series of alkyl chlorides (C, to C,) was decomposed on a barium sulphate catalyst [184] and the rate data were correlated by the Taft equation. Large negative values of p* were obtained, viz. -34.3 at 220°C and -40.3 at 280°C. Similarly, for a series of three alkyl bromides (ethyl, propyl and isopropyl) on silicaalumina, aluminaTABLE 9 Deuterium kinetic isotope effects in the dehydrochlorination o f 2-chloropropane ~ 8 3 1
_ _ _ _ _ ~
Catalyst
kH/kD
for the pair
CH3CHCICH3 CH3CDCICH3
CH3CHCICH3 CD3CHClCD3 1.8 1.75 1.8 1.6
References P P . 385-398
304 TABLE 10 Order of reactivities in the dehydrochlorination of chloroethanes o n different catalysts
I Reactant
Product
CHCl=CH:! CHCI=CH, CCl2=CH2 CC12=CH1 CICH=CHCI CCl2=CHC1
Cl2 CH-CH3 ClCH24HZCI C13C-CH 3 Cl;, CH-CHZCI
ClZCH-CHC12
Order of reactivity on
SrO and
A1203
A1203-Si02
6 5 3 2 4 1
3 5 1 6 4 2
2 5 1 6 3
4
boron oxide and alumina at 300”C, a p* value of about -20 was found [172]. This is in accordance with similar correlations of substituent effects on the rate of pyrolysis of alkyl chlorides and bromides at 400°C and 350°C ( p * = -23.5 and -22.3 respectively) [185]. The unusually high p* values can be attributed to the fact that the substituents were defined in both cases as R in RX instead of the more correct R ’ in R‘CH2CH2X and t o the high polarity of the transition states. The complex influence of both the reactant structure and catalyst structure is evident from attempts t o correlate the reactivities of a series of chloroethanes (see Table 10) on basic catalysts (SrO and BaO), on alumina and on strongly acidic catalysts (A1203-Si02, A1203-B203) with the delocalisability of hydrogen atoms [l86]. The delocalisability of an atom is a quantum chemical reactivity index calculated by the simple LCAO-MO method according t o Fukui et al. [187]. For basic catalysts, a good linear relationship was observed between the rate corrected by the number of equivalent hydrogen atoms and the delocalisability of the hydrogen atoms in the @position t o chlorine atoms for the nucleophilic abstraction. However, no correlation was obtained for acidic catalysts and alumina. The reactivity of a halogen compound in dehydrohalogenation over solid catalysts also depends on its steric arrangement. This was shown by studying the dehydrohalogenation of the rigid molecules l-bromo-1,2diphenylethylene [ 1881 and l-chloro-1,2-diphenylethylene [ 1711 on catalysts of the type of metal salts and metal oxides: the cis-compound was always more reactive than the tram-derivative.
H
,c6H5
C6H5,
/c=c\ >
.
Br
H
cis
/c=c Br
\
trans
C6H,
The work with 1-bromo-2-chloroethane allowed the influence of the nature of the halogen on its reactivity to be observed as either vinyl bromide or vinyl chloride are formed. The ratio of the chloride t o the bromide in the products changed with the nature of the catalyst, being around 0.1 for sulphates of Ni, Co, Mn, Cu, Zn and for silica-alumina, 0.6 for alumina and 5 for KOH-Si02 [ 1791. (c) Direction o f elimination
A large amount of work has been devoted t o the problem of the rules governing the dehydrohalogenation of halogen derivatives which can form several olefins. Although some regularities may be observed, the general picture is clouded by the following facts. (a) The direction of elimination depends very strongly on the nature of the catalyst because on different catalyst types, different mechanisms operate. (b) The nature of a catalyst (and the mechanism) may be changed by the hydrogen halide which is produced by the reaction [ 1893. (c) The composition of the product (and the participation of different mechanisms) is temperaturedependent and, as various catalysts sometimes differ appreciably in activity, a comparison of selectivities under the same conditions is impossible in a broader series of catalysts. (d) The composition of the product may be changed by a secondary isomerisation of the olefins formed. The point (a) is demonstrated by the data in Table 11for 2-halobutanes which give three products: 1-butene, cis-2-butene and trans-2-butene. The differences in selectivities can be even larger than indicated by these data. The dehydrochlorination of l,l,Z-trichloroethane yields 1,2dichloroethylene (I) and trans- and cis-l,2-dichloroethylene(11). On silica-alumon alumina, 0.30 and on KOHina, the value of the ratio 1/11 was SiOz, 1 0 [66]. The data in Table 11show some facts, the mechanistic consequences of which will be discussed in Sect 2.4.4. The 2-alkene/l-alkene ratio for the catalytic reaction differs significantly from that for the homogeneous decomposition. On all catalysts, this ratio is higher for the 2-bromo- than for the 2-chloroderivative; therefore the orientation also depends on the nature of the halogen. On some catalysts, both ratios (the 241- and cis/ trans) are equal or approximately the same as the equilibrium values, but on other catalysts, significant differences appear. The influence of temperature on the ratio of the products of the dehydrochlorination of 2-chlorobutane is seen from Fig. 6 [ 1901. Other examples may be found in the literature [190,194,195]. Some ratios are almost temperature-independent while some show large changes. Moreover, the data from various sources differ sometimes appreciably (cf. refs. 190 and 914 for 2-chlorobutane and refs. 66 and 195 for 1,1,2-trichloroethane). This might be caused by secondary isomerisation on strongly acidic catalysts of the olefins first formed such a reaction was proved at References p p . 385-398
306 TABLE 11 Examples of product ratio obtained from 2-halobutanes on different catalysts Catalyst
Temp. (“C)
2-Chlorobutane MgS04 Bas04
2-Butene 1-Butene
cis-2-Butene trans-2-Butene
Ref.
Li2SO4 K2S04 CUO CaO None
200 240 24 0 270 345 200 150 350
6.7 8.1 6.1 5.7 8.1 2.8 9.0 1.5
0.58 0.43 0.69 2.15 1.07 1.8 2.0 0.53
190 190 190 190 190 19 1 191 192
2-Bromobutane MgS04 Bas04 A12 (so413 &SO4 KzS04 Si02 KO H-S i 02
200 180 210 230 300 150 160
11.5 10.1 8.1 6.7 9.0 5.3 3.2
0.53 0.58 0.62 2.0 1.04 1.00 0.94
190 190 190 190 190 193 193
Equilibrium composition [I 911 100 13.3 200 6.7 300 4.3 400 3.3
0.43 0.58 0.62 0.67
A12 (s04
)3
least for MgS04 [190]. The water content of metal salts as dehydrohalogenation catalysts also influences the selectivity [ 1961 and consequently the thermal history of the solids must play a role. Therefore, all selectivity data should be judged with caution, especially in cases where the over-all conversion of the haloalkane was high and the catalyst illdefined. Further information about the direction of dehydrohalogenation was obtained with threo- and erythro-2deutero-3-bromobutanes on Si02 and Si02-KOH catalysts [ 193,1951. The determination of deuterium content in the butenes formed allowed the estimation of the extent of the synand anti-eliminations. The values of the deuterium kinetic isotope effect showed that the C,-H (or C,-D) bond is split in the rate-determining step. Over KOH-Si02, the anti-elimination was preferred, but at 3OO0C, syn-elimination was the peferential reaction mode. With SiOz, synelimination was favoured under all conditions. Further stereochemical studies with di- or tri-haloalkanes corroborated the general picture of a strong dependence of the direction of elimination on the nature of the catalyst. The data in Table 1 2 may serve as an exam-
307 I
I
200
100
300 Temperature
400
("C)
Fig. 6. Dependence of the selectivity ratios of the dehydrochlorination of 2-chlorobutane on temperature for different catalysts [ 1901.
ple and additional results were reported for 2,3-dibromobutanes on alkalised silica [ 1981. Large changes of stereoselectivity with the catalyst nature were also demonstrated in the dehydrohalogenation of 1,1,2-trichloroethane t o cis- and trans-dichloroethylene [ 66,172,179,199,2001 and 1,2dihalopropanes [ 1901.
TABLE 12 Ratio of cis- to trans-2-chloro-2-butenes from the dehydrochlorination of meso- and dZ-2,3-dichlorobutanes on different catalysts [ 197 ] Catalyst
K2C03
BaC03 NiC03 LiCO3 CaC12 Ca3(P04)2
Temp.
cisltrans ratio from
("C)
meso
dl
250 330 250 330 350 220 220
11.5 49 6.7 3.0 0.90 0.28 0.20
0.075 0.53 0.23 0.67 0.90 0.30 0.16
References p p . 385-398
308 2.4.4 Mechanism
The hypothesis of a continuous transition of the elimination mechanism from the extreme E l through concerted E2 t o the other extreme E l c B with the change of reactant structure and catalyst nature, described in Sect. 2.1, can be easily adopted for dehalogenation also. The data summarised in Sect. 2.4.3 show some inconsistencies but the over-all picture is clear. This can be demonstrated for some selected examples. Typical basic catalysts are strontium oxide and alkalised silica (studied by the group of Mochida and Yoneda) o r potassium carbonate (studied by the Noller group). With SrO and NaOH-SO2, a large deuterium kinetic isotope effect was observed for 1,2-dibromoethane [ 1721, which shows that the C,-H bond is split in the ratedetermining step. The trans/cis olefin selectivities in the reaction of 1,2-dichloropropane and 1,1,2-trichloropropane [179,189] on SrO and alkalised silica give the greatest values indicating the high stereoselectivity of the elimination. The same results are obtained in experiments with meso- and dl-dihalobutanes on NaOHSiOl [ 1981 and K2C03 [ 1971. All this indicates an E2-type mechanism, very probably an E2cB mechanism in which the fission of the C,-X bond is slightly preceded by the fission of the C,-H bond. However, the structure of the reactant may shift the timing of the two steps of the elimination in the direction of the extreme E l c B mechanism. Evidence for this can be seen in the lower value of the deuterium kinetic isotope effect for the dehydrohalogenation of 1,1,2,2-tetrachloroethane(1.2) than of 1,2dibromoethane (1.6) on KOH-Si02 [66]. The cause of this transition is the increased acidity of hydrogen atoms in the compound with more halogen atoms. Typical acidic catalysts are silica-alumina, transition metal sulphates o r chlorides, calcium phosphate etc. They are characterised by low deuterium kinetic isotope effects and low stereoselectivity (see Tables 8, 11 and 12). These results correspond t o the E2cA or E l mechanisms, between which a transition may be observed due t o the influence of the structure of the reactants, i.e. according t o the polarity of the C,-X and CD-H bonds. Again, the reactions of 1,2dibromoethane and 1,1,2,2-tetrachloroethane yielded the evidence. The deuterium kinetic isotope effect on silica-alumina was 1.0 for the dibromo-derivative, which indicates a pure E l mechanism, whereas for the tetrachloroderivative, the value of 1.5 was found. Alumina is a catalyst which shows intermediate behaviour and over which the concerted E2 mechanism is accepted [66] with slight transition either t o the E2cA or E2cB mechanisms according t o the structure of the reactant. Salts of strong acids and bases also show similar intermediate properties. The concerted or partly concerted mechanisms require two-site adsorption and because the mechanisms are ionic, the active centre must consist of a pair of an acidic and a basic site. Metal salts fulfil this
309 condition; their relative activities should depend on factors like acid and base strength of the surface atoms, lattice spacing, concentration of suitable site pairs etc. Another important factor is the contamination of the surface by adsorbed species like catalytic poisons [ 1791, water [ 1961 or product hydrogen halide [ 1891. Therefore all studies which attempted to find correlation between some property of the catalysts and their activity ended with inconclusive results (cf. refs. 67, 172 and 194). The most reliable data concern dehydrochlorination on zeolites containing different monovalent and divalent cations. A good linear relationship was obtained between the activation energy of the first-order reaction and the electrostatic field of the cation [180] and between the cis/trans selectivity and acid strength [ 2001. Both correlations are in agreement with the proposed shifts in mechanism. The catalysts which operate by means of an E2 mechanism give a high proportion of reaction products which are formed by the anti-elimination. This fact has been discussed in Sect. 2.1 and only few remarks need t o be added here. Quantum chemical calculations [73] on the transition state of the dehydrochlorination of chloroethane, initiated by an attack of a basic species, confirmed the preference of the anti-elimination over the syn-mode. On the contrary, calculations on the transition state for non-catalytic (homogeneous) thermal elimination [ 201,2021 confirmed the syn-elimination path. 2.5 DEALKYLATION BY CRACKING
2.5.1 Types of cracking reactions
In the chemical and petroleum industries, the term cracking is used t o describe a chemically complex process in which the decomposition of larger hydrocarbon molecules into smaller fragments plays a dominant role but is accompanied by a number of other reactions (isomerisation, cyclisation, polymerisation, disproportionation etc.). In this section, under catalytic cracking, only the primary fission of a C--C bond, which yields an alkene and a fragment with a C-H bond in the place of the former C-C! bond
I
R-C-C-H
I
=
R-H
I I + C=C I I
(A)
will be considered. The group R must therefore be alkyl or aryl. Reaction (A) resembles other olefin-forming elimination, not only formally but also in all general features. Of course, there are some special characteristics of the cracking reaction which are due t o the nature of the eliminated groups, R. The general chemistry of the cracking over solid catalysts was studied Rcfcrenccs p p . 385-398
310 intensively in the short period during the forties which followed the introduction of the cracking process and of aluminosilicate catalysts into the production of petrol (for a review see ref. 203). Two types of cracking reactions are of interest for the present treatment: (i) dealkylation of alkylaromatic compounds and (ii) fission of a saturated hydrocarbon chain. The first reaction is characterised by the fission of the bond between an aromatic ring and an alkyl group, e.g. O ( f 4 - H
= O
I I H + C=C I I
(B)
The second type of cracking can take place at any C-C bond of a saturated hydrocarbon with the exception of the bonds t o the terminal CH3 groups, e.g. / CH3*HZ*H2*H3
+ CH2=CH2
CH3*H,*H2-CH,-CH,
The cracking reactions are catalysed by strong acids, both liquid or solid, like sulphuric acid, aluminium trichloride, aluminosilicates (including zeolites) and similar two- or three-component oxide mixtures. Solid oxide catalysts are more suitable than other acids because they withstand the necessary high temperatures better. A close parallel between the action of inorganic Brqjnsted and Lewis acids, like sulphuric acid, hydrofluoric acid and aluminium trichloride on the one side, and aluminosilicates or natural clays on the other, was recognised very early [l]. Compositions like Al,03-Si0,, A1203-BZ03, MgO-SiO, and A1,03-Zr02
311 were the first synthetic catalysts employed, besides acid-washed natural clays, for cracking studies (for an early review see ref. 2). Later, preparations containing BF3 on an alumina carrier were used in laboratory work [ 204,2051.
2.5.3 Experimental kinetic results
( a ) Formal rate equations Several rate equations, some of them in the integrated form, have been used for a kinetic description of cracking. A critical comparison of them yielded the consistent kinetic model [ 2061 discussed below. Cracking is a reaction of the type
A=R+S where A denotes the starting compound, R the alkane or arene, and S the alkene formed. Most authors agree [204-2181 that the reaction is best represented by the simple Langmuir-Hinshelwood type rate equation, eqn. (3) (p. 280). Some authors have carefully tested it in the differential form [ 204,209,210,216-2181 , others followed the procedure given by Frost [ 2071 who integrated eqn. (3) and after some algebraic manipulation obtained the expression
where F/W denotes the space velocity and x the conversion. ( Y ~ . and OF are complex parameters containing rate coefficients and/or adsorption coefficients
where P is the total pressure in a flow reactor into which the pure compound A is injected. The Frost equation is evaluated [ 205,207,208,213, 2141 by plotting ln(1 - x ) - ' against x and taking the value of the intercept o F as a measure of reactivity or catalyst activity. The rate equation, eqn. (3), has been simplified for catalytic cracking on the basis of various assumptions. When all components are strongly adsorbed, the term C,K,p, in the denominator is much greater than 1and can be omitted. Then the expression [ 212,2191
r
=
R c ferciirrs
kKAPA- -~ KAPA + KRPR + K S P S 911.
385-398
(10)
312 is obtained. When only one component is strongly adsorbed, for example the alkene in the cracking of an alkane (cf. ref. 206), a simple expression can be used, i.e.
The drastic simplification based on the assumption that all reaction components are weakly adsorbed, which leads t o the first-order rate equation, has been found acceptable by some authors [ 206,220-2231. Only a few authors have arrived at different conclusions [224,225]. For example, the empirical rate equation, which fitted the data on cumene cracking, was
where C and C' are constants. In the presence of catalytic poisons, that is substances which are strongly adsorbed on the catalyst surface, the rate equation, eqn. (3), has t o be expanded by adding the term K p P p t o the denominator
Table 13 collects values of adsorption coefficients of some compounds determined by means of this equation from experiments on the cracking of cumene on an aluminosilicate catalyst. ( b ) Structure and reactiuity
The published data give a clear picture of effect of structure on reactivity in the cracking of alkylaromatic compounds. The reactivity increases with the size of the alkyl group, and with its branching, as the studies of TABLE 13 Adsorption coefficients (in bar-') of various substances on a aluminosilicate catalyst at 420°C [ 2 2 6 ]
_~____
Compound
KP
Cumene Benzene Naphthalene Methanol Acetone Phenol Py rid i ne Quinolint?
0.685 0.274 1.03 X lo2 1.28 X lo2 5.22 X lo2 7.32 X lo2 1.67 x 105 1.24 X lo6
313
Fig. 7 . Hammett correlation of the cracking of 1,l-diarylethanes on a kaolin catalyst at 500OC. (Data by May et al. [220].)
- 0.3
-0.2
by
-O*’
Fig. 8 . Taft correlation of the cracking of alkylphenols o n an aluminium fluoroborate catalyst at 425OC (Data by Schneider et al. [204].) References P P . 385-398
314 the cracking of alkylbenzenes on aluminosilicate catalysts [ 210,213,218, 2271 and of alkylphenols on an aluminium fluoroborate catalyst [204] have shown. Substituents on the aromatic nucleus influence the cracking of an alkyl group in alkylbenzenes in a similar way t o aromatic electrophilic substitution; this follows from the experiments with substituted ethylbenzenes [ 2161, isopropylbenzenes [ 2281, diphenylmethanes (a special case of elimination yielding substituted benzenes and toluenes) [220] and diphenylethanes [ 2291. When the series of compounds contain a sufficient number of derivatives, a linear free energy relationship can be applied for correlation of the rate data [ 125,227,229,2301. Two examples of such correlations are presented in Figs. 7 and 8. A summary of all correlations together with the values of the reaction parameters p and p' from the Hammett and Taft equations is given in Table 14. Other linear correlations were also sought, e.g. of adsorption coefficients of alkylbenzenes against the bond strength alkyl-CH, [ 2101 of of the apparent activation energy against enthalpy change of carbonium ion formation [ 2301. The most successful was the correlation of the dealkylation rate against the enthalpy of hydride ion abstraction from the corresponding alkane [ 2271. Less information is available about the cracking of alkanes. Three sources [ 212,232,2331 confirm that, in the series of straight-chain alkanes, the rate increases with the molecular weight. The data could be correlated by the Taft equation when the molecule was divided into the reaction centre and such substituents as R - C H 2 - C H 3 [ 1251. The reactivity of hexane isomers was studied at 550°C on an A1203-Si02-Zr0,
TABLE 14 Linear free energy relationships for catalytic cracking React ants
Catalyst
P or P"
Ref.
Hammett equation Ethylbenzenes 2-Propyl benzenes
Aluminium fluoroborate Aluminosilicate
-9.5 -2.0 -1.0 -4.2 -3.6 -5.0 -2.8 -2.4
216,125 213,125 221,228 230 230 231,125 220,125 229
-9.7 -22.6 -5.0 -18.8 -22.4
232,125 210,125 227 204,125 204,125
Aluminosilicate + alkali Zeolite Kaolin Aluminosilicate
1,l-Diphenylethanes
T a f t equation Alkanes A1kylbenzenes
Aluminosilicate Aluminosilicate
o-Alkylphenols p-A1kylphenols
Aluminium fluoroborate Aluminium fluoroborate
~
_
_
~
_
.
_
_
_
_
~~
_
.~
~
315 catalyst [ 2341 ; the conversion increased in the order 2,2dimethylbutane
< hexane < 2-methylpentane < 3-methylpentane < 2,3-dimethylbutane.
There is disagreement in the literature on the influence of reactant structure on activation energy. Some authors (e.g. refs. 212, 213 and 227) have found different activation energies for different reactants but their treatment of the rate data was rather simplified and the rate coefficients obtained were not separated from adsorption coefficients. However, in the studies where the kinetic analysis was more detailed and a true rate coefficient was calculated, the same activation energy was determined for all members of a series of alkylbenzenes [210] and a series of alkylphenols [ 2041. 2.5.4 Mechanism
The mechanism of catalytic cracking was very soon recognised t o differ basically from that of thermal cracking, where free radicals are the active intermediates [ 31. The basis for the interpretation of the processes taking place on the surface of cracking catalysts were (i) the observation that these catalysts possess centres with high acid strength [235,236], (ii) the discovery of a strong similarity between the catalytic action of strong acids and aluminosilicates and (iii) Whitmore's theory of carbonium ion reactions [ 41. The distribution of products suggested the formation of carbonium ions on the surface, which then undergo further transformations according to the well known rules for the homogeneous acid-catalysed reactions [237]. However, while it was quite easy t o explain the way in which various products are formed from the starting hydrocarbon and even qualitatively the order of experimentally observed reactivities of isomers or homologues, for a long time it has been controversial as t o how the carbonium ions are formed. The hydride ion abstraction
suggested by some authors (e.g. ref. 2) seemed improbable t o others, who preferred the formation of the carbonium ions by addition of a proton from the surface t o an alkene present in the feed as an impurity [ 31 I I I I C=C + H' = -C+I t
A'
The experimental evidence for the second hypothesis was the observed increase of the cracking rate of alkanes after addition of small amounts of alkenes t o the feed. Both early theories assumed the continuation of the cracking reaction by intermolecular transfer of the charge from the products t o fresh starting molecules, that is like an ionic chain mechanism with the catalyst acting only as an initiator. The problem was further clouded by the fact that two types of acid centres exist on the surface, the Br@References p p . 385-398
316 sted sites (protonic) and Lewis sites (uncoordinated surface metal ions). In spite of great efforts, it was not possible t o find out the part which they play in the catalysts of the cracking and their relative proportions on the working surface. Only quite recently, an attempt was made t o explain the mechanism of catalytic cracking within the framework of the theory of olefin-forming eliminations outlined in Sect. 2.1 [238]. Before going into details, the acidity of the cracking catalysts will be considered. The close connection between the number of acidic sites and the activity was soon discovered [2,208]. Further studies revealed that the sites can have various acid strengths, which was attributed t o the difference in the nature of Brqhsted and Lewis sites. The experiments with catalysts partially poisoned by alkali showed that the sensitivity to the amount of the poison depends on the structure of the cracked hydrocarbon. The conversion of tert-butylbenzene decreased more slowly with increasing alkali content than did the conversion of isopropylbenzene [ 2391. The pretreatment of a zeolite catalyst with ammonia had no effect on the cracking of tert-butylbenzene, whereas the cracking of isooctane was strongly diminished [240]. Also, the change in total acidity caused by variation in the ratio of A1203t o SiOz had a smaller influence on the cracking of tert-butylbenzene than of 2-butylbenzene [238]. This interplay of catalyst acidity and reactant basicity is in excellent agreement with the dynamic model of elimination mechanisms presented in Sect. 2.1. The strongly acid centres are preferentially poisoned and the remaining sites are weak for the activation of the less polar molecules, e.g. isooctane. The alkylbenzenes, especially tert-butylbenzene, need a smaller activation for the reaction. The linear correlation, with a negative slope, of the reactivity of alkylbenzenes against the enthalpy of hydride ion abstraction from an alkane [227,240] supports the view that the ratedetermining step is the splitting off of the alkyl group. The predominance of trans-2-butene in the products of the cracking of 2-butylbenzene at lower temperatures indicates an at least partially concerted mechanism (E2) [238]. The cracking of alkylbenzenes can be treated as a case of aromatic electrophilic substitution (for recent views on this type of reaction see ref. 241) where the attacking agent is either a proton from a surface Brqinsted site or a coordinatively unsaturated surface cation acting as a Lewis site (cf. ref. 238)
I/
PH -C \
317 or I/ C-H
\/
-C I M' I
I
/
-C-C-H
+ J L + i M
.__c
The next step is the fission of the bond between the aromatic nucleus and the alkyl group; assistance of a basic site (surface oxygen atoms -0-, -0- or - O H ) which takes the leaving proton from the C, atom is very probable. The last step is the desorption of the arene. Alkanes are less basic than arenes and therefore, according t o the general rules for elimination reactions (see Sect. 2.1),we may expect a strong adherence to the E l mechanism. The suggested carbonium ion formation from alkanes by the action of a strong acidic site [2,237] has been doubted [ 3 ] but has obtained support from studies of the interaction of alkanes with very strong liquid acids such as HF-SbF, or FSO3-SbF5 [ 2421. The abstraction of a hydride ion from an alkane by these systems, which can act either as Brqhsted or Lewis acids according to the conditions, takes place even at very low temperatures. A complete separation of a carbonium ion from the hydride ion is very probably not necessary. It has been shown [73] by MO calculations that any attack by a charged species on an atom bonded t o a carbon atom causes activation of the bonds from a 0-carbon atom t o the substituents. In this way, the splitting of the C,-C, bond can be induced by adsorption of the alkane on a strongly acidic site. The preferential cracking of a saturated hydrocarbon chain in 0-positions t o the position where a carbonium ion might be formed was observed early and named the 0-rule by Thomas [ 2 ] . The question remains open as t o which type of acidic centre is able t o activate an alkane molecule. The fact that an aluminosilicate catalyst is poisoned for the cracking of alkanes by irreversibly adsorbed ammonia suggests a Lewis site [ 2401, viz.
&?
The charge which develops on the y-carbon atom after the rupture of the Cp-C, bond can either be neutralised by a hydride ion forming an alkane or the carbonium ion-like residue may react further, i.e. isomerise, crack etc. References p p . 385-398
318 The electrophilic displacement mechanism for catalystic cracking is further confirmed by the negative slopes of the Hammett and Taft correlations (Table 14) and has been supported by MO calculations in which electrophilic, nucleophilic and free radical mechanisms were compared [ 2281. The formal kinetic description of the cracking of alkanes and alkylbenzenes is in a good agreement with the above mechanistic considerations and structure effects on the rate. For the reaction of alkylbenzenes, the Langmuir-Hinshelwood type rate equation (3), which is based on the assumption of an adsorption equilibrium preceding the fission of the Carom.-Caliph.bond as the ratedetermining step (see ref. 21O), was adequate. On the other hand, the order of reactivities of isomeric hexanes (see above) [234] indicates that, with alkanes, the first step, the attack on the hydrocarbon molecule by an acidic site, might be the rate-determining process; for this reaction, the simple first-order rate equation was sufficient either in the original form or in the form expanded by a term expressing the retardation by alkenes [eqn. ( l l ) ] (cf. ref. 206). 2.6 DEHYDROSULPHIDATION
2.6.1 T y p e s of dehydrosulphidation reactions and catalysts Dehydrosulphidation of thiols and sulphides is directly analogous to dehydration of alcohols and ethers. Thiols (mercaptans) yield both alkenes and sulphides RSH = alkene + H2S
2 RSH
= RSR
+ H2S
and the ratio of the reactions depends mainly on the temperature [243]. Sulphides give alkenes and thiols which can further split off hydrogen sulphide [ 2441 RSR = alkene + RSH The reactions are reversible and the equilibrium is less favourable for the decomposition of a thiol t o an alkene and H2S than of the corresponding alcohol t o an alkene and H 2 0 under the same conditions; data on equilibria in the reactions of propene with water [245] and with hydrogen sulphide [ 2461 indicate that the equilibrium constant of propene hydration is smaller than that for propene sulphidation by approximately two orders of magnitude. The same catalysts as for dehydration are suitable for the dehydrosulphidation, i.e. alumina, silica-alumina, zeolites, metal oxides and metal sulphides. (For a comparison of their activities, see ref. 247.)
319 2.6.2 Experimental kinetic results
No detailed kinetic study of thiol or sulphide dehydrosulphidation has been reported, but first-order kinetics have been assumed for measurements with a pulse flow reactor and found acceptable [248,249]. Effects of structure on reactivity have been studied several times. The sulphides are more stable than the thiols [248,250]. In both series of thiols and of sulphides, the reactivity increases with the inductive effect of the alkyl group [248,251,252], in accordance with other elimination reactions. A linear relation between the logarithm of the rate coefficient and the enthalpy change on carbonium ion formation from the corresponding alkanes has been observed [ 2481. As Fig. 9 shows, linear correlation of the same rate data by means of the Taft equation is also possible. 2.6.3 Mechanism By analogy with other elimination reactions and on the basis of observed structure effects on the rate, the E l and E2-like mechanisms may also be accepted for the dehydrosulphidation. Sugioka and Aomura [ 248, 249,2531 have proposed mechanisms which correspond t o the above designations. Their results on ethanethiol decomposition over a series of
- 0.2
-0.1
6*
0
Fig. 9 . Taft correlation of the dehydrosulphidation of alkanethiols (line 1, p' = -33) and of dialkylsulphides (line 2 , p * = -38). (Data by Sugioka and Aomura [248].) References p p . 385-398
320
2
6
lo
x
14
Fig. 10. Dependence of the dehydrosulphidation rate of ethanethiol over zeolite catalysts with different cations on the electronegativity of the cation x (after ref. 249).
metal-exchanged Y zeolites [249] indicate a change in timing of the fission of the C-S and C,-H bonds with the nature of the metal. The reaction rate was influenced to different degrees by the addition of pyridine; n o influence was observed with NaY and the largest influence with A1Y. A linear relationship was found between the relative decrease and the electronegativity of the metal. The dependence of the rate coefficient on the electronegativity of the metal was volcano-shaped with a maximum for Zn2+and Cd2+(Fig. 10).
3. Addition reactions In this section, additions t o the multiple bonds C=C, C-C and C=O as well as t o the epoxide bond C-C are described. Sections 3.1-3.4 deal ‘0’ predominantly with the reverse reactions t o olefin-forming eliminations which are the subject of Sect. 2, viz.
I
I
C=C + HX I I
I
= -C-C-
I
I I H X X can be halogen, OH, O-alkyl, alkyl and aryl. The published data on
321 reverse reactions t o deamination and dehydrosulfidation (X = NR2, SH and SR) is so small, especially with respect to kinetics, that a review of these reactions is not included. All additions of HX t o the carbon-arbon double bond, treated here as a part of heterogeneous catalysis, can also be catalysed homogeneously by BrBnsted or Lewis acids like sulphuric acid o r aluminium chloride. Because of the thermodynamics, the additions require relatively low reaction temperatures and, sometimes, elevated pressures for good preparative results and the homogeneous alternative has had much more attention. Since the introduction of some heterogeneous catalytic processes into industrial practice (hydration of alkenes, alkylation of aromatic hydrocarbons by alkenes), the kinetics and mechanism of these types of additions have been studied. Several concepts developed originally for the elucidation of homogeneous additions have been accepted for the heterogeneous case but sometimes with doubtful results. Relatively little use has been made of the knowledge of elimination mechanisms over solids, which must have common features with addition mechanisms, because, in most cases, the same catalysts are active for processes in both directions. In Sect. 3.5, the heterogeneous aldol reaction I -C-H
\
I
I
+ C=O =-C--C+-H
I I I I is reviewed. This acid- o r base-catalysed addition t o the carbon-oxygen double bond is a well known example of homogeneous catalysis and a vast literature on its kinetics and mechanism exists. The heterogeneous catalysis of the reaction has less practical significance, but its study helps in the understanding of the general features of acid-base catalysis over solids. 3.1 HYDRATION
Hydration means, in general, addition of the elements of water to a substance. Most of these reactions are nm-catalytic or homogeneously catalysed processes. In this section, only hydration of olefins t o alcohols, of acetylene t o acetaldehyde, and of alkene oxides t o glycols will be treated, since they are typical reactions where the application of solid catalysts has become important.
3.1.1 Hydration of olefins to alcohols Hydration of olefins I I C=C + H 2 0 = <-CI \ I H is the reverse reaction t o \
References p p . 385-398
I
(A) I OH olefin-forming dehydration of alcohols dealt with
322 in Sect. 2.2. The reaction is exothermic with AHo ranging from about -38 t o -46 kJ mol-' in the vapour phase [ 2541 and may be accompanied by the formation of the corresponding ether, aldehyde and olefin oligomers. In contrast t o the formation of olefins from alcohols, which has never become a large-scale industrial operation but has attracted much theoretical interest, the preparation of alcohols by hydration of olefins is one of the most important petrochemical processes (see, for example, refs. 255-260); the kinetics and the mechanism of olefin hydration, however, have been less thoroughly investigated. Originally, the hydration of olefins t o alcohols was carried out with dilute aqueous sulphuric acid as the catalyst. Recently, the direct vapour phase hydration of olefins with solid catalysts has become the predominant method of operation. From the thermodynamic point of view, the formation of alcohols by the exothermic reaction (A) is favoured by low temperatures though even at room temperature the equilibrium is still in favour of dehydration. To induce a rapid reaction, the solid catalysts require an elevated temperatue, which shifts the equilibrium so far in favour of the olefin that the maximum attainable conversion may be very low. High pressures are therefore necessary t o bring the conversion to an economic level (Fig. 11).To select an optimum combination of reaction conditions with respect t o both rate limitation and equilibrium limitation,
I
I
I
I
225
250
275
300
50 xeq
40
30
20
10
200
325
Tem p ero tu r e ("C) Fig. 1 1 . Calculated equilibrium conversion, x e q (%), in the vapour phase hydration with equimolar mixture of ethylene and water (after ref. 2 6 7 ) .
the thermodynamics of reaction (A) were thoroughly investigated together with the catalyst activity studies, The equilibrium data can be found in the literature [ 261-2781. The catalytic hydration of olefins can also be performed in a threephase system: solid catalyst, liquid water (with the alcohol formed dissolved in it) and gaseous olefin [258,279,280]. The olefin conversion is raised, in comparison with the vapour phase processes, by the increase in solubility of the product alcohol in the excess of water [ 2581. For these systems with liquid and vapour phases simultaneously present, the equilibrium composition of both phases can be estimated together with vapourliquid equilibrium data [ 2811. For the three-phase systems, ion exchangers, especially, have proved to be very efficient catalysts [ 260,2801. With higher olefins (2-methylpropene), the reaction was also performed in a two-phase liquid system with an ion exchanger as catalyst [282]. It is evident that the kinetic characteristics differ according to the arrangement (phase conditions), i.e. whether the vapour system, liquid-vapour system or two-phase liquid system is used. However, most kinetic and mechanistic studies of olefin hydration were carried out in vapour phase systems.
