Metal-Oxygen Clusters: The Surface and Catalytic Properties of Heteropoly Oxometalates (Fundamental and Applied Catalysis)

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Metal–Oxygen Clusters The Surface and Catalytic Properties of Heteropoly Oxometalates

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FUNDAMENTAL AND APPLIED CATALYSIS Series Editors: M. V. Twigg Johnson Matthey Catalytic Systems Division Royston, Hertfordshire, United Kingdom

M. S. Spencer Department of Chemistry Cardiff University Cardiff, United Kingdom

CATALYST CHARACTERIZATION: Physical Techniques for Solid Materials Edited by Boris Imelik and Jacques C. Vedrine CATALYTIC AMMONIA SYNTHESIS: Fundamentals and Practice Edited by J. R. Jennings CHEMICAL KINETICS AND CATALYSIS R. A. van Santen and J. W. Niemantsverdriet DYNAMIC PROCESSES ON SOLID SURFACES Edited by Kenzi Tamaru ELEMENTARY PHYSICOCHEMICAL PROCESSES ON SOLID SURFACES V. P. Zhdanov METAL–OXYGEN CLUSTERS: The Surface and Catalytic Properties of Heteropoly Oxometalates John B. Moffat PRINCIPLES OF CATALYST DEVELOPMENT James T. Richardson SELECTIVE OXIDATION BY HETEROGENEOUS CATALYSIS Gabriele Centi, Fabrizio Cavani, and Ferrucio Trifir`o

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Metal–Oxygen Clusters The Surface and Catalytic Properties of Heteropoly Oxometalates John B. Moffat University of Waterloo Waterloo, Ontario, Canada

Kluwer Academic Publishers New York, Boston, Dordrecht, London, Moscow

eBook ISBN: Print ISBN:

0-306-47378-X 0-306-46507-8

©2002 Kluwer Academic Publishers New York, Boston, Dordrecht, London, Moscow Print ©2000 Kluwer Academic / Plenum Publishers New York All rights reserved No part of this eBook may be reproduced or transmitted in any form or by any means, electronic, mechanical, recording, or otherwise, without written consent from the Publisher Created in the United States of America Visit Kluwer Online at: and Kluwer's eBookstore at:

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PREFACE TO THE SERIES

Catalysis is important academically and industrially. It plays an essential role in the manufacture of a wide range of products, from gasoline and plastics to fertilizers and herbicides, which would otherwise be unobtainable or prohibitively expensive. There are few chemical- or oil-based material items in modern society that do not depend in some way on a catalytic stage in their manufacture. Apart from manufacturing processes, catalysis is ®nding other important and everincreasing uses; for example, successful applications of catalysis in the control of pollution and its use in environmental control are certain to increase in the future. The commercial importance of catalysis and the diverse intellectual challenges of catalytic phenomena have stimulated study by a broad spectrum of scientists, including chemists, physicists, chemical engineers, and material scientists. Increasing research activity over the years has brought deeper levels of understanding, and these have been associated with a continually growing amount of published material. As recently as sixty years ago, Rideal and Taylor could still treat the subject comprehensively in a single volume, but by the 1950s Emmett required six volumes, and no conventional multivolume text could now cover the whole of catalysis in any depth. In view of this situation, we felt there was a need for a collection of monographs, each one of which would deal at an advanced level with a selected topic, so as to build a catalysis reference library. This is the aim of the present series, Fundamental and Applied Catalysis. Some books in the series deal with particular techniques used in the study of catalysts and catalysis: these cover the scienti®c basis of the technique, details of its practical applications, and examples of its usefulness. An industrial process or a class of catalysts forms the basis of other books, with information on the fundamental science of the topic, the use of the process or catalysts, and engineering aspects. Single topics in catalysis are also treated in the series, with books giving the theory of the underlying science, and relating it to catalytic practice. We believe that this approach provides a collection that is of value to v

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PREFACE TO THE SERIES

both academic and industrial workers. The series editors welcome comments on the series and suggestions of topics for future volumes.

Royston and Cardiff

Martyn Twigg Michael Spencer

PREFACE

Heterogeneous catalysts, as their name implies, are employed in multiphase systems, the ¯uid phase most frequently being gaseous. Although heteropoly oxometalates or metal±oxygen cluster compounds appear in a variety of forms, the present book is largely restricted to those with Keggin structure. In their acidic form, these present relatively small surface areas to the reactant molecules in the ¯uid phase. However, as described in Chapter 7, heteropoly oxometalates with microporous=mesoporous structures and hence relatively large surface areas can be prepared. Equally interesting and important, the structures and largely the morphology of the heteropoly oxometalates, which are monophasic, one-component systems, can be retained while the catalytic functionality is altered by variations of the elemental compositions of the anions. Because this allows the study of compositional effects on the catalysis without the intrusion of other variable effects, it is particularly useful for fundamental studies in catalysis. Although, with the simplest acidic forms, the syntheses are reasonably straightforward, in contrast the preparation of the salts and those solids containing atoms of more than three elements in the anion must be performed carefully with thorough characterization of the products, particularly elemental analyses. The relative quantities of the preparative elements are not necessarily re¯ected in the resulting solids. Not surprisingly, the stabilities of these materials are often less than desired, although again these vary with both the nature of the cation and the anionic composition. The use of high-surface-area supports has been investigated for the purpose of both augmenting the interphase interactions as well as increasing the stabilities of the supported heteropoly oxometalates. As the reader will discover, vii

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PREFACE

agreement on these aspects has not always been forthcoming and additional work is obviously required. In spite of the low surface areas and the absence of a pore structure in the acidic forms of the heteropoly oxometalates, polar molecules have been shown to be capable of penetrating into the structures, that is, between the cations and anions. Although this does not increase the real surface area available to reactant molecules in the ¯uid phase, nevertheless the portion of the solid available to these molecules is effectively increased. Although the crystal structure of the solids is retained during such diffusional phenomena, the lattice constants, not surprisingly, are found to vary. Thus, the heteropoly oxometalates can be considered as three-dimensional intercalates. The author is grateful to the many undergraduate and graduate students and postdoctoral fellows, too numerous to name, but whose identities are revealed in the references, who have contributed to the work. The author is also grateful to Catherine Van Esch and Beverley Winkler who translated the author's handwriting into legible typewritten form and endured numerous changes during the writing process. Last but not least, I thank my wife, Eleanor, for her enduring patience and support.

CONTENTS

CHAPTER 1. A BRIEF LOOK AT THE EARLY HISTORY OF HETEROPOLY OXOMETALATES . . . . . . . . . . . . . .

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References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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CHAPTER 2. SYNTHESIS . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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2.1. Introduction . . . . . . . . . . . . . . . . . . . 2.2. The 12-Heteropoly Acids . . . . . . . . . . 2.2.1. Tungsten-Containing Anions . . . 2.2.2. Molybdenum-Containing Anions 2.2.3. Other Heteropoly Acids . . . . . . 2.2.4. Mixed Addenda . . . . . . . . . . . . 2.3. Various Syntheses . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . .

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CHAPTER 3. CHARACTERIZATION . . . . . . . . . . . . . . . . . . . . . .

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3.1. Spectroscopy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1.1. Infrared Spectroscopy . . . . . . . . . . . . . . . . . . . 3.1.2. Photoacoustic FTIR Spectroscopy . . . . . . . . . . . 3.1.3. Nuclear Magnetic Resonance . . . . . . . . . . . . . . 3.1.4. Electron Paramagnetic Resonance and Electronic Spectroscopy . . . . . . . . . . . . . . . . . . . . . . . . .

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3.1.5 X-ray Photoelectron Spectroscopy . . . . . . . . . . 3.1.6 Scanning Tunneling Microscopy and Tunneling Spectroscopy . . . . . . . . . . . . . . . . . . . . . . . . 3.2. Electrochemical Methods . . . . . . . . . . . . . . . . . . . . 3.3. Elemental Analysis . . . . . . . . . . . . . . . . . . . . . . . . 3.3.1. HPLC . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3.2. Ion Chromatography. . . . . . . . . . . . . . . . . . . 3.3.3. EDTA Titration . . . . . . . . . . . . . . . . . . . . . . 3.4 Other Techniques . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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CHAPTER 4. STRUCTURE AND BULK PROPERTIES . . . . . . . . .

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4.1. Anions . . . . . . . . . . . . . . . . . . . . . . 4.2. Cations and Crystallographic Structure . 4.3. Langmuir Films . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . .

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29 34 37 39

CHAPTER 5. STABILITY . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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5.1. Thermal Stability . . . . . . . . . . . . . . 5.2. pH Stability . . . . . . . . . . . . . . . . . . 5.2.1. Molybdate=Phosphate System . 5.2.2. Tungstate=Phosphate System . . 5.3. Regeneration in the Presence of Water References . . . . . . . . . . . . . . . . . . .

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CHAPTER 6. SUPPORTED HETEROPOLY ACIDS AND THEIR DERIVATIVES. . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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6.1. Supports . . . . . . . . . . . 6.1.1. Carbon . . . . . . . . 6.1.2. Titanium Dioxide . 6.1.3. Silica . . . . . . . . . 6.1.4. Alumina . . . . . . . 6.1.5. MgF2 . . . . . . . . . 6.1.6. SiC . . . . . . . . . . 6.1.7. ZrO2 . . . . . . . . . 6.1.8. MCM-41 . . . . . .

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CONTENTS

6.1.9. Zeolite Y . . . . . . . . . 6.1.10. Heteropoly Salts. . . . . 6.1.11. Clays . . . . . . . . . . . . 6.1.12. Polymers. . . . . . . . . . 6.2. Formation of Heteropoly Acids References . . . . . . . . . . . . . .

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CHAPTER 7. MICROPOROSITY . . . . . . . . . . . . . . . . . . . . . . . . .