( a ) Catalysts The best catalysts for olefin hydration are not necessarily those which have proved most satisfactory for the reverse reaction. Some of the successful hydration catalysts are not typical dehydration catalysts. The more obvious reasons are: (i) different adsorption characteristics of the catalyst is desirable, e.g. stronger adsorption of olefin relative to alcohol. (ii) under the conditions used for the hydration, ether formation cannot be suppressed as readily as in the dehydration, (iii) at high pressures, the olefins tend to polymerise much more than at the low pressures used for the dehydration. The first solid catalysts used for the direct olefin hydration were inorganic acids (H3P04,H2SO4,H3B03)supported on a porous material (“solid acids”). Heteropolyacids (e.g. silicotungstic), acid clays, zeolites and several acidic oxides were also used. Tungstic oxide ( W 0 3 ) , alone or mixed with other oxides, unreduced or in a reduced state (so-called “blue-oxide”), was often reported to be an efficient catalyst. Metal phosphates and metal sulphates were also investigated in several studies. Comprehensive reviews [ 254,2571 have been published, summarising the literature concerning the different types of solid catalysts used for olefin hydration. The use of synthetic ion exchanger resins as hydration catalysts has also been reviewed [283]. These catalysts are more effective than the inorganic solids [280, 284,2851 and allow the use of lower temperatures where the position of the equilibrium of hydration is more favourable. It is apparent that all the catalysts cited are acidic in nature. The relation between the acidity and activity of catalysts was investigated and Rcfercnces P P . 385-398
324 demonstrated, e.g. for boron phosphate [ 2861 and cation exchange resins [ 2841. With solid H3P04 catalysts, the ethylene hydration rate was found to be linearly dependent on the acidity of the supported phosphoric acid [287,288], and the rate decreased in the course of the reaction due t o the dilution of the acid by the reactant water [287]. The relation between acidic properties and catalytic activity was thoroughly investigated with metal sulphates and other acidic catalysts [ 289,2901. For ethylene hydration [290], a good linear correlation was found for the series of metal sulphates between the initial reaction rate and the number of centres of acid strength -8.2 < Ho < -3.0, whereas there was no correlation with the number of centres of other acid strength. For the hydration of propene [ 2891, which is considered to be more basic than ethylene, the best correlation was found with centres of -3.0 < H o < +1.5. Both Brqjnsted and Lewis acid sites are assumed t o be active, but a conversion of Lewis t o BrQnsted sites is probable under the reaction conditions (presence of water). With Si0,-supported metal sulphates (MX), the formation of Br#nsted centres according t o the reaction
I -Si-OH I
I
+ MX = -%OM + H' + XI
is possible. With cation exchanged zeolites, proton formation by interaction of water vapour with the cation seems likely, viz.
M"+
+ H 2 0 = [M(OH)]("pl)' + H'
For ethylene hydration [ 2901, a correlation between the catalytic activity of cation-exchanged zeolites and the electronegativities of the cations was established. It may be concluded from all these results that the presence of acid centres is unavoidable for a catalyst t o be active in olefin hydration. The possible role of basic centres is less clear; they might participate in a fast step which follows the ratedetermining step. ( b ) Experimental kinetic results
The several attempts, published in the literature, t o describe the kinetics of vapour phase olefin (mostly ethylene) hydration can be classified into two groups according t o the basic model used. One model, for reactions catalysed by phosphoric acid supported on solids, treats the kinetics as if the process were homogeneous acid catalysis and takes into account the acid strength of the supported acid. Thus, a semiempirical equation for the initial reaction rate [ 2881 f J
=
kh OfA
B
was used for ethylene hydration at 250-330°C
and 10-80 bar; ho is the
325 acid strength of phosphoric acid, f A the fugacity of ethylene and aB the activity of water. For the same reaction under similar conditions and on supported H3P04, another equation was applied [ 2873. I t is based on Taft’s mechanism for homogeneously catalysed olefin hydration [ 291,2921 [see scheme (B), p. 3271 according t o which water enters into the reaction scheme in a fast step which follows the ratedetermining step and therefore appears in the rate expression in a negative term, viz. =
khO(pA
-PR/PBKp)
(14)
where p A ,pB and pR are the partial pressures of olefin, water and alcohol, respectively. Since. the reactant water can dilute the phosphoric acid and change its acid strength ho (defined by log ho = -Ho where Ho is the Hammett acidity function), a relationship between the acid strength and the partial pressure of water was derived, which after substitution into eqn. (14)gives the expression (k’/Pg5)(PA- P R / P B K p ) (15) which was also used for reactor optimisation purposes [ 293-2951. The other approach is based on the Langmuir-Hinshelwood kinetics. In all the work using this approach, the surface reaction of adsorbed olefin and water was found, or postulated, to be the ratedetermining step. The corresponding rate equation has the form =
and fitted well the hydration data on W03-Si02 (265-305°C) [277], metal sulphates (160-220°C) [ 289,2901, boron and chromium phosphates (260-320°C) [ 296,2971 and solid phosphoric acid [ 2981. Speculatively, a very similar model was also proposed for 2-butene hydration over boric and phosphoric acids on alumina [299]. Whereas in the reaction over W03-SiOz the reactants were assumed t o be adsorbed t o an equal extent on the catalyst surface [ 2771, a preferential adsorption of water was found on boron and chromium phosphates 1296,2971. On metal sulphates [289], the adsorption of all three reaction components was so weak that eqn. (16) could be reduced to a second-order rate term for hydration and a first-order term for dehydration; viz. +
t
r = k PAPB- k PR Only in one case (the hydration of 2-methylpropene over a H,S04-SiOz catalyst [ 2781) was the so-called “Rideal mechanism” proposed as a preferable model and expressed by the rate equation for single-site adsorption with retardation by the product alcohol, viz.
R r fcrc,n CL’S u p . 3 8 5-39 8
326 With ion exchangers as catalysts for olefin hydration, special attention was paid t o transport problems within the resin particles and t o their effects on the reaction kinetics. In all cases, the rate was found to be of the first order with respect to the olefin. The role of water is more complicated but it is supposed that it is absorbed by the resin maintaining it in a swollen state; the olefin must diffuse through the water or gel phase t o a catalytic site where it may react. The quantitative interpretation depends on whether the reaction is carried out in a vapour system, liquidvapour system o r two-phase liquid system. In the vapour system [284, 2851, the amount of water sorbed by the resin depends on the H 2 0 partial pressure; it was found a t 125--170°C and 1.1-5.1 bar that 2-methylpropene hydration rate is directly proportional to the amount of sorbed water
(17) where pA is the partial pressure of olefin and u B the volume of water sorbed per equivalent of resin. It is believed that this effect was caused by changes in the ionisation of the resin acid groups, by the dependence of the olefin diffusivity on the degree of the resin swelling and by the influence of water upon the solubility of the olefin in the resin. The results are consistent with the Wheeler model for simultaneous diffusion through a spherical particle and a first-order reaction. In the liquid-gas [ 2801 or liquid-liquid [ 2821 systems where the pressure was high enough t o maintain water in the liquid state at the temperatures used, water was present in such a large excess within the resin phase, relative to the olefin concentration, that reaction rates independent of water concentration could be found; an order of one with respect to olefin (propene and 2-methylpropene) was again observed. High reaction rates of 2-methylpropene hydration [282] were found t o result from a high diffusivity of the olefin within the resin catalyst. The analysis of the diffusion model used revealed that this finding was consistent with the transport mechanism involving the surface diffusion of 2-methylpropene in an adsorbed state, There have been no kinetic studies performed with the aim of comparing the reactivity of olefins of different structures. Only for the threephase hydration over reduced tungsten oxide [ 2791 are conversion data reported; from these, the reactivity order at 270°C, C3H, > C4HB (mixture of isomers) 3 C2H4, could be estimated. However, at temperatures above 300"C, a higher yield of alcohol can be obtained from ethylene than from the other two olefins. 2-Methylpropene is reported to be about one hundred times more reactive than n-butenes in vapour phase hydration at 115°C over a H2S04-Si02 catalyst [278]; for the butenes, the reactivity order cis-2-> 1->truns-2- has been established. The addition of water t o alkenes proceeds according t o the Markownikoff rule, i.e. 2-propanol from propene [280,279,289], 2-butanol from 1-and 2-butenes =kpA%
327 [ 278,279,2991 and 2-methyl-2-propanol from 2-methylpropene [ 278,282, 2851 are formed.
( c ) Mechanism
It has been assumed that, with phosphoric acid-based catalysts where the active component is liquid aqueous phosphoric acid adsorbed in the pores of the support, the reaction probably follows the scheme proposed by Taft [ 291,2921 for the hydration of olefins in aqueous solution, viz.
(fast)
I I - H-C-C-O'-H I I I
n-complex
L
+H2O &
-H20
4 5 2 0
+H20
(fast)
H
(rate-de termining step)
I I H-C-C-OH
I
+ H30'
I
(fast)
Adopting this mechanism and making some simplifying assumptions, Gelbsthein et al. [287] derived the kinetic equation, eqn. (14), or its alternative form, eqn. (15), mentioned in Sect. 3.1.1.(b) which fitted the experimental data on H3P04-Si02 catalyst well. The Langmuir-Hinshelwood rate equation (16) used by several authors for other types of catalyst was interpreted by Tanabe and Nitta [ 2901 on the basis of their results with NiSO,. The authors assume that the surface reaction of adsorbed ethylene and water molecules, which was found t o be the slowest process, may be written in greater detail as
It has been concluded from deuterium exchange experiments, using ethylene and heavy water, that the addition of an adsorbed proton t o adsorbed ethylene is the actual rate-determining step. It can be seen that the two schemes differ, mainly in that the latter includes dissociative adsorption of water on the surface of the catalyst and does not specify the adsorption of ethylene, but they are consistent in that they assume the formation of a carbonium ion as the ratedetermining step. 3.1.2 Hydration of acetylene t o acetaldehyde
The reaction CH=CH + H 2 0 = [CH2=CHOH] = CH3CH0 References p p . 385-398
328 is exothermic (M"= -150 to -167 kJ mol-' [300,301]). The isomerisation of the intermediate vinyl alcohol to acetaldehyde makes the over-all process practically irreversible since the equilibrium between the enol and carbonyl forms is strongly shifted in favour of the latter. The formation of acetaldehyde from acetylene is accompanied by several side reactions, the most important of them being the aldolisation of acetaldehyde with subsequent dehydration t o crotonaldehyde, and the polymerisation of acetylene. For a long time, only a liquid phase process was employed industrially for the hydration of acetylene t o acetaldehyde; mercury salts in acidic solution were used as catalysts. Only recent reports can be found in the literature (e.g. ref. 300) on the industrial utilisation of the direct vapour phase hydration of acetylene over solid catalysts. It has been reported that solid acids and oxides or salts of different metals can catalyse the vapour phase hydration of acetylene. Most typical are phosphoric acid and phosphates of bivalent metals, such as Zn or Cd. Organic ion exchangers and synthetic zeolites exchanged for Zn2+,Cd2+, Hg2+ and Cu2+ions were also employed. A survey of inorganic catalysts [254] or of organic ion exchangers [283] catalysing the hydration of acetylene or its derivatives can be found in literature. The temperatures reported in the kinetic studies range from 260 to 350°C. In most of the investigations, the hydration rates were found t o be of the first order with respect to acetylene [300,302-3041. With zinc phosphate [ 3031, cadmium-calcium phosphate [ 3001 and cation-exchanged zeolites [ 3041, the rates were independent of the concentration of water. Thus the simple kinetic equation
r = kpA (18) where p A is the partial pressure of acetylene, is valid in these cases. However, with phosphoric acid supported on carbon [ 3021, an increase of water partial pressure caused a decrease of the reaction rate. This was explained in a similar way as in ethylene hydration over a H3P04-Si02 catalyst [ 2871, i.e. by dilution of the supported acid with the water present in the reactant mixture (p. 325). The acidities, h o ,corresponding t o different partial pressures of water could be evaluated and it was established that the reaction rate could be then expressed by the equation (19) r = kh#A = k'p, Thus, the partial pressure of water is not really involved in the rate expression for the reaction catalysed by supported phosphoric acid, as for the other mentioned catalysts [ 300,303,3041 [ eqn. (IS)]. It was proposed [302] to explain this form of reaction kinetics on the basis of a homogeneous mechanism. The authors assume that the reaction proceeds in the film of phosphoric acid containing dissolved acetylene and they adopt the reaction scheme of Taft [291,292] for the hydration of
329 olefins in such a way that a carbonium ion CH2=C+His formed from a .rr-complex of acetylene; the other steps are formally analogous t o those in scheme (B). In the catalysis by zinc phosphate [303] instead of vinyl carbonium ion, the formation of a corresponding carbonium ion H--C'=C-H
I
Zn' is assumed which then reacts with water in the same way as the ion CHI = C'H. Another explanation was offered for the C d - C a phosphate catalyst [ 3051 ; the participation of acidic groups of the phosphate in the reaction mechanism was assumed, since a dependence of the activity of the catalyst on its acidity had been found. With cation-exchanged zeolites [ 3041, the first-order kinetics [eqn. ( I S ) ] is explained by the degeneration of the Langmuir-Hinshelwood equation for monomolecular transformation of adsorbed acetylene in the rate-determining step
r=
~KAPA
1+ C K i p i when the adsorption of all the reaction components is weak, i.e. CKipi << 1. Quite another type of kinetics was found for a zinc phosphate-phosphoric acid-activated carbon catalyst at 350°C [ 3061. The rate equation
where p B is the partial pressure of water, found experimentally, can be interpreted by assuming that the reaction occurs when acetylene in the gas phase attacks adsorbed water and a steady state is established in which the rate of the removal of water by chemical reaction is equal t o the rate of the adsorption of water. However, the same form of rate equation would result if it were supposed that water vapour attacks adsorbed acetylene. The latter interpretation would be more consistent with the mechanism based on the Taft scheme, according to which non-adsorbed water reacts with a carbonium ion formed from the n-bonded unsaturated compound. 3.1.3 Hydration of alkene oxides t o glycols
The reaction R-CH--CHz + H20=R--CH(OH)-CH20H \ ' 0 is exothermic and irreversible. Most investigations were carried out with ethylene oxide ( R = H) for which AHo = -96 kJ mol-' and log Kp 11 at 25°C and = 3 a t 300°C [ 3071. Usually, the reaction does not stop at the Refererices P P . 385-398
330 stage of monoglycol but this can react further with other molecules of alkene oxide and di, tri- and polyglycols may be formed. Other side reactions which may accompany the main reaction are isomerisation of the alkene oxide t o the corresponding aldehyde or the polymerisation t o ether polymers. In homogeneous liquid media, the hydration is catalysed by acids or bases. With solid catalysts, the reaction may be performed with both reactants either in the vapour [285,307-3111 or in the liquid [312-3141 phase. In the former case, temperatures of about 12O-25O0C, and in the latter case of about 25-90°C, were used. The reaction can also be conducted under conditions which establish a heterogenous system of vapour and liquid phases [ 3081. The solid catalysts for the alkene oxide hydration can also be of an acidic or basic nature. Inorganic solids such as SiO,, A1203,T h o 2 , supported H3P04, silica-alumina and molecular sieves did not prove t o be efficient catalysts [ 307,3111. Good results were obtained with silver oxide on an alumina carrier [307]. More intensively than the inorganic solid catalysts, however, organic polymer ion exchangers were investigated and used as catalysts for the alkene oxide hydration [285,308-3141. The catalytic activity of cation exchangers depended on their acid strength : sulphonated resins were much more active than those of the phosphonic o r carboxylic acid type [ 308,309,313,3141. Anion exchangers were found t o be less active than cation exchangers [313]; a strong base (with an amine group), however, proved to be a good catalyst [ 3081. Formal kinetic investigations (performed only with acidic ion exchange catalysts) revealed, in most cases, the first-order rate law with respect t o the alkene oxide [285,310,312] or that reaction order was assumed [ 309,3111. Strong influence of mass transport (mainly internal diffusion in the polymer mass) was indicated in several cases [285,309, 310,312,3141. The first-order kinetics with respect t o alkene oxide is in agreement with the mechanism proposed for the same reaction in homogeneous acidic medium [ 309,315-3171, viz. CH,\ rH2;O+HfCH2
1
CH,
P'H+
C'H2
I"'"/
CH20H
CHzOH
CH20H
+ H'
(C)
if the rate-controlling step is the decomposition of a protonates substrate t o a carbonium ion. This mechanism represents a classical S N 1 substitution; however, an alternative mechanism was proposed [ 317,3181 which is today most widely accepted for homogeneous acid catalysis, viz.
331 This mechanism does not involve the formation of a free carbonium ion but rather a nucleophilic displacement on a carbon atom in the oxonium complex. The reaction thus becomes a limiting case of a SN2substitution. There are, however, very few reliable kinetic measurements available which would allow the effect of water on the reaction kinetics to be determined and thus distinguish between the two alternative mechanisms in the catalysis by ion exchangers. The role of water is rather complex here: it is not only a reaction partner but it also greatly modifies the catalyst as has already been mentioned in connection with ion exchanger-catalysed hydration of olefins [Sect. 3.1.1.(b)). Water solvates the resin acid groups, causes swelling of the polymer and thus strongly affects the transport situation in the catalyst particle [285,310]. Therefore, the true reaction order with respect to water could not be established. It has been found rather empirically that the reaction rate is proportional to the volume of water retained by the resin catalyst [285,310], as in the hydration of olefins [eqn. (17)]. With ethylene oxide, the volume of sorbed water could be expressed by the BET isotherm which was substituted into eqn. (17) thus relating the reaction rate also to the water partial pressure, pB, ViZ.
where a is a constant and pB,sis saturation vapour pressure of water. The strong effect of transport processes in the polymer particle was further evidenced by the low values of the activation energy [310,312] and by the dependence of reaction rate on particle size [310], on the outer surface area of the resin particles [312), on the crosslinking degree of the copolymer [314] and on the cations introduced [314]; moreover, the lower rates observed with ion exchangers than with an equivalent of sulphuric acid [312,314] indicated that only a part of the acid groups in the polymer was easily accessible. The macrokinetic situation may therefore be characterised as follows: the resin particle is in a more or less swollen state (in the liquid system as well as in contact with water vapour) retaining a considerable amount of water. The alkene oxide is adsorbed on contacting the resin surface, probably by protonation; direct adsorption measurements [ 309 J revealed the occurrence of both physical adsorption and chemisorption. The alkene oxide then reacts on the surface with water according to mechanism (C) or (D) or diffuses within the swollen resin mass [ 3121 to the internal acid groups where it reacts in the same way. With an inorganic catalyst (Ag,0/A1203 [ 3071) a simple surface process is assumed for the vapour phase hydration: water adsorbed on the surface of silver oxide reacts with gaseous ethylene oxide to form adsorbed glycol which is then desorbed; this is obviously an oversimplification of the actual mechanism. References p p . 385-398
3.2 HYDROHALOGENATION
3.2.1 Types of hydrohalogenation reactions and catalysts The addition of hydrogen halides t o unsaturated organic compounds is called here hydrohalogenation in order t o stress that it is the reverse reaction t o dehydrohalogenation. Two types of the hydrohalogenation reaction have to be considered, the addition to a carbon-carbon triple bond --C=C-
+ HX = -CH=CX-
(A)
and the addition t o a double bond I I <=C-
I
I
+ HX = -CH-CX-
(B)
For both reactions, low temperatures and elevated pressures are favourable from the point of view of equilibria but high conversons of acetylene can be obtained even at atmospheric pressure and around 150°C. The same catalysts are active for the hydrohalogenation as for the dehydrohalogenation (see Sect. 2.4). Metal halides are preferred [ 319,3201 but other metal salts [ 3211 and alumina [ 3221 may be also used. When metal halides are the active components of the hydrohalogenation catalysts, they should have the same halogen as the reactant HX because exchange is easy [ 322,3231. For reaction (A), other catalysts are more active than for reaction (B). This difference and the less favourable equilibrium of the second step allow the reaction of acetylene with hydrogen chloride to be stopped at the stage of vinyl chloride. HgC12 is the proven catalyst for this industrially important process, with active carbon serving as the carrier. When a mixture of HgCl, and ZnC12 is used, both vinyl chloride and 1,l-dichloroethane are obtained [ 3201. 3.2.2 Experimental kinetic results The reported rate equations for the hydrohalogenation of acetylene, ethylene, propene and vinyl chloride are summarised in Table 15. Of special interest is the last entry; it is based on a model which assumes two types of active centres, the first one for the adsorption of acetylene, the second for the adsorption of hydrogen chloride and vinyl chloride. The addition of HX to alkenes proceeds according t o the Markownikoff rule, i.e. the halogen is attached to the more substituted carbon atom [321-3231. The reactivity order of butenes was found t o be dependent on the nature of the catalyst. Over MgS04, the order was isobutene > trans-2-butene > l-butene > cis-2-butene but with CaCI, , the reactivity decreased in the order isobutene > cis-2-butene > l-butene > truns-2butene [321]. Propene is more reactive than ethylene [318]. Earlier reports that tert-butylchloride is formed from l-butene and hydrogen
333 TABLE 15 Rate equations for hydrohalogenation reactions @ A , p B and pc denote partial pressures of the unsaturated compound, hydrogen halide and halogenated hydrocarbon, respectively.) ~
Reactants
Catalyst
r = k P d B -k ' p c 1-Butene, 2-butene (Y isobutene + HCI
MgSO, or CaC12
25-150
324
HgCl2
75-180
325* 326**
r = kpQgf5 Acetylene + HCI
Temp. ("C)
2 56-302 80-140
Ref.
327 328 329,330
Acetylene + HCI
A1203
256-302
Ethylene + HCI Propene + HCl Vinyl chloride + HCI
HgC12 ZrOCl2-Si02 A1203 ZnCl2
121-293 44-80 102-148
327 33 1 332 333 334
75-125
335
r = ~KAKBPAPBI(+ ~ KBPB Acetylene + HCl
+
KCpC)(1+ KAPAI
HgCl2
* a = 0.65-0.89, b = 0.25-0.38. ** a = 0.80, b = 0.37. chloride [336,337] were corrected later by the same group of authors [ 3211 ; 2-chlorobutane is the sole product as expected.
3.2.3 Mechanism Only a few speculations concerning the mechanism of the hydrohalogenation reactions can be found in the literature. They are mostly based on the better knowledge of the reverse reaction, the decomposition of haloalkanes. It seems evident that catalytic hydrohalogenation must also involve transfer of paired valence electrons and that a free radical-like process is highly improbable. The existence of .rr-complex intermediates has been suggested [ 321,336,3371 but the hypothesis lacks experimental evidence. Adsorption studies revealed that both reactants can be adsorbed on active catalysts. For the reaction of acetylene with hydrogen chloride, the References PP. 385-398
334 activity order of a series of metal halides was the same as the adsorptivity order for both components, HgC12 being the best catalyst and adsorbent [321,338]. 3.3 ALKYLATION BY OLEFINS
3.3.1 Types of alkylation reactions and catalysts Alkylation is a very broad reaction type and it can, depending on the nature of the alkylating agent, proceed either as a substitution or as an addition reaction. The alkylation by substitution of, for example, aromatic hydrocarbons, phenols or amines is based on the reaction with alkyl halides or alcohols. Some evidence indicates that, at least partly, the alkylation proceeds through the intermediate formation of alkenes from the alkylating agent when the reaction is conducted at atmospheric pressure and at high temperature. In this section we shall deal only with the reverse reaction t o dealkylation by cracking (Sect. 2.5), that is with additions of alkanes and aromatic compounds to the carbon-carbon double bond. The former reaction is described only in a single paper [ 3391; the formation of 2,2,4-trimethylpentane from isobutane and isobutene CH3 CH3 I I CH3+H + CH2=C-CH3 = CH,--C--CHZ--CH-CH3 I I I 1 CH, CH3 CH3 CH3 was catalysed by the strongly acidic rare-earth-exchanged crystalline aluminosilicates (zeolites) in the temperature range 25-lOO"C, and in all its features resembled the same reaction catalysed by H2S04 or HF. The alkylation was accompanied by polymerisation of isobutene, hydride transfer reactions etc. The reaction of aromatic compounds with alkenes giving alkylaromatic compounds has obtained more attention. A typical transformation is the alkylation of benzene by lower alkenes, e.g.
Not only benzene and its alkyl derivatives can be used as the aromatic component but also naphthalene [ 3411, phenol [ 340,3411 and thiophene 1341,3421. Low-molecular weight alkenes, C1 to C,, cyclohexene and dodecene have served as alkylating agents. Only strongly acidic solids can catalyse the heterogeneous alkylation of aromatic compounds. Amorphous aluminosilicates were the first catalysts
335 used for this purpose [342-3441 but they are distinctly less active than the zeolites [ 341,345-3501. Solid phosphoric acid [ 3431, iron phosphate [343], aluminium oxide activated by BF3 [351] or B203 [352] and sulphonated polymers [ 340,353-3551 are also suitable catalysts. The activity of the Al2O3-SiO2 catalysts can be enhanced by adding CC14 or other organic halogen compounds into the feed [ 3561. Both liquid phase and gas phase alkylations over solid catalysts under atmospheric and elevated pressures have been described. In order to achieve high conversions, it is necessary to operate at low temperature and pressures above atmospheric.
3.3.2 Experimental kinetic results Only a few authors have attempted t o describe the rate of the alkylation by kinetic equations [ 349,352,355-3571. Table 1 6 shows the conditions of experiments and the rate equations applied by various groups of authors. N o conclusions can be made on the basis of this small set obtained with different approaches. More information is available about orientation, when a second alkyl group is introduced into the aromatic ring, and about relative rates. As might be expected, propene reacts more easily than ethylene [ 342,3461 and isobutene more easily than propene [ 3421. Normal butenes are sometimes isomerised in the process; practically the same product composition, consisting mainly of 2,2,4-trimethylpentane, is obtained in the alkylation of isobutane whether the olefin component is isobutene or 2-butene [339]. In the alkylation of aromatic hydrocarbons, this side reaction is negligible. Toluene is more reactive than benzene [ 3581 as is phenol [ 3481. Alkyl groups and,hydroxyl direct the new entering alkyl group mostly into the TABLE 16 Rate equations for heterogeneous catalytic alkylation of aromatic hydrocarbons Reactants
Conditions and catalyst
Benzene + ethylene Benzene + propene
Gas phase, 300"C, 20 bar, A1203-Si02 Liquid phase, 55"C, 1 bar, ion exchanger Liquid phase, 1O-5O0C, 1 bar, A12 03-B2 0 3 Gas phase, 423-483"C, 1 bar, Al203-Si02
2-Propylbenzene + propene a
Rate equation
a
Indices denote: A aromatic hydrocarbon, B alkene, R product(s), S solvent. The second term in the denominator describes t h e swelling of the polymer.
References p p . 385-398
Ref.
357
356 352
349
336 ortho and para positions [ 342,351,3581. The ortho/para ratio decreases with the size of the alkyl group: CH3 > C2H5> 2-C3H, [347]; it is also influenced by the nature of the catalyst [ 343,3591. The metdpara ratio determined for the ethyiation of toluene at low conversions (in order to supress the consecutive isomerisation of the ethyltoluenes formed) fitted the Brown selectivity relationship for electrophilic aromatic substitution well [ 3581. This relationship correlates relative reactivity of toluene with respect t o benzene and the ratio of reactivities of the meta and para positions of toluene [ 3601 ; many homogeneous substitution reactions conform t o it. However, there are some contradictory reports on the composition of the products of toluene alkylation or benzene dialkylation at high conversions. In some cases, compositions corresponding to the thermodynamic equilibrium between ortho, meta and para isomers were found, and in other cases, kinetic control of orientation, giving mostly the ortho + para substitution, prevailed. Consecutive isomerisation of the ortho and para isomers t o the more stable meta isomer seems t o be the cause of the disagreement. More active catalysts gave more meta derivatives than the less active ones [343] and increasing the temperature has the same effect [ 3511. 3.3.3 Mechanism The observed effects of structure on rate and on orientation, confirmed by the Brown selectivity relationship, show that there is no basic difference between heterogeneous catalytic alkylation of aromatic compounds and homogeneous electrophilic aromatic substitution, cf. nitration, sulphonation etc. This agreement allows the formulation of the alkylation mechanism as an electrophilic attack by carbonium ion-like species formed on the surface from the alkene on Brqhsted acidic sites. The state of the aromatic compound attacked is not clear; it may react directly from the gas phase (Rideal mechanism ) [348] or be adsorbed weakly on the surface [ 3591. It seems that other acidic sites are the most efficient for the alkylation of aromatic compounds than for the reverse reaction, the cracking of alkylaromatic compounds [ 3611. For the forward process, a linear correlation was observed between the activity of decationized Y zeolites and the number of acidic sites corresponding t o H o < + 3.3, whereas for the cracking, the sites corresponding t o Ho < -3.0 correlated with the activity. 3.4 ADDITION OF ALCOHOLS TO ALKENES
The reaction
R’ R3 R’ R3 I I I I C=C + R’OH = H-C-C-0-R’ 1 1 I I R2 R4 RZ R4
R’ R3 I I + R’-O-C-C-H I I RZ R4
337 which is an old method for the preparation of alkyl-tert-butylethers by the action of isobutene on an alcohol or glycol catalysed by H2S04 [362, 3631 has only recently been conducted with heterogeneous catalysts. A great number of patents deal with suitable solid catalysts and conditions for this reaction. Ion exchange resins seem t o be the best catalysts, but aluminosilicates may be also used though with lower selectivity [ 3641. The principal side reactions are oligomerisation of the alkene and dehydration of the alcohol t o the ether. The addition itself is the reverse reaction to the first step of the dehydration of ethers t o olefins (see Sect. 2.2). Two recent papers report the main features of the heterogeneously catalysed addition of alcohols t o alkenes [ 364,3651. The reaction proceeds both in the liquid and gas phase [ 3641, and the temperature must be kept well under 150°C with respect t o the position of the equilibrium [ 3641. The reactivity of isobutene and 2-methyl-1-butene is much higher than that of propene, 2-butene and 3-methyl-1-butene [364,365]. 2-Methyl-1-butene reacts faster than 2-methyl-2-butene [ 3651. The reactivity of alcohols with isobutene decreases in the order methanol > ethanol > 1-propanol > 1-butanol [ 3651. The initial rate in the liquid phase reaction is zero-order in methanol, first-order in isobutene and about third-order in the sulphonic groups of the ion exchanger [ 3651. A comparison of catalysis by an ion exchanger and anhydrous p-toluenesulphonic acid indicated an efficiency of between 5 and 8, depending on temperature, according t o the Hammett definition [366]. The observed structure effects are similar, as in the reaction catalysed by sulphuric acid. On this basis and with the notion of the strong acidity of the heterogeneous catalysts used, it is possible to assume a mechanism similar to olefin hydration (Sect. 3.1) or alkylation (Sect. 3.3). Olefin protonation by the catalyst seems to be the first step, which is followed by the interaction with the nucleophile, in this case the alcohol. 3.5 ALDOL CONDENSATION AND RELATED REACTIONS
3.5.1 Types of reaction
In this proper sense, aldol condensation includes reactions producing P-hydroxyaldehydes or P-hydroxyketones by self condensation or mixed condensation of aldehydes and ketones; these reactions are, in fact, additions of a C-H bond activated by the carbonyl t o the C=O bond of the other molecule, viz.
R4 \
R3
I
CH-C=O +
I
R5' References p p . 385-398
R'
\
R1OH R4 R3 \I
I I
I
c=o= c - c - c = o
I R2
/ R2
R5
338 where R’ to R5 are H, alkyl or aryl. Depending on the conditions, the reaction stops either at the stage of the P-hydroxy carbonyl compound or, if R4 or R5 are hydrogen atoms, this is further transformed (dehydrated) t o a ,P-unsaturated aldehydes or ketones
R1 R4 R3 \ I I C--C--C=O = C=C--C=O + H 2 0 I / I R2 R2 H R L O HR4 R3
\I
I
I
In several cases, the intermediate hydroxy compound may be formed in an undetectable amount so that the reaction appears as a direct formation of a ,P-unsaturated carbonyl compound by condensation of two carbonyl compounds R1 R1 R4 R3 \ \ I I R4CH2-C=0+ C=O = C=C-C=O + H 2 0 I I R2 R2
R3
I
In a broader sense, the term aldol condensation has sometimes been applied t o many so-called “aldol-type” condensations involving reactions of an aldehyde or ketone with a substance containing a mobile hydrogen, namely R4(or X) CH2Y +
R1 OH R4(or X) \I I C=O = /C--CHY
R1\
I
RZ R4(or X)
R1 \ I = ,C=CY
RZ
+ HzO
RZ where X or Y is an activating group such as COOR, CONHR, CN, NO*, S02C H3;in a similar way t o the true aldol condensations, these reactions may produce a hydroxy compound or its dehydration product. The condensation represented by reaction (D) is very probably involved as the first step in the synthesis of bis( ary1)alkanes from carbonyl compounds and substituted benzenes (e.g. phenol)
This reaction has become industrially important as a large scale process for the production of bisphenol A [ 2,2-bis(p-hydroxyphenyl)propane] from acetone and phenol.
cl,
TABLE17 Equilibrium data for some aldolic reactions Reactants
Product
Reaction scheme
Reaction conditions
Equilibrium constant
Acetaldehyde
fi-Hydroxybutyraldehy de
(A)
Aqueous medium, 25°C
5.7 x 10'4 x 102
Acetaldehyde
Crotonaldehyde
(C)
Vapour phase, 25°C
6.2-32.5
Acetone
Diacetone alcohol
(A)
Liquid phase, 20-30" C
-10-2 -104
Vapour phase
0.19 (300 K ) 2.2 (600 K )
Acetaldehyde + benzaldeh yde a
Cinnamaldehyde
(C)
A€r
Ref.
(kJ mot-')
-41.2
367
-9.7
a
368 C
369 ~
_
_
Calculated from thermochemical data in ref. 367. Evaluated from experimental data. Calculated from free energy data in ref. 368, obtained by t h e group contribution method.
w
w
co
Aldol condensations are reversible and slightly exothermic reactions. The values of equilibrium constants and reaction enthalpies of some aldol reactions reported in the literature are listed in Table 17. Aldol condensations were originally carried out in the liquid phase and catalysed homogeneously by acids or bases; this way of operation is still predominant. Solid-catalysed aldol reactions can also be performed in the liquid phase (in trickle or submerged beds of catalyst), but in many cases vapour phase systems are preferred; the factors determining the choice are the boiling points and the stability of the reactants at elevated temperatures. At higher temperatures, the formation of (Y ,P-unsaturated aldehydes or ketones [reactions (B) and (C)] is preferred t o aldol (ketol) formation [reaction (A)]. A side reaction, which may become important in some cases, is the self-condensation of the more reactive carbonyl compound if a mixed condensation of two different aldehydes or ketones is occurring. The Cannizzaro reaction of some aldehydes or polymerisation t o polyols or other resin-like products can also accompany the main reaction.
3.5.2Catalysts The solid substances catalysing aldol condensations are similar to the homogeneous catalysts in that they may be of an acidic or basic nature, but the latter are preferred. Alkali and alkaline earth metal hydroxides o r phosphates (supported or unsupported), Ca- or Sr-exchanged zeolites and anion exchange resins are typical examples of efficient base catalysts. As acidic catalysts, cation exchanger resins and zeolites in the hydrogen form were used; calcium hydrogen phosphate was also assumed t o act as an acidic catalyst. Condensations of carbonyl compounds with simple aromatic compounds giving bis(ary1)alkanes [reaction (E)] represent a particular case where acidic ion exchange resins are the most successful catalysts. The use of ion exchange resins (mainly of basic types) as catalysts for various types of liquid phase aldol condensation and related reactions has been reviewed [ 370,3711. The effect of the basicity of aldol condensation catalysts on their activity was thoroughly investigated by Malinowski et al. [ 372-3791. The observed linear dependence of the rate coefficients of several condensation reactions on the amount of sodium hydroxide contained in silica gel (Figs. 1 2 and 13) supported the view that the basic properties of this type of catalyst were actually the cause of its catalytic activity, though the alkali-free catalyst was not completely inactive. The amphoteric nature of the catalysis by silica gel, which can act also as an acid catalyst, was demonstrated [380]. By a stepwise addition of sodium acetate t o a HN0,-pretreated silica gel catalyst the original activity for acetaldehyde self-condensation was decreased to a minimum (when an equivalent amount of the base was added); by further addition of sodium acetate, the activity increased again because of the transition t o a base
341 I
I
I
1
I
i
0.8
0
OD 2
OD4
0.00
0.06 mN13
Fig. 1 2 . Dependence of apparent rate coefficient, k (sec-I), on sodium content, m N a (mol Na per 100 g cat), in silica gel catalysts for the vapour phase,condensation of formaldehyde with (1)acetaldehyde, (2) acetone, ( 3 ) acetonitrile, a t 275OC [ 3721.
mNCi Fig. 1 3 . Dependence of apparent rate coefficient, k (sec-I), on sodium content, mNa (mol Na per 100 g cat), in silica gel catalysts for the vapour phase condensation of acetaldehyde with (1) formaldehyde, ( 2 ) acetaldehyde, ( 3 ) benzaldehyde, at 3OO0C [376]. References PP. 385-398
342 type of catalysis. The groups -Si-ONa in the catalyst were assumed t o be the active sites [372,378]. The effect of different alkali metal ions was compared and the activity order Na < K < Cs established [377]. When ion exchange resins were employed as catalysts in liquid phase condensations, a dependence of the activity on the base strength of their functional groups was observed. In self-condensations of some aldehydes and ketones, only the strongly basic ion exchangers, such as Amberlite IRA-400 were active [ 381,3821. In a more detailed study [ 3831, an activity decrease of the functional groups in the order CN > OH > C1 was found; the acetate group was found t o be inefficient. A non-linear relatinship between the exchange capacity and the reaction rate was observed [384], from which it could be concluded that only part of the hydroxyl groups of the resin was accessible and took part in the catalytic action. In benzaldehyde-acetophenone condensation [ 3701, a decrease of the rate coefficient with increasing degree of crosslinking of the resin was found. With acidic ion exchangers, the same effect of the resin degree of crosslinking was observed [ 385,3861 for bisphenol A synthesis [reaction type (E)]. These results indicate that the accessibility of the resin functional groups and the transport phenomena may become important factors in the kinetics of condensation reactions catalysed by ion exchangers (see also ref. 387), as will be discussed in greater detail in Sects. 4.1.3, 4.2.1 and 4.2.2 in connection with esterification and hydrolysis.
3.5.3 Experimental kinetic results
(a) Formal rate equations As with homogeneous aldol reactions, simple power-type rate equations have been frequently used t o describe the kinetics of solid-catalysed condensations. For several liquid phase reactions, second-order kinetics was established, viz. r
=
kcAcB
Examples are the formation of diacetone alcohol from acetone. [reaction type (A)] catalysed by barium or strontium hydroxide at 20-30°C [368] o r by anion exchange resin at 12.5-37.5"C [387], condensation of benzaldehyde with acetophenone [type (C)] catalysed by anion exchangers at 25-45°C [ 3701 and condensation of furfural with nitromethane [type (D)] over the same type of catalyst [ 3841. The vapour phase self-condensation of acetaldehyde over sodium carbonate or acetate at 50°C [388], however, was found to be first order with respect t o the reactant. Langmuir-Hinshelwood-type equations were applied in some cases. The kinetics of t h e vapour phase condensation of acetaldehyde with formaldehyde t o acrolein at 275-300°C over sodium-containing silica gel
343 [ 373-3751 was interpreted by means of the equation
where the subscripts A and B denote the reactants and R and S the products (acrolein and water). The equation was derived on the assumption that acetaldehyde is adsorbed on a basic site and reacts with formaldehyde from the gaseous phase (by analogy with the mechanism of homogeneously catalysed aldol condensation). The equation was slightly modified for the self-condensation of acetaldehyde on the same catalyst [ 3791. On the basis of a series of kinetic studies of aldolisation reactions in which either the hydrogen donor (acetaldehyde, acetone, acetonitrile) [ 3721 or the hydrogen acceptor [ 3761 (formaldehyde, acetaldehyde, benzaldehyde) and the catalyst basicity [ 372-374,378,3791 were systematically varied (Figs. 1 2 and 13), the authors have concluded [376] that the rate coefficient is proportional t o the catalyst basicity, K B , the acidity of the hydrogen donor, K H , and t o a factor Y which is related t o the oxygen basicity of the acceptor, viz.
k = /3KBKH Y (22) where p is the proportionality constant. A similar relation has been proposed and discussed on the basis of a kinetic study of benzaldehyde condensation reactions [ 3701. For the liquid phase ketolisation of acetone t o diacetone alcohol over barium hydroxide-sodium silicate or barium hydroxidesodium hydroxide-borax at 12--40°C [ 3891, the equation
r=
k(cA - CR/K) 1 + K A c ~+ K ~ c ,
from a set of seven Langmuir-Hinshelwood equations, was found t o give the best fit t o the experimental data. It is consistent with a model, according to which the desorption of the product is the ratedetermining step. The kinetic data for phenolacetone condensation t o bisphenol A [reaction type (E)] in the liquid phase at 9 1 ° C over sulphonated ion exchanger [ 3851 were best represented by the equation
kcAck r=(1+ KAcA + KBcB + K s p s + K ~ C , ) ’ where A is acetone, B phenol, S water and I methylcyclohexane (added as a non-polar substance). The equation corresponds t o a model, in which the surface reaction is the ratedetermining step and the reactants, water and methylcyclohexane are competitively adsorbed. This semi-empirical model had to be slightly refined in order t o be in closer agreement with the reaction mechanism proposed (Sect. 3.5.4). References p p . 385-398
344 ( b ) Structure and reactivity
With regard t o the effect of the structure of the hydrogen donor, Malinowski et al. [372] obtained, for vapour phase reactions with formaldehyde over Na-modified silicagel, the reactivity order (see Fig. 1 2 ) CH3CH0 > CH3COCH3 > CH3CN. The authors pointed out that the acidities of these substances decreased in the same order and suggested that the rate coefficient be expressed as a function of the hydrogen donor acidity [eqn. (22)]. Kraus [ 1251, using Malinowski’s data, has demonstrated that there is a very good linear relationship between log k for these reactions and the acidities expressed as the pK of the hydrogen donor. In the liquid phase reactions of substituted acetophenones with benzaldehyde over an anion exchanger [ 370 3 , the order of the effect of substituents on the reactivity was found to be p-OCH3 < rn-OCH3 < no substituent < p-F < p-Br < rn-Br. With another group of hydrogen donors, the reactivity order malonitrile > benzyl cyanide > acetophenone > ethyl cyanoacetate was established. Contrary t o the above view [ 3721, these authors [ 3701 came t o the conclusion that the reaction rate was not an unequivocal function of the hydrogen donor activity. The reactivity of hydrogen acceptors in vapour phase reactions with acetaldehyde over alkalised Si02 [376] decreased in the order (see Fig. 13) CH20 > CHJCHO > C,H,CHO. It has been assumed that it is the basicity of the acceptor oxygen which affects the reaction rate, since oxygen basicity is a measure of the ease with which a proton adds to oxygen and, subsequently, of the ease with which the carbonyl carbon attaches an anion [see scheme (F)]. The positive charge on the acceptor carbonyl carbon is dependent on the nature of the closest neighbour of the carbonyl group. Since formaldehyde has no substituent, its carbonyl group has the largest positive charge. In acetaldehyde, hyperconjugation reduces the densities of negative and positive charges on the oxygen and the carbonyl carbon, respectively, and in benzaldehyde, the mesomeric and steric effects suppress the basicity of the carbonyl oxygen and the positive charge on the carbonyl carbon still more. The lower reactivity of benzaldehyde with respect t o acetaldehyde was found also in the vapour phase aldolisation over lithium phosphate [ 3901. Over the same catalyst, the reactivity order in the self-condensations of aldehydes could be estimated as CH3CH0 > CH3CH2CH0>> (CH,),CHCHO. The reactivity of isobutyraldehyde in the self-condensation was almost undetectable, probably due to steric hindrance on the &-carbon, but this substance was able t o react as a hydrogen acceptor with cyclohexanone. With propionaldehyde over a calcium hydroxide catalyst, a Cannizzaro-type reaction occurred t o some extent simultaneously with the aldolisation [ 3901. This unexpected result was also recorded by other authors [391], who established that the tendency t o aldolisation decreased, and the tendency t o the Cannizzaro reaction increased, with
345 increasing chain length of the aldehyde; both tendencies appear to be in agreement with the results above.
3.5.4 Mechanism In the work concerning the mechanism of solid-cataiysed aldol reactions, the analogy between the homogeneous and heterogeneous mechanisms is usually assumed [370,372-3751. The mechanism of base-catalysed condensations, which has received much attention (cf. ref. 371), may be pictured in general as
R4--CH2Y+ B * R4-€-)HY
H R' I I R4-C-C-0'-) I I Y R2
+ B(+)H
(1)
H R' I I + B'+'H + R4-C--C-0H + B I I Y R2
(3)
using the notation of scheme (D). The hydroxy compound formed can be dehydrated to an unsaturated product by a carbanion elimination ElcB mechanism [see Sect. 2.11
H R' I I R4-C-C-0H I 1 Y R2 R' I R4--C(-)+--OH
I
Y
I
R2
R' + B * R4--C'-'&H
I Y
I
+ B'"H
R2
R' I + B'+'H =+R4--C=C + H 2 0 + B I I Y R2
It is not necessary that the acid B(+'H involved in the aldolisation (F) and dehydration (G) mechanisms must be just the conjugated form of the basic catalytic site B acting in steps (1)and (4), but it is possible that a cooperation of basic and acidic sites originally present on the surface of the catalyst as pairs of suitable mutual distance, occurs, though the basicity of the catalyst is the main property determining its catalytic activity. The activating group Y can stabilise the carbanion RC'-)HY or R4Y C(-)C(R'R2)-OH; in the case of true aldolisation when the activating group is a carbonyl, R3C=0,the stabilisation can occur via a keto-enol tautomerReferences p p . 385-398
346 ism
R3 I R4C(-)H-C=0 R4
R3 I + R4CH=C-O-)
R'
\
C(-)-&oH I I R3-C R2 \\0
R4
+
\
R'
I
C-C-OH 0 1 R3--C RZ &(-)
The mechanism presented by schemes (F) and (G) is consistent with the particular mechanisms suggested by different authors for the condensations catalysed by solid basic catalysts [ 370,372,376,3921, and seems to be supported by the effects of the structure of either hydrogen donors or hydrogen acceptors [Sect. 3.5.3.(b) and eqn. (22)]. For acid-catalysed aldol condensations (which are less frequent), the homogeneous mechanism [371] can again be accepted. The enol form of the hydrogen donor interacts with the protonated form of the hydrogen acceptor, viz. R3 I R4RSCH--C =+R4R5C=C
P3
\
\\
OH
0
R1\
C=O I R2
R1\
+ H' *
C")--OH
I
R2
R1\ C'"-OH I R2
(Hf
P3 * R1oH \I R4 I I C--C!-C=O
+ R4R5C=C \
OH
R3
/ R2
I R5
+ H'
The dehydration of the aldol (ketol) proceeds more rapidly with acidic than with basic catalysts, and this is the reason why, with the former, the CY $-unsaturated carbonyl compounds are the products most frequently encountered. The dehydration follows one of the elimination mechanisms discussed in Sect. 2.1, depending on the particular nature of the used catalyst and on the temperature. With solid acid catalysts, it is possible that the protonation of the carbony1 compound does not lead to a fully ionic form, but that by interaction with surface protons, a hydrogen-bonded intermediate can be formed. Such a mechanism, not differing in the other features from scheme (H), was proposed for acetaldehyde self-condensation over a CH 3 / CaHPO, catalyst [ 3931; a surface complex P--O...H..-O--C was
'H
347 assumed. The formation of hydrogen-bonded intermediates on the surface active sites is not unlikely t o occur even in the base-catalysed reactions [372]; the first step of scheme (F) would then lead t o the formation of an intermediate surface complex of the type R4YCH...H.-B . The mechanism proposed for the acid-catalysed synthesis of bis(ary1)alkanes [ 3941 [scheme (E)] follows the main features of the aldolisation scheme (H). The protonised form of the carbonyl compound reacts by an electrophilic attack with the quinonoid structure of the aromatic molecule (e.g. phenol), viz.
The intermediate tertiary carbinol could not be detected (with the exception of bis(trifluoromethyl)(hydroxyphenyl)carbinol from hexafluoroacetone and phenol [ 3951) and reacts readily with another molecule of phenol; this second stage of the reaction is, in fact, an alkylation of phenol by the tertiary carbinol, or by the carbonium ion formed from it, by a common carbonium ion alkylation mechanism (Sect. 3.3). The mechanism represented by scheme (J) was accepted for heterogeneous catalytic reactions and only slightly modified for the reactions catalysed by hydrogen forms of zeolites [395] or for bisphenol A synthesis catalysed by sulphonated ion exchangers [ 385,3961. In the former case, a Rideal-like mechanism was assumed, in which the chemisorbed (by a H-bond) conjugated acid of the carbonyl compound reacts with the nonadsorbed aromatic molecule [395]. In the latter case, acetone is considered to be chemisorbed by hydrogen bonding of its carbonyl group to a resin -SO,H group; chemisorbed acetone then reacts with phenol from the surrounding non-polar matrix of the resin [385]. The tertiary carbonium ion intermediate is assumed in both cases [385,395] t o react in References P P . 385-398
348 the chemisorbed form by an electrophilic attack on the other aromatic molecule.
4. Substitution reactions Substitution reactions represent a very large group of transformation of organic molecules. Substitutions at a saturated carbon atom as well as on the aromatic nucleus are classified according t o mechanism into nucleophilic and electrophilic and are well known under the notation SN and SE. Ester formation and hydrolysis, which are, in fact, also substitutions, namely in the carboxyl group, are usually treated separately in monographs and textbooks (see, for example, refs. 397, 398). In homogeneous media, the kinetics and mechanism of all these substitution reactions have been very thoroughly studied. In heterogeneous catalysis, however, the situation is different: intensive kinetic and mechanistic investigations of esterification and hydrolysis have been performed, while papers concerning the classical aliphatic and aromatic substitutions catalysed by solid substances are scarce and the results diffuse. The substitutions for the OH group in alcohols by amino groups, aryl and halogens can serve as examples. These reactions, catalysed by acidic or basic solid catalysts, frequently proceed by an eliminationaddition pathway with intermediate formation of carbonium ions or highly polarised species and follow the mechanisms which were treated in Sects. 2 and 3 concerning elimination and addition reactions. This is another reason, besides the lack of systematic studies on kinetics and mechanism, why we have not included the solidcatalysed aliphatic and aromatic substitutions in this chapter. Section 4 is therefore limited t o heterogeneous esterification, transesterification and hydrolysis (including hydrolysis of compounds other than only esters) where the extensive published literature permits a review of their kinetics and mechanisms. 4.1 ESTERIFICATION AND TRANSESTERIFICATION
4.1.1 T y p e s of reactions and catalysts In a broader sense, the term esterification may include all reactions in which esters, both of organic and inorganic acids, are formed. We sha!l limit the discussion in this section, however, t o ester formation from organic carboxylic acids and alcohols RCOOH + R’OH = RCOOR’ + H,O
(A)
In the literature concerning solid-catalysed esterifications, kinetic studies of other ester-forming reactions are scarcely reported. Reactions of
organic esters with alcohols (alcoholysis)
RCOOR" + R'OH
RCOOR' + R"OH
(B) are similar to reaction (A) with respect t o the mechanism and catalysts used. They are called transesterifications and will be included in this chapter. A special case of this substitution reaction of esters, hydrolysis (R' = H), will be dealt with separately (Sect. 4.2.1) because of its specific importance. Reactions (A) and (B) are reversible and this must be taken into account in deriving rate aquations; the effect of the reverse reaction is frequently suppressed by working at low conversion or in excess of one of the reactants. For illustration, esterification equilibrium data for some pairs of acids and alcohols in liquid phase are given in Table 18. For vapour phase esterifications, different values can be found, as is evident, for example, by comparing the equilibrium constant for acetic acidethanol esterification in the liquid phase at 155°C ( K , = 3.96) with that determined in the vapour phase at 150°C (KP = 30.9-33.6 [400]). This discrepancy may be due to different activities (fugacities) of reactants in liquid and vapour phases or to errors in determining the conversions caused by adsorption of some gaseous reactant on the solid catalyst [400]. Since the equilibrium constants for esterification of a given acid with homologous alcohols do not usually differ significantly, the equilibrium constants for transesterification, which are given by the ratio of esterification equilibrium constants of the acid with the two alcohols in question =
- Kst.1
Ktrans --
Kst.2
will not be far from unity in most cases (see some experimental values in Table 19). TABLE 18 Equilibrium constant, K,, and limit conversions, xeq, of esterification [ 399 ] (155"C, initial molar ratio of reactants 1 : 1 . ) ~
Acetic acid with various alcohols Alcohol
KC
xeq
2-Methyl-1-propanol with various acids Acid
K C
(%)
Methanol Ethanol 1-Propanol 2-Propanol I-Butanol 2-Butanol Ally1 alcohol
5.24 3.96 4.07 2.35 4.24 2.12 2.18
References p p . 385-398
69.6 66.6 66.85 60.5 67.3 59.3 59.4
xeq
(%)
Formic Acetic Propionic Butyric Isobutyric Methylethylacetic Trimethylacetic Benzoic
3.22 4.27 4.82 5.20 5.20 7.88 1.06 7.00
64.2 67.4 68.7 69.5 69.5 73.7 72.65 72.6
TABLE 1 9 Equilibrium constants ( K , o r Kp) of transesterification. Ester
Alcohol
Phase
Temperature ("C)
K , or K ,
Ref.
Ethyl acetate
Methanol
Vapour
Methanol 1-Butanol
Liquid Liquid
0.365 0.550 0.790 0.924 1.35 1.3
401
Ethyl acetate Ethyl acrylate
170 200 220 240 60 100
402 403,404
The physical properties of most acids (esters) and alcohols allow the reaction t o be carried out either in the liquid or in the vapour phase. In the liquid phase, t h e effects of solvents and of transport phenomena may play a more important role than in the vapour phase. On the other hand, the side reactions (mainly the ether and/or olefin formation from the alcoTABLE 20 Reactants and inorganic catalysts used in kinetic studies of esterification (transesterification) Acids (Ester)
Alcohols
Catalysts
Temperature ("C)
Ref.
CH3COOH CH3COOH
C2 H5 OH C2 H5 OH
425 350
405 406
C2-€, normal and branched
250
126
CH3COOH CH3COOH CH3COOH CH3COOH CH3COOH CH3COOH CH3COOH and C3H7COOH CH3COOH CH3COOH
CI-C4 primary, secondary, tertiary CZHsOH C2 Hs OH C2 Hs OH n-C3H7 0 H n-C3H,OH C2 H5 OH Cz H5 OH and n-C4HgOH n-CsH, 7 0 H n-C4 H, OH
Si02-A1203 (SA) Si02-A1203 (SA) Al203-B203 (AB) Na-poisoned SA Na-poisoned AB O3 Na-poisoned SA
Si02 SiO, Si02 SiO, SiOz wo 3 /A12 0 Bauxite
200-260 246-286 150-270 170-230 200-260 255
407 408 409 410 41 1 412 413
218-241 155-197
414 415
CH3COOH HCOOH CH3COOCzHs
Cz HSOH C2 HSOH CHjOH
120-140
416 417 401
a
3
Bauxite Acid activated Korvi earth H3P04/C Ca-metaphosphate SiOz -~.._.I_.~__
a
Transesterification.
351 hol) are usually unimportant in the liquid phase, whereas they may become significant in the vapour phase because of the higher temperatures used. In homogeneous media, esterifications are catalysed by acids. Similarly, for kinetic studies of heterogeneously catalysed esteifications, the use of solids of acidic character is reported. These catalysts may be divided into two groups: (i) inorganic acid catalysts and (ii) organic polymer-based ion exchangers in the acid form. With inorganic catalysts, the majority of kinetic studies were performed in the vapour phase, whereas the main use of ion exchanger resins is for liquid phase processes.
4.1.2 Reactions catalysed b y inorganic catalysts The reactants and inorganic catalysts used in kinetic studies of heterogeneous catalytic esterifications (transesterifications) are summarised in Table 20. As can be seen, no systematic comparative study with more than one catalyst (with the exception of paper [ 4061 ) has been performed by any one worker. The greatest attention was paid t o silica gel [4074111. The reactants were usually low molecular weight acids and alcohols; a typical pair of reactants is acetic acid-ethanol. Only in one study [ 1261 was the structure of the reactants systematically varied in order to establish the effect on the reactivity.
( a ) Formal kinetics For a formal kinetic description of vapour phase esterifications on inorganic catalysts (Table 21), Langmuir-Hinshelwood-type rate equations In were applied in the majority of cases [ 405-408,410-412,414,4151. some work, purely empirical equations [413] or second-order power lawtype equations [ 401,4091 were used. In the latter cases, the authors found that transport phenomena were important: either pore diffusion [ 4011 or diffusion of reactants through the gaseous film, as well as through the condensed liquid on the surface [ 4091, were rate-controlling. The Langmuir-Hinshelwood-type rate equation correspond, in most cases, t o the assumption of surface reaction being the rate-determining step [ 405,406,410-412,414,4151. However, the details of the model differ in individual cases: either one (acid or alcohol) [410,411,415] or both reactants [ 405,406,4141 are assumed to be adsorbed. Of rate-determining steps other than surface reaction, only adsorption of the acid is reported [407,408,415,416]. In several cases with silica gel, an activating effect of water was observed; this was described either by including a Langmuir isotherm for adsorption of water into the rate equation on the assumption that each adsorbed water molecule creates an additioiial active site [ 410, 4111, or by including an empirical function L = rn f (pR)for the number, L, of active sites (pRis partial pressure of water) [ 407,4081. References p p . 385-398
TABLE 21 Equations reported as best fitting esterification (transesterification) data o n inorganic catalysts Catalyst
Equation
a
Rate-determining step
Ref.
Si02-Al203 Bauxite
1
S R between molecularly adsorbed reactants
405,4 1 4
Si02-Al20 A1203-B203 A1203
I
S R between molecularly adsorbed reactants
406
SR between molecularly adsorbed alcohol and acid reacting from t h e vapour phase
4 10,411
SiO2
(water increases t h e number of sites) SiOz Acid activated Korvi earth
i
AdsA (alcoho: reacts from t h e vapour phase, water increases the number of sites)
407,408
AdsA (alcohol reacts from the vapour phase; temperature, 176 and 197°C) S R between molecularly adsorbed acid and alcohol reacting from t h e vapour phase (temperature, 155°C)
415 415
412
Bauxite a
Power-law type equation (second order)
401 409
Empirical rate equation
413
Symbols: k,, k A = rate coefficients for surface reaction and adsorption of acetic acid, respectively; E = effectiveness factor; L = total number of active sites; p i , Ki = partial pressure and adsorption coefficients of reactant ( A = acid, B = alcohol, R = water, S = ester); KK = adsorption coefficient of water in the second layer; L' = concentration of active sites a t zero water partial pressure; m, a, b = constants. SR = surface reaction, AdsA = adsorption of the acid. Ref. 401 concerns transesterification ( A denotes the ester).
As appears from the examination of the equations (giving the best fit to the rate data) in Table 21, no relation between the form of the kinetic equation and the type of catalyst can be found. It seems likely that the equations are really semi-empirical expressions and it is risky t o draw any conclusion about the actual reaction mechanism from the kinetic model. In spite of the formalism of the reported studies, two observations should be mentioned. Maatman et al. [ 4101 calculated from the rate coefficients for the esterification of acetic acid with 1-propanol on silica gel, the site density of the catalyst using a method reported previously [ 4181. They found a relatively high site density, which justifies the identification of active sites of silica gel with the surface silanol groups made by Fricke and Alpeter [411]. The same authors [411] also estimated the values of the standard enthalpy and entropy changes on adsorption of propanol from kinetic data; from the relatively low values they presume that propanol is weakly adsorbed on the surface, retaining much of the character of the liquid alcohol.
( b ) Effect o f reactant structure on reactivity Mochida et al. [126] have compared the reactivity of eight alcohols with acetic acid and of seven carboxylic acids with ethanol over sodiumpoisoned silicaalumina at 250°C using a gas chromatographic technique; their results correlated by the Taft equation are shown in Figs. 14 and 15. The effect of the structure of both alcohols and carboxylic acids on their reactivity in heterogeneous catalysed esterification is very small. In contrast, the polar effect of substituents in the alcohol molecule in its intramolecular dehydration to olefin, accompanying the esterification, is much larger (Fig. 14) suggesting different ratedetermining steps in these two reactions. Figure 15 demonstrates the difference in the steric effect of substituents in carboxylic acid molecules in heterogeneous and homogeneous esterifications; however, extrapolation of the published data [ 4191 for homogeneous esterification t o higher temperatures indicates that the difference would become less significant if closer temperature for both types of catalysis were used [ 1261. No other systematic studies of the structure-reactivity relations for esterifications on solid inorganic catalysts have been reported. Only Fricke and Alpeter [411] compared their own results obtained for 1-propanol with those of other authors concerning methanol and ethanol in esterification of acetic acid over silica gel. The rate drops sharply with the increasing chain length (in contrast to the findings of Mochida et al. [126] and in agreement with Heath [420]), and the kinetic model changes from a dual site one (with both reactants adsorbed) for methanol to a single site one for 1-propanol (only alcohol is adsorbed). An interpretation which could explain this change is that steric hindrance by the larger molecules of 1-propanol prevents adsorption of the acid; this changes the kinetics and lowers the reaction rate. Reference6 p p . 385-398
354 4
I
I
I
1
3
2 L
F
4
1
0
I
- 03
I
- 0.2
I
I
- 0.1
bsc
0
Fig. 14. Taft correlation with polar substituent constants ( u * ) of the vapour phase esterification of acetic acid with alcohols ( 0 )and of the olefin formation from alcohols (0) over Na-poisoned silica-alumina at 25OoC [126]. 1,Methanol; 2, ethanol; 3, l-propanol; 4, 1-butanol; 5 , 2-methyl-1-propanol; 6, 2-propanol; 7 , 2-butanol; 8, 2-methyl2-propanol.
Fig. 15. Taft correlation with steric substituent constants (E,) in the vapour phase esterification of carboxylic acids with ethanol over Na-poisoned silicaalumina at 25OoC ( 0 ) [126] and in homogeneous acid-catalysed esterification at 4OoC (0)[419]. Acids: 1, acetic, 2, propionic, 3, butyric, 4, isobutyric, 5, isovaleric, 6, pivalic, 7, 2ethylbutyric.
355 (c) Mechanism
The type of catalyst must be taken into account when considering the reaction mechanism. Mochida e t al. [ 4061 investigated several oxide or mixed-oxide catalysts for esterification of acetic acid with ethanol. For silica-alumina, alumina-boria and sodium-poisoned silicaalumina, the authors assumed that active sites for esterification were the protonic sites. Kinetic experiments revealed that the more basic ethanol was more strongly adsorbed than acetic acid, and poisoning experiments with organic bases led to the conclusion that esterification proceeded on even weaker sites. The authors suggested a mechanism (similar t o that used for homogeneous esterification [421,422]) for which the attack of the oxonium ion, formed from acetic acid, on the adsorbed ethanol was ratedetermining. They have excluded the possibility of carbonium ion participation on the basis of a number of arguments, including the formation of ethyl thioacetate and no ethyl acetate from acetic acid and ethanethiol over the silicaalumina catalyst. However, the investigations of the mechanism of olefin and ether formation from alcohol (Sect. 2.2) revealed the importance of basic sites. It is feasible that, for esterification also, pairs of acidic and basic sites might be necessary. A mechanism similar t o that proposed by Mochida for the above-mentioned group of catalysts, though not so explicitely formulated, might also be valid for acetic acid-ethanol esterification over a H3P04/C catalyst [ 4161. According t o the author, the adsorbed acid in a polymolecular film on the surface of the catalyst reacts with protonated molecules of the adsorbed ethanol. These mechanisms are in formal agreement with kinetic equations assuming surface reaction between molecularly adsorbed reactants; besides the group of catalysts used by Mochida e t al. [ 4061, such equations were also found t o fit the kinetic data for silicaalumina [405] and bauxite [414] (see Table 21). An activating effect of water was observed for the catalyst H3P04/C [ 4161 and for silica gel; for the latter, the effect seems to be more general since it has been established by several authors [ 407,408,4111. A plausible explanation of the promoting effect of water on silica gel was suggested by Fricke and Alpeter [411]. The authors assume that water is adsorbed on silica gel in two layers. In the first layer, it is adsorbed strongly, hydrating the surface according t o the recation
The silanol groups formed may act as active sites, since it has been obReferences PP. 385-398
served [423] that surface hydroxyl groups can act as adsorption sites for molecules having electron donor atoms. This idea is in harmony with the mechanism proposed by Mochida et al. [126,406] for esterification on other oxide or mixed-oxide catalysts. The concentration of surface silanol groups can be expressed as a function of water partial pressure and introduced into the rate equation [see Sect. 4.1.2.(a)]. As to the second layer of water, it is assumed to be less strongly adsorbed and to cause the free catalyst surface to decrease by competition with adsorption of reactants. This second, inhibiting effect is expressed by the corresponding Langmuir term K k p , in the denominator of the rate equation (see Table 21). The proposed mechanism of the effect of water can be supported by two other findings: (i) the calculations of Maatman et al. [410] revealed that the active sites could be identified with surface silanol groups [Sect. 4.1.2.(a)] and (ii) independent studies of other authors [ 424-4261 showed that silica gel could actually adsorb two layers of water; the first layer is strongly chemisorbed whereas the second is less strongly adsorbed and retains much of the character of free water. The standard enthalpy and entropy changes on adsorption determined from kinetic adsorption coefficients, KR and Kk,for the first and second layer, respectively [411], are consistent with this observation. It follows from all the above considerations that the acidic character of the surface is necessary for the esterification reaction. This view is supported by the parallel found by some workers [ 405,4061 between the rate of esterification and that of other typical acid-catalysed reactions. A linear correlation was established between the rate of acetic acid-ethanol esterification and that of deisopropylation of isopropylbenzene on a series of silica-alumina, alumina-boria and alumina catalysts [ 4061 ; a similar relation was found between the rate coefficient of the same esterification reaction and the cracking activity of a series of silicaalumina catalysts prepared in a different way [ 4051.