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7.1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7.2. Microporosity of Salts Prepared from the Monovalent Cations of the Group Alkali Metals and Ammonium and Related Cations . . . . . 7.2.1. Surface Areas and Pore Structures from the Analysis of Nitrogen Adsorption±Desorption Isotherms . . . . . . . . . . . . . 7.2.2. The Source of the Microporosity . . . . . . . . . . . . . . . . . . . . 7.2.3. The Dependence of Morphological Properties of Salts Prepared from Group 1 Monovalent Cations on the Source of the Cations and the Stoichiometry of the Salt . . . . . . . . . . 7.3. Microporosity of Salts of other Cations. . . . . . . . . . . . . . . . . . . . . 7.3.1. Microporosity of Salts Prepared from the Monovalent Cations of the Group 11 and 13 Elements . . . . . . . . . . . . . . 7.3.2. Stoichiometric and Nonstoichiometric Salts of Groups 11 and 13 and Their Morphological Properties . . . . . . . . . . . . . . . . 7.4. Pore Structures from 129Xe NMR . . . . . . . . . . . . . . . . . . . . . . . . 7.5. Sorption and Diffusion in Metal±Oxygen Cluster Compounds . . . . . 7.6. Ion Exchange and Structure Retention . . . . . . . . . . . . . . . . . . . . . 7.6.1. Cation Exchange and Microporosity . . . . . . . . . . . . . . . . . . 7.6.2. Crystal Structure and Morphology Retention on Ion Exchange . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7.7. Argon Adsorption . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7.8. Divalent Salts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7.9. Reactions on the Microporous Salts and Shape Selectivity . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

97 98 98 101 105 110 110 111 116 119 121 121 125 135 136 139 141

CHAPTER 8. THE TWO FUNCTIONS: ACIDITY AND OXIDATION±REDUCTION . . . . . . . . . . . . . . . . . . . 143 8.1. Acidity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 143 8.1.1. Electrical Conductivity . . . . . . . . . . . . . . . . . . . . . . . . . . . 143 8.1.2. Calorimetry. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 144

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8.1.3. Theoretical . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8.1.4. Solid±Liquid Phase Titrations. . . . . . . . . . . . . . . . . . . . 8.1.5. Adsorption±Desorption . . . . . . . . . . . . . . . . . . . . . . . . 8.1.6. Spectroscopy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8.1.7. Temperature-Programmed Desorption . . . . . . . . . . . . . . 8.1.8. Probe Reactions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8.2. Oxidation±Reduction and the Properties of the Anionic Oxygen Atoms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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CHAPTER 9. ACID-CATALYZED PROCESSES . . . . . . . . . . . . . . . 175 9.1. Methanol Conversion to Hydrocarbons . . . . . . . . . . . . . . . . . . . 9.1.1. Heteropoly Acids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.1.2. Metallic Salts of 12-Tungstophosphoric Acid . . . . . . . . . . 9.1.3. Ammonium 12-Tungstophosphate . . . . . . . . . . . . . . . . . . 9.1.4. Mechanistic Studies with Photoacoustic FTIR Spectroscopy 9.1.5. Alkylammonium Oxometalates . . . . . . . . . . . . . . . . . . . . 9.2. Ethanol . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.3. Propene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.4. Propanol . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.5. Butane. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.6. Butene. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.7. Isobutane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.8. Isobutene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.9. Butanol . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.10. Alkylation Processes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.10.1. Isobutene and Methanol. . . . . . . . . . . . . . . . . . . . . . . . 9.10.2. Isobutene and 2-Butene . . . . . . . . . . . . . . . . . . . . . . . . 9.10.3. Benzene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.11. Friedel±Crafts Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.12. C5±C8 Alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.13. C6±C8 Alkenes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.14. Ring Expansion of Methylcyclopentane and Ring Contraction of Cyclohexane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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175 175 178 180 182 186 189 192 193 194 196 201 206 208 209 209 212 214 215 216 217

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CHAPTER 10. OXIDATION PROCESSES . . . . . . . . . . . . . . . . . . . 227 10.1. Methane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 227 10.1.1. Methane with N2 O and O2 as Oxidants . . . . . . . . . . . . . 227

CONTENTS

10.2. 10.3. 10.4. 10.5. 10.6. 10.7. 10.8. 10.9. 10.10. 10.11. 10.12.

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10.1.2. The Effect of Gas Phase Additives in the Conversion Methane on Heteropoly Oxometalates . . . . . . . . . . Ethane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Propane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Butane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Isobutane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Methacrolein . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Isobutyric Acid . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . n-Pentane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Butadiene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1-Butene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Cyclohexane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Ammoxidation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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248 252 259 260 263 267 268 283 284 284 285 286 286

CHAPTER 11. ENVIRONMENTALLY RELATED PROCESSES . . . 289 11.1. Conversion of Nitrogen Oxides . . . . . . . 11.1.1. The Heteropoly Acids . . . . . . . . 11.1.2. Ammonium 12-Tungstophosphate 11.2. Hydrodesulfurization . . . . . . . . . . . . . . 11.3. Polymerization . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . .

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289 289 296 302 303 304

INDEX . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 307

1 A BRIEF LOOK AT THE EARLY HISTORY OF HETEROPOLY OXOMETALATES

Although some disagreement on the topic appears in the literature, the ®rst heteropoly compound can be traced to the work of Berzelius who in 1826(1) prepared ammonium 12-molybdophosphate as a yellow precipitate from the addition of ammonium molybdate to phosphoric acid. The heteropoly molybdates of Cr3‡ and Fe3‡ were reported in 1854 by Struve as double salts.(2) By 1862 Marignac had prepared 12-tungstosilicic acid and provided analytical compositions.(3) Although hundreds of polyoxometalates were synthesized during the next half-century, little progress was made in understanding their structures. In 1908 Miolati attempted to provide a structural interpretation for these materials by application of ionic coordination theory.(4) This theory was further developed by Rosenheim(5,6) whose contributions continued for the next quartercentury. In 1929 the Miolati±Rosenheim theory was criticized by Pauling who suggested a cage structure of MoO6 octahedra joined by corners into a shell enveloping the PO4 3ÿ ion(7) Although Pauling's proposal was a step in the right direction, Hoard,(8) employing X-ray diffraction techniques, was unable to provide support for this theory. In 1933 Keggin(9) provided the ®rst de®nitive information on a heteropoly compound by showing from X-ray diffraction that the WO6 octahedral units in H3PW12O40
5H2O were connected by both shared edges and corners. His work was con®rmed by Bradley and Illingworth in 1936(10) from their studies of H3PW12O40
29H2O. Although these results were based on powder X-ray diffraction, they were largely supported by the singlecrystal experiments of Brown and co-workers the results of which were reported in 1977.(11) Three hydrates of 12-tungstophosphoric acid H3PW12O40
nH2O, where n ˆ 6,(11) 21,(12) 29,(13) have been studied by single-crystal X-ray and 1

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CHAPTER 1

neutron diffraction and with the 6-hydrate the protons were found to be hydrogenbonded with two water molecules. Additional discussion of the syntheses and structures of these materials can be found in the chapters on these topics. More detailed reviews of the history of these materials together with extensive discussions of a variety of structural and chemical properties are available.(14±49)

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26.

J. Berzelius, Pogg. Ann. Phys. Chem. 6, 369, 380 (1826). H. Struve, J. Prakt. Chem. 55, 888 (1854). C. Marignac, J. Prakt. Chem. 77, 417 (1862). A. Miolati and R. Pizzighelli, J. Prakt. Chem. 77, 417 (1908). A. Rosenheim and J. Jaenicke, Z. Anorg. Allg. Chem. 100, 304 (1917). A. Rosenheim, in: Handbuch der Anorganischen Chemie (R. Abegg and F. Auerbach, eds.), Vol. 4, Part 1, p. 977, Hirzel Verlag, Leipzig (1921). L. Pauling, J. Am. Chem. Soc. 51, 2868 (1929). J. L. Hoard, Z. Kristillogr. 84, 217 (1933). J. F. Keggin, Nature 131, 968 (1933); 132, 351 (1933); Proc. R. Soc. London Ser. A 144, 75 (1934). A. J. Bradley and J. W. Illingworth, Proc. R. London Ser. A 157, 113 (1936). G. M. Brown, M. R. Noe-Spirlet, W. R. Busing, and H. A. Levy, Acta Crystallogr. B 33, 1038 (1977). M. R. Noe-Spirlet, G. M. Brown, W. R. Busing, and H. A. Levy, Acta. Crystallogr. A 31, S80 (1975). M. R. Spirlet and W. R. Busing, Acta Crystallogr. B 34, 907 (1978). J. W. Mellor, in: Comprehensive Treatise on Inorganic and Theoretical Chemistry, Longmans, Green, London; Vol. 9, p. 757 (1929) (Vanadates); Vol. 11, p. 535 (1931) (Molybdates), p. 659 (Molybdophosphate), p. 762 (Tungstates), p. 862 (Tungstophosphates). Gmelins: Handbuch der Anorganischen Chemie. System No. 53 (Molybdenum), System No. 54 (Tungsten), Verlag Chemie, Berlin (1935). A. Morette, in: Nouveau Traite de Chimie MineÂrale (P. Pascal, ed.), Masson, Paris (1958), Vol. 12, p. 207 (Vanadates); R. Rohmer, pp. 468, 600 (Niobates and Tantalates); L. Malaprade, Vol. 14, p. 656 (Molybdates), p. 903 (Heteropolyanions); A. ChreÂtien and W. Freundlich, p. 835 (Tungstates). D. L. Kepert, Prog. Inorg. Chem. 4, 199 (1962). P. Souchay, Polyanions et polycations, Gauthier Villars Editeur, Paris (1963). G. A. Tsigdinos, Heteropoly Compounds of Molybdenum and Tungsten. Climax Molybdenum Company Bulletin Cdb-12a (Revised), Ann Arbor, MI (1969). P. Souchay, Ions MineÂraux CondenseÂs, Masson, Paris (1969). L. C. W. Baker, in: Advances in the Chemistry of the Coordination Compounds (S. Kirschner, ed.), p. 604, Macmillan Co., New York (1961). K. F. Jahr and J. Fuchs, Angew. Chem. Int. Ed. Engl. 5, 689 (1966). H. T. Evans, Jr., Perspect. Struct. Chem. 4, 1 (1971). D. L. Kepert, The Early Transition Elements, Academic Press, New York (1972). P. Souchay, M. Boyer, and F. Chauveau, Kg. Tek. Hoegsk. Handl. 259 (1972). [Contributions to Coordination Chemistry in Solution, Stockholm, p. 159 (1972)]. D. L. Kepert, in: Comprehensive Inorganic Chemistry (A. F. Trotman-Dickerson et al., eds.), Vol. 4, p. 607, Pergamon Press, Oxford (1973).