4.1.3 Reactions catalysed by organic polymer-based cation exchangers Organic cation exchangers are copolymers (mostly of the styrenedivinylbenzene type) with acid functional groups, such as -S03H. The close chemical analogy between these groups and inorganic acids used as homogenous catalysts (e.g. H,SO,) led to the idea of using organic ion exchangers as solid catalysts in proton-catalysed reactions, such as esterification. The fact that homogenously catalysed esterifications are carried out in the liquid phase determined the method of operation in the majority of esterification studies using ion exchanger catalysts: only a few kinetic studies of esterification and transesterification in the vapour phase with these catalysts were performed. The same fact also influenced the approach t o the kinetic analysis of esterifications catalysed by ion exchangers. As for homogeneous reactions, power law-type rate equations
357 (second- or pseudo-first-order) have been used in almost all published work, in spite of the fact that the presence of the solid catalyst may introduce changes into the kinetic relationships and complicate the analysis by the effects of slow diffusion of reactants to the solid surface and/or through the polymer mass. In order t o account for the well known sorption and swelling properties of polymer ion exchangers, Helfferich’s model [ 4271 is frequently used for liquid phase reactions. According to this, the pore liquid of the resin, where the reaction occurs, is treated as a homogeneous system and the reactant is assumed t o be distributed according t o a distribution coefficient
between the pore liquid and the supernatant solution; creSand csol are the respective reactant concentrations. The concentrations c,,, in the pore liquid are introduced instead of cSo1into the second-order kinetic law. In Helfferich’s relation it is implicitly assumed that the diffusion of the reactants through the pore liquid is fast enough so that equilibrium between the pore liquid and the supernatant solutions exists. Several papers, however, report a particle size effect on the esterification rate [4284331 indicating diffusional limitations; attempts t o describe,quantitatively the transport situation in the polymer mass are less frequent (e.g. ref. 434). For the liquid phase kinetic studies of esterification, with a few exceptions [ 402,435-4371 only the standard (non-porous, see Sect. 1.2.5) ion exchangers were used. The macroreticular (porous) ion exchangers with a large inner surface area are prefered for vapour phase reactions, especially in more recent studies [ 436-4431. The authors claimed that diffusion was not the limiting process under their conditions. This observation cannot be generalised, however, and even with vapour phase reactions and macroreticular polymers, the possibility of transport limitations through the pores o r the polymer mass cannot be excluded a priori. As with inorganic solid catalysts, the most extensively studied system was acetic acid-ethanol [ 428,432,434,444-4481.Other alcohols used in kinetic studies were methanol [ 430,449,4501,2-propanol [ 4381, l-butanol [ 429,431,433,451-4581,ally1 alcohol [ 4591,l-pentanol [ 4341 and ethyleneglycol [ 4601 ; besides acetic acid, the reactions of formic [ 4501, propionic [443,461],salicylic [ 430,4491,benzoic [ 453-4571 and oleic acids [430,451-4531 and of phthalic anhydride [462]have been reported. Investigation of a greater variety of reactants is reported in only one paper [463]:six alcohols (C4, C, and C,) and five acids (mainly dicarboxylic were studied. Transesterification kinetic studies were performed with ethyl formate [437,439,441],isobutyrate [ 437,439-4411 acetate [ 402, 435-437,439-4421, methoxyacetate [ 4411 and acrylate [ 403,404,464, 4651 ; the alcohols used were methanol [ 402,435,437,439-442,4501, References PP. 385-398
358 2-methoxyethanol [441], l-propanol [435-437,439-4411, l-butanol [ 403,4041, 3-methyl-l-butanol [ 464],2,2-dimethyl-l-propanol [ 437,439, 4411 and ally1 alcohol [ 4641. (a) Formal
kinetics
As has been already said, the majority of esterification (transesterification) kinetic data measured in the liquid phase were treated by using second- or pseudo-first-order rate equations. The concentrations of reactants were corrected in some cases [ 403,404,434,449,454-457,4641by using the Helfferich distribution coefficient, h. Some authors have recognised the oversimplification involved in the Helfferich model, which takes into account the concentration differences inside and outside the polymer particle but neglects the heterogeneous character of the chemical interaction between the liquid reactants and the functional groups of the solid catalysts. Interaction of such a type can be described by an adsorption isotherm; the Langmuir isotherm can be used, because all the functional groups of a given ion exchanger are chemically identical. Bochner et al. [ 4491 have developed a rate equation for salicyclic acid-methanol esterification (Dowex 50W catalyst) using this approach and assuming that all the reaction components can be associated with the protons of the -S03H groups in the polymer (competitive chemisorption). By using some simplifications, they obtained eqn. (24) in which it was assumed that the reaction between chemisorbed salicyclic acid molecules and methanol molecules in the pore liquid was the rate-determining step, r=
kbl[SA] [CH,OH] 1 + bz[HzOl
where bl and b2 are the association affinities for salicyclic acid (SA) and water, respectively. Equation (24) is, in fact, a Langmuir-Hinshelwood-type equation. Similar models with a single site surface reaction as the ratedetermining step were used for other liquid phase esterifications [ 448,451 1. Experimental data for the l-butanol-oleic acid system were best fitted by eqn. (24) [452] or eqn. (25) [451]
where Ki are adsorption coefficients and N i mole fractions of the reaction components (A = acid, B = alcohol, R = water, S = ester) and K is the thermodynamic equilibrium constant. Langmuir-Hinshelwood equations, with a surface reaction as the ratedetermining step, were also found suitable for liquid phase transesterifications [435-4371. With ethyl acetate and l-propanol in dioxan as a sol-
359 vent, the data obeyed eqn. (26)
where A is ester, B alcohol and I solvent. It is assumed that the ester reacts from the liquid phase with the adsorbed alcohol; the solvent competes with the alcohol for the active sites. With cyclohexane, there is no interaction of this inert solvent with the active sites and both reactants, alcohol and ester, are adsorbed, as it has been found for the same reaction in the vapour phase [439] [eqn. (27)].
If methanol, as a less basic alcohol than 1-propanol, is used in dioxan as solvent, then, contrary t o the model giving rise t o eqn. (26), the ester is adsorbed (protonated) and alcohol reacts directly from the liquid phase, and thus
Considerably fewer kinetic studies were performed with reactants in the vapour phase than in the liquid phase. The second-order rate equation was only used for acetic acid-ethanol esterification at 130°C and 175°C on a KU-2 standard ion exchanger [ 444,4451. A semiempirical second-order rate equation with slight inhibiting effect of reaction products, viz.
kPAPB 1 + aps ibPR was proposed for the vapour phase esterification of acetic acid with butanol over an oxidised phenol-formaldehyde carbon [ 4581. Other equations. For the vapour authors used Langmuir-Hinshelwood-type phase esterification of acetic acid with 2-propanol [ 4381 or ethanol [ 4461 and of propionic acid with ethanol [443], a dual site model with both reactants adsorbed was found to fit the experimental data r=
Herrman [ 4461 established an order-of-magnitude agreement between the values of the adsorption coefficients obtained by direct measurements of adsorption of alcohol and water vapour and those evaluated from kinetic data as KB and KR , which, in the author’s opinion, supported the physical meaning of these constants. The physical meaning of the Langmuir-Hinshelwood model was also examined by means of several transesterification reactions in the vapour phase at 120°C on a macroreticular ion exchanger [439,440,442]. The References p p . 385-398
TABLE 22 Values of the parameters of rate equation ( 2 7 ) for vapour phase transesterification catalysed with macroreticular ion exchanger at 120"C [ 4 3 9 ] Reaction components
Rate coefficient, h (mol kg-' h - l )
B
A Ethyl acetate Ethyl acetate Ethyl acetate Ethyl acetate Ethyl isobutyrate Ethyl formate
Methanol l-Propanol 2,2-Dimethylpropanol l-Propanol l-Propanol l-Propanol
Adsorption coefficients KA
KB
(bar-')
(bar-')
2625 284 80.4
1.3 1.0 0.9
0.3 2.8 7.6
284 55.8 37050
1.0 0.55 0.4
2.8 3.1 1.8
kinetics were expressed by the rate equation (27) and the resulting values of the parameters are summarised in Table 22. As can be seen, the adsorption coefficients of ethyl acetate obtained by kinetic analysis of its reaction with three differently reactive alcohols have very closely similar values; the same is true for the adsorption coefficients of l-propanol in its reaction with ethyl acetate and isobutyrate. Thus, the adsorption coefficients do not depend on the nature of the second reaction partner and are not empirical constants valid only for one particular reaction but they characterise the compound in question more generally. Kinetic analyses of systems where two esters compete for one alcohol or two alcohols for one ester were also performed [440]. In spite of the fact that more reactants are present in the reaction system and are adsorbed competitively on the surface of the catalyst, good agreement between the values of the parameters of eqn. (27), determined by two independent methods (kinetic analysis of isolated and competitive reactions), was again obtained. The results of both studies [439,440] demonstrate the applicability of the Langmuir-Hinshelwood model to reactions catalysed by ion exchangers.
( b ) Effect o f reactant structure
A quantitative correlation of structural effects of four esters and four alcohols in the vapour phase transesterification on a macroreticular ion exchanger at 120°C was made using the Taft equation [ 4411. The authors found that rate coefficients [from eqn. (27)] yielded better correlation with steric (E,) than with polar (u*) parameters, while there was no significant difference between the correlations of the adsorption coefficients of alcohols, K,, with both parameters. The correlations with E , yielded the slopes 1.4 and 0.6 for the reactivity of the esters and the alcohols, respectively, and - 0 . 4 for the adsorptivity of the alcohols. The observed
TABLE 23 Ratios of initial transesterification rates on the least and the most cross-linked ion exchanger [ 4 3 7 ] (M = macroreticular polymer, S = standard (non-porous) polymer.) Reactants Ester
Liquid phase (in dioxan, 52°C)
Vapour phase (120°C)
Ma
Sb
Ma
A1c oh 01 Sb
Ethyl formate Ethyl acetate Ethyl isobutyrate
1-Propanol 1-Propanol 1-Propanol
2.18 1.01 1.31
78.9 62.2 24.1
0.44 0.36 0.21
9.28 8.33 9.20
Ethyl acetate Ethyl acetate Ethyl acetate
Methanol 1-Propanol 2,2-Dimethylprop ano1
2.22 1.01 1.02
16.5 62.2 82.1
0.21 0.36 0.11
1.58 8.33 16.0
.-
a
Resins with 10 and 60% DVB were compared. Resins with 2 and 50% DVB were compared.
decrease in the reactivity of alcohols with the increasing chain length is analogous to that discussed by Fricke and Alpeter [411] for esterification over a silica gel catalyst. The importance of steric hindrance was also stressed in a comparative esterification study in the liquid phase [ 4631. As will be discussed later [Sect. 4.1.3.(c)], the reactivity in reactions catalysed by ion exchangers depends on the degree of crosslinking of the polymer. If these effects were different for reactants of different structures, one could expect changes in the relative reactivities (selectivities) of reactants on varying the degree of crosslinking. Table 23 summarises the ratios of initial transesterification rates [437] on the least and the most cross-linked ion exchanger (from 2 to 60% divinylbenzene content) for the same reactants as in Table 22. It is apparent that the selectivity in transesterification, e.g. of ethyl acetate with methanol and 2,2dimethylpropanol on standard ion exchangers, would increase five times (82.1/ 16.5 = 4.98) in the liquid phase, or ten times (16/1.58 = 10.1) in the vapour phase, in favour of methanol in going from the lowest t o the highest degree of crosslinking. With macroreticular ion exchangers, the effect would be reversed. When reactants of large molecular size, such as oleic acid, were reacted over an ion exchanger catalyst, a direct proportionality between the reaction rate and the surface area of the catalyst was found [433]. The authors explain the result by assuming that, for the bulky reactant molecules, only acid groups at or near the surface of the catalyst particle can be effective catalysts. The efficiency of the catalyst (the rate coefficient with the resin compared to the rate coefficient with the same stoichiometric amount of dissolved inorganic acid) was found to be considerably References p p . 385-398
362 lower for oleic than for benzoic acid [ 4531. The structure of the reactants can affect the relative adsorptivities of ester and alcohol and perhaps, according t o the view of Setinek and Rodriquez [ 4351, also the kinetic mechanism, as already discussed in Sect. 4.1.3.(a) [see eqns. (26) and ( 2 8 ) ] . It was found in transesterification of ethyl acrylate in the liquid phase over a non-porous KU-2 catalyst [ 4641, that the structure of the alcohol influenced the value of the limiting sorption of alcohol by the ion exchanger, the logarithm of this value being a linear function of the dielectric constant of the alcohol. As the second-order rate coefficients yielded the same sequence as the limiting sorption values, viz. ally1 alcohol > l-butanol > 3-methyl-1-butanol, Filippov et al. [ 4641 assumed a relation between the dielectric constant and the reactivity of the alcohols.
( c ) Effect o f ion exchanger properties With one exception [447], only sulphonated resins were used as catalysts in kinetic studies of esterification and transesterification, the resins being almost exclusively styrene-divinylbenzene copolymers; in one case, a sulphonated phenol-formaldehyde resin was also used [ 4331. The main factors determining the catalytic activity are (i) the concentration of functional groups in protonated form (-S03H groups) and (ii) the degree of crosslinking of the copolymer (characterised by the divinylbenzene content).
(i) Concentration of ucid groups. It is generally accepted that the protons of the acid functional groups in the polymer are responsible for the catalytic activity in esterification. According t o expectation, the activity was always found t o drop when the number of acid groups decreased [428, 455,456,4601. However, this dependence is not linear and the activity of the remaining protons in a partially neutralised catalyst is lower than in a catalyst fully in the H'-form [ 428,455,4561. This observation indicates that the dependence of the reaction rate on the proton concentration is of an order higher than unity and that more than one proton or protonated species might participate in the formation of the activated complex in the ratedetermining step. The activity of the remaining protons is dependent, not only on their concentration, but also on the nature of the cation used for neutralisation. For the esterification of acetic acid with ethylene glycol on a KU-2 catalyst, the inhibiting effect of cations increased in the order Na' < K' < Ca2' < Li+< A13+[460]; from data on the esterification of benzoic acid with 1-butanol over Dowex 5 0 W catalyst [ 4551, the order Mg2' < Na' < Ba2' < Cs' can be estimated. However, the activation energy (-75 kJ mol-') was found t o be independent of the type of cation and the degree of neutralisation [ 4551. The changes in activity are attributed by the authors t o the changes in activation entropy, which decreases with
363 the increasing concentration of inorganic ions in the resin. Since, with higher ion concentrations, the sorption of benzoic acid also increases, the authors assume that the activation entropy is the more reduced the stronger is the adsorption of this reactant and the more restraints upon the transition state. A similar view was expressed in a similar study performed with the cations Li', Na', K', A13', C2HSNH: and (C2H,),NH', [456].In the vapour phase esterification of propionic acid with ethanol over a macroreticular catalyst [443],the effect of cations was not quite consistent with that found in the liquid phase reactions: an increasing inhibiting effect with decreasing cation valency was found (Me3' < Me2+ < Me').
(ii) Degree of polymer crosslinking. The degree of crosslinking of the copolymer is controlled by the amount of divinylbenzene (DVB) added t o the copolymerisation mixture, The effect of crosslinking upon the catalytic activity of the ion exchanger is dependent on the type of resin (standard or macroreticular) and on the conditions of its use (liquid or vapour phase). With the exception of macroreticular catalyst in vapour phase reactions [ 436,4371, the catalysts with a higher degree of crosslinking always exhibited lower activity in esterification or transesterification than (see the less crosslinked ones [ 404,428,429,431,436,437,445,451,4571 also Table 24). The difference in the effects of crosslinking in liquid and vapour phase processes, and especially the different behaviour of standard and macroTABLE 24 Effect of degree of crosslinking of the resin and comparison of heterogeneous and homogeneous rates (a) Esterification of acetic acid with ethanol at 75°C [ 4 2 8 ] Catalyst
Reaction rate (mol equiv-' h-')
Ion exchanger with DVB content (%) 20
1
22.8
82.8
H2S04
108
( b ) Transesterification of ethyl acetate with 1-propanol at 52°C [ 4 3 6 ]
Catalyst
Reaction rate (mol equiv-' h - l ) a
50
25
15
8
2
0
p-toluenesulphonic acid
0.19
0.25
0.46
1.88
4.07
8.50a
9.40
Ion exchanger with DVB content (%)
Extrapolated value.
References p p . 385-398
364
0
10
20
30
40
50 "la
60 DVB
Fig. 1 6 . Effect of degree o f crosslinking ( % DVB) of standard (non-porous) ion exchanger on initial transesterification rate, ro (mol kg-1 h-1 ), of ethyl acetate with 1-propanol [436]. ( 1 ) Liquid phase at 52OC; initial composition (mole%),0.4 ethyl acetate, 0.4 1-propanol, 0 . 2 dioxan (solvent). ( 2 ) Vapour phase at 12OoC; partial pressure of reactants, 0 . 5 bar (ester-alcohol ratio 1:l).
reticular catalysts, was demonstrated (Figs. 1 6 and 17) for transesterification of ethyl acetate with 1-propanol [436]. From Fig. 16, it follows that, with standard ion exchangers, the effect of crosslinking is greater for liquid than for vapour phases processes. This can be explained by the swelling of the polymer, which increases with decreasing crosslinking and is more extensive in the liquid phase. The swelling of standard polymers is very probably the reason why in the liquid phase the reaction rates are higher than in the vapour phase in spite of the lower temperatures used (see also ref. 428). The behaviour of macroreticular ion exchangers (Fig. 17) is different from that of the standard ones. The porosity (pores are mainly about 1 5 nm in diameter) of the macroreticular type, and correspondingly the internal surface area, are higher with higher crosslinked polymers. This would lead to higher reaction rates, but the effect of increasing surface area is compensated in the liquid phase by decreased swelling. Consequently, the catalytic activity of the ion exchangers is only slightly dependent on the degree of crosslinking (Fig. 17, curve 1).On the other hand, for vapour phase reactions, swelling is much less important and, since the surface area is larger with higher crosslinked polymers, the rate increases with the
365
Fig. 1 7 . Effect of degree o f crosslinking (W DVB) of macroreticular ion exchanger o n initial transesterification rate, ro (mol kg-1 h - l ) , o f ethyl acetate with 1-propanol [436].( 1 ) Liquid phase. (2)Vapour phase (reaction conditions the same as in Fig. 16). Dotted curves (1’) and ( 2 ’ ) represent reaction rates o n a surface o f area equal t o the surface area of the ion exchanger with 10%DVB.
degree of crosslinking (Fig. 17, curve 2). If, however, the rates are related to unit surface area (dotted curves in Fig. 17), they decrease with the degree of crosslinking. The crosslinking effect also depends on the structure of the reactants [437] as already discussed [Sect. 4.1.3.(b)]. In liquid phase reactions, the importance of the swelling properties and the related sorption capacities for the catalytic activity of ion exchangers was demonstrated. The rate coefficient of 1-butanolacetic acid esterification [431] decreased with the degree of crosslinking in the same manner as did the water sorption capacity and the solvation coefficient of l-butanol. A similar effect was found for the transesterification of ethyl acrylate with 1-butanol [ 4041. Davini and Tartarelli [ 4571 found an increase of activation energy and activation entropy with degree of crosslinking of the resin in the liquid phase esterification of benzoic acid with 1-butanol. This finding is in contradiction to the view of Bernhard and Hammett [366] according to which the resin structure imposes the more severe restraints upon the transition state the higher is its degree of crosslinking. Nevertheless, Davini and Tartarelli [ 4571 try to explain their observation by reference t o their sorption data. References pp. 385-398
It appears from the relations between the catalytic activity, the structure of the reactants, the degree of crosslinking and the swelling and sorption properties of the resin that all these factors might influence the accessibility of the reactants to the acid groups of the ion exchanger. Thus, the rate of transport of reactants (diffusion) t o the active groups may also become an important factor in the kinetics, especially with larger reactant molecules and more highly crosslinked polymers.
( d ) Effect o f water and solvents (i) Effect o f water. An inhibiting effect of water on esterification (greater than that which corresponds t o the reversibility of the reaction) was observed for both liquid [ 428,431,433,449,4521and vapour phase reactions [438,445,446].In the liquid phase where the resin is in a swollen state, the effect can be ascribed to the interaction of water molecules (in competition with the reactants) with the active protons of the catalyst rendering them less active or inactive. The effect of water has generally been considered qualitatively. However, the competition was expressed quantitatively in salicylic acid-methanol esterification [ 4491 by a Langmuir-Hinshelwood model and eqn. (24)is a simplified expression based on this model. In a similar manner, the inhibition by water was expressed by the Langmuir-Hinshelwood equation (29)for vapour phase esterification of acetic acid with 2-propanol [ 4381 or ethanol [ 4461. With vapour phases processes, however, some other effects were observed. Setinek and Ber6nek [439]found an activating effect of water in acetic acid-ethanol esterification when an ion exchanger, predried in vacuo at a higher temperature, was used. Cunningham et al. [438] observed a loss in catalyst weight (initial water content was 15-2096) and a decrease of the catalytic activity when the esterification was carried out in a flow of nitrogen at a partial pressure higher than 0.6 bar; the authors ascribed the effect t o a partial dehydration or “deswelling” of the resin, which decreased the site accessibility. Herrman [ 4461 reported an increase in the effective rate coefficient when the pressure was increased or the temperature decreased; he suggested that the number of active sites might change with the equilibrium amount of water in the ion exchanger which increases with pressure and decreases with temperature. All the cited experimental observations [ 438,439,4461could have a common explanation: the first effect of water vapour contacting a dry catalyst may consist in slight swelling of the resin thus increasing the amount of accessible sulphonic groups; this effect is reversible and dependent on the partial pressure of the water and on the temperature. With a stabilised resin with a water content above a certain limit, the accessibility of the active groups does not change significantly, however, and the effect of water then consists of competition for the active groups with other reactants, as mentioned above.
367 (ii) Effect of solvents. Because liquid phase esterifications or transesterifications are usually carried out in excess of one of the reactants (most frequently of alcohol) which serves also as a solvent, the influence of other solvents has not been thoroughly investigated. Nevertheless, the available results indicate that there may exist a relationship between the dielectric constant and the polar properties of the solvent, and the reaction rate. The solvents investigated can be classified into three groups: (a) non-polar (heptane, benzene, toluene), (b) polar solvents negatively charged at the oxygen atom and capable of solvating cations (tetrahydrofuran, dipropyl ether, dioxan) and (c) polar solvents not capable of solvating cations (nitropropane, nitroethane, nitromethane). The effect of group (a) (non-polar solvents) was examined in the esterification of benzoic acid with 1-butanol over a Dowex-W X2 catalyst [ 4541. The solvent affected both the Helfferich distribution coefficients and the esterification rates. Dielectric constants, corresponding t o the composition of the pore liquid, were estimated and the kinetic data related t o the polar properties of the medium within the catalyst. In Fig. 18 are plotted specific rate coefficients versus the reciprocal value of dielectric constant of the pore liquid. The slope of the correlation is positive as for
-5
-C -6
3
5
4
6
104/ DT Fig .18.Dependence of specific rate coefficient, h of the esterification benzoic acid + 1-butanol at 3OoC o n the reciprocal of the dielectric constant, D, of the liquid reaction medium [ 4541. ( 1 ) Homogeneously catalysed reaction. ( 2 ) Reaction catalysed by Dowex-W X2 ion exchanger. T,Absolute temperature; 00,without s o l v e n t ; o l , in heptane; a m , in benzene; AA, in toluene. R e f e r e n c e s p p . 385-398
368 the data on homogeneous catalysis, which is consistent with a reaction mechanism involving a positive ion and a dipolar molecule. The effect of solvents of groups (b) and ( c ) was investigated in transesterification of ethyl acrylate with l-butanol over a KU-2-8 catalyst [465]. It was established for group (b) and for n-heptane that the higher the sorption capacity of the resin for a given solvent, the lower was the rate of reaction in that solvent. This can be explained by the formation of a hydrogen bond between the polar aprotic solvent and the -S03H group, e.g. (C3H,),=O,\
--.H-SO3-
(C)
which lowers the activity of the catalyst, The sorption of non-polar heptane is negligible and the rate of reaction in thia solvent is the highest of the series. A linear dependence of the rate coefficient on the reciprocal of the dielectric constant was found, which indicates that the polarity of the medium determines its ability to solvate the active species which are assumed to be the protonated alcohol molecules, C4H90Hf, thus affecting the reaction rate. An interaction of the type above [scheme (C)], i.e. the formation of hydrogen bond between the oxygen of the solvent and the -S03H group of the catalyst, was proposed by Rodriguez [435,466] to explain the lower reaction rates observed in transesterification of ethyl acetate with l-propanol in dioxan compared with those in cyclohexane [see also eqn. (26)]. As t o solvents in group (c) (njtroalkanes), which hardly solvate cations, higher reaction rates were observed than with other solvents [465]; the rates increased with increasing dielectric constant and the sorption capacity of the resin for the solvent. This result indicates that a solvation of the anions -SO; by the nitroalkane molecules might occur by the formation of a hydrogen bond -S06,-....H6+-CH2N02 ; consequently, the activity of the protons of the sulphonic groups increases and the transesterification rate is enhanced.
( e ) Mechanism It follows from the dependence of the catalytic activity of ion exchangers on the amount of functional groups in the H+-form [Sect. 4.1.3.(c)] and on their acid strength [447], that acidity is the essential property of esterification catalysts and plays a decisive role in the reaction mechanism (see also ref. 438). However, lower specific rates were observed in catalysis by ion exchangers than if the reaction was catalysed by an equivalent amount of p-toluenesulphonic, sulphuric or hydrochloric acids [ 431,453, 454,460,462,4631. The difference between the homogeneous and heterogeneous catalytic reaction rates decreases, however, with decreasing degree of crosslinking of the polymer [Sect. 4.1.3.(c)] and by extrapolation to
369
zero divinylbenzene content, both rates were found t o be approximately equal (Table 24). Therefore, the mechanisms of both homogeneous and heterogeneous reactions might be expected t o be the same and the reduced reaction rate with ion exchanger catalysts could be simply ascribed to a lower number of sulphonic acid groups of the resin accessible to reactants. An equal effect of the dielectric constant of solvents, found for both types of catalytic esterification (Fig. 18), further supports the similarity of both mechanisms. The homogeneous acid-catalysed esterification has been thoroughly studied and a lot of information on its mechanism has been accumulated. The assumed close relation to the heterogeneous mechanism can, therefore, be made use of in mechanistic considerations about the esterification catalysed by ion exchangers. In the homogeneous mechanism, the reaction is assumed to start by protonation of one of the reactants, either ester (mechanisms denoted as AAcl and AAc2 [397,398]) or, less frequently, alcohol (mechanism AALl). It seems likely that protonation of reactants is an important step in esterification catalysed by ion exchangers, too. This follows from all that has been said above about the effect of the acidic properties of ion exchangers on their catalytic activity and is further supported by the effect of the dielectric constant of solvents (Fig. 18), which indicates that the reaction mechanism involves a positive ion and a dipolar molecule [ 4541. In most homogeneously catalysed esterifications, the protonated species is considerdd to be formed from the acid (or from the ester in transesterification) [ 397,398,4671. Starting with this view, Bochner et al. { 4491 assumed a mechanism for heterogeneous esterification [see eqn. (24)] in which the reaction between protonated salicylic acid and methanol in solution was the ratedetermining step. In oleic acid esterification with 1-butanol [452], an analogous mechanism was indicated [see eqn. (25)] by the kinetic analysis. On the other hand, the sorption of the acid was considered to be nonionic and found to decrease with increasing concentration of functional groups of the ion exchanger in the H+-form [455]. The amount of the alcohol sorbed was higher than that of the acid [455] or that of the ester in transesterification [ 403,404,464,4651. The rate coefficients of transesterification varied in the same way as did the amount of the sorbed alcohol when different alcohols [ 4641 and solvents [ 4651 were used or when the degree of crosslinking of the ion exchanger was altered [404]; such a relation did not exist between the reaction rate and the sorbed amount of the ester. These results could be interpreted on the basis that it is the alcohol that is preferentially adsorbed and protonated; the formation of the ion ROH; and its role as a reactive intermediate were explicitely formulated [ 404,4651. Two arguments, however, that weaken this interpretation may be raised: (i) the amount of a reactant sorbed in the bulk of the ion exchanger References p p . 385-398
370 is not necessarily a measure of the chemical interaction of the reactant with Lhe sulphonic acid groups; (ii) chemical interaction of the --S03H groups with the alcohol can consist not only in the proton transfer to the alcohol [scheme (a)] but can also take place through the formation of hydrogen bond between the hydroxyl group of the alcohol and the oxygen atom of the sulphonic group acting as a basic site [scheme (b)].
OH
RS-H /\o=s=o/ ' \ It has also been suggested that an alcohol molecule can be bound t o two o r more sulphonic groups by combined interaction of both types; simultaneous interaction with two -S03H groups, for example, is illustrated by the scheme (c).
OH
0 ''
Interactions of this types are assumed t o be involved in the mechanism of alcohol dehydration over sulphonated ion exchangers [ 112,131,138,4684701 from spectroscopic evidence [ 138,4711. Therefore, it cannot be ruled out that in esterification also more than one active site (-S03H group) may, in general, be involved in the reaction mechanism. Alcohol can be adsorbed on one or more active sites according t o schemes (a)-(c), and acid be protonated as represented by scheme (d). This assumption seems t o be supported by the dependence of the esterification rate on the concentration of sulphonic groups which is of an order higher than unity [see Sect. 4.1.3.(c)]. In a particular case, the adsorption of either the acid (ester in transesterification) or of the alcohol may prevail, and, in an extreme case, only one reactant would be adsorbed, the other reacting from a non-adsorbed state. These considerations can be formulated by schemes (e)-(g) below (R' = H o r alkyl). Scheme (e), in which the acid (ester) is protonated and alcohol reacts in non-adsorbed state, corresponds t o the mechanisms AAcl or AAC2 proposed for homogeneous esterification and hydrolysis; with ion exchanger catalysts, the mechanism (e) was assumed to be operating in the liquid phase esterification of salicyclic acid with methanol [ 4491 and in t h e transesterification of ethyl acetate with the same alcohol in dioxan as
371 solvent [435]. Scheme (f), according t o which alcohol is adsorbed and the acid (ester) reacts from the non-adsorbed state, was proposed for the reaction of ethyl acetate with 1-propanol in dioxan [435] and may perhaps also be operating in cases where protonation of the acid (ester) was not considered to be a likely step [404,454,457,465]. Mechanism AAll the first step of which is protonation of the alcohol, is rare even in homogeneous catalysis. There is no reason t o suppose that it would be more probable in heterogeneous catalysis. A more general scheme with both reactants adsorbed is that represented by scheme (g). It was considered probable in the liquid phase transesterification in a non-polar solvent (cyclohexane) [435] and it may correspond to all vapour phase esterifications and transesterifications where the rate equations from the kinetic analysis (see eqns. (27) and (29) suggested the involvement of dual- or triple-sites.
4.2 HYDROLYSIS
4.2.1 Hydrolysis of esters Hydrolysis of esters, producing acids and alcohols RCOOR' + H 2 0 = RCOOH + R'OH
(D)
is the reverse reaction t o esterification [scheme (A)] treated in Sect. 4.1. The conversion in such a reversible reaction may be limited by thermodynamic equilibrium (see equilibrium constants of esterification in Table 18). Kinetic studies of ester hydrolysis, however, have been almost exclusively performed in an excess of water which causes the reaction t o be practically irreversible. Although, in homogeneous media, ester hydrolyses are catalysed both by acids and bases, for kinetic investigations of heterogeneously catalysed hydrolysis, the use of solids with acidic properties only is reported, most of them being organic polymer ion exchangers. No side reactions, accompanying the main reaction (D), were observed under References p p , 385-398
37 2 the conditions used. The analogy with homogeneously catalysed ester hydrolysis determined the method of operation (liquid phase, mainly batch reactors) and the kinetic approach. With solid catalysts, the kinetics of reaction (D) were also assumed to be of the second order (see, for example, refs. 472, 473); in excess of water, the order degenerates and pseudo-first-order kinetic equations were used in most cases. In two papers concerning the hydrolysis of ethyl acetate over sulphonated styrenedivinylbenzene copolymers as catalysts [ 448,4511, LangmuirHinshelwood single site models with surface reaction as the ratedetermining step were applied [see eqn. (25)]. (a) Catalysis by inorganic catalysts The activities of six mixed oxides, Si02-A1203, SiO, -Zr02, SiO, TiOz, A1203-Zr02, TiOz-ZrO, and A1203-Ti02, in the hydrolysis of methyl and ethyl acetates at 25 and 35°C were compared and related to the acidic properties of the catalysts [474]. It was found that the pseudofirst-order rate coefficients for both esters and both temperatures were proportional t o the number of acid centres of pK, < -3.0 as determined by the l-butylamine titration of the used catalysts. Since the points for HC1 lay on the same straight line as did the points €or inorganic catalysts, it was believed that the acid centres of pK, < -3.0 of these catalysts behaved even in the presence of water, very much like hydrochloric acid or any other mineral acid. This view was further supported by the fact that, by poisoning the catalysts with dicinnamalacetone (pK, = -3.0), the hydrolysis was fully suppressed. The values of the observed activation energies for all solid catalysts and HC1 (24.7-25.5 and 33.9-35.5 kJ mol-' for methyl and ethyl acetate, respectively) were closely similar and indicate a similarity of the mechanism of hydrolysis with all catalysts and again support the view that there is no substantial difference in the action of dissolved HCl and solid acid catalysts in ester hydrolysis. Specific rate coefficients (related to unit amount of acid centres) were approximately the same for solid catalysts as well as for HC1 [474]. However, when a montmorillonite clay activated by adsorption of protons on its surface was used as the catalyst in ethyl acetate hydrolysis [475], a higher specific rate coefficient (about 1.8 times at 25°C) was found for the reaction catalysed by adsorbed protons than by dissolved acid, this result being explained by the authors by an increase of activation entropy in the former case. (b) Catalysis by organic ion exchangers
The majority of kinetic studies of ester hydrolyses using this type of catalyst were performed with sulphonated styrenedivinylbenzene copolymers. Only in a few cases [476-4781 was the use of phenol-formalde-
373 hyde polycondensates with either -S03H or -€OOH groups reported. The main interest in kinetic investigations with organic ion exchanger catalysts was concentrated on the effects of reactant structure, solvents, the degree of cross-linking of the polymer and partial neutralisation of acid groups with different cations. These effects are usually interrelated and t o discuss each of them separately is of little value. Since, with ion exchangers, the specific properties of the solid acid are of major interest, the efficiency of the catalyst for a given reaction was defined [366] by the relation 4
khet
=-
khom
where khet is the rate coefficient for the reaction catalysed by a solid resin and khom that for the same reaction catalysed by a dissolved inorganic acid, such as HC1, both coefficients being related t o the same amount of -S03H groups or protons in a unit reaction volume.
(i) Effect of reactant structure and solvent. It has been shown that the effect of the ester structure on both the absolute values of the rate coefficients and the efficiencies is strongly dependent on the properties of the reaction medium used (addition t o water, of solvents which are capable of solvating cations). If the reactions are carried out in acetone-water mixtures [366,479-4811, a decrease of the reactivity as well as of the. efficiency (this being always less than unity) with increasing length of the aliphatic chain in the acyl or in the alkoxy groups was observed. Similar trends were established for the structure effects using water-dioxan mixtures [ 482-4841 o r in excess of the reacting ester [ 485-4871. When, on the other hand the reaction was carried out in excess of water without any other solvent added, other effects of the ester structure on the reactivity and the efficiency of the catalyst were observed [476,488, 4891. The efficiencies, q, were higher than unity (see also ref. 490), which means that the resin-catalysed reaction was faster than that catalysed by HCl; the q values increased with increasing chain length of the alkyl group [476], contrary t o what was found with mixtures of water and other solvents. Several attempts have been made t o explain the variations in efficiency of ion exchanger catalysts for different esters and reaction media. Hammett et al. [ 366,479,4881 suggested that the difference in efficiency for different esters arises from a difference in the magnitude of the loss in internal entropy of the ester molecule which accompanies its fixation on the resin catalyst in the formation of the transition state. It can be shown that the ratio of efficiencies for two esters, 1and 2, is given by
*
-RT ln(q*/q1) = (G:,het - G2,horn) - (G:,het - G:,hom) where the G*’s are the standard free energies of the transition states and References p p . 385-398
37 4
I
01 40
I 60
I
I 100
ao
- 2.c
- 1.5 0-
0" -1.0
- 0.5
0 40
I 60
I 80
I
100
Fig. 1 9 . The efficiency, q , of t h e Amberlite IR-120 ion exchanger catalyst in ester hydrolysis as a function of t h e entropy, S , of t h e parent hydrocarbon RH or R'H of the substituents. ( a ) Hydrolysis (at 25-45OC) of methyl esters RCOOCH3 [ 3661 : 1, acet a t e ; 2, chloroacetate; 3, benzoate; 4, cyclopentanecarboxylate; 5 , phenylacetate; 6, a-naphthylacetate; 7 , l-octanoate. ( b ) Hydrolysis (at 35OC) of acetates CH3COOR' [480]:1 , methyl; 2, ethyl; 3 , cyclopentyl; 4 , cyclohexyl; 5 , 1-butyl; 6 , 2-pentyl; 7 , 1-pentyl; 8, 1-hexyl; 9 , 1-octyl.