THE EARLY HISTORY

3

27. G. A. Tsigdinos, Ind. Eng. Chem. Prod. Res. Dev. 13, 267 (1974). 28. G. A. Tsigdinos, in: Methodicum Chimicum (F. Korte, editor-in-chief), Vol. 8 (K. Niedenzu and H. Zimmer, eds.), Chapter 32, p. 552, Academic Press, New York (1976). 29. T. J. R. Weakley, Struct. Bonding 18, 131 (1974). 30. L. P. Kazanskii, P. A. Torchenkova, and V. I. Spitsyn, Usp. Khim. 43, 1337 (1974). 31. L. Barcza and M. T. Pope, J. Phys. Chem. 79, 92 (1975). 32. G. A. Tsigdinos and C. J. Hallada, Cdb-14, Climax Molybdenum Co., Ann Arbor, MI (1969). 33. K. H. Tytko and O. Glemser, Adv. Inorg. Chem. Radiochem. 19, 239 (1976). 34. G. A. Tsigdinos, Top. Curr. Chem. 76, 1 (1978). 35. L. P. Kazanskii, M. A. Fedotov, I. V. Potapova, and V. I. Spitsyn, Dokl. Chem. 244, 36 (1979). 36. L. P. Kazanskii, E. A. Torchenkova, and V. I. Spitsyn, Russ. Chem. Rev. 43, 525 (1974). 37. M. T. Pope and B. W. Dale, Q. Rev. Chem. Soc. 22, 527 (1975). 38. V. I. Spitsyn, L. P. Kazanskii, and E. A. Torchenkova, Sov. Sci. Rev. B Chem. Rev. 3, 111 (1981). 39. M. T. Pope, Heteropoly Oxometalates, Springer-Verlag, Berlin (1983). 40. I. V. Kozhevnikov and K. I. Matveev, Appl. Catal. 5, 135 (1983). 41. V. W. Kay, W. G. Klemperer, C. Schwartz, and R.-C. Wang, in: Surface Organometallic Chemistry: Molecular Approaches to Surface Catalysis NATO ASI Ser. 231, 173 (1988). 42. M. Misono, Catal. Rev.-Sci. Eng. 29, 269 (1987). 43. J. B. Moffat, Rev. Chem. Intermed. 8, 1 (1987). 44. J. B. Moffat, Chem. Eng. Commun. 83, 9 (1989). 45. M. T. Pope and A. MuÈller, Angew. Chem. Int. Ed. Engl. 30, 34 (1991). 46. Polyoxometalates: From Platonic Solids to Anti-Retroviral Activity (M. T. Pope and A. MuÈller, eds.), Kluwer, Dordrecht (1994). 47. I. V. Kozhevnikov, Catal. Rev.-Sci. Eng. 37, 311 (1995). 48. R. J. J. Jansen, H. M. van Veldhuizen, M. A. Schwegler, and H. van Bekkum, Recl. Trav. Chim. Pays-Bas 113, 115 (1994). 49. T. Okuhara, N. Mizuno, and M. Misono, Adv. Catal. 41, 113 (1996).

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2 SYNTHESIS

2.1. INTRODUCTION Although, as noted in the chapter on history, synthesis of the heteropoly oxometalates can be traced back to as early as 1826, as Tsigdinos has pointed out, the products were often not fully characterized and the purity of the product may often be suspect.(1) Tsigdinos provided a critical evaluation of preparative procedures in 1976(2) and the monograph by Pope updated this to 1983(3) with a later but more concise review with MuÈller(4) and a historical perspective by Baker and Glick.(5) Most recently, Was® and Johnson have published a review of syntheses and procedures for the tungsten-containing heteropoly oxometalates.(6) Of course, as is well known the methods available to assess the nature and purity of products and hence the validity of a synthetic method have burgeoned in the last several decades with obvious advantages to synthetic chemists. It is neither practical nor appropriate to attempt a complete review of the heteropoly oxometalates and the methods by which they have been prepared. Pope has noted that heteropoly anions have been prepared with more than 65 elements as the central atom (in Pope's terminology the heteroatom) while the elements that serve as peripheral metal atoms appear to be more restrictive, apparently requiring certain ionic radius and charge and the ability to form d±p peripheral metal±oxygen bonds.(3) Accordingly, those heteropoly oxometalates that have been employed as heterogeneous catalysts or appear to have potential applications for such purposes will be discussed. These are largely of Keggin structure although some with Dawson structure have been examined as catalysts. That which follows here is intended to illustrate the synthesis of a few of the myriad of species that have been prepared to date. 5

6

CHAPTER 2

2.2. THE 12-HETEROPOLY ACIDS 2.2.1. Tungsten-Containing Anions Probably more heterogeneous catalytic studies have involved heteropoly anions with phosphorus as central atom and Keggin structure than any other heteropoly oxometalates. Almost inevitably the peripheral metal atoms are tungsten, molybdenum, and vanadium or combinations of these: PW12 O40 3ÿ PMo12 O40 3ÿ PW12ÿx Mox O40 3ÿ PW12ÿx Vx O40 …3‡x†ÿ PMo12ÿx Vx O40 …3‡x†ÿ Although somewhat tangential to the present purposes, it is not uninteresting to note that, where 12 vanadium atoms are imagined to be present, the charge on the anion would be ÿ15. Not surprisingly, this has not been synthesized, although [PV14O42]9ÿ which has an a-Keggin structure with two additional vanadium atoms occupying trans-related VO5 trigonal bipyramids has been prepared.(3,7) In addition to the aforementioned phosphorus compounds, those with Keggin structure and silicon as a central atom have also been studied for their catalytic properties. As will be discussed later, the use of these materials undoubtedly relates to their relatively high stabilities. 12-Tungstophosphoric acid can be prepared from appropriate tungstates and phosphates, for example, sodium tungstate and disodium phosphate, respectively, in acidi®ed aqueous solutions(2,3,6,8±12): 12WO4 2ÿ ‡ HPO4 2ÿ ‡ 23H‡ ! …PW12 O40 †3ÿ ‡ 12H2 O The control of temperature and pH has been shown to be important.(3) Pope cites the example of the use of excess H3PO4 which after boiling in the appropriate aqueous solution yields a mixture of tungstophosphates including [P2W18O62]nÿ . As will be discussed in the chapter on stability, the heteropoly acids are well known to be pH-sensitive. A combined 31P NMR and FTIR spectroscopic study has shown that, with H3PW12O40 in aqueous solution, at pH < 8 a single NMR peak appears at ÿ10:0 ppm which converts to 3.5 ppm at pH > 8.(13) The former is attributed to PW11O39 7ÿ, which exists as the major species down to pH 2.0. Although the heteropoly anions can be isolated from solution as their salts, the acids themselves can be isolated from solution by ether extraction, often referred to as the etherate method.(2,3,14) When excess ether is added to a very strongly acidi®ed solution of the aforementioned reactants, three layers form, an upper

SYNTHESIS

7

ether layer, an aqueous layer, and a heavy oily etherate.(3) The lowest layer is shaken with excess ether to remove any residual aqueous solution and separated again. Water is added to decompose the etherate and the ether is removed. Crystallization occurs as the water is evaporated from the aqueous solution. Because evaporation by heating causes reduction of the acid, removal of the ether by purging with a clean inert gas is recommended.(6) 12-Tungstoarsenic acid (H3AsW12O40, abbreviated as HAsW) has been prepared from arsenic pentoxide, sodium hydroxide, sodium tungstate, and hydrochloric acid and characterized by direct current plasma emission, infrared, X-ray powder diffraction, differential thermal analysis, nitrogen adsorption± desorption, and 75As NMR.(15,16) The preparation, modi®ed from that of Malik et al.,(17) involved the dissolution of arsenic pentoxide in aqueous sodium hydroxide solution which was added to a sodium tungstate solution of appropriate concentration and volume, followed by the addition of HCl to pH 0.5. The etherate extraction method was used to separate HAsW. A variety of tungsten-containing anions of Keggin structure can be prepared by similar methods, including those with silicon, boron, cobalt, iron, and carbon as central atoms. The ammonium salts of various heteropoly anions were synthesized by Illingworth and Keggin(18) and of that with boron as the central atom by Signer and Gros.(19) Brown has prepared a number of tungsten-containing acids and salts(20) including chromium,(21) manganese,(22) and cobalt.(23)

2.2.2. Molybdenum-Containing Anions 12-Molybdophosphoric acid may be prepared in a similar manner to that for 12-tungstophosphoric acid using, for example, sodium molybdate and phosphate(2): 12MoO4 2ÿ ‡ PO4 3ÿ ‡ 27H‡ ! H3 PMo12 O40 ‡ 12H2 O Ether extraction followed by recrystallization from water at room temperature may again be employed to obtain a yellow crystalline solid, H3PMo12O40
29H2O. A ®ne precipitate of decomposition products may form during the initial recrystallization and should be removed by ®ltration.(2)

2.2.3. Other Heteropoly Acids Methods for the preparation of heteropoly acids with various central singleelement and peripheral metal atoms can be found in Refs. 2 and 3.