37 5 the subscripts refer either t o heterogeneous or t o homogeneous systems. Since the differences in activation energy between the two systems were found [479] (for reactions in 70% aqueous acetone) to be small (of about -5.4 k J mol-I) and were little, if at all, affected by the chain length of the ester, it follows that it is the entropy difference that is controlling the efficiency of the resin for different esters (values of -23.0 and -37.6 J mol-' K-' were observed [479] for methyl acetate and ethyl butyrate respectively). It seems reasonable that the loss in entropy should be the smaller, the smaller is the entropy of the ester. This hypothesis has a surprising degree of quantitative validity, as is shown in Fig. 19, where, in (a), the efficiencies, 4, of the hydrolysis of seven methyl esters, RCOOCH,, are compared with the entropy, S, at 298 K of their simplest structural analogues, the gaseous compounds RH and, in (b), the efficiencies in the alkyl acetate series CH,COOR' are plotted against the S values of the parent hydrocarbon, R'H. A drop in efficiency with increasing degree of crosslinking of the resin was also observed [366], this effect being much greater with the larger molecules of ethyl hexanoate than with methyl acetate. The results are considered by Hammett et al. to be consistent with their hypothesis: the structure of the more highly crosslinked resin imposes more severe restraints in the formation of the transition state than does the structure of the less crosslinked resin. Bernhard and Hammet [ 4881 tried t o explain the apparent discrepancy between the results [366,479] for 70% aqueous acetone and those with water (Table 25). They found that, with water, the activation entropies were, in contrast t o those using aqueous acetone, higher for the resin than for HCl. The observation is consistent with the following hypothesis. In the mixed solvents, the electrically charged transition state is subjected t o restraints arising from solvation, which are stronger than those acting on the ester. Still greater restraints imposed by the resin network lead t o a still greater decrease in the entropy of activation with increasing chain length. In water solution, however, the solvation of the polar group in the ester itself is so powerful that no further important restraints are imposed on the internal motion of the ester by the solvation of the charged transiTABLE 25 Effect of ester structure o n the efficiency of ion exchanger in water [ 4 7 6 ] (Catalyst: Amberlite IR-100; temperature: 25°C.) Ester ___.
Methyl acetate Ethyl acetate 1-Butyl acetate Benzyl acetate
Efficiency ~
~
_
_
-~ ._ _
References p p . 385-398
_
_
_
1.8 2.3 10 20 ~
-
__-
376 tion state or by the resin network. Consequently, the entropy change involved in the conversion of the ester t o the transition state is no longer more negative for esters of a greater chain length. Another approach attempts t o explain the different effect of the ester structure in different reaction media simply by the changing ability of the esters t o be absorbed by the resin. Qualitatively, this approach was used [476] t o interpret the results for water and aqueous acetone and a similar idea was suggested for the hydrolysis of dicarboxylic acid esters in water-dioxan mixtures [ 482,4831. Quantitative interpretation was based [481,489] on Helfferich’s model [427]. It follows from eqn. (30) and from the relation -
(31) resulting from the Helfferich model [eqn. (23)] ( h h e t is the rate coefficient corresponding to the concentration of ester in the pore liquid) that hhet = h h e t h
-
4 =-
hhet
(32)
hhom
The experimental evidence of the validity of eqn. (32) was illustrated by a plot of the experimental q versus data; the slope obtained, z h e t / h h o m = 1, means that the reactivity of the ester in the pore liquid (i;h,t) of the resin is the same as in the H2S04-catalysedreaction ( h h o m ) for all the esters investigated [ 4891. Thus, the effect of ester structure is believed to consist only of influencing the distribution coefficients, A , of the ester between the pore liquid and the supernatant solution, which is in accordance with Helfferich’s model (see also ref. 491). Investigating the effect of acetone concentration in acetone-water mixtures on ethyl, 1-propyl and 2-propyl acetate hydrolysis with Dowex X50 containing 10% DVB, Tartarelli e t al. [481] found that the rate coefficients and the efficiencies decreased with increasing acetone concentration (cf. ref. 477). The inversion of the efficiencies, q , from positive to negative values took place at about 45-5576 acetone concentration (see also refs. 478 and 482). The authors [481] plotted the rate coefficients found at different acetone concentrations for the three esters in the homogeneous and heterogeneous reactions and the corresponding distribution coefficients as hhomh versus k h e t and obtained a straight-line dependence with the slope equal t o unity. This means that k h o m x = h h e t ; from this and the relation (31), it again follows that jr;het = h h o m , in agreement with the results of Tartarelli e t al. [489] and with Helfferich’s model.
(ii) Partial neutralisation of ion exchanger. The two theories, that of Hammett and that based on the Helfferich concept, were also used t o interpret the effect on the hydrolysis rate of the partial neutralisation of acid
377 groups of the ion exchanger by different cations. When metal ions (Na+, MgZ+,Ba2+)and NH4' were used for neutralisation, the dependence of the rate coefficient on the proton concentration in the resin was found t o be first order [ 489,4921. In another study [ 4931, a slight decrease of the specific rate coefficient with increasing degree of neutralisation with Na' ions was reported. When cations of larger dimensions such as ethylenediammonium ions or other quarternary ammonium ions with larger organic substituents were used for the neutralisation, the specific rate coefficients were found to change with the degree of neutralisation [ 492,494,4951, this effect being more pronounced in the case of an ester with a longer acyl chain [ 4921. The effect of partial neutralisation by cations of the last mentioned type was, in general, negative. An unexpected enhancement of the specific rate coefficient ( h h e t )was, however, established when the neutralising quarternary ion and the ester had some prominent structural features in common [ 4941 (e.g. hexadecyltrimethylammonium ion in the hydrolysis of ethyl hexanoate, an ester with a rather long aliphatic chain). The interpretation presented by Riesz and Hammett [494] follows from the old principle that like dissolves like. In the case above, this is t o be interpreted in the sense that increasing incorporation into the resin of longchain aliphatic structures lowers the standard free energy of the transition state for the hydrolysis of an ester containing similar structures relative t o standard free energies of the transition states of esters of a different structure. The alternative approach to the problem is based on Helfferich's model. Tartarelli et al. [495] measured the distribution coefficients, A, and specific rate coefficients, k h e t , of the hydrolysis of 1-propyl acetate and ethyl 1-hexanoate; the ion exchanger catalyst was neutralised t o a different degree with benzyldimethylhexadecyl- and trimethylbenzyl-ammonium ions. The specific rate coefficients for the pore liquid of the resin, h h e t , were calculated according t o eqn. (31).The distributioncoefficients, A, of the aliphatic ester were increased and the coefficients hhet lowered when the ion exchanger contained ions with long chain aliphatic groups (benzyldimethylhexadecylammonium ions) and the reverse was true when it contained trimethylbenzylammonium ions (without any long aliphatic chain). The authors explain the results by assuming that the matrix of the resin interacts with the ester being sorbed, involving its reactive group. Hence, the stronger these intereactions are, the higher is the distribution coefficient but the lower is the reactivity.
(iii) The degree of crosslinking and diffusion. The efficiency of the ion exchanger catalyst in ester hydrolysis, as in esterification (Sects. 4.1.3.(b) and (c)], decreases as the degree of crosslinking of the resin increases [ 366,483,484,488,490,4931, this effect being more pronounced with larger ester molecules. Bernhard and Hammett [366] tried t o explain the References p p . 385-398
378 phenomenon by restraints imposed by the resin on the activated state (see p. 375). However, Goldenshtein and Freidlin [ 483,4841 advanced another idea, relating the influence of crosslinking simply to geometric effects of the resin network. Smaller molecules are able t o penetrate into the polymer mass whereas the reaction of the larger ones can take place only on the outer surface of the resin particles, the permeability of the resin being dependent on its degree of crosslinking. A similar idea was put forward [433] in the discussion of the effects of the catalyst surface area on the esterification of oleic acid, and will also appear in the discussion of other hydrolytic reactions [Sect. 4.2.2(b)]. This “sieve effect’’ cannot be considered statically as a factor that only determines the amount of accessible acid groups in the resin in such a way that the boundary between the accessible and non-accessible groups would be sharp. It must be treated dynamically, i.e. the rates of the diffusion of reactants into the polymer mass must be taken into account. With the use of the Thiele’s concept about the diffusion into catalyst pores, the effectiveness factors, Thiele moduli and effective diffusion coefficients can be determined from the effect of the catalyst particle size. The apparent rates of the methyl and ethyl acetate hydrolysis [490] were corrected for the effect of diffusion in the resin by the use of the effectiveness factors, the difference in ester concentration between swollen resin phase and bulk solution being taken into account. The intrinsic rate coefficients, h i n t r ,
TABLE 26 Effect of diffusion and adsorption in methyl and ethyl acetate hydrolysis in water at 40°C (Adsorption coefficients, K A , (dimensionless), effective diffusion coefficients, D , f , (cm2 sec-‘) and intrinsic rate coefficients, k i n t r , (cm3 equiv-’ min-’) for Amberlite ion exchanger catalysts of different degrees of crosslinking [490].) Ester
DVB in resin
KA
D,f
X
lo6
kintr
(%I Ethyl acetate
Methyl acetate
4
8 10 12
1.23 1.07 1.02 1.02
4 8 10 12
1.oo 0.88 0.83 0.85
0.513 0.250
32.6 32.5 32.0 32.3
0.852 0.452
33.1 35.5 38.0 35.3
From hintr and data for inorganic acid, the mean values of intrinsic efficiency q i n t r were calculated: for ethyl acetate qintr = 1.28 and for methyl acetate q i n t r = 1.45.
379 were calculated according t o the formula
where khet is the measured specific rate coefficient, 7 the effectiveness factor (ratio of diffusion limited reaction rate t o the non-limited one) and K A the ratio of ester concentration in the volume of the swollen resin to that in the bulk solution. As can be seen from Table 26, the intrinsic rates of hydrolysis are independent of the degree of crosslinking, whereas the adsorption of ester by the resin and its diffusivitity in the resin decrease with crosslinking. It might, therefore, be assumed that the lower hydrolysis rates observed with higher crosslinked polymers in this [ 4901 as well as in other studies [ 366,483,484,4881 were very probably due t o increased diffusion resistance and decreased adsorption capacity of the more tightly crosslinked resins. I t is not impossible that some effects of solvents, ester structure, or introduced cations, interpreted either by changes of activation free energy [366,479,488,492] o r reactant concentration in the pore liquid [476, 481,489,4951, could also be satisfactorily explained by the changes in the accessibility of catalytically active groups and in the diffusivity of reactants. Higher hydrolysis rates and resin efficiencies were always observed, for example, in reaction media in which the resins swell more, e.g. in water [ 476,481,4951 and alcohols [478]. In solvents in which the resins are less swollen (acetone, dioxan, esters), the rates were lower [ 366,477, 479-481,485-487,4951. The lower rates observed in solvents with negatively charged oxygen atoms, such as acetone and dioxan, might also be due, at least partially, t o the competition of the solvent with the reacting ester for the --S03H groups via an interaction illustrated by scheme (C) (p. 368), as was presumed [ 435,465,4661 for transesterification.
(iu) Mechanism. N o special investigations of the mechanism of the ion exchanger-catalysed hydrolysis of esters were reported. In most papers, the analogy with the mechanism of acid-catalysed hydrolysis in homogeneous medium [ 397,3981 is impIicitly assumed. Haskell and Hammett [479] pointed out that, in resin-catalysed hydrolysis, as in that catalysed by strong aqueous acids, the mechanism involves a positively charged transition state in which a proton has been transferred t o the ester molecule from the oxonium ion of the acid. The view seems t o be in harmony with the acidic properties of ion exchanger catalysts and with the dependence of their activity on the strength of the acid groups in the polymer [476] as well as on their concentration [489,492,494,495]. Some analogies established between homogeneous and heterogeneous hydrolyses in the structure effects of the ester [366,478-480,4881 and in the effects of solvents [478,481] can be considered to support the concept of the identity of both mechanisms. R e f e r e n c e s p p . 385-398
380 4.2.2 Other hydrolyses
In comparison with ester hydrolysis, little attention has been paid to kinetic investigation of other solid-catalysed hydrolytic reactions, though they could be numerous with respect to the types of functional groups to be hydrolysed. With inorganic solid catalysts, some studies were devoted to the hydrolysis of carbonic acid derivatives, aryl chlorides and compounds with metalloid-hydrogen bonds. With organic ion exchangers, the reactions studied served rather as a model for the investigation of the catalytic action of, and transport phenomena in, the polymer catalysts. A typical hydrolytic reaction used for this purpose is sucrose inversion; limited attention was paid t o acetal hydrolysis. (a) Hydrolysis over inorganic catalysts
The kinetics of the hydrolysis of diethyl carbonate
CO(OC2HS)Z + H2O
=
C02 + 2 CzH50H
was investigated at 160-270°C in the vapour phase (excess of water) in a flow system [496]. About fifteen oxide or metal salt catalysts were compared. The first step in the formation of the activated complex over the majority of catalysts (those with ionic surfaces) is assumed t o involve reaction with the hydroxyl groups of water dissociated on the surface. The activated complex may be pictured as a semi-ionic system of water and diethyl carbonate polarised by the dissociative adsorption of water on an ionic surface. The hydrolysis of carbonyl sulphide
COS + H2O = C02 + H2S in the vapour phase in flow systems was investigated [497-4991. Firstorder kinetics with respect t o carbobyl sulphide and zero-order with respect to water (used in excess) were found. Catalysts with basic properties exhibited high activity whereas those with acidic properties were almost inactive [ 4971 ; alumina, being a good catalyst, is assumed t o act as a basic catalyst. Addition of sodium hydroxide t o the catalyst enhanced the activity considerably [ 498,4991. Alumina and cobalt molybdate supported on alumina were found to be equally active catalysts [499]. The surface of the catalyst is believed t o be partially covered with hydroxyl ions which are formed by dissociative adsorption of water on basic sites [4981 HOH + B(s)+ BH'(s)+ OHwhere B(s) is the basic site on the surface of the catalyts. The kinetics of vapour phase hydrolysis of aryl chlorides
X,C,H,-Cl+
H20 = XnC6H,,0H
+ HCI
381 was studied on catalysts containing lanthanum, cerium and other rareearth phosphates [500]; X was a methyl group and the number, n, of the substituents on the aromatic ring was 0, 1 or 2. The relative reactivity increased from chlorobenzene t o chloroxylenes. The reaction was inhibited at high conversions by the product HCl. The hydrolysis occurs via a reaction of adsorbed aryl chloride and a surface hydroxyl group to give a phenolic product and a surface chloride. The reaction of steam with the surface chloride produces HCl and regenerates the surface hydroxyls. Hydrolytic reactions of compounds with Si-H or B-H bonds in the liquid phase with great excess of water were investigated; zero reaction order with respect t o HzO was found. In the hydrolysis of para-substituted phenyldimethylsilanes at 20°C + H,O p-XC6H4Si(CH3)*H
= p-XC,H,Si(CH,),OH
+ H,
a Pd/C catalyst was used and the effect of substituents (X = H, C1, CH3, CH,O and CzH,O) on the reaction rate (first order with respect to the silane) was correlated by the Hammett equation [ 5011. The value of the p constant was +0.73, which indicated a nucleophilic substitution mechanism with a slightly polarised transition complex; very probably a SN2 mechanism was operating. The hydrolysis of the borohydride anion BH, + 4 H,O
=
B(OH), + 4 H2
was catalysed by rhodium, ruthenium and iron at 20-60"C in 1 N NaOH solution [ 502-5041. The reaction was formally of a fractional order (0 < n < 1) with respect t o the borohydride anion. The variation of the rate with initial concentration of borohydride led t o the conclusion that the monomolecular surface reaction of adsorbed BH, is the ratedetermining step according to a Langmuir-Hinshelwood mechanism. ( b ) Hydrolysis catalysed b y ion exchangers
Sucrose inversion, a typical proton-catalysed irreversible reaction C,zH,ZOl, + HZO sucrose
H+
=
C ~ H U O+ , C6H1206 glucose
fructose
has been investigated by several authors with the use of ion exchanger catalysts, mainly in order t o get specific information on the catalytic action of this type of solid catalyst when a large reactant molecule is used. The reaction in excess water was studied a t temperatures in the range 25-100°C (the order with respect t o sucrose concentration was found to be one). The effects of the particle size, degree of crosslinking and partial neutralisation of sulphonic groups indicated the strong influence of the diffusion of the reaction components into or out of the resin particle on References p p . 385-398
382 TABLE 27 Effect of degree of crosslinking of the polymer catalyst o n the rate coefficient ( k ) of sucrose inversion a t 50°C [505]
DVB
k
(%)
(min-')
1
199.2 110.3 26.3 3 .O 0.7
4
10 15 20
X
lo4
the reaction rate. With increasing particle size, the reaction rate or the first-order rate coefficients decreased [ 505-5071 . Table 27 illustrates the effect of the degree of crosslinking of the resin; the effect on sucrose inversion rates is much greater than was observed for esterification and transesterification (p. 363) or for ester hydrolysis (p. 377). The difference can be attributed to the larger size of sucrose molecules compared with that of acids and esters used in the other reactions mentioned. Figure 20,
Fig. 20. Effect of degree of crosslinking (% DVB) of a standard ion exchanger on the diffusivities, D,f (cm2 min-I), and t h e selectivity ratio, S = k e f S / k e f A c ( k , f = effective rate coefficient, S = sucrose, Ac = ethyl acetate). Data were obtained by rate measurements and Wheeler-Thiele analysis of simultaneous sucrose and ethyl acetate hydrolysis a t 70°C [508].
383 presenting the results obtained by rate measurements of sucrose inversion and ethyl acetate hydrolysis proceeding simultaneously in the same reaction system [ 5081, clearly demonstrates the situation. The decrease of the diffusivity values with increasing degree of crosslinking, obtained by Wheeler-Thiele analysis, is much steeper for sucrose, the reactant with the larger molecule, than for ethyl acetate. This is the reason for the changes of the selectivity ratio k e f S / k e f A cwith degree of crosslinking in favour of the reactant with smaller molecules. The diffusion of the large sucrose molecule may be so slow that a large proportion of the sulphonic acid groups inside the polymer become inaccessible and cannot participate in the reaction [506,508], as was also assumed for esterification (p. 361) and ester hydrolysis (p. 377). Other results also confirm the important role of internal diffusion. Experimental activation energies (67-75 kJ mol-') of the sucrose inversion catalysed by ion exchangers [ 506-5091 were considerably lower than those of a homogeneously catalysed reaction (105-121 kJ mol-') [ 505, 506,5081 and were close t o the arithmetic average of the activation energy for the chemical reaction and for the diffusion in pores. The dependence of the rate coefficient on the concentration in the resin of functional groups in the H'-form was found t o be of an order lower than unity. A theoretical analysis based on the Wheeler-Thiele model for a reaction coupled with intraparticle diffusion in a spherical bead revealed [ 510,5111 that the dependence of the experimental rate coefficient on acid group concentration should be close to those found experimentally (orders, 0.65 and 0.53 for neutralisation with Na' and K' ions respectively [511] or -0.5 with Na' ions [510]). Intrinsic rate coefficients, k i n t r , of sucrose inversion catalysed by Dowex 50W-X8 were evaluated [ 5061 by correcting the experimentally observed coefficients, k O b s using , the relation
hkintr 7 where X is the experimentally determined absorption coefficient of sucrose in the resin and 77 is the catalyst effectiveness factor calculated by Wheeler-Thiele analysis from rate data on catalyst beads of different size. Good agreement between the calculated effective diffusivities of sucrose and those obtained by non-equilibrium sorption measurements demonstrated that the diffusion model used was reasonable. The intrinsic rate coefficients of the reaction in the resin phase were only 60% as large as those determined in 3 N benzenesulphonic, p-toluenesulphonic or hydrochloric acid solution. According t o Gilliland et al. [506] one possible explanation of this discrepancy is that steric hindrance prevents close approach or favorable alignment of a sucrose molecule at an active site for inversion. A similar idea was suggested by Murakami and Mori [ 5121, who assume that only a number of protons smaller than one tenth of those present in the resin grains may effectively catalyse the sucrose kobs =
R e f e r e n c e s p p . 385-398
384 inversion because protons are much too close to each other on the resin surface compared with the dimensions of a sucrose molecule. The hydrolysis of acetals acetal + H,O
=
2 alcohol + aldehyde (ketone)
is another acid-catalysed reaction that was used to investigate transport phenomena in ion exchange resins [ 508,5131. First-order rate coefficients of Dowex 5OW-catalysed reaction ( k r e s ) of a series of acetals at 20°C were plotted against the rate coefficients of reactions catalysed by dissolved HCl ( k h o m ) . As Fig. 21 shows, the hydrolysis rate of less reactive acetals is controlled by a chemical reaction: the slope is near unity and the coefficients k,,, and khom are directly proportional t o one another kres = C k h o m
The hydrolysis of more reactive acetals is influenced by the rate of intraparticle diffusion, since the slope was found to be close to 0.5, which is in agreement with eqn. (33) derived from the Wheeler-Thiele model for a
-1
I
I
I
I
I
I
I
-2
I -1
I
-3
0
:-2 a! m 0
-3
log
khorn
Fig. 21. Hydrolysis of acetals at 2OoC o n a Dowex 5 0 W X10 resin catalyst [513 1. Rate coefficients of the resin-catalysed reaction (kres) versus rate coefficients of the reaction catalysed by dissolved inorganic acid (hhom). 1, Formaldehyde dimethylacetal ; 2 , formaldehyde diethylacetal; 3, formaldehyde di-2-propylacetal; 4 , acetaldehyde ethyleneacetal; 5 , acetone ethyleneacetal; 6 , acetaldehyde dimethylacetal; 7 , acetalde0.5. hyde diethylacetal. The slope for acetals 1-3 is 1 , for the acetals 3-7
-
-
385 TABLE 28 Effect of acetal structure o n the diffusivity (D,f) in a bead of ion exchanger [ 5 0 8 ] (Catalyst: Dowex 5OW-X12; temperature: 20°C.) Acetal
DefX lo6 (cm2 min-’)
Acetone ethyleneacetal Acetaldehyde dimethylacetal Acetaldehyde diethylacetal Benzaldehyde di( 2-buty1)acetal
10.0 6.6 2.6 0.008 -.
first-order reaction coupled with intraparticle diffusion in spherical beads
(33) The transition from the kinetic t o the internal diffusion region was also shown by the change in the activation energy [ 508,5131. The rate coefficient for an acetal with a much larger molecule, benzaldehyde di( 2-butyl) acetal, was, however, so low that it could not be correlated with the corresponding khom by the above relationship. This can be explained by a steric effect of this bulky reactant whose diffusion is much more hindered than that of smaller molecules (“sieve effect”, see also pp. 361 and 378). Effective diffusion coefficients of some acetals were estimated [ 5081 from the kinetic data for Dowex 5OW-X12 catalyst using the Wheeler-Thiele model (Table 28). The great difference (by about three orders of magnitude) between the diffusivity of benzaldehyde di-(2-buty1)acetal and that of the other acetals is evident. kres
=C’(hh~m~ef)~’~
References 1 2 3 4 5
6 7 8 9 10 11 12 13
A.V. Frost, Zh. Fiz. Khim., 1 4 (1940) 1313. C.L. Thomas, Ind. Eng. Chem., 4 1 (1949) 2564. B.S. Greensfelder, H.H. Voge and G.M. Good, Ind. Eng. Chem., 4 1 (1949) 2573. F.C. Whitmore, Chem. Eng. News, 2 6 (1948) 668. K. Tanabe, Solid Acids and Bases, Kodansha, Tokyo and Academic Press, New York, 1 9 7 0 . A. Eucken and E. Wicke, Naturwissenschaften, 32 ( 1 9 4 5 ) 161. H. Pines and J. Manassen, Adv. Catal., 1 6 ( 1 9 6 6 ) 49. F. Figueras, A. Nohl, L. d e Mourgues and Y. Trambouze, Trans. Faraday SOC.,67 (1971) 1155. H. Noller and W. Kladnig, Catal. Rev., 13 (1976) 1 4 9 . H. Pines and W.O. Haag, J. Am. Chem. SOC.,8 2 (1960) 2471. H.P. Boehm, Adv. Catal., 1 6 (1966) 1 7 9 . C. Okkerse, in B.G. Linsen (Ed.), Physical and Chemical Aspects of Adsorbents and Catalysts, Academic Press, London, 1 9 7 0 , p. 213. A.V. Kiselev and V.I. Lygin, Infrakrasnye Spektry Poverkhnostnykh Soedinenii i Adsorbirovanykh Veshchestv, Nauka, Moscow, 1972.
1 4 V. Ya. Davydov, L.T. Zhuravlev and A.V. Kiselev, Zh. Fiz. Khim., 38 (1964) 2047. 1 5 V. Ya. Davydov, A.V. Kiselev and L.T. Zhuravelv, Trans. Faraday Soc., 60 (1964) 2254. 16 J.B. Peri, J. Phys. Chem., 70 (1966) 2937. 17 J.B. Peri and A.L. Hensley, J. Phys. Chem., 72 (1968) 2926. 18 O.G. Armistead, A.J. Tyler, F.H. Hambleton, S.A. Mitchell and J.A. Hockey, J . Phys. Chem., 7 3 (1969) 3947. 1 9 K.G. Miesserov, J. Catal., 1 3 (1969) 169. 20 O.M. Dzhigit, A.V. Kiselev, N.N. Mikos-Avgul and K.D. Sherbakova, Dokl. Akad. Nauk SSSR, 70 (1950) 441. 21 E. Borello, A. Zecchina and C. Morterra, J. Phys. Chem., 7 1 (1967) 2938. 22 M.J.D. Low, Y.E. Rhodes and P.D. Orphanos, J . Catal., 40 (1975) 236. 23 H. Jeziorowski, H. Knozinger, W. Meye and H.D. Miiller, J. Chem. Soc., Faraday Trans. I, 69 (1973) 1744. 24 B.C. Lippens and J.J. Steggerda, in B.G. Linsen (Ed.), Physical and Chemical Aspects of Adsorbents and Catalysts, Academic Press, London, 1970, p. 171. 25 J.B. Peri, J. Phys. Chem., 69 (1965) 220. 26 J.B. Peri, J. Phys. Chem., 69 (1965) 211. 27 J.B. Peri, J. Phys. Chem., 69 (1965) 231. 28 G.M. Schwab and H. Kral, Proc. 3rd Int. Congr. Catal., Amsterdam, 1964, NorthHolland, Amsterdam, 1965, Vol. I, p. 433. 29 M. Yamadaya, K. Shimomura, T . Konoshita and H. Uchida, Shokubai, 7 (1965) 313; see ref. 5, p. 48. 30 N.S. Figoli, S.A. Hillar and J.M. Parera, J. Catal., 20 (1971) 230. 31 A.A. Baushkin and A.V. Uvarov, Dokl. Akad. Nauk SSSR, 110 (1956) 581. 32 R.O. Kagel, J. Phys. Chem., 7 1 (1967) 844. 33 E.V. Ballou, R.T. Barth and R.A. Flinn, J. Phys. Chem., 65 (1961) 1639. 34 V.A. Chernov and T.V. Antipina, Kinet. Katal., 7 (1966) 739. 35 H. Gerberich, F.E. Lutinski and W.K. Xall, J. Catal., 6 (1966) 209. 36 Yu. A. Bitepazh, Zh. Obshch. Khim., 17 (1947) 199. 37 R.S. Hansford, Ind. Eng. Chem., 39 (1947) 849. 38 H.A. Benesi, J. Phys. Chem., 6 1 (1957) 970. 39 A.P. Ballod and K.V. Topchieva, Usp. Khim., 20 (1951) 161. 40 G.A. Oblad, T.H. Millikan and G.A. Mills, Adv. Catal., 3 (1951) 199. 41 L.B. Ryland, M.W. Tamele and J.N. Wilson, in P.H. Emmett (Ed.), Catalysis, Reinhold, New York, 1960, Vol. VII, p . 1. 42 J.B. Peri, J. Catal., 41 (1976) 227. 4 3 P.B. Venuto and P.S. Landis, Adv. Catal., 18 (1968) 259. 44 P.B. Weisz and V.J. Frillete, J. Phys. Chem., 64 (1960) 382. 45 J.A. Rabo, P.E. Pickert, D.N. Stamires and J.E. Boyle, Actes 2Bme Congr. Int. Catal., Paris, 1960, Technip, Paris, 1961, Vol. 11, p. 2055. 46 H. Hattori and T. Shiba, J. Catal., 1 2 (1968) 111. 47 J.A. Rabo, C.L. Angel1 and V. Shomaker, Proc. 4th Int. Congr. Catal., Moscow, 1968, AkadBmiai Kiad6, Budapest, 1971, Vol. 11, p. 96. 48 Ya. I. Isakov, A.L. Klyachko-Gurvich, A.T. Khudiev, Kh. M. Minachev and A.M. Rubinshtein, Proc. 4th Int. Congr. Catal., Moscow, 1968. Akademiai Kiado, Budapest, 1971, Vol. 11, p. 123. 49 A. BielaAski and J. Datka, J. Catal., 32 (1974) 183. 50 K. Tanaka and A. Ozaki, J. Catal., 8 (1967) 1. 51 I. Mochida, A. Kato and T. Seiyama, J. Catal., 22 (1971) 23. 52 G. Zundel, H. Noller and G.M. Schwab, Z. Elektrochem., 66 (1962) 129. 53 G. Zundel and H. Metzger, Z. Phys. Chem. (Frankfurt am Main), 59 (1968) 225.
54 G. Zundel, Hydration and Intermolecular Interaction, Academic Press, New York, 1969. 55 K. Kochloefl, M. Kraus and V. BaZant, Proc. 4th Int. Congr. Catal., Moscow, 1968, AkadCmiai Kiad6, Budapest, 1971, Vol. 11, p. 490. 56 C.L. Kibby and W.K. Hall, J. Catal., 29 (1973) 144. 57 M. Boudart, Kinetics of Chemical Processes, Prentice-Hall, Englewood Cliffs, 1968, p. 102. 58 M. Balikovi and L. Berinek, Collect. Czech. Chem. Commun., 42 (1977) 2352. 59 K.R. Bakshi and G.R. Gavalas, J. Catal., 38 (1975) 312. 60 K.R. Bakshi and G.R. Gavalas, J . Catal., 38 (1975) 326. 6 1 K. Setinek and L. Berinek, Collect. Czech. Chem. Commun., 38 (1973) 3790. 62 S.L. Kiperman, Vvedenie v Kinetiku Geterogennykh Kataliticheskikh Reaktsii, Nauka, Moscow, 1964. 63 M. Kraus, Wiss. Z. Tech. Hochsch. Chem. Carl. Schorlemmer Leuna-Merseburg, 1 4 (1972) 327. 64 D.V. Banthorpe, Elimination Reactions, Elsevier, Amsterdam, 1963. 65 A.F. Cockerill, in C.H. Bamford and C.F.H. Tipper (Eds.), Comprehensive Chemical Kinetics, Vol. 9, Elsevier, Amsterdam, 1963, p. 163. 66 I. Mochida, Y. Anju, A. Kato and T. Seiyama, J. Org. Chem., 39 (1974) 3785. 67 H. Noller, P. AndrCu and M. Hunger, Angew. Chem. Int. Ed. Engl., 10 (1971) 172. 68 K. Thomke, Proc. 6th Int. Congr. Catal.,The Chemical Society. London, 1977. 69 H. Knozinger, in S. Patai (Ed.), The Chemistry of the Hydroxyl Group, Interscience, London, 1971, p. 641. 70 J. Sedl5Eek and M. Kraus, React. Kinet. Catal. Lett., 2 (1975) 57. 7 1 H. Knozinger, H. Buhl and K. Kochloefl, J. Catal., 24 (1972) 57. 72 J. SedlBEek, J . Catal., in press. 73 J. SedlaEek, Collect. Czech. Chem. Commun., 42 (1977) 2027. 74 G.A. Mills and G.S. Hindin, J. Am. Chem. SOC.,72 (1950) 5549. 75 A.G. Oblad, G.S. Hindin and G.A. Mills, J. Am. Chem. SOC., 75 (1953) 4096. 76 P. Sabatier, Catalyse en Chimie Organique, Beranger, Paris, 2nd eqn., 1920. 77 W. Ipatiew, Ber., 37 (1904) 2986. 78 H. Adkins and P.P. Perkins, J. Am. Chem. SOC.,47 (1925) 1163. 79 W.J. Brey and K.A. Krieger, J. Am. Chem. SOC.,7 1 (1949) 3637. 80 J.C. Balaceanu and J.C. Jungers, Bull. SOC.Chim. Belg., 60 (1951) 476. 8 1 K.V. Topchieva, K. Yun-Pin and I.V. Smirnova, Adv. Catal., 9 (1957) 799. 82 J.B. Butt, H. Bliss and C.A. Walker, AICHE J., 8 (1962) 42. 83 I.M. Kolesnikov and G.M. Panchenkov, Zh. Fiz. Khim., 38 (1964) 96. 84 H. Knozinger and R. Kohne, J. Catal., 5 (1966) 264. 85 M.E. Winfield, in P.H. Emmett (Ed.), Catalysis, Vol. VII, Reinhold, New York, 1960, p. 93. 86 H. Knozinger, Angew. Chem., 8 (1968) 778. 87 B. Notari, Chim. Ind. (Milan), 5 1 (1969) 1200. 88 A. Ch. Bork and A.A. Tolstopyatova, Acta Physicochim. URSS, 8 (1938) 577. 89 H. Dohse and W. Kalberer, Z. Phys. Chem., Abt. B 5 (1929) 131. 90 H. Dohse, Z. Phys. Chem., Abt. B 6 (1930) 343. 91 A. Ch. Bork and A.A. Tolstopyatova, Acta Physicochim. URSS, 8 (1938) 591. 92 A. Ch. Bork and A.A. Tolstopyatova, Acta Physicochim. URSS, 8 (1938) 603. 93 A. Ch. Bork and S.V. Kirillova, Zh. Fiz. Khim., 25 (1951) 224. 94 K. Kochloefl, M. Kraus, Chou-Chin Shen, L. Berinek and V. BaZant, Collect. Czech. Chem. Commun., 27 (1962) 1199. 95 M. Kraus, Collect. Czech. Chem. Commun., 34 (1969) 699. 96 C.L. Kibby, S.S. Lande and W.K. Hall, J. Am. Chem. SOC.,94 (1972) 214.