8

CHAPTER 2

2.2.4. Mixed Addenda Heteropoly anions with so-called mixed addenda, that is, peripheral metal atoms of two or more elements, have been prepared. Although the structure of these will be discussed in more detail later, such species are capable of existing in various positional isomeric forms, depending on the location of the peripheral metal atoms or addenda. Kokorin prepared HPMo10V2,(24) and Courtin, Chauveau, and Souchay synthesized HPMo11V, HPMo10V2, and HPMo9V3.(25) Tsigdinos and Hallada prepared HPMo11V, HPMo10V2, HPMo9V3, and the sodium and ammonium salts of HPMo11V.(26) Although the latter authors made use of the method developed by the former authors, it is important to note that it was found that the acidity speci®ed for extraction into ether (5 N H2SO4) was ``totally insuf®cient to accomplish extraction'' and additional acid was required. The latter report provided detailed characterization data which were missing from the earlier studies. The polyoxometalates substituted by noble metal cations (M ˆ PdII or IrIV) with the formula (Bu4N)n [XW11MO39(OH2)] (X ˆ B, n ˆ 7, X ˆ Si or Ge, n ˆ 6, x ˆ P, n ˆ 5) were prepared by re®lling the vacant site of the lacunary precursor Kn XW11O39. Characterization was obtained from 183W NMR, UV±vis and IR spectra, and cyclic voltammetry.(27) It is of interest to note the synthesis and characterization of the ®rst 12tungstovanadate compound with Keggin structure, [Me4N]7[VW12O40]
15H2O, with tungsten in the peripheral metal positions and vanadium(IV) in the central atom position.(28) Titanium-substituted heteropolytungstates K7[BTiIVW11O40] and V IV K6[PV Ti W10O40] have been synthesized and characterized by elemental analyses and spectroscopic methods.(29)

2.3. VARIOUS SYNTHESES Peripheral metal±oxygen groups have been replaced by other metal±ligand groups. For example, PW12O40 3ÿ has been converted to H2PCoW11O40 5ÿ,(30) CpPTiW11O39 4ÿ,(31,32) CH3PSnW11O39 4ÿ,(32,33) and CpFe(CO)2PSnW11 O39 4ÿ.(34) The electrophilic O-alkylation of PMo12O40 3ÿ and PW12O40 3ÿ has been reported.(35) [C6H13)4N]2CH3OMo12PO39, [(n-C6H13)4N]2C2H5OMo12PO39, [(CH3)3O]3W12PO40, and [(C6H13)4N]2CH3OW12PO39 were synthesized. Cs6W5P2O23, Cs7W10PO36, and Cs7Na2W10PO37 have been prepared.(36) The synthesis and characterization of what is stated to be the ®rst Kegginbased true heteropoly dioxygen (peroxo) anion has been reported.(37) Exposure of a-[Co3‡ W11O39]9ÿ to H2O2 reduces the central atom Co3‡ to Co2‡ and each

SYNTHESIS

9

of the four unshared oxygen atoms surrounding the vacancy is replaced by a peroxide group, forming salts of the tetraperoxide anion b3 [(Co2‡ O4)W11O31(O2)4]10ÿ . The authors note that this type of reaction does not occur with the XW12 Keggin compounds. There has been, and continues to be, signi®cant interest in the use of heteropoly oxometalates in layered structures, at least in part as a result of the size of the anions. While relatively few applications of such layered materials in heterogeneous catalysis have been reported, some reference to these should be made. Layered double hydroxides pillaring by a-[XM12O40]nÿ have been investigated.(38) Complete intercalative ion exchange reaction occurred with aqueous solutions of a-[H2W12O40]6ÿ and a-[SiV3W9O40]7ÿ but no reaction was observed with a-[PW12O40]3ÿ and a-[SiW12O40]4ÿ . Three-dimensional tunnel structures in which the organic templates are sandwiched by two polyanions were obtained with the pyridazinium, pyrimidinium, and pyrazinium salts of HPMo.(39) Thermal decomposition of these compounds begins at room temperature with the release of water molecules. The self-assembly of a-[SiW12O40]4ÿ on Ag(III) surfaces from acidic aqueous solutions has been reported.(40) The spacing on the metal surface Ê ) matches the diameter of the anion. (10.2  0.5 A The incorporation of iron cations in the octahedral position in a Keggin anion has been demonstrated.(41) HPMo trapped in the supercages of Y zeolite was formed from MoO3 and H3PO4 in a slurry mixture of Y zeolite and deionized water. After washing in hot water, HPMo was found to remain in the Y zeolite.(42) Evidence for the formation of HSiMo with MoO3=SiO2 catalysts was obtained from IR and Raman spectra and indirectly from application to methane oxidation.(43) Subsequent work on the use of the same catalysts with a variety of loadings and either O2 or N2O as oxidant in the aforementioned oxidation process provided further evidence.(44) Keggin aluminotungstic species have been shown to be formed in a tungstate solution in which alumina has been suspended and to be deposited on the support during the preparation of WO3=Al2O3 catalysts by equilibrium adsorption.(45) HPW can be immobilized in silica matrices by means of a sol±gel technique that involves the hydrolysis of ethyl orthosilicate.(46) Uniform spherical colloidal particles of zirconium and thorium heteropoly tungstosilicate(47) and of thorium and cesium tungstophosphates(48) have been formed by aging aqueous solutions of salts of the corresponding metals and the appropriate heteropoly acid at temperatures up to 90  C. With the tungstosilicate anion the thorium ion substitutes for four protons whereas with zirconium hydrolyzed cationic zirconium species reacted with the heteropoly anion to produce nonstoichiometric solids. With the tungstophosphate anion the cesium

10

CHAPTER 2

compound was stoichiometric whereas the thorium heteropoly particles contained excess thorium. Hydrothermal syntheses have found recent application in the isolation of mixed-valence polyanion clusters.(49±53) Zubieta and co-workers have prepared a mixed-valence arsenic=molybdenum Keggin anion [H4AsIII 2 AsVMoV 8 MoVI 4 O40]ÿ .(51) A mixed-valence K6Mo3W9PO40
13H2O(52) and an organic salt of [PMo4.27W7.73O40]6ÿ with Mo and W present in both V and VI oxidation states(53) have been synthesized. The ®rst heteropoly compound with terminally bonded ¯uorine, [SiW11O39MF]5ÿ where M is Zr or Hf, has been recently prepared.(54) A new method to prepare silica-supported heteropolyanion catalysts has been proposed.(55) The synthesis of bifunctional rhodium±oxometalate catalysts was reported by Siedle and co-workers in 1987.(56) Iridium polyoxometalates were prepared and shown to activate aromatic C±H bonds.(57) The rhodium and iridium oxometalates were found to be ``chemically microporous,'' with molecules such as acetylene and cyclopentene capable of diffusing into the lattices.(58) The results of further studies by the same authors of [(Ph3P)2IrH2]3PW12O40 (A) and its reactions with organic molecules such as acetone, methanol, cyclooctane, benzene, and toluene were interpreted as the consequence of the dissolution of these latter species in the hydrophobic regions of the solid that are formed by the phenyl rings on the triphenylphosphine ligands.(59) Further reports on compound A and its properties have been published(60±63) and the platinum analogue has also been synthesized.(63) Aluminum 12-tungstophosphate has been reported to be prepared in high purity through electrochemical dissolution of an aluminum anode in aqueous HPW.(64) Aluminum reduces the Keggin anion which is reoxidized by atmospheric oxygen. This is of particular interest as single-crystal X-ray diffraction studies of various heteropoly salts prepared from divalent cations cast some doubt on the appropriateness of the usual method of synthesis when applied to such cations.(65±66)

REFERENCES 1. G. A. Tsigdinos, IEC Prod. Res. Dev. 13, 267 (1974). 2. G. A. Tsigdinos, in: Methodicum Chimicum (F. Korte, editor-in-chief), Vol. 8 (K. Niedenzu and H. Zimmer, eds.), Chapter 32, p. 552, Academic Press, New York (1976). 3. M. T. Pope, Heteropoly Oxometalates, Springer-Verlag, Berlin (1983). 4. M. T. Pope and A. MuÈller, Angew. Chem. Int. Ed. Engl. 30, 34 (1991). 5. L. C. W. Baker and D. C. Glick, Chem. Rev. 98, 3 (1998). 6. S. H. Was® and W. I. Johnson, Synth. React. Inorg. Met. Org. Chem. 29, 697 (1999). 7. R. Kato, A. Kobayashi, and Y. Sasaki, J. Am. Chem. Soc. 102, 6571 (1980).