97 T.V. Antipina and A.V. Frost, Vestnik Mosk. Gosud. Univ., ( 3 ) (1950) 81. 9 8 T.V. Antipina, V.I. Savushkina and A.V. Frost, Vestnik Mosk. Gosud. Univ., (3-4) (1946) 119. 99 T.V. Antipina and A;V. Frost, Vestnik Mosk. Gosud. Univ., (8) (1951) 69. 1 0 0 K.V. Topchieva and Yun-Pin, Vestnik Mosk. Gosud. Univ., ( 2 ) (1953) 89. 101 N.C. R o y and N.K. Bose, Trans. Indian Inst. Chem. Eng., 11 (1958-1959) 1. 1 0 2 J.R. Laible, Thesis, University of Wisconsin, 1959. 1 0 3 W.C. Johnson, Thesis, University of Wisconsin, 1960. 104 D.N. Miller and R.S. Kirk, AICHE J., 8 (1962) 183. 1 0 5 M. Kraus and K. Kochloefl, Collect. Czech. Chem. Commun., 32 (1967) 2320. 106 L. Zanderighi, A. Greco and S. Carrd, J. Catal., 33 (1974) 327. 107 J.B. Moffat and A S . Riggs, J. Catal., 42 (1976) 388. 1 0 8 N.C. Roy and N.K. Bose, Indian Chem. Eng. Trans., 6 (1964) 36. 109 S. Carrd, U. Santangelo and A. Fusi, Chim. Ind. (Milan), 48 (1966) 229. 110 H. Bremer and U. Werner, Chem. Tech. (Leipzig), 21 (1969) 353. 111 J.C. Gottifredi, A.A. Yeramian and R.R. Cunningham, J . Catal., 12 (1968) 245. 112 K. JeMbek, V. BaZant, L. Berdnek and K. Setinek, Proc. 5th Int. Congr. Catal., Miami Beach, 1972, North-Holland, Amsterdam, 1973, p. 1193. 113 H. Knozinger and H. Buhl, Ber. Bunsenges. Phys. Chem., 7 1 (1967) 73. 114 L. Kabel and L.N. Johanson, AICHE J., 8 (1962) 621. 1 1 5 D. Ka116 and H. Knozinger, Chem. Ing. Tech., 39 (1967) 676. 116 H. Knozinger and E. Ress, Z. Phys. Chem. (Frankfurt am Main), 54 (1967) 136. 117 H. Knozinger, H. Buhl and E. Ress. J. Catal., 1 2 (1968) 121. 118 H.J. Solomon, H. Bliss and J.B. Butt, Ind. Eng. Chem. Fundam., 6 (1967) 325. 119 E.R. Haering and A. Syverson, J. Catal., 3 2 (1974) 396. 120 K. Thomke, J. Catal., 44 (1976) 339. 1 2 1 P. Bautista, M. Hunger and H. Noller, Angew. Chem. Int. Ed. Engl., 7 (1968) 140. 122 H. Knozinger and A. Scheglila, Z. Phys. Chem. (Frankfurt am Main), 63 (1969) 197. 1 2 3 K. Kochloefl and H. Knozinger, Proc. 5th Int. Congr. Catal., Miami Beach, 1972, North-Holland, Amsterdam, 1973, p. 1171. 124 J.E. Stauffer and W.L. Kranich, Ind. Eng. Chem. Fundam., 1 (1962) 107. 1 2 5 M. Kraus, Adv. Catal., 17 (1967) 75. 126 I. Mochida, Y. Anju, A. Kato and T. Seiyama, Bull. Chem. SOC.Jpn., 44 (1971) 2326. 127 D. Dautzenberg and H. Knozinger, J. Catal., 33 (1974) 142. 128 H. Knozinger and H. Buhl, Z. Naturforsch. Teil B, 24 (1969) 290. 129 E.E. Eliel, Stereochemistry of Carbon Compounds, McGraw-Hill, New York, 1962. 130 A.J. Lundeen and W.R. van Hoozer, J. Org. Chem., 32 (1967) 3386. 131 B.C. Gates and L.N. Johanson, AICHE J., 17 (1971) 981. 132 K. Thomke and H. Noller, Proc. 5th Int. Congr. Catal., Miami Beach, 1972, NorthHolland, Amsterdam, 1973, p. 1183. 1 3 3 G.K. Boreskov, Yu.M. Shchekochikhin, A.D. Makarov and V.N. Filimonov, Dokl. Akad. Nauk SSSR, 156 (1964) 901. 134 V.I. Yakerson, L.I. Lafer and A.M. Rubinshtein, Dokl. Akad. Nauk SSSR, 1 7 4 (1967) 111. 135 R. ReYicha and K. Kochloefl, Collect. Czech. Chem. Commun., 34 (1969) 206. 136 H. Knozinger and H. Stolz, Ber. Bunsenges. Phys. Chem., 74 (1970) 1056. 137 A.V. Deo, T.T. Chuang and I.G. Dalla Lana, J. Phys. Chem., 7 5 (1971) 234. 1 3 8 R. Thorton and B.C. Gates, Proc. 5th Int. Congr. Catal., Miami Beach, 1972, North-Holland, Amsterdam, 1973, p. 357. 139 Le Nhu Thanh, K. Setinek and L. Berinek, Collect. Czech. Chem. Commun., 37 (1972) 3878.
140 141 142 143 144 145 146 147 148 149 150 151 152 153 154 155 156 157 158 159 160 161 162 163 164 165 166 167 168 169 170 171 172 173 174 175 176 177 178 179 180 181 182 183 184
J . Sedlirek and M. Kraus, Collect. Czech. Chem. Commun., 41 (1976) 248. H. Pines and C.N. Pillai, J . Am. Chem. Soc., 83 (1961) 3270. E.I. Heiba and P.S. Landis, J. Catal., 3 (1964) 471. J . R . Jain and C.N. Pillai, J . Catal., 9 (1967) 322. J.E. Dambrowski, J.B. Butt and H. Bliss, J. Catal., 18 (1970) 297. H. Pines and C.N. Pillai, J . Am. Chem. Soc., 82 (1960) 2401. L. Beranek, M. Kraus, K. Kochloefl and V. BaZant, Collect. Czech. Chem. Commun., 25 (1960) 2513. W.S. Brey and D.S. Cobbledick, Ind. Eng. Chem., 5 1 (1959) 1031. J . Pagek, J. Tyrpekl and M. Machovi, Collect, Czech. Chem. Commun., 31 (1966) 4108. J . Koubek, Thesis, Institute of Chemical Technology, Prague, 1975. S. Ogasawara, H., Hamaya and Y. Yamamoto, Kogyo Kagaku Zasshi, 74 (1971) 385. J. P a h k , Collect. Czech. Chem. Commun., 28 (1963) 1007. J.P. Catry and J.C. Jungers, Bull. Soc. Chim. Fr., (1964) 2317. M. Kikry Ebeid and J. Pagek, Collect. Czech. Chem. Commun., 35 (1970) 2166. S. Ogasawara, K. Hamaya and Y. Kitajima, J. Catal., 25 (1972) 105. P. Hogan and J . P a k k , Collect. Czech. Chem. Commun., 38 (1973) 1513. P. Hogan and J. Pagek, Collect. Czech. Chem. Commun., 39 (1974) 2696. J. Medena, J.J.G.M. van Bokhoven and A.E.T. Kniper, J. Catal., 25 (1972) 238. H.A. Benesi, 3. Catal., 28 (1973) 176. J. Koubek, J. Volf and J. P d e k , J. Catal., 38 (1975) 385. N.S. Kozlov, V.A. Polikarpov, V.A. Tarasevich and L.I. Moiseyenok, Dokl. Akad. Nauk BSSR 1 9 (1975) 150. H. Dunken, P. Fink and E. Pilz, Chem. Tech. (Leipzig), 18 (1966) 490. K. Hirota, K. Fueki and T . Sakai, Bull. Chem. Soc. Jpn., 35 (1962) 1545. M. Tanaka and S. Ogasawara, J. Catal., 25 (1972) 111. Y.P. Hsieh, J . Catal., 2 (1963) 211. J . SedlhEek and J. Koubek, to be published. E.E. Fleck, J. Org. Chem., 1 2 (1947) 708. J.B. Senderens, Bull. Soc. Chim. Fr., 3 (1908) 827. H. Biltz and E. Kiippers, Ber., 37 (1904) 2413. S. Okazaki and M. Komata, Nippon Kagaku Kaishi, (1972) 1615. S. Okazaki and M. Komata, Nippon Kagaku Kaishi, (1973) 459. P. AndrBu, E. Schmitz and H. Noller, Z. Phys. Chem. (Frankfurt am, Main) 42 (1964) 270. I. Mochida, Y. Anju, H. Yamamoto, A. Kato and T . Seiyama, Bull. Chem. SOC. Jpn., 44 (1971) 3305. P. AndrBu and J . Madrid, Acta Cient. Venez. Suppl., 24 (1973) 169. F. Asinger, Chemie und Technologie der Paraffin-Kohlenwasserstoffe, AkademieVerlag, Berlin, 1959, p. 248. G.M. Schwab and H. Noller, Z. Elektrochem., 58 (1954) 762. A. Heinzelmann, R. Letterer and H. Noller, Monatsh. Chem., 102 (1971) 1750. K. Setfnek, M. d e Goldwasser, 0. Rodriguez and P. AndrBu, An. Quim., 67 (1971) 1045. J.J. Prinsloo, P.C. Van Berge and J. Zlotnick, J. Catal., 32 (1974) 466. I. Mochida, A. Kato and T. Seiyama, J. Catal., 18 (1970) 33. W. Kladnig and H. Noller, J. Catal., 29 (1973) 385. I. Hadzistelios, H.J. Sideri-Katsanou and N.A. Katsanos, J. Catal., 27 (1972) 16. A. Lycourghiotis, N.A. Katsanos and I. Hadzistelios, J. Catal., 36 (1975) 385. M. Hunger, Acta Cient. Venez. Suppl., 24 (1973) 155. F.J. Lopez, P. AndrBu, 0. Blassini, M. Pa& and H. Noller, J. Catal., 18 (1970) 233.
185 186 187 188 189 190 191 192 193 194 195 196 197 198 199 200 201 202 203 204 205 206 207 208 209 210 211 212 213 214 215 216 217 218 219 220 221 222 223
M. Kraus, Chem. Ind. (London), (1966) 1263. I. Mochida, J. Take, Y. Saito and Y. Yoneda, J. Org. Chem., 32 (1967) 3894. K. Fukui, H. Kato and T. Yonezawa, Bull. Chem. SOC.Jpn., 34 (1961) 1111. P. AndrBu, M. Calzadilla, J. Baumrucker, S. Perozo and H. Noller, Ber. Bunsenges. Phys. Chem., 7 1 (1967) 69. I. Mochida and Y. Yoneda, J. Catal., 33 (1974) 474. P. Andrbu, M. Rosa-Brusin, C. Sanchez and H. Noller, Z. Naturforsch. Teil B, 22 (1967) 814. H. Noller, W. Low and P. AndrBu, Ber. Bunsenges. Phys. Chem., 6 8 (1964) 663. H. Heydtmann and G. Rinck, Z. Phys. Chem. (Frankfurt a m Main), 36 (1963) 75. M. Misono, J. Catal., 30 (1973) 226. P. AndrBu and J. Madrid, J. Catal., 4 1 (1976) 343. M. Misono and Y. Yoneda, Chem. Lett., (1972) 551. S. Villalba and H. Noller, Z.Phys. Chem. (Frankfurt am Main), 90 (1974) 113. H. Noller, P. AndrBu, E. Schmitz, S. Serain, 0. Neufang and J. Gir6n, Proc. 4th Int. Congr. Catal., Moscow, 1968, Akadbmiai Kiad6, Budapest, 1971, p. 420. M. Misono and Y. Yoneda, J. Org. Chem., 3 3 (1968) 2161. R.B. Valitov, A.A. Malikov, B. Je. Prusenko, A.N. Ikhasnov, G.G. Garifzjanov and G.M. Panchenkov, Khim. Prom. (Moscow), (1974), 823. I. Mochida and Y. Yoneda, J. Org. Chem., 3 3 (1968) 2161. I. TvaroSka, V. Klimo and L. Valko, Tetrahedron, 3 0 (1974) 3275. P.C. Hiberty, J. Am. Chem. SOC.,97 (1975) 5975. V. Haensel, Adv. Catal., 3 (1951) 179. P. Schneider, M. Kraus and V. BaZant, Collect. Czech. Chem. Commun., 26 (1961) 1636; 27 (1962) 9. V.A. Chernov and T.V. Antipina, Kinet. Katal., 4 (1963) 595. H.H. Voge, in P.H. Emmett (Ed.), Catalysis, Vol. VI., Reinhold, New York, 1958, p. 463. A.V. Frost, Vestn. Mosk. Gosud. Univ., (3-4) (1946) 11. A.P. Ballod, I.V. Pacevich, A.S. Feldman and A.V. Frost, Dokl. Akad. Nauk SSSR, 78 (1951) 509. T.E. Corrigan, J.C. Garver, H.F. Raseand R.S. Kirk, Chem. Eng. Prog., 49 (1953) 603. H.F. Rase and R.S. Kirk, Chem. Eng. Prog., 50 (1954) 35. C.D. Prater and R.M. Lago, Adv. Catal., 8 (1956) 294. G.M. Panchenkov and A.S. Kazanskaya, Zh. Fiz. Khim., 32 (1958) 1779. Ch. D. Georgiev and B.A. Kazanskii, Izv. Akad, Nauk SSSR, Otd. Khim. Nauk, (1959) 491,499. A.E. Pinsker, Khim. Prom. (Moscow), (1963) 660. W.F. Pansing, J. Phys. Chem., 69 (1965) 392. P. Strnad and M. Kraus, Collect. Czech. Chem. Commun., 30 (1965) 1136. K.V. Topchieva, B.V. Romanovskii and V.I. Timoshenko, Kinet. Katal., 6 (1965) 471. M. Rosa-Brush, M. d e Goldwasser and M. HBjek, Acta Cient. Venez. suppl., 24 (1973) 120. G.M. Panchenkov, Zh. Fiz. Khim., 22 (1948) 209. D.R. May, K.W. Saunders, E.L. Kropa and J.K. Dixon, Discuss. Faraday SOC., 8 (1950) 290. Sh. Battalova, G.M. Panchenkov and K.V. Topchieva, Zh. Fiz. Khim., 26 (1952) 903. G.M. Panchenkov and V.V. Krasinchev, Dokl. Akad. Nauk SSSR, 94 (1954) 891. I.M. Kolesnikov, V.A. Tulupov and G.M. Panchenkov, Zh. Fiz. Khim., 40 (1966) 182.
224 L. d e Mourges, M. Perrin, F. Sattonay and Y. Trambouze, Bull. Soc. Chim. Fr., (1965) 843. 225 P. Andrbu, C. Villamizar and G. Cavalleto, Acta Cient. Venez. Suppl., 24 (1973) 126. 226 R.W. Maatman, R.M. Lago and C.D. Prater, Adv. Catal., 9 (1957) 531. 227 I. Mochida and Y. Yoneda, J. Catal., 7 (1967) 386. 228 R.M. Robertsand G.M. Good, J. Am. Chem. Soc., 7 3 (1951) 1320. 229 Y. Fukui, H. Takaoka, J. Ishii, K. Hirai and T. Takahashi, Kogyo Kagaku Zasshi, 69 (1962) 82. 230 I. Mochida and Y. Yoneda, Shokubai, 6 (1964) 281. 231 G.M. Schwab and G. Mandr6, Z. Phys. Chem. (Frankfurt am Main), 47 (1965) 22. 232 J.L. Franklin and D.E. Nicholson, J . Phys. Chem., 60 (1956) 59. 233 B.S. Greensfelder and H.H. Voge, Ind. Eng. Chem., 37 (1945) 515. 234 G.M. Good, H.H. Voge and B.S. Greensfelder, Ind. Eng. Chem., 39 (1947) 1032. 235 M.W. Tamele, Discuss. Faraday SOC.,8 (1950) 270. 236 T.H. Milliken, G.A. Mills and A.C. Oblad, Discuss. Faraday SOC.,8 (1950) 279. 237 R.C. Hansford, Adv. Catal., 4 (1952) 1. 238 P. Andrbu, G. Martin and H. Noller, J. Catal., 21 (1971) 255. 239 I. Mochida and Y . Yoneda, J. Catal., 7 (1967) 393. 240 A.A. Beglaryan and B.V. Romanovskii, Kinet. Katal., 17 (1976) 367. 241 E. Berliner, Prog. Phys. Org. Chem., 2 (1964) 263. 242 D.M. Brower and H. Hogeveen, Prog. Phys. Org. Chem., 9 (1972) 179. 243 P. Sabatier and A. Mailhe, Compt. Rend., 150 (1910) 1570. 244 I.N. Tits-Skvortsova, S. Ja. Levina, A.I. Leonova and Je. A. Karasova, Zh. Obshch. Khim., 21 (1951) 242. 245 J.R. Kaiser, H. Beuther, L. Moore and R.C. Odioso, Ind. Eng. Chem. Prod. Res. Dev., l ( 1 9 6 2 ) 296. 246 F.T. Barr and D.B. Keyes, Ind. Eng. Chem., 26 (1934) 1111. 247 A.V. Mashkina, T.S. Suchareva and V.I. Chernov, Neftekhimiya, 6 (1966) 619. 248 M. Sugioka and K. Aomura, Bull. Jpn. Pet. Inst., 1 5 (1973) 136. 249 M. Sugioka and K. Aomura, Nippon Kagaku Kaishi, (1973) 1279. 250 J.C. Elgin, Ind. Eng. Chem., 22 (1930) 1290. 251 N.V. Imadze, A.K. Getmanova and T.A. Danilova, Neftekhimiya, 7 (1967) 105. 252 N.V. Shelemina, T.A. Danilova and Kh. Khaumbaev, Neftekhimiya, 8 (1968) 777. 253 M. Sugioka and K. Aomura, Bull. Jpn. Pet. Inst., 17 (1975) 51. 254 M.E. Winfield, in P.H. Emmett (Ed.), Catalysis, Vol. 7 , Reinhold, New York, 1960, pp. 158-182. 255 C.R. Nelson and M.L. Courter, Chem. Eng. Prog., 50 (1954) 526. 256 M.A. Dalin, Khim. Nauka Prom., l ( 1 9 5 6 ) 259. 257 A.F. Millidge, in S.A. Miller (Ed.), Ethylene and its Industrial Derivatives, Benn, London, 1969, pp. 690-801. 258 Y. Onoue, Y. Mizutani, S. Akiyama, Y. Izumi and H. Ihara, Bull. Jpn. Pet. Inst., 1 5 (1973) 50. 259 Y . Onoue, Kagaku Kogyo, 26 (1975) 355. 260 W. Neier and J. Woellner, Erdoel Kohle Erdgas Petrochem. Brennst.-Chem., 28 (1975) 19. 261 E.R. Gilliland, R.C. Gunnes and V.O. Bowles, Ind. Eng. Chem., 28 (1936) 370. 262 C.W. Smart, H. Burrows, K. Owen and O.R. Quayle, J. Am. Chem. SOC., 6 3 (1941) 3000. 263 H.M. Stanley, J.E. Youell and J.B. Dymock, J. SOC. Chem. Ind. London, 53 (1934) 205. 264 F.M. Majewski and L.F. Marek, Ind. Eng. Chem., 30 (1938) 203.
265 M.P. Appleby, J.V.S. Glass and G.F. Horsley, J. SOC. Chem. Ind. London, 56 (1937) 279. 266 R.H. Bliss and B.F. Dodge, Ind. Eng. Chem., 29 (1937) 19. 267 A.A. Vvedensky and L.F. Feldman, Zh. Obshch. Khim., 1 5 (1945) 37. 268 C.S. Cope and B.F. Dodge, AIChE J . , 5 (1959) 10. 269 Yu. M. Bakshi, A.I. Gelbshtein and M.I. Temkin, Dokl. Akad. Nauk SSSR, 126 (1959) 314. 270 Yu. M. Bakshi, A.I. Gelbshtein and M.I. Temkin, Dokl. Akad. Nauk SSSR, 132 (1960) 157. 271 J. Muller and H.I. Waterman, Genie Chim., 7 8 (1957) 173. 272 G.M. Barrow, J. Chem. Phys., 20 (1952) 1739. 273 F.J. Saunders and B.F. Dodge, Ind. Eng. Chem., 26 (1934) 208. 274 R.S. Aries, Can. Chem. Process Ind., 32 (1974) 1004. 275 M.E. Dyatkina, Zh. Fiz. Khim., 2 8 (1954) 377. 276 H. Kuribayashi and M. Kugo, Kogyo Kagaku Zasshi, 69 (1966) 1930. 277 C.V. Mace and C.F. Bonilla, Chem. Eng. Prog., 50 (1954) 385. 278 K. Deguchi, I. Yuhara and S. Ogasawara, Nippon Kagaku Kaishi, (1974) 319. 279 R.C. Zabor, R.C. Odioso, B.K. Schmid and J.R. Kaiser, Actes 2dme Congr. Int. Catal., Paris, 1960, Technip, Paris, 1961, Vol. 2, 2601. 280 J.R. Kaiser, H. Beuther, L.D. Moore and R.C. Odioso, Ind. Eng. Chem. Prod. Res. Dev., 1 (1962) 296. 281 C.S. Cope, J. Chem. Eng. Data, 11 (1966) 379. 282 V.P. Gupta and W.J.M. Douglas, AIChE J., 1 3 (1967) 883. 283 H. Matyschok and S. Ropuszynski, Wiad. Chem., 26 (1972) 783. 284 P.K. Lashmet, Ph.D. Thesis, University of Delaware, 1962; Diss. Abstr., 23 (1963) 2462. 285 A.B. Metzner, P.K. Lashmet and J.E. Ehrreich, Actes 26me Congr. Int. Catal., Paris, 1960, Vol. l, T e c h n ip , Paris, 1961, p. 735. 286 M. Giorgini and A. Lucchesi, Chim. Ind. (Milan), 48 (1966) 24. 287 A.I. Gelbshtein, Yu. M. Bakshi and M.I. Temkin, Dokl. Akad. Nauk SSSR, 132 (1960) 384. 288 R. Ono and T. Sugiura, Sekiyu Gakkai Shi, 1 4 (1971) 984; Chem. Abstr. 76 (1972) 986 58. 289 Y. Ogino, J. Catal., 8 (1967) 64. 290 K. Tanabe and M. Nitta, Bull. J p n . Pet. Inst., 1 4 (1972) 47. 291 R.W. Taft, Jr., J. Am. Chem. SOC., 74 (1952) 5372. 292 R.W. Taft, Jr., E.L. Purlee, P. Riesz and C.A. de Fazio, J. Am. Chem. SOC., 77 (1955) 1584. 293 Zh. A. Bril and V.M. Platonov, Khim. Prom. (Moscow), 44 (1968) 112. 294 Zh. A. Bril, V.M. Platonov and M. Ya. Klimenko, Khim. Prom. (Moscow), 45 (1969) 332. 295 B.S. Bouden, A.I. Gelbshtein, G.I. Gagarina and G.G. Goryacheva, Neftepererab. Neftekhim. (Moscow), ( 5 ) (1968) 33. 296 M. Giorgini, P.F. Marconi, G. Monzani, R. Simula and R. Tartarelli, J. Catal., 24 (1972) 521. 297 M. Giorgini, P. Davini, P.F. Marconi, R. Falleni and R. Tartarelli, Ann. Chim. (Rome), 62 (1972) 624. 298 H. Kuribayashi and M. Kugo, Kogyo Kagaku Zasshi, 69 (1966) 1935. 299 C.B. Dale, C.M. Sliepcevich and R.R. White, Ind. Eng. Chem., 4 8 (1956) 913. 300 G.V. Pipko, Khim. Prom. Ukr., (4) (1969) 31. 301 W.B. Innes, in P.H. Emmett (Ed.), Catalysis, Vol. 2, Reinhold, New York, 1955, p. 1. 302 E.N. Tsybina, A.I. Gelbshtein, A.A. Arest-Yakubowich and M.I. Temkin, Zh. Fiz. Khim., 3 2 (1958) 856.
303 304 305 306 307
E.N. Tsybina, A.I. Gelbshtein and M.I. Temkin, Zh. Fiz. Khim., 32 (1958) 995. G. Gut and K. Aufdereggen, Helv. Chim. Acta, 57 (1974) 441. Yu. A. Gorin, I.K. Gorn and S.F. Lokteva, Kinet. Katal., 10 (1969) 322. A.J. Szonyi and W.F. Graydon, Can. J. Chem. Eng., 40 (1962) 183. R.R. Cartmell, J.R. Galloway, R.W. Olson and J.M. Smith, Ind. Eng. Chem., 40 (1948) 389.
308 309 310 311
L.M. Reed, L.A. Wenzel and J.B. O’Hara, Ind. Eng. Chem., 48 (1956) 205. G.E. Hamilton and A.B. Metzner, Ind. Eng. Chem., 49 (1957) 838. A.B. Metzner and J.E. Ehrreich, AIChE J., 5 (1959) 496. L.E. Wing, Ph.D. Thesis, Ohio State University, 1968; Diss. Abstr. B, 29 (1969)
312 313 314 315 316 317
D.F. Othmer and M.S. Thakar, Ind. Eng. Chem., 50 (1958) 1235. N.S. Rabovskaya, Zh. Fiz. Khim., 33 (1959) 2467. N.G. Polyanskii and N.L. Potudina, Neftekhimiya, 3 (1963) 706. F.A. Long and J.G. Pritchard, J. Am. Chem. Soc., 78 (1956) 2663. J.G. Pritchard and F.A. Long, J. Am. Chem. Soc., 78 (1956) 2667. R. Landau and R.E. Lidov, in S.A. Miller (Ed.), Ethylene and its Industrial Derivatives, Benn,London, 1969, p. 568. R.E. Parker and N.S. Isaacs, Chem. Rev., 59 (1959) 737. J.P. Wibaut, J.J. Diekmann and A.J. Rutgers, Rec. Trav. Chim., 47 (1928) 477. J.P. Wibaut and J. van Dalfsen, Rec. Trav. Chim., 5 1 (1932) 636. P. AndrBu, H. Noller, M. Pakz and G. SanchBz, Z. Phys. Chem. (Leipzig), 246
3301.
318 319 320 321
(1971) 145. 322 P. AndrBu, A.J.Schwab and H. Noller, An. Quim., 65 (1969) 965. 323 P. Andrku, A.J. Schwab and H. Noller, Z. Phys. Chem. (Frankfurt am Main), 73 (1970) 132. 324 P. Andreu, H. Noller and M. Paez, An. Quim., 65 (1969) 921. 325 S. Kondo, M. Shindo, S. Wakano and H. Sakurai, Kogyo Kagaku Zasshi, 57 (1954) 909. 326 D.W. Weston and J.B. Agnew, Indian Chem. Eng., 15 (1973) 37. 327 K.V. Topchieva and S.A. Venyaminov, Kinet. Katal., 4 (1963) 450. 328 A.I. Gelbshtein, G.G. Shcheglova and A.A. Khomenko, Kinet. Katal., 4 (1963) 625. 329 H. Bremer, H. Lieske and K. Hertwig, Wiss. Z. Tech. Hochsch. Chem. Carl Schorlemmer Leuna-Merseburg, 1 4 (1972) 355. 330 K, Hertwig, R. Adler and K.D. Henkel, Chem. Tech. (Leipzig), 26 (1974) 639. 331 T. Watanabe, Kogyo Kagaku Zasshi, 62 (1959) 125. 332 G. Thodos and L.F. Stutzman, Ind. Eng. Chem., 50 (1958) 413. 333 L.E. Swabb and H.E. Hoelscher, Chem. Eng. Prog., 48 (1952) 564. 334 R.G. Rinker and W.H. Corcoran, Ind. Eng. Chem. Fundam., 6 (1967) 333. 335 R.D. Wesselhoft, J.M. Woods and J.M. Smith, AICHE J., 5 (1959) 361. 336 H. Noller, R. Letterer and P. AndrBu, Z, Phys. Chem. (Frankfurt am Main) 40 (1964) 334. 337 P. Andrku, R. Letterer, W. Low, H. Noller and E. Schmitz, Proc. 3rd Int. Congr. Catal., Amsterdam, 1964, North-Holland, Amsterdam, 1965, p. 857. 338 A.I. Gelbshtein, M.I. Siling, G.A. Sergeeva and G.G. Shcheglova, Kinet. Katal., 4 (1963) 149. 339 F.W. Kirsch, J.D. Potts and D.S. Barmby, J. Catal., 27 (1972) 142. 340 N.G. Polyanskii, Usp. Khim., 3 1 (1962) 1046. 341 P.B. Venuto, K.A. Hamilton, P.S. Landisand J.J. Wise, J. Catal., 5 (1966) 81. 342 W.M. Kutz and B.B. Corson, J. Am. Chem. SOC.,68 (1946) 1477. 343 W.M. Kutz, J.E. Nickels, J.J. McGovern and B.B. Corson, J. Org. Chem., 16 (1951) 699. 344 A.A. O’Kelly, J . Kellett and J . Plucker, Ind. Eng. Chem., 39 (1947) 154.
345 Kh. M. Minachev, Ya. I. Isakov, V.I. Garanin, L.I. Piguzova, V.I. Bogomolov and A.S. Vitukhina, Neftekhimiya, 5 (1965) 676. 346 Kh. M. Minachev, Ya. I. Isakov, V.I. Garanin, L.I. Piguzova and A.S. Vitukhina, Neftekhimiya, 6 (1966) 47. 347 Kh. M. Minachev and Ya. I. Isakov, Neftekhimiya, 6 (1966) 694. 348 P.B. Venuot, L.A. Hamilton and P.S. Landis, J. Catal., 5 (1966) 484. 349 I.M. Kolesnikov, G.M. Panchenkov and Ye. A. Morozov, Zh. Prikl. Khim., 39 (1966) 415. 350 B.V. Romanovskii, D.T. Allakhverdieva and K.V. Topchieva, Kinet. Katal., 9 (1968) 1384. 351 K. Matsura, T. Watanabe, A. Suzuki and M. Itoh, J. Catal., 26 (1962) 127. 352 Yu. I. Kozorezov, J.G. Gerasimenko and T.S. Novozhilova, Kinet. Katal., 16 (1976) 238. 353 I.G. Gakh, F.G. Gabdpakhmanov and Ye. P. Babin, Zh. Obshch. Khim., 34 (1964) 2807. 354 J.M. Kapura and B.C. Gates, Ind. Eng. Chem. Prod. Res. Dev., 1 2 (1973) 62. 355 Yu. I. Kozorezov and O.A. Lazareva, Zh. Fiz. Khim., 48 (1974) 2755. 356 R.B. Wesley and B.C. Gates, J. Catal., 34 (1974) 288. 357 Y. Fukui, H., Takaoka, J. Ishii, K. Hirai and T. Takahashi, Bull. Jpn. Pet. Inst., 9 (1967) 19. 358 P.B. Venuto, J. Org. Chem., 32 (1967) 1272. 359 K.A. Becker, H.G. Karge and W.D. Streubel, J. Catal., 28 (1973) 403. 360 L.M. Stock and H.C. Brown, Adv. Phys. Org. Chem., 1 (1963) 35. 361 Y. Morita, T. Kimura, F. Kato and M. Tamagawa, Bull. J p n . Pet. Inst., 14 (1972) 192. 362 A. Reychler, Bull. Soc. Chim. Belg., 21 (1907) 71. 363 T.W. Evans and K.R. Edlund, Ind. Eng. Chem., 28 (1936) 1186. 364 K. Setfnek, Z. Prokop, M. Krausand Z. Zitng, Chem. Prumysl, 27 (1977) 451. 365 F. Ancilloti, M. Massi Mauri and E. Pescarollo, J. Catal., 46 (1977) 49. 366 S.A. Bernhard and L.P. Hammett, J. Am. Chem. Soc., 7 5 (1953) 1798. 367 J.P. Guthrie, Can. J. Chem., 52 (1974) 2037. 368 T.V. Subramanian, B. Jagannadhaswamy and G.S. Laddha, Indian J. Technol., 11 (1973) 385. 369 S. Malinowski, W. Kiewlicz and E. Soltys, Bull. SOC.Chim. Fr., (1963), 439. 370 Z. Csuros, Gy. Defik, M. Haraszthy-Papp and L. Prihradny, Acta Chim. Acad. Sci. Hung., 55 (1968) 411. 371 A.T. Nielsen and W.J. Houlihan, in A.C. Cope (Ed.), Organic Reactions, Vol. 16, T h e Aldol Condensation, Wiley, New York, 1968. 372 S. Malinowski, S. Basinski, S. Szczepanska and W. Kiewlicz, Proc. 3rd Int. Congr. Catal., Amsterdam, 1964, North-Holland, Amsterdam, 1965, p. 441. 373 S. Malinowski and S. Basinski, J. Catal., 2 (1963) 203. 374 S. Malinowski and S. Basinski, Bull. Acad. Pol. Sci., Ser. Sci. Chim., 11 (1963) 55. 375 S. Basinski and S. Malinowski, Rocz. Chim., 38 (1964) 635. 376 W. Kiewlicz and S. Malinowski, Bull. Acad. Pol. Sci., Ser. Sci. Chim., 17 (1969) 259. 377 S. Basinski and S. Malinowski, Rocz. Chem., 38 (1964) 843. 378 S. Malinowski, S. Basinski and S. Szczepanska, Rocz. Chem., 38 (1964) 1361. 379 W. Kiewlicz and S. Malinowski, Rocz. Chem., 44 (1970) 1895. 380 Z. Czarny, Chem. Stosow. Ser. A, 7 (1963) 609. 381 A. Narebska, Rocz. Chem., 37 (1963) 663. 382 M.J. Astle and J.A. Zaslowsky, Ind. Eng. Chem., 44 (1952) 2867. 383 K. Ueno and Y. Yamaguchi, Kogyo Kagaku Zasshi, 55 (1952) 234.