SYNTHESIS

11

8. H. Wu, J. Biol. Chem. 43, 189 (1920). 9. A. G. Scroggie, J. Am. Chem. Soc. 51, 1057 (1929). 10. J. C. Bailar, in: Inorganic Syntheses, Vol. I (H. S. Booth, ed.), p. 132, McGraw±Hill, New York (1939). 11. E. O. North, in: Inorganic Synthesis, Vol. I (H. S. Booth, ed.), p. 129, McGraw±Hill, New York (1939). 12. H. S. Booth, in: Inorganic Synthesis, Vol. I (H. S. Booth, ed.), p. 127, McGraw±Hill, New York (1939). 13. G. B. McGarvey and J. B. Moffat, J. Mol. Catal. 69, 137 (1991). 14. E. Drechsel, Ber. Chem. Ges. 20, 1452 (1887). 15. G. B. McGarvey and J. B. Moffat, J. Colloid Interface Sci. 125, 51 (1988). 16. G. B. McGarvey and J. B. Moffat, J. Magn. Reson. 88, 305 (1990). 17. W. V. Malik, S. K. Srivastava, and S. Kumar, Talanta 23, 323 (1976). 18. J. W. Illingworth and J. F. Keggin, J. Chem. Soc. 575 (1935). 19. R. Signer and H. Gros, Helv. Chim. Acta 17, 1076 (1934). 20. D. H. Brown and J. A. Mair, J. Chem. Soc. 2597 (1958). 21. D. H. Brown, J. Chem. Soc. 3322 (1962). 22. D. H. Brown, J. Chem. Soc. 4408 (1962). 23. D. H. Brown, Spectrochim. Acta 19, 585 (1963). 24. A. I. Kokorin, J. Gen. Chem. USSR 24, 697 (1954). 25. M. P. Courtin, F. Chauveau, and P. Souchay, Compt. Rend. 258, 1246 (1964). 26. G. A. Tsigdinos and C. J. Hallada, Inorg. Chem. 7, 437 (1968). 27. H. Liu, W. Sun, B. Yue, S. Jin, J. Deng, and G. Xie, Synth. React. Inorg. Met. Org. Chem. 27, 551 (1997). 28. M. I. Khan, S. Cevik, and R. Hayashi, J. Chem. Soc. Dalton Trans. 1651 (1999). 29. R. Muragesan, P. Sami, T. Jeyabalan, and A. ShunMugasundaram, Transition Met. Chem. 23, 583 (1998). 30. A. Komura, M. Yayash, and H. Imanaga, Bull. Chem. Soc. Jpn. 49, 87 (1976). 31. R. K. C. Ho and W. G. Klemperer, J. Am. Chem. Soc. 100, 6772 (1978). 32. W. H. Knoth, J. Am. Chem. Soc. 101, 759 (1979). 33. F. Zonnevijlle and M. T. Pope, J. Am. Chem. Soc. 101, 2731 (1979). 34. W. H. Knoth, J. Am. Chem. Soc. 101, 2211 (1979). 35. W. H. Knoth and R. L. Harlow, J. Am. Chem. Soc. 103, 4265 (1981). 36. W. H. Knoth and R. L. Harlow, J. Am. Chem. Soc. 103, 1865 (1981). 37. J. Server-Carrio, J. Bas-Serra, M. E. Gonzalez-Nunez, A. Garcia-Gastaldi, G. B. Jameson, L. C. W. Baker, and R. Acerete, J. Am. Chem. Soc. 121, 977 (1999). 38. T. Kwon and T. J. Pinnavaia, Chem. Mater. 1, 381 (1989). 39. M. Ugalde, J. M. GutieÂrrez-Zorrilla, P. Vitoria, A. Luque, A. S. J. WeÂry, and P. RomaÂn, Chem. Mater. 9, 2869 (1997). 40. M. G. B. Zhong, W. G. Klemperer, and A. A. Gewirth, J. Am. Chem. Soc. 118, 5812 (1996). 41. A. Oszko, J. Kiss, and I. Kricsi, Phys. Chem. Chem. Phys. 1, 2565 (1999). 42. S. R. Mukai, T. Masuda, I. Ogino, and K. Hashimoto, Appl. Catal. A 165, 219 (1997). 43. S. Kasztelan, E. Payen, and J. B. Moffat, J. Catal. 112, 320 (1988). 44. M. A. Banares, J. L. G. Fierro, and J. B. Moffat, J. Catal. 142, 406 (1993). 45. X. Carrier, J.-B. d'Espinose de la Caillerie, J.-F. Lambert, and M. Che, J. Am. Chem. Soc. 121, 3377 (1999). 46. Y. Izumi, Rev. Chem. Intermed. 24, 461 (1998). 47. A. Koliadima, L. A. Perez-Maqueda, and E. Matijevic, Langmuir 13, 3733 (1997). 48. L. A. Perez-Maqueda and E. Matijeric, Chem. Mater. 10, 1430 (1998). 49. M. I. Khan, Q. Chen, and J. Zubieta, Inorg. Chem. 31, 1556 (1992).

12

CHAPTER 2

50. M. I. Khan, Q. Chen, D. P. Gosborn, H. Hope, S. Parker, and J. Zubieta, J. Chem. Soc. 114, 3241 (1992). 51. M. I. Khan, Q. Chen, and J. Zubieta, Inorg. Chem. 32, 2924 (1993). 52. A. Leclaire, M. M. Borel, J. Chardon, and B. Raveau, Mater. Res. Bull. 30, 1075 (1995). 53. J.-C. P. Gabriel, R. Nagarajan, S. Natarajan, A. K. Cheetham, and C. N. R. Rao, J. Solid State Chem. 129, 257 (1997). 54. H. W. Roesky and R. Siefken, Z. Anorg. Allg. Chem. 624, 171 (1998). 55. J.-M. Tatibouert, C. Montalescot, and K. BruÈckman, Appl. Catal. A 138, L1 (1996). 56. A. R. Siedle, C. G. Markell, P. A. Lyon, K. O. Hodgson, and A. L. Roe, Inorg. Chem. 26, 219 (1987). 57. A. R. Siedle, R. A. Newmark, K. A. Brown-Wensley, R. P. Skarjune, L. C. Hadlad, K. O. Hodgson, and A. L. Roe, Organometallics 7, 2078 (1988). 58. A. R. Siedle, R. A. Newmark, W. B. Gleason, R. P. Skarjune, K. O. Hodgson, A. L. Roe, and V. W. Day, Solid State Ionics 26, 109 (1988). 59. A. R. Siedle, R. A. Newmark, M. R. V. Sahyun, P. A. Lyon, S. L. Hunt, and R. P. Skarjune, J. Am. Chem. Soc. 111, 8346 (1989). 60. A. R. Siedle, New J. Chem. 13, 719 (1989). 61. A. R. Siedle and R. A. Newmark, J. Am. Chem. Soc. 111, 2058 (1989). 62. A. R. Siedle and R. A. Newmark, Organometallics 8, 1442 (1989). 63. A. R. Siedle, R. A. Newmark, P. A. Lyon, S. L. Hunt, and V. W. Day, Inorg. Chim. Acta 259, 241 (1997). 64. A. R. Siedle, T. E. Wood, M. L. Brostrom, D. C. Koskenmaki, B. Montez, and E. Old®eld, J. Am. Chem. Soc. 111, 1665 (1989). 65. G. B. McGarvey and J. B. Moffat, Catal. Lett. 16, 173 (1992). 66. G. B. McGarvey, N. J. Taylor, and J. B. Moffat, J. Mol. Catal. 80, 59 (1993).

3 CHARACTERIZATION

For methods of investigation of heteropolyanions in solution, the reader should consult the monograph by Pope.(1) As the reader will appreciate, investigative techniques employed for homogeneous systems are frequently inapplicable for solids and their surfaces. However, spectroscopic measurements have found useful application for both systems. Almost inevitably, reports on the surface and catalytic properties of the heteropoly oxometalates contain the results of characterization studies. With discussions of the aforementioned reports found elsewhere in this monograph, the present chapter illustrates methods of characterization by reference to selected publications.

3.1. SPECTROSCOPY 3.1.1. Infrared Spectroscopy Infrared spectroscopy may well be the most frequently employed technique for the characterization of heteropoly oxometalates, particularly as a consequence of the identifying ®ve or six peaks of the Keggin structure. It is not surprising, therefore, that many of the publications in which catalytic studies are reported also have made use of infrared spectroscopy for demonstrations of the presence of the Keggin structure. In the following, a selected group of publications that have dealt primarily with infrared spectra are discussed.(2±35) One of the ®rst, if not the ®rst, infrared investigations of the heteropoly oxometalates was reported by Sharpless and Munday who examined the ammonium salts of 12-tungsto- and 12-molybdophosphoric, tungstoboric, tungsto- and molybdosilicic, tungsto- and molybdoarsenic, molybdomanganic and molybdotitanic acids in potassium bromide disks and provided band assignments.(2) 13

14

CHAPTER 3

Brown obtained the ultraviolet, visible and near-infrared spectra of 12tungstochromic(III)(3) and 12-tungstomanganic(IV)(4) acids in aqueous solution as well as a variety of heteropolytungstic acids and their salts in the form of KCl or KI disks for the near-infrared and far-infrared regions, respectively.(5) What appears to be the ®rst detailed investigation of the molybdovanadophosphoric acids (H4PMo11VO40, H5PMo10V2O40, and H6PMo9V3O40) and their sodium and ammonium salts was published by Tsigdinos and Hallada in 1968.(6) The visible and ultraviolet spectra were obtained with aqueous solutions whereas the infrared spectra employed Nujol mulls. Since 1968, many infrared spectra have been published on these materials with those from the French group probably the most frequently cited.(12±14,26±27) 3.1.2. Photoacoustic FTIR Spectroscopy The ®rst applications of photoacoustic (PAS) FTIR spectroscopy to heteropoly oxometalates were reported in 1984.(29±35) The parent 12-tungstophosphoric ‡ salts have been examined at various acid and its Al3‡ , NH‡ 4 , and Na temperatures. The calorimetric mode of detection of PAS is particularly advantageous with strongly scattering and=or opaque materials for which conventional photometric methods are not suitable. With the heteropoly acids whose surface areas are relatively small, necessitating the use of Nujol or KBr with conventional methods, PAS is particularly advantageous because the solid sample may be studied in its pure form. Not surprisingly, however, the technique is not without its disadvantages.(30,34) 3.1.2.1. H3PW12O40 The PAS spectra of 12-tungstophosphoric acid (H3PW12O40, abbreviated as HPW) show broad bands assignable to water (3200 cmÿ1 ) at room temperature and the oxonium ion, H3O‡, at 1710 cmÿ1 as well as a group of ®ve or six bands below 1100 cmÿ1 , characteristic of the Keggin Unit (KU) (PW12O40 3ÿ ) (Fig. 3.1a).(29±35) After heating at 200  C the band at 1710 cmÿ1 vanishes while a weak band attributable to the binding vibration of lattice water is evident at 1640 cmÿ1 (Fig. 3.1b). Weak bands assignable to the corresponding asymmetric and symmetric stretching vibrations appear at 3060 and 2980 cmÿ1 , respectively, with a third band at 3160 cmÿ1 , probably the ®rst overtone of the bending fundamental. The position of the stretching vibrations indicates the presence of hydrogen bonding. The band at 2240 cmÿ1 may originate from crystal water (1±2 molecules KUÿ1 by weight loss). The development of structure in the KU bands at 1080 and 980 cmÿ1 indicates that some distortion of the anion has occurred. The former band is attributed to the triply degenerate asymmetric stretching vibration of the central PO4 tetrahedron and the

CHARACTERIZATION

15

Figure 3.1. PAS-FTIR spectra of 12-tungstophosphoric acid after heating at various temperatures ( C) in vacuo.(29±35)

latter to a stretching vibration involving the central W atom and the isolated terminal O (W±Ot).(2±14) The appearance of two new bands at 1120 and 1069 cmÿ1 is indicative of the complete resolution of the aforementioned degeneracy and implies a lowering of symmetry from Td to C2v .

16

CHAPTER 3

After heating at 350  C the band at 2240 cmÿ1 appears to intensify and then disappears after 450  C, suggesting the complete elimination of lattice water at the higher temperature. The diminution of the background absorption is believed to result from the loss of water and=or protons. Continuous absorption in the infrared has been related to proton mobility in adsorbed layers on solid surfaces. Although the KU bands remain, broadening occurs and absorption below 1000 cmÿ1 is observed, indicative of partial decomposition. 3.1.2.2. (NH4)3PW12O40 The PAS spectrum of ammonium 12-tungstophosphate remains relatively invariant on heating up to 350  C (Fig. 3.2a).(29,31,32) The background absorption with this salt is much weaker than that found for the parent acid. The characteristic bands for NH4 ‡ are 3200 and 1420 cmÿ1 , attributed to the triply degenerate asymmetric stretching (n3 ) and bending (n4 ) fundamentals, respectively. Additional bands, assigned to combinations and overtones, are found at 3280 cmÿ1 (n1 ‡ n5 ), 3220 cmÿ1 (n3 ), 3060 cmÿ1 (n2 ‡ n4 ), and 2820 cmÿ1 (2n4 ), where n1 is the symmetric stretching fundamental, n2 is the doubly degenerate bending fundamental, and n5 is a lattice mode. The absence of structure in the KU bands here is indicative of an undistorted anion with Td symmetry. The assignments, based on those in Ref. 14, are summarized in Table 3.1. The absence of bands at 2240 and 1640 cmÿ1 indicates that little or no lattice water is present. As the ammonium salt is heated (Fig. 3.2b±d), loss of bands characteristic of NH4 ‡ begins at approximately 450  C and a broadening and shifting of the KU bands with the P±O and W±O stretching bands shifting to lower and higher frequencies, respectively. However, the anion structure is partially retained at 500  C. 3.1.2.3. AlPW12O40 In addition to the anion bands, broad bands at 3200, 2240, and 1640 cmÿ1 indicate the presence of residual lattice water most of which disappears by 450  C, although the anion appears stable even above 500  C but not at 690  C (Fig. 3.3).(29,31,32) 3.1.2.4. Na3PW12O40 The PAS spectrum of the sodium salt presents an interesting contrast to that of the parent acid as well as the NH4 ‡ and Al salts (Fig. 3.4).(29,31,32) The strong

CHARACTERIZATION

17

Figure 3.2. PAS-FTIR spectra of (NH4)3PW12O40 after heating at various temperatures ( C) in vacuo.(29,31,32)

sharp bands at 3590, 3537, and 1640 cmÿ1 are assigned to the stretching [asymmetric (n1 ), symmetric (n3 ) and bending (n3 )] fundamentals of water, respectively. Bands in the range 2200 1200 cmÿ1 are believed to result from residual impurities from the preparative reagents. The anion appears to be stable up to 500  C.

18

CHAPTER 3

TABLE 3.1. PAS FTIR Bands for (NH4)3PW12O40(29,31,32) Assignmenta NH4 ‡ (cmÿ1 ) 3200 1420

Triply degenerate asymmetric stretching (n3 ) fundamental Bending fundamental (n4 )

PW12O40 3ÿ (cmÿ1 ) 1083 987 891 815 596

P±O stretching W±Ot stretching Stretching of W±O±W bridges between corner-sharing WO6 octahedra Stretching of W±O±W bridges between edge-sharing octahedra P±O bend

a

After Ref. 14.

3.1.3. Nuclear Magnetic Resonance Prior to the general availability of MAS NMR, much of the work, not surprisingly, was focused on solutions of the heteropoly oxometalates. Pope has a concise summary of the state of the art up to 1983.(1) A later review appeared in 1994.(36) Not unnaturally, workers in heterogeneous catalysis have made signi®cant use of MAS in their investigations. For purposes of illustration, reference will be made here to selected publications that demonstrate the use of MAS NMR for such purposes. Silica-supported HPMo has been investigated with 31P Magic Angle Spinning (MAS) NMR.(37) Bulk HPMo, that is in unsupported form, was found to have a 31P NMR peak at ÿ8 ppm. A silica-supported HPMo catalyst of 23% loading after heating at 200  C for 2 h also had a 31P NMR peak at ÿ8 ppm. Heating of the sample at 350, 450, and 550  C for 16 h each as well as use in a methane oxidation reaction at 570  C for 4 h produced the same peak position. The full widths at half-maximum (FWHM) were 1.0±1.2 Hz with the aforementioned supported samples. After heating at 600 and 730  C each for 16 h the peak position shifted to ÿ5 and FWHM values were recorded as 1.8 and 9.4 Hz, respectively. The values for the bulk HPMo are consistent with the 31P NMR spectra reported earlier.(1) The shift in the 31P peak position and FWHM after heating at 600  C is indicative of the destruction of the anion of HPMo. Samples with loadings between 1.2 and 39%, after heating at 350  C for 2 h showed 31P peak positions of ÿ8 ppm regardless of the loading and FWHM values of 1.0± 1.5 Hz up to 31% increasing to 4.1 Hz for a loading of 39%. Evidently the catalyst stability is not dependent on the loading of the SiO2-supported samples. NMR has been employed to study species of heteropoly oxometalates which exist at room temperature in aqueous solutions of various pH.(38) Further details can be found in the chapter on stability.

CHARACTERIZATION

19

Figure 3.3. PAS-FTIR spectra of AlPW12O40 after heating at various temperatures ( C) in vacuo.(29,31,32)

The 31P and 29Si MAS NMR spin±lattice relaxations have been proposed as a probe for the dispersion of HPMo and HSiMo, respectively, on silica.(39,40) With HPMo=SiO2 the 31P mean relaxation time T 1* remains essentially constant as the loading increases up to approximately 1 mol HPMo per 160 mol of SiO2 (10 Mo wt %) (Fig. 3.5). For further increases in loading up to 3.5 mol HPMo per 160 mol SiO2, T 1* increases linearly. Although 3 mol HPMo per 160 mol SiO2 is

20

CHAPTER 3

Figure 3.4. PAS-FTIR spectra of Na3PW12O40 after heating at various temperatures ( C) in vacuo.(29,31,32)

calculated to correspond to monolayer coverage, the discontinuity in T 1* occurs at a much lower value. The authors argue that there are two forms of HPMo, that anchored on the SiO2 surface and corresponding to macroscopic HPMo, the former exhibiting a short T 1* while the latter has a value (20 s) of T 1* similar to that for unsupported HPMo. The discrepancy between the value of T 1* at the discontinuity and that for the expected monolayer is attributed to the difference between chemical and physical HPMo=support interactions.

CHARACTERIZATION

21

Figure 3.5. Solid-state 31P mean relaxation time T 1* of HPMo=SiO2.(39) Reproduced by permission of The Royal Society of Chemistry.

Palladium and HPW in the presence of H2 have been studied with 1H MAS and broad-line NMR.(41) The catalyst was prepared from an aqueous solution of Pd(NO3)2 and HPW evaporated to dryness at 353 K. Reduction of the solid with H2 increased the intensity of a signal at 9.1 ppm. Heating to 333 K in the presence of H2 resulted in increase of the linewidths of the aforementioned 9.1-ppm signal due to protons and a peak at 19.8 ppm due to hydrogen atoms adsorbed on metallic palladium, the latter of which shifted toward that of acidic protons, suggesting that interconversion occurs between the hydrogen atoms on the palladium and the protons in HPW. The authors conclude that the presence of H2 and palladium metal modi®es the mobility of the protons and hence the acid strength. 51 V NMR, both in solid state and in aqueous solution, has been employed for a number of years to assist in the characterization of heteropoly oxometalates.(42±46) What appears to be the ®rst application of 51V NMR to these materials was introduced by O'Donnell and Pope in 1976.(42) Since that time the technique has found wide usage among inorganic chemists and more recently those interested in the catalytic properties of vanadium-containing Keggin anions.(46) 31P NMR and 51V NMR have been used to obtain information on the nature and localization of vanadium in molybdovanadophosphoric acids before and after use in oxidation processes.(46) Vanadium localized in the secondary structure was found to exchange with the molybdenum atoms of the

22

CHAPTER 3

Keggin anion during the reaction. In the oxidation of butane to maleic anhydride, vanadium ions in the secondary structure are more selective at low conversion, while V ions in the peripheral metal positions of the anion are more selective at higher conversions. 1 H MAS NMR has been applied to examine the variation of chemical shifts with loading of HPW=SiO2 catalysts.(47) A single relatively sharp peak with a chemical shift that increases with loading is observed (Fig. 3.6). A combination of 31P NMR and FTIR has shown that on HPW=TiO2, the acid is present in ®ve forms: a bulk salt phase, two weakly bound intact Keggin species, a strongly bound partially fragmented Keggin unit, a strongly bound highly fragmented and a pure phosphate phase formed from the decomposition of the Keggin structure.(48) Adsorption of the Keggin anion on the TiO2 surface results in an increase in the acidity. The ®rst reports of 75As NMR spectra of heteropoly anions were published in 1990.(49) With potassium hexa¯uoroarsenate as an external standard, the chemical shifts and linewidths for HAsW and HAsMo were 291 ppm, 214 Hz and 337 ppm, 183 Hz, respectively. The 75As chemical shifts for 12-tungstoarsenic and 12-molybdoarsenic acids show a dependence on the nature of the peripheral metal atom that is similar to that shown by the 31P chemical shifts for

Figure 3.6. 1H NMR chemical shift of HPW=SiO2 at various loadings.(47) Reproduced by permission of Baltzer Science Publishing.