384 R.U. Usmanov and Kh. R. Rustamov, Uzb. Khim. Zh., 8 (1) (1964) 82; 9 ( 2 ) (1965) 57. 385 R.A. Reinicker and B.C. Gates, AIChE J., 20 (1974) 933. 386 Z.N. Verkhovskaya, M. Ya. Klimenko, L.B. Vystavkina, E.M. Zalesskaya and L.N. Makarova, Khim. Prom. (Moscow), 44 (1968) 490. 387 A. Basinski and A. Narebska, Rocz. Chem., 35 (1961) 1673, 1685. 388 T. Kawaguchi and S. Hasegawa, Tokyo Gakugei Daigaku Kiyo, Dai-4-Bu, 22 (1970) 47; Chem. Abstr., 74 (1971) 6853. 389 .G.B. Rao, S.H. Ibrahim and N.R. Kuloor, Indian J. Technol., 5 (1967) 119. 390 F.M. Scheidt, J. Catal., 3 (1964) 372. 391 F. Wolf and A. Losse, J. Prakt. Chem., 313 (1971) 137, 145. 392 J.F. Vitcha and V.A. Sims, Ind. Eng. Chem., Prod. Res. Dev., 5 (1966) 50. 393 N.A. Rozenberg and Yu. A. Gorin, Zh. Prikl. Khim. (Leningrad), 43 (1970) 2301. 394 H. Schnell and H. Krimm, Angew. Chem., 75 (1963) 662. 395 P.B. Venuto and P.S. Landis, J. Catal., 6 (1966) 237. 396 Kh. Ya. Maksudova, Kh. R. Rustamov and R.U. Usmanov, Uzb. Khim. Zh., 13 (1969) 27. 397 C.K. Ingold, Structure and Mechanism in Organic Chemistry, Cornell Univ. Press, Ithaca, N.Y., 1953, pp. 767-782. 398 A.J. Kirby, in C.H. Bamford and C.F.H. Tipper (Eds.), Comprehensive Chemical Kinetics, Vol. 10, Elsevier, Amsterdam, 1972, pp. 57-146. 399 P.H. Groggins (Ed.), Unit Processes in Organic Synthesis, McGraw-Hill, New York and London, 1947, pp. 621-623. 400 R.E. Kirk and D.F. Othmer (Eds.), Encyclopedia of Chemical Technology, Vol. 5, Interscience, New York, 1950, p. 794. 401 V. Kumar and M. Satyanarayana, Ind. Eng. Chem., Process Des. Dev., 10 (1971) 289. 402 S. Renganathan and N. Krishnaswamy, J. Appl. Polym. Sci., 18 (1974) 2527. 403 V.O. Reikhsfeld, N.A. Filippov and G.P. Yavshits, Zh. Prikl. Khim. (Leningrad), 43 (1970) 101. 404 V.O. Reikhsfeld, N.A. Filippov and G.P. Yavshits, Zh. Prikl. Khim. (Leningrad), 46 (1973) 2223. 405 A. Ballesteros Olmo, J.F. Garcia de la Banda and E. Hermana Tezanos, An. Real Soc. Espan. Ffs. Quim., Ser. B, 63 (1967) 179. 406 I. Mochida, Y. Anju, A. Kato and T. Seiyama, J . Catal., 2 1 (1971) 263. 407 R.A. Buckley and R.J. Alpeter, Chem. Eng. Prog., 47 (1951) 243. 408 C. Venkateswarlu, M. Satyanarayana and M.N. Rao, Ind. Eng. Chem., 50 (1958 973. 409 H.T. Hoering, Don Hanson and O.L. Kowalke, Ind. Eng. Chem., 35 (1943) 575. 410 R. Maatman, P. Mahaffy and R. Mellema, J. Catal., 35 (1974) 44. 411 A.L. Fricke and R.J. Alpeter, J. Catal., 25 (1972) 33. 412 N.R. Amundson and M. Stusiak, PB Rep. 136 776, US Govt. Res. Rep., 3 1 (1959) 135: Chem. Abstr., 54 (1960) 11671. 413 R.A. Fitz, Ph.D. Thesis, Ohio State University, 1956. 414 Y. Venkatesham, Ph. D. Thesis, Ohio State University, 1953; Diss. Abstr., 20 (1959) 240. 415 M.K. Husain, Chem. Age India, 1 9 (1968) 183; Chem. Abstr., 69 (1968) 76127. 416 A.M. Chashchin., Sb. Tr. Tsentr. Nauchn.-Issled. Proektn. Inst. Lesokhim. Prom., (1963) 1 2 ; C h e m . Abstr., 62 (1965) 7629. 417 A. Afridzhanov and Ya. Yu. Gafurdzhanov, Tr. Tashk. Politekh. Inst., 90 (1972) 45; Chem. Abstr., 83 (1975) 78094. 418 R.W. Maatman, Catal. Rev., 8 (1973) 1. 419 M.S. Newman (Ed.), Steric Effects in Organic Chemistry, Wiley, New York, 1956, p. 205.
420 C.E. Heath, Ph. D. Thesis, University of Wisconsin, 1956. 421 E.S. Gould, Mechanism and Structure in Organic Cehmistry, Holt, New York, 1959, p. 314. 422 L.F. Fieser and M. Fieser, Introduction t o Organic Chemistry, Heath, 1957, p. 138. 423 V.N. Filimonov, Opt. Spektrosk., 1 (1956) 490. 424 M. Baverez and J. Bastick, Bull. Soc. Chim. Fr., (1967) 4070. 425 J.H. Anderson and K.A. Wickersheim, Surface Sci., 2 (1964) 252. 426 S. Kurosaki, J. Phys. Chem., 5 8 (1954) 320. 427 F. Helfferich, J. Am. Chem. Soc., 76 (1954) 5567. 428 T.I. Andrianova, Kinet. Katal., 5 (1964) 724. 429 S.Y. Kerke and M.G. Rao, Trans. Indian Inst. Chem. Eng., (Oct.) (1969) 115. 430 B. Satyanarayana and Y.B.G. Varma, Indian J. Technol., 8 (1970) 312. 431 N.G. Polyanskii and G.V. Gorbunov, Zh. Prikl. Khim. (Leningrad), 45 (1972) 2282. 432 Lu Meng-Kun, Lu Shau-Zou and Fu Mou-Ying, Hua Hsueh, 1 (1968) 7; Chem. Abstr., 7 1 (1969) 100 984. 433 C.L. Levesque and A.M. Craig, Ind. Eng. Chem., 40 (1948) 96. 434 M. Tomescu, N. Demian and V. Pgrgusanu, Rev. Chim. (Bucharest), 22 (1971) 519. 435 K. Setfnek and 0. Rodriguez, Vses. Konf. Mekhan. Katal. Reak., Moscow, 1974. 436 0. Rodriguez and K. Setfnek, J. Catal., 39 (1975) 449. 437 K. Setfnek, Collect. Czech. Chem. Commun., 42 (1977) 979. 438 A.A. Yeramian, J.C. Gottifredi and R.E. Cunningham, J. Catal., 1 2 (1968) 257. 439 K. Setfnek and L. Beranek, J . Catal., 17 (1970) 306. 440 L. Zanderighi, K. Setinek and L. Beranek, Collect. Czech. Chem. Commun., 35 (1970) 2367. 441 L. Beranek, K. Setfnek and M. Kraus, Collect. Czech. Chem. Commun., 37 (1972) 2265. 442 K. Setfnek and L. Beranek, Collect. Czech. Chem. Commun., 38 (1973) 3790. 443 N. Takamiya, T. Fukuda, J. Horii and S. Murai, Nippon Kagaku Kaishi, (1975) 2069. 444 T.I. Andrianova and B.P. Bruns, Kinet. Katal., 1 (1960) 440. 445 T.I. Andrianova, Kinet. Katal., 5 (1964) 927. 446 A.J. Herrman, Ph.D. Thesis, University of Washington, 1955. 447 N.M. Lebedeva, M.A. Shaburov, A.M. Chashchin, K.M. Saladze and L.V. Parshukova, Gidroliz. Lesokhim. Prom., 24 (5) (197 1) 10. 448 K.C. Oosterhout, Ph.D. Thesis, Pennslyvania State University, 1967 ; Diss. Abstr. B, 28 (1967) 2406. 449 M.B. Bochner, S.M. Gerber, W.R. Vieth and A.J. Rodger, Ind. Eng. Chem. Fundam., 4 (1965) 314. 450 R. Linarte Lazcano and F.B.L. Rivera, Rev. Inst. Mex. Pet., 6 ( 3 ) (1974) 31. 451 S. Bhatia, K. Rajamani, P. Rajkhowa and M.G. Rao, Ion Exch. Membranes, 1 (1973) 127. 452 R. Tartarelli, G. Stoppato, F . Morelli, A. Lucchesi and M. Giorgini, J . Catal., 1 3 (1969) 108. 453 R. Tartarelli, G. Stoppato, F. Morelli, A. Lucchesi and M. Giorgini, J. Catal., 18 (1970) 219. 454 R. Tartarelli, F. Piladi, G. Stoppato and A. Lucchesi, Chim. Ind. (Milan), 53 (1971) 150. 455 P. Davini and R. Tartarelli, Chim. Ind. (Milan), 53 (1971) 1119. 456 P. Davini and R. Tartarelli, Ann. Chim. (Rome), 62 (1972) 444. 457 P. Davini and R. Tartarelli, Chim. Ind. (Milan), 54 (1972) 133.
458 459 460 461 462 463
464 465 466 467 468 469 470 471 472 473 474 475 476 477 478 479 480 481 482 483
484 485
486 487
488 489 490 491 492 493 494 495
S.S.Stavitskaya and D. Strazhesko, Ukr. Khim. Zh., 41 (1975)1314. Lee Min-Dar, Lee Wei-Kuo and Tsay Quei-Tsu, Hua Hsueh, (1965)52. Ya. Yu. Gafurdzhanov and Kh. R. Rustamov, Uzb. Khim. Zh., 11 (1967)37,58. O.N. Sharma, G.D. Nageshwar and P.S. Mene, Indian J. Technol., 11 (1973)360. J. Fodor and Z. Hajos, Acta Chim. Acad. Sci. Hung., 10 (1956)141. Huang Ting-Chia, Wan Jassen and Tsai Lung-Hai, J. Chin. Chem. SOC.(Taipei), 14 (2)(1967)72;Chem. Abstr., 68 (1968)104295. N.A. Filippov, G.P. Yavshits and V.O. Reikhsfeld, Zh. Prikl. Khim. (Leningrad), 43 (1970)467. V.O.Reikhsfeld, N.A. Filippov and G.P. Yavshits, Zh. Prikl. Khim. (Leningrad), 46 (1973)2453. 0. Rodriguez, Ph.D. Thesis, Czechoslovak Academy of Sciences, Prague, 1974. J. Hine, Physical Organic Chemistry, McGraw-Hill, New York, 1962,2nd edn., pp. 275-280,285-287. B.C. Gates and L.N. Johanson, J. Catal., 14 (1969)69. B.C. Gates, J.S. Wisnouskas and H.W. Heath, Jr., J . Catal., 24 (1972)320. L. Beranek, Catal. Rev. Sci. Eng., 16 (1977)1. E.Knozinger, Inaugural Dissertation, University of Munich, 1966. A. Nakamura, Kogyo Kagaku Zasshi, 72 (1969)626. Kim Yong Ho, Kim Zin U and Zo In Yeng, Chosun Kwahakwon Tongbo, (2) (1964)29;Chem. Abstr., 62 (1965)11649. S.P. Walvekar and A.B. Halgeri, J. Indian Chem. SOC.,50 (1973)246. L. Delvaux and H. Laudelout, J. Chim. Phys., 61 (1964)1153. C.W. Davies and G.G. Thomas, J. Chem. SOC.,(1952)1607. Kh. R. Rustamov, Uzb. Khim. Zh., 2 (1961)32. L.M. Goldenshtein and G.N. Freidlin, Zh. Prikl. Khim. (Leningrad), 38 (1965) 2538. V.C. Haskell and L.P. Hammett, J. Am. Chem. SOC.,71 (1949)1284. S.Affrossman and J.P. Murray, J. Chem. SOC.B, (1968)579. R. Tartarelli, F. Morelli, M. Giorgini and A. Lucchesi, Chim. Ind. (Milan), 50 (1968)528. L.M. Goldenshtein and G.N. Freidlin, Zh. Prikl. Khim. (Leningrad), 37 (1964) 2540. L.M. Goldenshtein and G.N. Freidlin, Zh. Prikl. Khim. (Leningrad), 38 (1965) 1345. L.M. Goldenshtein and G.N. Freidlin, Zh. F’rikl. Khim. (Leningrad), 38 (1965) 2541. M.B. Ordyan, Ya T. Eidus, L.A. Sarkisyan and A.E. Akopyan, Arm. Khim. Zh., 19 (1966)632. M.B. Ordyan, Ya. T. Eidus, R. Kh. Bostandzhyan and A.E. Akopyan, Arm. Khim. Zh., 20 (1967)873, M.B. Ordyan, Ya. T. Eidus, R. Kh. Bostandzhyan and A.E. Akopyan, Arm. Khim. Zh., 23 (1970)511. S.A. Bernhard and L.P. Hammett, J. Am. Chem. SOC.,75 (1953)5834. R. Tartarelli, G . Nencetti, M. Baccaredda and S. Cagianelli. Ann. Chim. (Rome), 56 (1966) 1108. H. Teshima, T. Sumi and N. Morita, Nippon Kagaku Kaishi, (1972)711. F.Wolf, Papers presented a t t h e Conference on Ion Exchange Processes in Industry, 1969,SOC.Chem. Ind., London, 1970,pp.285,291. S.A. Bernhard, E. Garfield and L.P. Hammett, J . Am. Chem. SOC.,76 (1954)991. H. Teshima, T. Sumi and N. Morita, Nippon Kagaku Kaishi, (1972)2045. P. Riesz and L.P. Hammett, J. Am. Chem. SOC.,76 (1954)992. R. Tartarelli, A. Lucchesi and B. Stoppato, J. Catal., 19 (1970)310.
496 497 498 499 500 501 502 503 504 505 506 507 508 509 510 511 512 513
R.W. Sauer and K.A. Krieger, J. Am. Chem. SOC.,7 4 (1952) 3116. S. Namba and T. Shiba, Kogyo Kagaku Zasshi, 7 1 (1968) 93. Z.M. George, J. Catal., 3 5 (1974) 218. M. ZdraZil, Chem. Prum., 24 (1974) 549. R.J. Rennard, Jr. and W.L. Kehl, Am. Chem. Soc. Div. Pet. Chem. Prepr., 14 ( 4 ) (1972) E47;Chem. Abstr., 8 0 (1974) 145041. V.P. Milishkevich, A.A. Ignat’ev and A.V. Kargin, Zh. Obshch. Khim., 41 (1971) 2498. A. Yu. Prokopchik and J. Valsiuniene, Liet. TSR Mosklu Akad. Darb. Ser. B, (4) (1974) 11;Chem. Abstr., 8 3 (1975) 85426. A. Yu. Prokopchik, J. Valsiuniene and A. Norgailaite, Liet. TSR Mosklu Akad. Darb. Ser. B, (5) (1974) 17;Chem. Abstr., 8 4 (1976) 185407. M. Salkauskas, H. Zielys, S. Raudys and J. Valsiuniene, Liet. TSR Mosklu Akad. Darb. Ser. B, ( 5 ) (1974) 3 ; Chem. Abstr., 8 4 (1976) 185444. G. Bodamer and R. Kunin, Ind. Eng. Chem., 4 3 (1951) 1082. E.R. Gilliland, H.J. Bixler and J.E. O’Connell, Ind. Eng. Chem. Fundam., 1 0 (1971) 185. E.W. Reed and J.S. Dranoff, Ind. Eng. Chem. Fundam., 3 (1964) 304. F. JirBEek, J. Horlk and Z. PotuEkovl, Chem. Prum., 24 (1974) 385. N. Lifshutz and J.S. Dranoff, Ind. Eng. Chem. Process Des. Dev., 7 (1968) 266. J.J.R. De Aimeida and D. Glasser, S. Afr. J. Sci., 6 9 (1973) 109. T.J. McGovern and J.S. Dranoff, AIChE J., 16 (1970) 536. Y. Murakami and 0. Mori, Kogyo Kagaku Zasshi, 69 (1966) 588. F. JirlEek and J. HorBk, Chem. Prum., 22 (1972) 337.
399
Index acrylonitrile, from ammoxidation of C3H6,135,136,155,164-174 acetaldehyde, aldol condensation of, 339, activated complex, see transition state activation energy, and ammoxidation, 341-344,346 -, from CzH2 + HzO, 327-329 166,168,172,223 -, in ammoxidation of C3H6, 167 -, and catalytic cracking, 315 -, in heterogeneous oxidation of CzH4, -, and heterogeneous eliminations, 281, 126,128,133,134,140,157 287,290 -, in heterogeneous oxidation of C4H8, -, and heterogeneous oxidation, of aro187 matics, 199-202, 206, 209, 218, 219 acetals, heterogeneous hydrolysis of, 384, _ ,_ , of C2H4,128,130,133,134 _ , _ , of C3&,144,156,161,238 385 - ,_ , o f C4H8, 176,180,184,185,187acetic acid, adsorption of, 248 -, and dehydration of alcohols, 268, 293 189 -, esterification with, 349-357, 359, -, -, of CHJOH, 226 362,363,365,366 -, --, of NH3, 229, 243 -, from heterogeneous oxidation, of -, -, o f S O z , 2 3 1 C2H4, 1 3 3 , 1 3 4 , 1 4 4 , 1 4 7 , 1 6 0 , 1 6 2 -, of catalytic dehydrohalogenation, 302 -, -, O f C4H8,194 -, of esterification, 362, 365 _ ,- , of PhMe, 205 -, of hydrogenation, of CzH2, 60, 108 -, hydrogenation in, 49 -, -, Of CzH4,29,30,108 -, -, Of C3H6, 3 7 , 3 8 , 1 0 1 acetone, adsorption of, 312 -, -, of C4H,5,81,108 -, aldol condensation of, 339, 341-344, -, -, Of C4H8, 3 9 , 4 0 , 4 3 347 -, from heterogeneous oxidation of -, -, of MeCFCH, 69 -, of hydrolysis, 372, 375, 383, 385 C3H6.135, 136,153,157,162-164 -, of isomerisation, of C4H8, 181 -, hydrolysis in, 373, 375, 376, 379 -, -, of C2H40, 128 -, reaction + PhOH, 338 acetonitrile, and aldol condensation, 341, active sites, 2, 1 0 3 -, and acid-base catalysis, 263, 264, 343,344 266, 268, 269, 272, 273, 275, 292-,from ammoxidation of C3H6, 165294, 299, 300, 308, 315-318, 336, 167,169 343,345,353,355,356,370,371 acetophenone, aldol condensation of, -, and heterogeneous oxidation, 146342,344 149,163,168,181,195,247-251 acetylene, adsorption of, 50-54, 6 3 adjacent-site mechanism, see Langmuir-, hydration of, 327-329 -,hydrogenation of, 4, 8, 9, 10, 13, 55Hinshelwood mechanism adsorption coefficients, 5, 13, 14, 16, 41 68,108,110,111 -, hydrohalogenation of, 332, 333 -, and deamination, 297 acrolein, from heterogeneous oxidation, -, and ester hydrolysis, 378 of C3H6, 135-159, 163, 172, 237, -, and esterification, 352, 360 238,242,246,251 -, for aluminosilicate catalyst, 312 -, -, Of C4H8, 187 adsorption isotherms, 5, 6 --,heterogeneous oxidation of, 159, 165, -, and C2H40 + H20.331 167,173 -, and esterification, 358 acrylic acid, from heterogeneous oxida- -, for C2H2 on metals, 51, 6 3 tion of C3H.5, 135-137, 140, 144, -, for CzH4 on metals, 18-20 147,159,160 allene, see propadiene A
400 alloys, as hydrogenation catalysts, 107109 allyl, see pi-ally1 allyl alcohol, esterification with, 349, 357,358, 362 -, hydrogenation of, 105 allyl bromide, and heterogeneous oxidation Of C3H6,147 allyl carbocation, and heterogeneous oxidation of C3H6, 139, 140 alumina, as carrier, for hydration catalysts, 325, 330, 331 -, -, for hydrogenation catalysts, 3, 1820, 30, 32-34, 36, 37, 40-45, 47, 49, 50, 52, 58-66, 70-78, 81-85, 100.102,109-111 -, -, for oxidation catalysts, 133, 134, 154, 157, 163, 185, 221, 223, 234, 242 -, catalyst for alkylation, 335 -, catalyst for cracking, 310, 311, 314, 316,317 -, catalyst for deamination, 296-300 -, catalyst for dehydration, 279, 282286,288--290,293-295 -, catalyst for dehydrohalogenation, 301-305, 308 -, catalyst for dehydrosulphidation, 318 -, catalyst for esterification, 350, 352356 -, catalyst for hydrohalogenation, 332, 333 -, catalyst for hydrolysis, 372, 380 -, nature of, 266-269 aluminium chloride, reaction + silica, 265,268 aluminium fluoroborate, as cracking catalyst, 313, 314 aluminium phosphate, and dehydration, 295 aluminium sulphate, and cracking, 310 -, and dehydrohalogenation, 306, 307 ammonia, adsorption of on catalysts, 248,266,293,298,316,317 -, and heterogeneous oxidation, of aromatics, 221-223 _ , _, of C3H6,149,165--174 -, heterogeneous oxidation of, 123, 125, 165,227-230,236, 243 aniline, and deamination, 296-299 anthracene, heterogeneous oxidation of, 218,219 anthraquinone, from anthracene oxidation, 218, 219
antimony, and amtroxidation, 166, 169173,222,237 -, and heterogeneous oxidation, 240, 243,252 - ,_ , O f C3H6, 154-158, 161, 239, 246, 247 -, -, of C4Hs, 178, 179,186-190 _ ,- , Of C6H6, 203 -, -, o f N H 3 , 2 3 0
_ ,- , of PhMe, 204 -, free energy of decomposition of oxide, 232 arsenic, and heterogeneous oxidation, 1 6 0 , 1 6 1 , 1 6 9 , 241 Auger spectroscopy, 103
B band theory of metals, and catalysis, 107 barium chloride, catalyst for dehydrochlorination, 301 barium hydroxide, and aldol condensation, 342, 343 barium phosphate, and dehydration, 295 barium sulphate, and dehydrohalogenation, 303, 306, 307 -, as catalyst support for hydrogenation, 70,72 benzaldehyde, aldol condensation of, 339,341-344 -, from heterogeneous oxidation of aromatics, 201, 204, 205, 207, 221, 222 benzaldehyde di( 2-butyl)acetal, hydrolysis of, 385 benzene, adsorption of, 312 -, alkylation of, 334-336 -, esterification in, 367 -, from heterogeneous oxidation, of C3H6,135,155,161 -, -, of PhMe, 205, 207-210 -, hydrogenation in, 49 -,hydrogenation of, 16, 104, 105, 109, 110 -, oxidation of, 196-204, 219 benzoic acid, and alumina catalysts, 268 -, esterification with, 349, 357, 362, 363,367 -, from heterogeneous oxidation of PhMe, 201,204,205 benzonitrile, from ammoxidation of PhMe, 221,222
401 C4H8, 175, 179-181, 191-193, 248, benzoquinone, from heterogeneous oxi249 dation of aromatics, 198, 201, 203, -, heterogeneous oxidation of, 246, 248 205 benzyl acetate, hydrolysis of, 375 -,hydrogenation of, 10, 57, 58, 81-92, benzyl alcohol, dehydration of, 282, 285 109,110 benzyl cyanide, and aldol condensation, butanols, esterification with, 349, 350, 344 354,357-359,362, 365,367-369 beryllium oxide, and heterogeneous oxi- -, from heterogeneous oxidation of i-C4Hs, 195 dation, 242 bismuth hydroxide, and heterogeneous -, heterogeneous dehydration of, 283oxidation of C3H6, 1 4 1 286,288-290,292-295 bismuth oxide, and ammoxidation, 164, -, reaction + olefins, 337 n-butenes, alkylation by, 335 166-169,171-173,222 -, and heterogeneous oxidation, 124, -, heterogeneous oxidation of, 124, 146, 174, 179-195, 239, 244, 248, 249, 240-242,248-251 253 -, -, of aromatics, 202--206, 209, 210 -, hydration of, 325, 326 -, -, O f C3H6, 135,140--151,154,157162, 237, 244-247, 250-252 -, hydrochlorination of, 332, 333 -, hydrogenation of, 10, 38-48, 109, -, -, of C4H8, 175-184,189,194,195 -I -, of CsHio, 1 9 5 111 -, -, of NH3, 229, 230 -, isomerisation of, 25, 26, 49 -, and oxygen exchange, 235, 238, 239 -, reaction + ROH, 337 -, free energy of decomposition, 232 butyl acetate, hydrolysis of, 374, 375 bisphenol A, production of, 338, 342, t-butylacrolein, from heterogeneous oxi343,347 dation, 1 9 5 boehmite, as catalyst support, 81, 83, 84, butylamine, deamination of, 296 110 butylbenzene, catalytic cracking of, 316 bond dissociation energy, and heterogene- butynes, hydrogenation of, 70-73, 109 ous acid-base catalysis, 276, 277 butyric acids, esterification with, 349, 350,354,357 boric acid, and deamination, 296 boric oxide, and alkylation, 335 -, and cracking, 310 C -, and dehydrohalogenation, 303, 304 -, and esterification, 350, 355, 356 borohydride anion, hydrolysis of, 381 cadmium, and C2H2 + HzO, 328,329 boron phosphate, and dehydration, 284 -, and heterogeneous oxidation, 183, -, and hydration, 324, 325 185, 242 boron trifluoride, and alumina catalysts, calcium, and aldol condensation, 340, 268,311,335 344,346 bromobutanes, dehydrobromination of, -, and C4Hs + HCl, 333 302,306 -, and C2Hz + H20, 328,329 l-bromo-2-chloroethane, dehydrohalo- -, and dehydration, 295 genation of, 305 -, and dehydrochlorination, 301-303, l-bromo-l,2-diphenylethane,dehydrobro306,307 mination of, 304 -, and esterification, 350 bromoethane, dehydrobromination of, -, and heterogeneous oxidation, 132, 303 150,169 bromopropanes, dehydrobromination of, Cannizarro reaction, and aldol condensa303 tions, 340, 344 Brown selectivity relationship, and heter- carbon, as catalyst carrier, for esterificaogeneous alkylation, 336 tion, 350, 355 buta-l:2-diene, hydrogenation of, 77-81 -, -, for hydration, 328, 329 buta-l:d-diene, adsorption on Pd, 5 3 -,-,for hydrogenation, 3, 48, 49, 73, -, from heterogeneous oxidation of 74,94
402 carbon, as catalyst carrier - continued
-, -, for hydrohalogenation, 332 -, -, for hydrolysis, 381 -, -, for oxidation, 133, 164
-,
catalyst for dehydrohalogenation, 301 carbon-13, and hydrogenation, 21 carbon-14, and heterogeneous oxidation, 139,140,210,211,215 -,and hydrogenation, 13, 18, 20, 51, 63, 64 carbon dioxide, hydrogenation of, 7 carbon monoxide, and hydrogenation of C2H2,64 -, heterogeneous oxidation of, 1 0 8 carbon tetrachloride, and alkylation catalysts, 335 -, reaction + silica, 265 carbonyl sulphide, hydrolysis of, 380 carriers, for hydrogenation catalysts, 3, 4 cerium, and heterogeneous oxidation, of C3H6,150,157,169 - ,_ , of C4Hg,183,184 -, and hydrolysis, 381 charcoal, see carbon chemisorption, and aldol condensation, 347,348 -, and esterification, 356 -, and heterogeneous oxidation, 129, 130,231,246 -,and hydrogenation, 2, 3, 5, 6, 16, 17, 97,103,104 chlorine, and heterogeneous oxidation of C2H4,130 -, reaction + silica, 265 chlorobenzene, hydrolysis of, 381 chlorobutanes, dehydrochlorination of, 301,302,305-307 l-chloro-1,2-diphenylethylene,dehydrochlorination of, 304 chloroethane, dehydrochlorination of, 302,309 chloropropanes, dehydrochlorination of, 302,303 chlorotoluenes, hydrolysis of, 381 chloroxylenes, hydrolysis of, 381 chromium, and heterogeneous oxidation, Of C3H6,157,162 _ , _ , of C4Hg, 1 7 7 , 1 7 8 , 1 9 0 , 1 9 1 - ,_ , of CsHio, 196 -, -, Of C6H6, 201, 202 - ,_ , of MeOH, 225,226 _ ,_ ,and hydration, 325 -, hydrogenation of C2H4 on, 1 0 6 cinnamaldehyde, hydrogenation of, 8 , 9
&--trans isomerisation, and heterogeneous oxidation, 1 8 1 -, and hydrogenation, 24-26, 39--43, 46--49, 95-97 cobalt, and dehydrohalogenation, 305 -, and heterogeneous oxidation, of aromatics, 202, 217, 219, 221 -,-,Of C3H6, 153, 157, 159, 160, 162, 1 6 4 , 2 4 6 , 2 5 2 , 253 _ ,- ,Of C4H8,177,178,185, 190,195 -, -, of MeOH, 227 -, -, of NH3, 228, 229 -, and hydrolysis, 380 -, and oxygen exchange, 234, 235 -, catalyst for dimerisation of C2H4, 18 -, catalyst for hydrogenation of alkenes, 30, 31, 38, 106 _ ,- , of alkynes, 59, 60, 62, 69, 70, 72 - ,_ , of c - C ~ H 1~0,0 , 1 0 1 - ,_ , o f dienes, 74-77, 81, 82, 84-86, 88,91-93 -, catalyst for isomerisation of n-C4Hs, 47 conductance, and heterogeneous oxidation, 144, 243, 244 coordination number, of surface atoms in metals, 104 copper, and conductance of oxides, 243 -, and dehydrohalogenation, 305, 306 -, and heterogeneous oxidation, of aromatics, 202, 217 -, -, Of C3&, 1 3 5 , 1 3 8 , 1 4 1 , 1 5 6 , 1 5 7 - ,_ , of C4Hs,179,190 -, -, of NH3, 228,229 -, and oxygen mobility, 234, 242 -, catalyst for hydrogenation, 2, 107, 108 , of alkenes, 7, 30, 107, 109 _ , _ _ , _ , of alkynes, 59, 60, 62, 69, 72 -, -, O f C-C3H6, 1 0 0 , 1 0 1 _ , _ , Of C6H6, 109 -,-,of dienes, 81, 82, 84, 86, 88, 92, 93,109 -,free energy of decomposition of oxides, 232 cross-linking in resins, and esterification, 361-366 -, and hydrolysis, 373, 375, 377-379, 381-383 crotonaldehyde, and C2Hz + H20, 328 -, from heterogeneous oxidation of C4HS1174,175,191 crystal phases, in oxide catalysts, 142144, 151, 170, 179, 184, 186, 189, 214,266
403 crystal voltage, and oxide catalysts, 244246 cyclohexane, transesterification in, 359, 368,371 cyclohexanols, dehydration of, 283-285, 287,290 cyclohexenes, hydrogenation of, 95-97, 105 -, reaction + aromatics, 334 cyclohexyl acetate, hydrolysis of, 374 cyclohexylamine, deamination of, 296 cyclopentene, and Fe catalyst, 98 cyclopentyl acetate, hydrolysis of, 374 cyclopropanes, hydrogenation of, 7, 94, 100-103,105,110
D delocalisability, of H and dehydrohalogenation, 304 desorption, and catalysis, 5, 6, 29, 32, 33, 234 deuterium, and dehydration, 286, 290 -, and dehydrohalogenation, 303, 306, 308 -, and hydration, 327 -, effect on heterogeneous oxidation of C3H6,138,155,163 -, effect on silica, 265 -, reaction + alkenes, 10, 16, 17, 22-24, 27,28,32-39,43-45,47,48 -,reaction + alkynes, 55, 57, 58, 64, 65, 67, 68, 70, 71, 73, 74 -, reaction + c-C3H6, 100, 102 -, reaction + C6H6, 105 -, reaction + cycloalkenes, 96, 98 -, reaction + dienes, 78-80, 84, 86, 8891 2,3-dibromobutane, dehydrobromination of, 307,308 1,2-dibromoethane, dehydrobromination of, 3 0 3 , 3 0 8 2,3dichlorobutane, dehydrochlorination of, 307, 308 dichloroethanes, dehydrochlorination of, 300,302,304, 305 -, effect on heterogeneous oxidation of C2&,133 1,2-dichloropropane, dehydrochlorination of, 308 dicinnamalactone, and ester hydrolysis, 372
dielectric constant, of solvent and esterification, 367-369 diethylamine, deamination of, 296, 297, 299 diethyl carbonate, hydrolysis of, 380 diethylether, dehydration of, 285 diffusion, and esterification, 351, 357 -, and hydrolysis, 377-379, 381-383, 385 -, and ion-exchange catalysts, 271, 326, 331 2,3-dimethylbicyclo[ 2,2,2]oct-2-ene, hydrogenation of, 9 8 , 9 9 2,5-dimethylhexa-l:5-diene,from heterogeneous oxidation of C4Hs, 194, 195 2,5-dimethylhexa-2:4-diene, hydrogenatior, of, 94 3,4-dimethylhex-3-ene, isomerisation of, 49 dioxan, hydrolysis in, 373, 376, 379 -, transesterification in, 358, 361, 364, 367,368,370, 371 diphenylethanes, cracking of, 313, 314 diphenylmethane, cracking of, 314 dipropylamine, deamination of, 297 distribution coefficient, and esterification, 357, 358, 367, 376, 377 -, and hydrolysis, 383 dodecene, alkylation by, 334 durene, heterogeneous oxidation of, 220
E efficiency, of catalyst for hydrolysis, 373, 374,376 electron diffraction, and heterogeneous oxidation, 129 electronegativity of cations, and dehydrosulphidation, 320 -, and hydration, 324 electronic effect, and heterogeneous oxidation, 248-250 -, and hydrogenation, 3 , 1 0 6 , 1 0 9 , 1 1 2 electron microscopy, and catalysis, 103 electron spin resonance spectroscopy, and heterogeneous oxidation, 130, 151, 153, 157, 163, 169, 197, 200, 214, 241, 242, 244-247 enthalpy, see also heat -, of hydride ion abstraction and cracking, 314, 316 -, of reaction, of aldol condensation of MeCHO, 339 -, -, of CzH2 + H20, 328
404 enthalpy, of reaction - continued
N-ethyl-N-propyl-N-butylamine, deamination of, 298
entropy, and hydrolysis, 372, 374-376 -, and oxygen desorption, 234 -, of activation and esterification, 362, 363,365 -, of adsorption on oxide catalysts, 253 equilibrium constant, for aldol condensations, 339 -, for C2H4O + HzO, 329 -, for esterifications, 349, 350 ethane, and hydrogenation, 1 0 1 ethanethiol, dehydrosulphidation o f , 3 19,
F
_ , _ , o f C2H40 + H,O, 329 - ,_ , of olefins + H20, 322
320
ethanol, as solvent for hydrogenation, 99 -, dehydration of, 283-285, 288, 289, 293 -, esterification with, 349-351, 353-357,359,363, 366 -, reaction + C4H8, 337 ethyl acetate, hydrolysis of, 372, 3743 7 6 , 3 7 8 , 3 8 2 , 383 -, transesterification with, 350, 357, 358, 360, 361, 363-365, 368, 370, 371 ethyl acrylate, transesterification with, 3 5 0 , 3 5 7 , 3 6 2 , 365,368 ethylbenzene, cracking of, 314 -, heterogeneous oxidation of, 220, 221 ethyl butyrate, hydrolysis of, 375 -, transesterification with, 357, 360, 361 ethyl cyanoacetate, and aldol condensation, 344 ethylene, adsorption of o n metals, 1821, 23, 5 1 -, alkylation by, 335 -, and C2Hz adsorption, 50, 63, 64 -, from ammoxidation of CqHh. - _ . 167 -, heterogeneous oxidation of, 124-135, 236 -, hydration o f , 322, 324, 325, 327 -, hydrochlorination o f , 332, 333 -,hydrogenation of, 1, 3, 7-10, 13, 17, 18, 27-37, 44, 106, 107, 110, 111, 113 ethylene glycol, esterification with, 357, 362 ethylene oxide, hydration of, 329-331 -, in heterogeneous oxidation of C2H4, 1 26-1 2 8 , 1 3 0-1 33 ethyl formate, transesterification with, 357,361 ethyl hexanoate, hydrolysis o f , 375, 377