CHARACTERIZATION

23

HPW and HPMo. The magnitude of the measured linewidths demonstrates the stabilizing effect that the surrounding M3O13 trimetalate units exert on the tetrahedral arsenate species at the center of the Keggin anion. 3.1.4. Electron Paramagnetic Resonance and Electronic Spectroscopy Although considerable use has been made of electron paramagnetic resonance as well as electronic spectroscopy for the studies of heteropoly oxometalates,(1) only one reference will be noted here to illustrate the nature of the data that can be obtained with these techniques. The coordination geometry around Cu(II) in the tetrabutylammonium (TBA) salts, (TBA)4Hx [XW11CuO39], where X is P or B, was found to be square pyramical, with copper bonded to the ®ve oxygen donor atoms.(50) In contrast, with the potassium salt of the anion [XW11Cu(H2O)O39]nÿ , a tetragonally elongated pseudo-octahedral geometry was deduced. It should be noted that both solid- and solution-phase EPR spectra were obtained in this work. 3.1.5. X-ray Photoelectron Spectroscopy (XPS) Relatively few reports of the application of this technique have appeared, although this will undoubtedly change in the near future. The dispersion of HPMo=SiO2 has been estimated for various loadings.(37) A graph of the Mo3d=Si2p XPS intensity ratio versus the loading shows three identi®able parts. The ®rst, a linear section that extends up to 0.04 KU nmÿ1 (10 wt % HPMo) is approximately coincident with the line calculated for a monolayer dispersion and probably corresponds to single or small aggregates of the anion dispersed over the surface. For loadings higher than 10 wt % HPMo up to 23 wt % (0.10 KU nmÿ2 ) the measured Mo3d=Si2p intensity falls below the line for monolayer dispersion, indicating the appearance of larger aggregates and for higher loadings large aggregates are formed. After use in methane oxidation the samples have similar XPS intensity curves, showing that no drastic changes in the dispersion have occurred. XPS has also been employed in an attempt to demonstrate that a small fraction of Al3‡ can be substituted by Fe3‡ in the octahedral positions of the Keggin anion.(51) XPS has been applied to bulk HPMo and that supported on silica and on zirconia±silica.(52) The surface Mo=Si ratio of HPMo=SiO2 increased slightly on heating from 383 K to 623 K, indicating that the Keggin anions remained essentially intact at these temperatures. However, after heating at 723 K the Mo=Si ratio increased signi®cantly, which the authors interpreted as resulting from the decomposition of the Keggin anion into a mixture of molybdenum and phosphorus oxides. The XPS Mo=Si ratio for HPMo=ZrO2±SiO2 at 383 K is

24

CHAPTER 3

signi®cantly higher than that for HPMo=SiO2, which is interpreted as resulting from a higher perturbation of the Keggin anion by the mixed support. High values of the XPS Mo=Zr ratios suggest a strong interaction between Mo and Zr. 3.1.6. Scanning Tunneling Microscopy (STM) and Tunneling Spectroscopy (TS) Ordered arrays of heteropoly acids and isopolyacids have been observed on highly oriented pyrolytic graphite (HOPG) by STM and TS.(53,54) Ordered twodimensional monolayer arrays were found with periodicities consistent with the molecular dimensions of the molecules. The periodicities of the arrays of Ê H3PMo12O40, H4PMo11VO40, and H8PMo10VCuO40 were approximately 11 A Ê .(55) On adsorption of whereas those of the K‡ and Cs‡ salts were 12±14 A pyridine the periodicity of H3PMo12O40 on HOPG increased by approximately Ê , the size of the pyridine molecule, to 16.5 A Ê and the array symmetry was 6A approximately hexagonal, indicating that the pyridinium ions were located at the interstitial positions in the array.(56) Cation-exchanged heteropoly acids formed two-dimensional ordered arrays on HOPG, whose periodicity was dependent on the nature of the cation.(57±59) Cesium substitution in HPMo results in a monotonic increase in the reduction temperature and the negative differential resistance peak shifts to larger negative voltages whereas substitution of copper has the opposite effect.(57±59) The aforementioned work has recently been extended to three additional structures, namely, those of Finke±Droege, Wells± Dawson, and Pope±Jeannin±Preyssler.(60) STM and TS have also been employed to study pyridine adsorption with well-ordered monolayers of H3PMo12ÿx WxO40 (x ˆ 0; 3; 6; 9; 12) on highly oriented pyrolytic graphite.(61) Exposure to pyridine increased the lattice constants of the two-dimensional heteropoly acid arrays by Ê as well as shifting the negative differential resistance peak approximately 6 A voltages of the acids to less negative values in the TS measurements. The earlier work on the adsorption of pyridine on HPW as studied by PAS FTIR showed that at room temperature the dimer ion salt (py2H)3PW12O40 forms, while above 100  C, pyridine is desorbed to generate the pyridinium salt, (pyH)3PW12O40.(30) Sorption of pyridine was found to be less facile than that observed with NH3 on HPW. Further discussion can be found in the section on acidity. 3.2. ELECTROCHEMICAL METHODS Polarography has been applied to heteropoly oxometalates in solution for the identi®cation of species as well as other related purposes.(1) For obvious reasons relatively little use has been made of this technique by those interested in heterogeneous catalysis. Recently, polarographic measurements on aqueous

CHARACTERIZATION

25

HCl=dioxan solutions of HPMo have been employed to show that because the wave heights decrease continuously from 280  C up to 380  C and collapse at 400  C, the PMo12O40 3ÿ anion apparently begins to decompose at 280  C.(62) The ®rst successful voltammetric studies of vanadium-containing polyanions were reported by Smith and Pope in 1973.(63) Cyclic voltammetry has been applied to the study of transition metalÿ substituted heteropoly anions [XW11Mm‡ (H2O)O39]…n‡4ÿm† [(X ˆ P, Si, B; M ˆ transition metal, Cr(III), Fe(III), Co(II, III), and Ni(II)] with K‡ as the ÿ cation and for comparison, the [XW11O39]…n‡4† anions.(62,64) In all cases the aisomers are considered. The transition metal-substituted heteropolytungstates were found to be more dif®cult to reduce than the corresponding ÿ [XW11O39]…n‡4† species. 3.3. ELEMENTAL ANALYSIS The most commonly cited methods for the determination of the elemental composition of heteropoly oxometalates, particularly where these are found in liquid solution, have undoubtedly been atomic absorption and what many authors have described, without elaboration, as ``chemical analysis,'' probably meaning that provided by a commercial laboratory. The reader is again referred to Pope's book for methods of characterization of heteropoly oxometalates in solution(1) as well as other sections of this chapter. 3.3.1. HPLC Relatively little work has been reported on the application of chromatographic methods, probably as a consequence of the large size and complex nature of the heteropoly anions. A reversed-phase liquid chromatographic technique that provides separations of highly charged ions such as PW11VO40 4ÿ, PW10V2O40 5ÿ, PW9V3O40 6ÿ, and P2W18Zn4(OH2)2O68 10ÿ was reported in 1989.(65) The method uses ion-interaction reagents to promote retention and the citrate ion as a competing ion to reduce tailing. Separations within 15 min were shown to be possible for a mixture of ions with charges from 4ÿ to 10ÿ . 3.3.2. Ion Chromatography An ion chromatographic method has been developed that permits the measurement of the concentrations of heteropoly anions and the determination of their elemental compositions.(66) The method depends on the decomposition of the anions into their simpler constituent ions by addition of an aqueous solution of lithium hydroxide and the subsequent separation and analyses of these ions.

26

CHAPTER 3

The stabilities of HPMo, HPW, HSiMo, and HSiW in aqueous solution have been studied at various values of pH by use of the ion chromatographic method.(67) Further discussion of these results can be found in the chapter on stability. 3.3.3. EDTA Titration Although EDTA has been widely employed in the titrimetric determination of phosphorus and tungsten, the procedure as previously employed was not suitable for heteropoly acids. Modi®cations of the procedure have been shown to produce satisfactory results for the latter materials.(68) 3.4. OTHER TECHNIQUES A number of additional characterization methods that do not ®t under the aforementioned categories are also available. Because the nature of the results obtained from most, if not all, of these is discussed elsewhere in this volume, here only the following listing will be given: (1) physisorptionÐprobably most employed for measurement of surface areas and pore structures; (2) chemisorptionÐfor surface site identi®cation and strength of binding of chemisorbed species; (3) temperatureÐprogrammed desorption, exchange, reduction±oxidation, reaction; and (4) ion exchange. REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19.

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28

CHAPTER 3

61. M. S. Kaba, M. A. Barteau, W. Y. Lee, and I. K. Song, Appl. Catal. A 194±195, 129 (2000). 62. C. Rocchiccioli-Deltcheff, A. Aouissi, M. M. Bettahar, S. Launay, and M. Fournier, J. Catal. 164, 16 (1996). 63. D. P. Smith and M. T. Pope, Inorg. Chem. 12, 331 (1973). 64. F. A. R. S. Couto, A. M. V. Cavaleiro, J. D. Pedrosa de Jesus, and J. E. J. Simao, Inorg. Chim. Acta 281, 225 (1998). 65. A. D. Kirk and R. G. Finke, Inorg. Chem. 28, 792 (1989). 66. A. JuÈrgensen, G. B. McGarvey, and J. B. Moffat, J. Chromatogr. 602, 173 (1992). 67. A. JuÈrgensen and J. B. Moffat, Catal. Lett. 34, 237 (1995). 68. H. Hayashi and J. B. Moffat, Talanta 29, 943 (1982).