Fermi energy, see free energy of electrons ferric chloride, catalyst for dehydrochlorination, 300 field emission microscope, 17, 18, 51 films, see metal films flowsystems, and cracking, 311 -, and dehydration, 283 -, and dehydrohalogenation, 302 -, and heterogeneous oxidation, 126, 146, 151, 154, 155, 158, 159, 161, 162, 167-169, 172, 173, 176-178, 180, 186-188, 191, 199, 200, 204, 205,214,215,221,238 -, and hydrogenation, 4 fluid bed reactor, 201 formaldehyde, and aldol condensation, 341-344 -, from heterogeneous oxidation, of C3H6,140,144,157, 242 -, -, of MeOH, 224, 249 formic acid, decomposition of on Pd/Au, 107,108 -, esterification with, 349, 350, 357 free energy, of adsorption, 8, 12, 1 3 -, of decomposition of oxides, 232 -, of electrons in oxides, 243, 244 -, of transition state in hydrolysis, 373 furan, from heterogeneous oxidation of n-C4H~,180,181,185,191-193 furfural, aldol condensation of, 342 G
gamma radiation, and heterogeneous oxidation, 132, 1 5 3 gas chromatography, and eliminations, 298 -, and esterification, 353 -, and heterogeneous oxidation, 130, 185 -, and hydrogenation, 39 geometric factors, and heterogeneous oxidation, 197 -, and hydrogenation, 3, 103-106 glyceride oils, hydrogenation of, 9 gold, and heterogeneous oxidation, 134, 225 -, as hydrogenation catalyst, 2, 47, 73, 81,83-85,107-110,113
405 H Hammett acidity function, and alkylation, 336 -, and hydration, 325 Hammett correlation, and cracking, 31 3, 314 -, and dehydration of alcohols, 288, 289, 292 -, and hydrolysis, 381 heat, see also enthalpy -, of adsorption, and dehydration, 294 -, -, of CdH8, 253 _ ,- , o f N H 3 , 298 -, of desorption of 0 2 and catalysis, 234 -, of formation of metal oxides and catalysis, 233 heptane, esterification in, 367, 368 heptanols, dehydration of, 283, 288 1,5-hexadiene, from heterogeneous oxidation O f C3H6, 135,147,160-162 hexafluoroacetone, reaction + PhOH, 347 hexanes, cracking of, 310, 314, 315 hexanols, dehydration of, 288, 289 hex-1-ene, adsorption of on Pt, 22 -, hydrogenation of, 48, 105 hexyl acetate, hydrolysis of, 374 hex-1-yne, hydrogenation of, 7 1 Hofmann rule, 277 A'.' -D-hornogonene, hydrogenation of, 73 hydride ion, and cracking, 314-317 hydrofluoric acid, and alumina catalysts, 268 hydrogen, chemisorption of, 3 -, effect on olefin adsorption, 21, 22 -, spillover and hydrogenation, 110, 111 hydrogen atoms, and adsorbed CzH4,17 --, and adsorbed C4H6, 5 3 -, and hydrogenation, 4 4 , 4 7 -, and olefin isomerisation and exchange, 25 hydrogen bonding, and aldol condensation, 346, 347 -, and deamination, 298 -, and esterification, 368, 370 -, and ion-exchange catalysts, 270, 291 -, and silica catalysts, 266 hydrogen chloride, and deamination, 296 -, and hydrolysis, 372, 373, 375, 381, 383,384 -, hydrochlorination by, 332, 333 hydrogen cyanide, from ammoxidation of C3H6,165,166,169,173
*
hydrogen sulphide, reaction + C3H6, 318 hydroxyapatite, dehydration on, 283, 289 hydroxyl groups, and hydrolysis, 381 -, on alumina, 266, 267,269, 285 -, on silica, 265, 266, 298 hydroxyl ion, and hydrolysis, 380 hyperconjugation, and acid-base catalysis, 277, 344
I indium, and heterogeneous oxidation, 161,195 infrared spectroscopy, and alumina catalysts, 266, 290 -, and eliminations, 293, 298 -, and heterogeneous oxidation, 130132, 148, 157, 160, 226, 239, 241, 251,252 -,and hydrogenation, 17, 21, 23, 53, 67, 113 -, and ion-exchange resins, 270 -, and silica catalysts, 265 ion-exchange resins, 270-272 -, for aldol condensation, 340, 342, 343, 347 -, for dehydration, 284, 285, 291 -, for esterification, 351, 356-371 -, for hydration, 323, 324, 326, 330 -, for hydrolysis, 371-379, 381-385 -, for ROH + alkenes, 337 iridium, adsorption of hydrocarbons on, 1 7 , 1 9 , 51 -, catalysis of hydrogenation, of alkenes, 30, 31. 33-38.40.42. . . . 43.46.47. . . . 107 -,-,of alkynes, 58, 60-62, 64-67, 69, 73-74 -, -, Of C-CsH6, 100, 1 0 1 -, --,Of C6H6,105 -, -, of cyclohexenes, 9 5 -, -, of dienes, 74--77, 81, 83, 84 -.-, catalysis of isomerisation of olefins, 25 iron, and alkylation, 335 -, and ammoxidation, 170-173, 222 -, and bonding of oxygen on, 230, 234 -, and heterogeneous oxidation, of aromatics, 202--204, 217, 218 -,-,Of C3H6, 151, 153-157, 162, 237, 239,245, 246,250,252, 253 -,-,of C4H8, 177, 178, 180, 183-185, 189--191,244 -, -, of C ~ H 1 0 , 1 9 6
406 iron, and heterogeneous oxidation - continued -, -. , of MeOH, 225-227,249 - ,_ , o f NH3, 228-230 -, and hydrolysis, 381 -, and oxygen exchange, 235 -, catalysis of hydrogenation, of alkenes, 30, 31, 38, 47, 48, 106 _ , _ , o f alkynes, 59, 60, 62, 69, 70 -, -, of c - C ~ H 1~0,0 , 1 0 1 -, -, of cyclohexenes, 9 5 , 9 8 -, -, of dienes, 74-77, 81, 8 2 -, free energy of decomposition of oxides, 232 isobutane, reaction + i-C4Hs, 334 isobutene, alkylation by, 334, 335 -, heterogeneous oxidation of, 124, 174-179,194,195,251 -, hydration o f , 323,325-327 -, hydrochlorination o f , 332, 333 -, reaction + ROH, 337 isobutyraldehyde, aldol condensation of, 344 isoprene, from heterogeneous oxidation, 195,196 isotope effect, see kinetic isotope effect
K kaolin, as cracking catalyst, 313, 314 kieselguhr, support for hydrogenation catalyst, 69, 1 0 2 kinetic isotope effect, and dehydration, 286,292,294, 295 -, and dehydrohalogenation, 303, 306, 308 Korvi earth, as esterification catalyst, 350,352 L Langmuir-Hinshelwood mechanism, for aldol condensation, 342, 343 -, for cracking, 311, 318 -, for esterification, 351, 358-360, 366 -, for heterogeneous acid-base catalysis, 271-273,280,284,285,302 -, for heterogeneous oxidation, 125127, 132, 134, 138, 155, 156, 173, 179, 187, 188, 194, 196, 197, 216, 224 -, f o r hydration, 325, 327, 329 -, f o r hydrogenation, 6, 7, 13, 15, 1 0 0 -, for hydrolysis, 372, 381
lanthanum, and heterogeneous oxidation, 183 -, and hydrolysis, 381 lead, and heterogeneous oxidation, 157, 1 6 9 , 1 8 3 , 184 -, as hydrogenation catalyst, 81 linear free energy relationships, and deamination, 297 -, and dehydration, 286, 288, 289 -, and dehydrohalogenation, 303, 304 -, and dehydrosulphidation, 319 -, and esterification, 354 -, and hydrocarbon cracking, 313, 314 lithium salts, and aldol condensation, 344 -, and dehydrohalogenation, 302, 303, 306,307 low energy electron diffraction, and heterogeneous oxidation, 130 -, and hydrogenation, 18, 54,103,112
M magnesium oxide, and cracking, 310
-, and dehydrobromination, 303 -, and heterogeneous oxidation, 241, 242
- ,_ , Of C3H6, 157 - ,_ , of C4H8,185,190,191 -, -, of CsHio, 196 magnesium sulphate, and dehydrohalogenation, 302, 303, 306 -, and hydrochlorination, 332, 333 magnetic susceptibility, and hydrogenation, 17 maleic anhydride, from heterogeneous oxidation, of aromatics, 196-206, 211-213,217,218, 220 - ,_ , Of C4H6, 246,248 -,-,of C4H8, 174, 175, 180, 182, 185, 191-193,248, 249,251 _ , _ , o f C s H i o . 196 malonitrile, and aldol condensation, 344 manganese, and dehydrohalogenation, 305 -, and heterogeneous oxidation, of aromatics, 202, 217 - , , of C3H6, 152, 157,159, 253 _ ,- , of C4H8, 185 _ , _ , of NH3, 228, 229 Markownikoff rule, 326, 332 mass spectrometry, and hydrogenation, 70
407 mass transport, and heterogeneous catalysis, 5, 6, 342, 351, 357, 384 mercuric chloride, as catalyst for hydrohalogenation, 332-334 mercury, effect o n hydrogenation, 46, 111 metal films, and hydrogenation, 3, 4, 23, 29, 30, 33, 34, 54, 102 methacrolein, from heterogeneous oxidation of i-C4Hs, 174-179 methacrylonitrile, from ammoxidation of i-C4&, 174-179 methanol, adsorption of, 312 -,dehydration o f , 282, 285, 289, 291, 294 -, esterification with, 349, 350, 353, 354,357-361, 366, 369, 370 -, heterogeneous oxidation of, 123, 125, 153,224-227,236, 249 -, hydrogenation in, 49 -, reaction + i-CqH8, 337 -, reaction + silica, 265 methyl acetate, hydrolysis of, 372, 374, 375,378 methylacetylene, hydrogenation of, 6870 methylamine, adsorption of, 298 methyl benzoate, hydrolysis o f , 374 o-methylbenzyl alcohol, and heterogeneo u s oxidation of o-xylene, 215 methyl bromide, and heterogeneous oxidation of C3H6, 156 2-methylbuta-l : 3-diene, hydrogenation of, 94 methylbutenes, heterogeneous oxidation of, 1 9 5 , 1 9 6 methylcyclohexane, and MezCO + PhOH, 343 methyl cyclopentanecarboxylate, hydrolysis of, 374 methyl elaidate, hydrogenation of, 49 methyl ethyl ketone, from heterogeneous oxidation of n-C4Hs, 175, 195 3-methyl-2-methylenebicyclo[ 2,2,2] octane, hydrogenation of, 98, 99 methylnaphthalene, heterogeneous oxidation o f , 220 methyl octanoate, hydrolysis o f , 374 methyl oleate, hydrogenation of, 49 methyl propyl ketone, from heterogeneo u s oxidation of C4H10, 196 microcrystallites, and catalysis, 104 microwave spectroscopy, and hydrogenation, 48
moderator (modifier) and heterogeneous oxidation, 127, 130, 156, 241 molecular sieves, as catalyst for deamination, 296 -, as catalyst for hydration, 330 -, as catalyst support, 110 molybdenum oxide, and oxygen exchange, 2 3 5 , 2 3 8 , 2 3 9 -, binding energy and, 234 -, for ammoxidation, 164, 166--169, 172,173,222,223 -, for heterogeneous oxidation, 124, 240-242,248-253 -, -, of aromatics, 199--206, 209, 217, 219,221 -, -, Of C3H6, 135, 140--154,157-160, 162-164, 237, 244--247, 250,251 -,-,of C.4H8, 175-185, 187, 189, 193, 195 _ ,_ , o f C5H1o1l95,196 -, -, of MeOH, 225-226 - ,_ , of NH3, 228-230 -, free energy of decomposition, 232 multiplet theory of catalysis, 103, 104
N naphthalene, adsorption of, 312
-, alkylation of, 334 -, heterogeneous oxidation o f , 124, 196, 217-2 19 naphthaquinone, from heterogeneous oxidation of naphthalene, 217, 218 nickel, and heterogeneous oxidation of aromatics, 202, 217, 221 -, -, O f C3H6, 157 -, -, of MeOH, 225, 227 -, as catalyst for hydrogenation, 2, 107, 108 -, -, of alkynes, 50, 52-54, 58-60, 62, 64,66,67,69-72 -,-,of CzH4, 15, 17-19, 21-23, 27, 30-34,106,110 _ ,- , o f C 3 H 6 , 3 8 -, -, of n-C4H8, 39, 40, 47, 48, 109 -, -, Of C6H6, 16,104, 105, 109 -, -, of cycloalkenes, 95-99 -, -, of cyclopropane, 100-102 -, -, of dienes, 74-86, 88, 91-94, 105, 109 -, as catalyst for isomerisation of olefins, 25,49 nickel carbonate, and dehydrochlorination, 307
408 nickel sulphate, and dehydrohalogenation, 302,303,305 -, and hydration, 327 niobium, and heterogeneous oxidation,
160,202,252 nitric oxide, and ammoxidation of C3H6,
-, as catalyst for olefin isomerisation, 25 -, clusters and hydrogenation, 106 oxygen, in lattice and heterogeneous oxidation, 236-239 -, o n Ag surface, 129,130 oxygen-18, and heterogeneous oxidation,
233-2 35,237-2 39
165
-, from heterogeneous oxidation o f NH3, -, -, O f C3H6, 145,146,160,163 -, -, of MeOH, 226,227 227,228 nitrogen, from heterogeneous oxidation of NH3, 227-229, 243 nitromethane, aldol condensation of, 342 nitrous oxide, effect o n heterogeneous oxidation of CzH4, 131 -, f r o m heterogeneous oxidation of NH3,
227,243
P palladium, and heterogeneous oxidation,
124,126,133-135,164
-, and hydrolysis, 381
N-profiles, and deuteration, 58, 79, 86, -, as catalyst for hydrogenation, 2,1 0 7 ,
91 nuclear magnetic resonance spectroscopy, and hydrogenation, 70,77
108,111,113 -, -, o f alkenes, 17-21, 23,30,31,33.-
38,40,42-45,47,49,107 -, -, of alkynes, 50-53, 59-67, 69-75,
109
0
-, -, Of C-C3H6, 100, 101
Ag7'0-octalin, hydrogenation of, 95,96, -, -, Of C6H6,16, 110 98 _ , _ , of cycloalkenes, 95,96,98,99 octanes, cracking o f , 316 _ , _ , o f dienes, 74-77, 81, 83,84, 86, octyl acetate, hydrolysis o f , 374 91-94 oleic acid, esterification with, 357, 358, -, as catalyst for olefin isomerisation, 25 361,362,369,378 particle size, and hydrogenation, 3, 104, qrder o f reaction, and ammoxidation, 105,113 165,166,168,171,223 partition function, and bimolecular sur-, and heterogeneous oxidation, of aroface reaction, 13-15 matics, 202,219 penta-l:3-diene, hydrogenation of, 92,93 -, -, o f C Z H ~128,133-135 , pentanols, dehydration o f , 283,288,289 -,-,Of C3H6, 144,151,153-155, 161, -, esterification with, 354, 357, 358,
162,164
-,
360-362
of C4H8, 176-178, 180,184-186, pentenes, hydrogenation of, 48,49 194 -, reaction + ROH, 337 pentyl acetate, hydrolysis of, 374 -, -, o f NH3, 229 pentynes, hydrogenation of, 71,73 -, and hydration, 326,328,330 phase, see crystal phase -, and hydrogenation, 7,8,14 phecanthrene, heterogeneous oxidation -, -, of alkenes, 29,30,42,47 o f , 219,220 -,-,of alkynes, 59,60,66,67,69,70, phenols, alkylation of, 334,335 72 -, cracking of, 312-314 -, -, O f C-C3H6, 100,102 -, -, of dienes, 74,81,85 -, reaction + Me*CO, 338,343,347 phenyldimethylsilanes, hydrolysis of, 381 -, and hydrolysis, 372,377,381 -, and MeOH + C4H8, 337 phenylethanols, dehydration of, 283, 284,288,292 osmium, as catalyst for hydrogenation, of alkenes, 30,33-36, 38,40,42,43,47 phosphoric acid, as alkylation catalyst, 335 _ ,- , of alkynes, 58,60-62,65-67,72-, as esterification catalyst, 355 74 -, -, of C-C3H6, 100, 101 --,as hydration catalyst, 323-325, 327--, -, of propadiene, 76,77 330 7 ,
409 phthalic anhydride, esterification with, 357 -, from heterogeneous oxidation, 196, 205,206,210-220 pi-ally1 species, in heterogeneous oxidation, 124, 136, 138-140, 145, 148, 150, 155-157, 161, 165, 174, 241, 251,252 -,in hydrogenation, 20, 25, 26, 38, 44, 47, 49, 87, 9 8 pi-complex and heterogeneous oxidation, 210,241, 248 -, and hydration, 327, 329 --, and hydrogenation, of alkenes, 22, 23, 105 - ,_ , of alkynes, 54-56 .-, -, of c - C ~ H1 ~0,0 , 1 0 2 --,-, of dienes, 79, 80, 87 -, and hydrohalogenation, 333 piperdine, and dehydration, 293 platinum, as catalyst for hydrogenation, 2,107,110 _ ,- , o f alkenes, 7, 18, 19, 21-23, 27, 30, 31, 33-38, 40, 42-44, 47, 49, 105,107,110,111 59-62, 64, -, -, o f alkynes, 50, 52-54, 65, 67, 69, 72, 74, 76, 77 _ ,- , O f C-C3H6, 100-102,105 _ ,- , of cycloalkenes, 95-99, 105 -,-,of dienes, 75-77, 83, 84, 86, 9094 -, as catalyst for olefin isomerisation, 25 -, as catalyst for oxidation, 225, 228 poisoning, of hydrogenation catalysts, 4, 8,17 polymerisation, and aldol condensation, 340 -, and C4H10 + C4&, 334 -, and CzH2 + H 2 0 , 328 -,and hydrogenation, 59, 60, 62, 68, 70, 75,81 pores in catalysts, 6, 133, 270 -, and esterification, 351, 357, 364 -, and hydrolysis, 376, 378, 383 -, and olefin elimination, 278 potassium oxide, and hydrogenation, 100 pre-exponential factor, and heterogeneous oxidation, of anthracene, 219 _ , - , of C2H4, 1 2 8 -, -, Of i-C4H8,176 _ ,- , 0 f S O z , 2 3 1 --,and hydrogenation on alloys, 109 propadiene, adsorption of on metals, 5 3 -, hydrogenation of, 74-76
propane, reaction + Dz on metals, 100, 101 propanols, dehydration of, 279, 283-286,288,289 -, esterification with, 349, 350, 353, 354,357-361,363--366,368,371 -, reaction + i-C4H8, 337 propene, adsorption of on Pt, 20, 21 -, alkylation by, 334, 335 -,heterogeneous oxidation of, 124, 135 e t seq., 237-239, 242-246,250-252 -, hydration of, 324,326 -, hydrochlorination of, 332, 333 -, hydrogenation of, 28, 36-38, 44 -, isomerisation of, 25, 48 -, reaction + H2O and H2S, 318 -, reaction + ROH, 337 propene oxide, from heterogeneous oxidation O f C3H6, 141, 157 propionaldehyde, aldol condensation of, 344 propionic acid, esterification with, 349, 350,357, 363 propyl acetate, hydrolysis of, 376, 377 propylamine, deamination of, 296 propylbenzene, adsorption of, 312 -, alkylation of, 335 -, cracking of, 314, 316 pulse technique, for elimination reactions, 281, 302, 319 -, for heterogeneous oxidation, 126, 127, 130, 136, 145, 146, 153, 154, 157, 158, 170, 171, 180, 181, 185, 187, 189, 190, 195, 209, 214, 225, 237,238 pumice, support for hydrogenation catalysts, 7, 34, 36, 37, 60, 62, 64, 68-70, 74,76,100,101,1G3,108 pyridines, adsorption of on oxide catalysts, 248, 249, 293, 294, 297-299, 312
Q quinoline, adsorption of, 312
R rake model, for heterogeneous oxidation, 124,125,201 rate coefficient, for aldol condensation, 341
410 rate coefficient - continued resonance, and intermediate in dehydra. -, f o r ammoxidation of C3H6, 166, 167, tion, 292 173 rhenium, as hydrogenation catalyst, 2, -, for esterification, 360, 367 30,31,101 -, for heterogeneous oxidation, of CzH4, rhodium, adsorption of C2H4 on, 18, 19 128 -, and oxidation of CzH4, 134 -, -, Of C3H6,145,167 -, catalyst f o r hydrogenation, of alkenes, -, -, Of C4H8, 1 8 2 , 1 8 7 , 1 8 9 , 1 9 2 30, 31, 33, 34, 36, 38, 40-44, 46, 47, -, -, Of C6H6, 1 9 9 , 2 0 3 49,107,111 -, -, of MeOH, 226 - ,_ , of alkynes, 50-53, 59-67,69,72-, -, of NH3, 229 75 -, -, of xylenes, 214,223 - ,_ , of c - C ~ H1~0,0 , 1 0 1 -, f o r hydrolysis, of acetal, 384 -,-,of cyclohexenes, 95, 96 -, -, o f esters, 378 -,-, o f dienes, 74-77, 81-84, 86, 88, -, -, of sucrose, 382 90,91,94 rate-controlling step, and acid-base catal- -, catalyst for hydrolysis, 381 ysis, 276 -, catalyst for olefin isomerisation, 25 -, and aldol condensation, 343 -, clusters and hydrogenation, 106 -, and cracking, 316, 318 Rideal-Eley mechanism, for aldol con-, and esterification, 351, 352, 355, 358 densation, 347 -, and heterogeneous oxidation, of aro- -, for alkylation, 336 matics, 209 -, for heterogeneous oxidation, 127, -, -, Of C 3 H 6 , 1 3 7 - 1 3 9 , 1 6 1 , 1 6 3 , 1 6 4 131,132,224 -, -, Of C4H8, 1 9 4 , 1 9 5 -, for hydration, 325 -, f o r hydrogenation, 6 -8, 14, 15 -, -, of MeOH, 225 -, -, o f SOz, 230 ruthenium, adsorption of C2H4 on, 19 -, and hydration, 325, 327, 330 -, and heterogeneous oxidation of C2H4, -, and hydrogenation, 4, 5, 67 134 -, and hydrolysis, 372, 381 -, catalyst f o r hydrogenation, of alkenes, rate law, for aldol condensation, 342, 343 30,31,33-36,38,40-43, 47,106 -, for alkylation, 335 -, -, o f alkynes, 53, 60--62, 65-67, 69, --,f o r deamination, 296, 297 7 2-7 4 -, for dehydration of alcohols, 283-285 -, -, of c - C ~ H 100,101 ~, -, for dehydrohalogenation, 302 -, -, of cyclohexenes, 95 -, for elimination reactions, 280 - ,_ , o f dienes, 74-77, 81, 8 2 -, for esterification, 352, 356, 358, 359 -, catalyst for hydrolysis, 381 -, catalyst for olefin isomerisation, 25 -, for heterogeneous oxidation, 125 -, -, of aromatics, 198, 206, 213 -, -, Of CzH4,128 C3H6, 138, 155 --,,___,, OOff C4H8, 176,179,182,187-189 S -, -, of MeOH, 224 -, -, of NH3, 228 salicylic acid, esterification with, 357, - ,_ , of SO*, 230, 231 358;366, 369, 370 -, for hydration, 324-326, 328, 329, Saytzeff rule, 277 331 scheelite type catalysts, for heterogene-, f o r hydrocarbon cracking, 311, 312 o u s oxidation, 250 -, for hydrochlorination, 333 -, for hydrogenation, 6 , 7 -,- , O f C3H6, 149--151,168 -- , -, O f C4H8,184 -, -, Of CzH2, 66, 67 selectivity, in ammoxidation, of aromat-, -, O f C3H6, 3 7 , 3 8 ics, 222 -, -, Of n-C4H8, 3 9 , 4 1 , 42 redox mechanism, 125, 126, 137, 138, -, -, of C3H~,164,167-169,171-173 -, in heterogeneous oxidation, of aromat152, 155, 156, 176, 178, 185-188, ics, 199-201, 203-205, 211, 213, 191, 197, 198, 206, 209, 213, 218, 215-218,221 219, 224,230, 244
411
-, -, of C~H4,127,131,133,134
stereospecificity, and dehydration, 292,
295 -, and dehydrohalogenation, 306, 307, 151--156,159,161,162,164,248 309 -,-,of C4H8,174,176-183, 185,186, -, and hydrogenation, 50,56,57,72,98, 189-193,195 -, -, of C5H10,195,196 112 steric effects, and aldol condensation, , , of MeOH, 224-227 _ , _ , of NH3,228,229 344 -,in hydrogenation, 11, 50,55,58,62, -, and deamination, 299 63, 69, 70, 72, 74-76, 85, 94-96, -, and dehydration, 287 -, and dehydrohalogenation, 304 112 -, and esterification, 353,361 sieve effect, 271,378,385 -, and heterogeneous oxidation, 253 silanes, reaction + silica, 265 -, and hydrolysis, 383 silica, and aldol condensation, 340-344 -, and cracking, 310,312,314,316,317 stop effect, in deamination, 300 strontium, and aldol condensation, 342 -, and deamination, 296,298,299 -, and dehydration, 283-286, 288,292, -, and ammoxidation, 169 294,295 -, and dehydrohalogenation, 302-304, -, and dehydrohalogenation, 302-308 308 -, and dehydrosulphidation, 318 styrene, from heterogeneous oxidation of --,and esterification, 350-355, 361 PhEt, 221 -, and hydrolysis, 372 sucrose, hydrolysis of, 380-384 -, carrier for hydration catalysts, 324---sulphur dioxide, heterogeneous oxidation 327,330 of, 123,125,230,231,236 -, carrier for hydrogenation catalyst, 3, sulphuric acid, and esterification, 363 15, 17-23, 29-31, 41-43, 46,50- -, and hydration, 323,326 52, 58, 59, 62, 63,74,76,78,100- -, and hydrolysis, 376 102,105,109,110 -, and ROH + C4H8,337 -, carrier for oxidation catalyst, 133.- surface adsorption sites, 13,15,16 135, 153, 154, 157, 159, 164, 166, surface defects, 3 173, 185, 201, 213, 219, 231, 234, surface potential, and CzH2 o n Ni, 54 241,242 -, and C2H4 o n Ni and Pd, 23 -, nature of, 264-266, 268,269 surface spectroscopy, and heterogeneous silver, and ammoxidation of C3H6,169 oxidation of C3H6,143 -,-,Of
_
C3H6,136,137,142,144,147,
I
-, catalyst for heterogeneous oxidation, of C2H4,124-133, 236
_ , _ , of MeOH, 224,225
T
-, catalyst for hydrogenation, 107,108
Taft correlation, and cracking, 313,314
silver oxide, and heterogeneous oxidation of C2H4,129 -, and hydration, 330,331 -, free energy of decomposition, 232 sodium, and aldol condensation, 340-
-, and dehydration of alcohols, 288,289,
344
--,and esterification, 350,353,354 -, and hydrolysis, 380,383 specific activity, of catalyst and hydrogenation, 30-32, 105-107 spinels, and catalysis of oxidation, 190 stannous nitrate, and heterogeneous oxidation of C3H6,141 steady state, and acid-base catalysis,
273,285
--, and CzH2 + H2,66,67
292-294
-, and dehydrohalogenation, 303,304 -, and dehydrosulphidation, 319
-, and esterification, 353,354,360 tantalum, and heterogeneous oxidation,
160,202
-, hydrogenation of
C2H4 o n , 106 tellurium, and heterogeneous oxidation, C3H6, 151, 152, 157, 159, 160, 246,252 -, -, of i-C4H8,179 Of
terephthalonitrile, from ammoxidation of p-xylene, 222,223 1,1,2,2-tetrachloroethane,dehydrochlorination of, 304,308
412 tetracyanoethylene, and dehydration of alcohols, 268, 293 2,2,5,5-tetramethylhex-3-ene, hydrogenation of, 3 thallium, and heterogeneous oxidation, of C3H6,136,161,250 _ , _ , Of C6H,5, 202 thiols, dehydrosulphidation of, 318-320 thiophene, alkylation of, 334 thoria, and elimination reactions, 287, 301 -, and hydration, 330 tin, and ammoxidation, 166, 170-173, 222 -, and free electrons in oxide, 243 -,and free energy of decomposition of oxides, 232 -, and heterogeneous oxidation, of aromatics, 202-204, 206, 220 -,-,Of C3H6, 152, 155-163, 237, 239, 246,248,250, 252 -, -, of C4H8, 179, 186--189, 191-195, 244 -,-, of CsHlo, 196 -, -, of NH3, 230 titanium, and ammoxidation, 172, 173, 222 -, and dehydration, 283, 286, 288, 292, 294 -, and heterogeneous oxidation, of aromatics, 202-204, 210, 214-21 7 -, -, of C Z H ~134 , -, -, O f C3H6, 1 5 4 , 1 6 1 , 1 6 2 , 246 _ , _ , of C4H8,185,194 -, -, of C ~ H 1 0 , 1 9 6 -, and hydrolysis, 372 -, and oxygen exchange, 235 -, free energy of decomposition of oxide, 232 tolualdehyde, and heterogeneous oxidation of xylenes, 206 208, 210-214, 216, 217 toluene, alkylation of, 335, 336 -, ammoxidation of, 221, 222 -, esterification in, 367 -, from heterogeneous oxidation of xylenes, 207, 208 -, heterogeneous oxidation of, 201, 202, 204-207,209,219 toluylnitrile, from ammoxidation of p-xylene, 222, 223 transient-response technique, 272 transition state, and bimolecular surface reaction, 13-15
-, -, -, -,
and dehydration, 287 and dehydrohalogenation, 309 and hydrolysis, 313, 375, 378-381 and oxygen exchange, 235 2,2,2-trichloro-l,1-bis(p-chloropheny1)ethane, dehydrochlorination of, 300 trichloroethanes, dehydrochlorination of, 302,304-307 1,1,2-trichloropropane, dehydrochlorination of, 308 triethylamine, deamination of, 297, 300 1,2,4-trimethylbenzene, heterogeneous oxidation of, 220 trimethylbut-1-ene, heterogeneous oxidation of, 195 tritium, as tracer in hydrogenation, 49, 111 tungsten, and ammoxidation of C3H6, 169 -, and esterification, 350, 352 -, and heterogeneous oxidation, 240, 242,252 _ , _ , of aromatics, 201-203, 205, 220, 221 _ ,- , O f C3H6,149,150,158-160,163 - ,_ , of CqHB, 183,194 _ , _ , of MeOH, 225 -, and hydration, 323, 325, 326 -, as hydrogenation catalyst, 2, 17, 51, 95,106 -, free energy of decomposition of oxides, 232
U ultraviolet spectroscopy, and heterogeneous oxidation, 148 uranium, and ammoxidation, 169, 170, 172,173,222 -, and heterogeneous oxidation, of aromatics, 202-204, 209, 210 -, -, OfC3&,163,166, 237, 239,247 _ , _ , of C4H8, 178, 1 7 9 , 1 8 9 , 1 9 0 -, free energy of decomposition of oxide, 232
V vanadium, and ammoxidation, 172, 173, 222,223 -, and binding energy of oxygen, 234 -, and heterogeneous oxidation, 240242, 247,248,252
413
_ , _ , of
acrolein, 159
_ , _ , o f aromatics, 196, 197, 199--204, 206, 209-21 1, 213-220 - ,_ , of C2H4,134,135 _ , _ , O f C3H6, 157,169,242,251 -,-,Of CqHs, 179,180,183,184,191194, 239, 244 _ , _ , O ~ C S H L196 O, _ , - , of MeOH, 226,227 -, -, of NH3,228,229 -, --, o f S02, 230, 231 -, and oxygen exchange, 235 -, free energy of decomposition of oxides, 232 vinyl alcohol, and C2H2 + H20, 328 vinyl chloride, and C2H2 + HCI, 332, 333
W Wacker process, 126, 1 3 3 water, and alumina catalysts, 266 -, and ammoxidation of C3H6, 167, 170, 171 -, and elimination reactions, 279, 281, 293, 294, 306, 309 -, and esterification, 351, 355, 356, 358, 366 -, and heterogeneous oxidation, of CzH4, 133,134 -,-,of C3H6, 137, 146, 148, 149, 159, 160,162-164,240 - ,_ , of CqHs, 1 7 8 , 1 8 0 , 1 8 6 , 1 8 9 -, -, Of C,H6, 200 -, and silica catalysts, 265 -, reaction + C2H2, 327-329 -, reaction + C2H40, 329-331 -, reaction + olefins, 318, 321-327 work function, and hydrogenation, 17
X X-ray analysis, 103, 143, 151, 202, 214 xylenes, ammoxidation of, 222, 223 -, from heterogeneous oxidation of C4Hs,194 -, heterogeneous oxidation of, 124, 196, 204,206,207, 210-217
Z
zeolites, and aldol condensation, 340, 347 -, and elimination reactions, 282, 302, 309,314,316, 318,320 -, and heterogeneous oxidation, 157, 229 -, as alkylation catalysts, 334-336 -, as hydration catalysts, 323, 324, 328, 329 -, nature of, 268, 269 zinc, and dehydration, 284 -, and dehydrohalogenation, 305 -, and heterogeneous oxidation, of C&, 157,161 -, -, of C4Hs,190,191 _ ,- ,of ~ ~ H196 I D ~ -,-,Of C6H6, 202 -, and hydration of C2H2, 328, 329 zinc chloride, as catalyst for hydrohalogenation, 332, 333 zirconium, and cracking, 310, 314 -,and dehydration, 283, 286, 288, 292, 294 -, and dehydrochlorination, 301 -, and heterogeneous oxidation of C6H6, 202 -, and hydrolysis, 372
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