4 STRUCTURE AND BULK PROPERTIES

4.1. ANIONS In view of the focus of this monograph, the discussion will be restricted to species of interest in catalytic studies and particularly heterogeneous catalysis. Thus, the so-called Keggin anions will receive special attention. For a complete description of the structures of both heteropoly and isopolyanions, at least as known up to 1983, the reader is referred to the book by Pope.(1) The ®rst X-ray powder diffraction investigation of HPW
5H2O was reported by Keggin in 1934.(2) The structures of the cesium salts of a variety of the heteropoly acids as determined from powder XRD were published a year later but the author was primarily concerned with the location and numbers of cations.(3) The ®rst single-crystal studies were not reported until 40 years later at which time a number of papers appeared within one or two years.(5±10) Although several isomers of the anion ®rst investigated by Keggin are now known,(1) that which has been most frequently employed in heterogeneous catalysis is referred to as having the a form with Td symmetry (Fig. 4.1). A central atom is bonded to four oxygen atoms arranged tetrahedrally. Surrounding the central tetrahedron are 12 octahedra with a peripheral metal element at each of their approximate centers, the latter in four groups each of which contains three edge-shared octahedra. The groups are connected by shared corners to each other and to the central tetrahedron (Fig. 4.2). The structures of HPW
29H2O,(4) of triclinic NaH2PW
xH2O (x ˆ 12± (5,6) and of triclinic HPMo
(13±14)H2O(6) have been determined. The 14), structure of HPMo
(29±31)H2O, believed to be tetragonal, has also been reported,(7,8) although this compound was later shown to be cubic and isomorphous with HPW
29H2O.(9) Single-crystal X-ray and neutron diffraction data for HPW
nH2O were ®rst reported in 1974 and published in 1975(9) and 1977.(10) This probably remains as 29

30

CHAPTER 4

Figure 4.1. Anion of Keggin structure. Large circles: central atom and peripheral metal atoms; small circles: oxygen atoms. Reproduced by permission of IUCr.

the most complete structure determination on the heteropoly oxometalates available to data because all hydrogen atom positions were included. It should be noted that what was considered to be the pentahydrate by Keggin(2) was found to be the hexahydrate.(10) It is remarkable that, considering the relative simplicity of the equipment available to Keggin(2) in 1934, the later single-crystal work showed that the structure of the PW12O40 3ÿ anion was semiquantitatively that found 40 years earlier.(10) At least from a historical viewpoint, it is of interest to note that Clark and Hall(11) describe the anion as a clathrate structure in which 12MoO5 pyramids share edges to form Mo12O36 that encloses a guest PO4 3ÿ ion. However, Brown and co-workers note that the space within the Mo12O36 or the W12O36 unit is not suf®ciently large to contain a PO4 3ÿ ion as a guest.(10) The guest=host discussion concerning the possibility that the phosphate of the heteropoly anion is a guest in the remainder of the anion as host has reemerged a decade or so later.(12±16) Pope

STRUCTURE AND BULK PROPERTIES

31

Figure 4.2. Keggin anion structure illustrating the trimetallic groups.(10) Reproduced by permission of IUCr.

and MuÈller(16) note that the polyoxometalate structures can be considered as pseudoclathrates, an example of which is [PMo12O40]3ÿ with [PO4]3ÿ in a neutral shell (Mo12O36) of edge- and vertex-shared MoO5 square pyramids. However, there is no evidence that the Keggin anions are capable of accepting guests as usually found with conventional clathrates.(16) Jeannin et al.(17) note that the Keggin anions can be viewed from two perspectives, focusing on the central atom and considering it as ligated by 12 octahedra or, in contrast, imagining a cage formed by 12 octahedra that is then ®lled up by the central atom. These authors argue that the second viewpoint is preferred because the cage of 12 octahedra may be constructed in the absence of a central atom but with two nonexchangeable protons in its stead, with the resulting stoichiometry given by [H2W12O40]6ÿ . In his 1983 book,(1) Pope noted that the central atom (he prefers to refer to these as heteroatoms) could be one of more than 65 elements of the periodic table while a smaller number have been used as peripheral metal atoms. Pope and MuÈller in their introduction to a collection of papers note that the predominant species of polyoxometalates are polyoxoanions of molybdenum(VI) and tungsten(VI).(16) These authors list some general principles controlling these structures. A number of these are arrangements of edge- and vertex-sharing MO6 octahedra, each with one or two unshared vertices that are terminal oxygens.(18,19) The periphery of the anions contain weakly bonded or nonbasic multiply bonded oxygen atoms which reduce the probability of polymerization of the anions.

32

CHAPTER 4

TABLE 4.1. Known Heteroatoms in Heteropoly Anions as Recorded by Pope(1) H Lia Na K Rb Cs

Be Mg Ca Sr Ba

Ti Zr Hf

Y La Ce Th

a

V Nb Ta Pr

Cr Mo W

Nd U

Mn Re Np

Fe Ru Os

Sm Pu

Co Rh

Ni

Cu Ag

Pt

Eu Gd Am Cm

Tb

Zn

B Al Ga In Tl Ho

Cf

C Si Ge Sn Pb Er

P As Sb Bi

S Se Te

I

Yb

Elements shown in italic type have been observed only as secondary heteroatoms, see Ref. 1.

Elements that have served as central or peripheral metal atoms as known up to the date of Pope's book are listed in Table 4.1. Bond lengths and internuclear separations for some Keggin anions are summarized in Table 4.2. Pope notes that most heteropoly anions with tungsten as the peripheral metal element have the aKeggin structure while a much smaller number of these with molybdenum in the peripheral position possess the Keggin structure. Table 4.3 provides a summary of these. Delgado, Dress, MuÈller and Pope (DDMP)(27) remark that the favorable combination of ionic radius and charge together with the accessibility of empty dorbitals for metal±oxygen p-bonding contributes to the ease with which molybdenum(VI) and tungsten(VI) form polyoxoanions. Others with similar properties include vanadium(V), niobium(V) and tantalum(V). DDMP also list hexavalent Tc, Re, Ru, Os, pentavalent Cr, Mo, W, Tc, Re, and tetravalent Ti, V, Cr, Mo, W as possible cluster formers. Ê )a of a-Keggin Anions, [XM12O40]nÿ (1) TABLE 4.2. Mean Dimensions (A X

M

X±Oa

M±Oa

M±Ot

P P Si

W Mo W

Si Ge H2 Co2‡

Mo Mo W W, COd

1.53 1.54 1.63 1.64 1.62 1.73 Ð 1.92

2.44 2.43 2.38 2.35 2.35 2.29 2.26 2.14

1.70 1.66 1.68 1.71 1.67 1.69 1.70 1.71

a

M


M 3.41b 3.41 3.42 3.38 3.36 3.35 3.32 3.25

Averaged to Td symmetry. Within M3 triplet, edge-shared MO6 octahedra. Between M3 triplets, corner-shared MO6 octahedra. d Crystallographically disordered; each M site contains on average 11=12 W ‡ 1=12 Co. b c

3.70c 3.70 3.68 3.68 3.70 3.74 3.69 3.73

Refs. 10 6 20 21 22, 23 24 25 26

STRUCTURE AND BULK PROPERTIES

33

TABLE 4.3. Heteropoly Molybdates and Tungstates with the a-Keggin Structure(1) [XW12O40]nÿ

[XMo12O40]nÿ

X ˆ H, H2, B, Al, Ga(III),a Si, Ge(IV), P(V), As(V), V(V),b Cr(III), Mn(IV),c Fe(III), Co(III), Co(II), Cu(II), Cu(I), Zn, Se(IV),c Te(IV),c Sb(III),c Bi(III)c X ˆ Si, Ge(IV), P(V), As(V), V(V),b Ti(IV),c Zr(IV),c In(III),c H2 d , Mod

a

Known as a lacunary (GaW11) anion only. In mixed addenda (V ‡ W, V ‡ Mo) anions only. Con®rmation desirable. d Existence questionable, see Ref. 1. b c

DDMP state the two general principles for the formation of polyoxometalate structures, with the most important ®rst. Each metal atom is contained in an MOn coordination polyhedron (CP), an octahedron in the case of anions with Keggin structure. The metal atoms are not centered in the CP but are shifted, due to MO p-bonding, toward the vertices at the periphery of the anion. MOn polyhedra are linked in various ways to form the structures thereby producing different types of faces on the surfaces. Although infrequently employed in heterogeneous catalysis, at least in part as a result of their instabilities, the b, g, d and E isomers of the Keggin anion are of some interest particularly in view of some recent studies concerning the relative stabilities of the a and b structures.(28) The b-Keggin structure is identical qualitatively to the a-Keggin unit but with one of the three edge-shared M3O13 triplets of the latter rotated by 60 .(29) The internuclear separations between two W atoms in the b form are shorter than those in the a structure and the W±O±W angles are smaller, which Pope suggests may account for the lower stability of the former structure.(1) Consequently, coulombic repulsions between peripheral metal atoms in the edge sharing should be more substantial than those in the cornersharing structures, possibly accounting for the relative stabilities of the a and b forms.(30) However, Pope has suggested, in view of the contentions of Kepert,(30) that the a and b structures should have similar stabilities.(31) In contradiction, because b isomers rearrange to form a isomers, the latter isomer is expected to be the more stable. Weinstock and co-workers have recently shown that the a and b forms of (AlIIIW12O40)5ÿ can exist in equilibrium with respect to the interconversion process and have reported that at 200  C the a isomer of the acid form of the aforementioned anion is more stable than the b form by 2:1  0:5 kcal molÿ1 under the conditions employed in their experiments.(28) Although both a and b isomers of [PMo12O40]3ÿ have been synthesized,(32) the b-[PW12O40]3ÿ has only recently been prepared.(33)

34

CHAPTER 4

Heteropoly anions of the type Xx‡ Zz‡ W4 O40 Hn …14-x-z-n†ÿ , in which one tungsten atom of the anion Xx‡ W12O40 …8-x†ÿ is replaced by an atom Z, were ®rst reported by Baker and co-workers.(34) Weakley has prepared salts of heteropoly anions containing two different heteroatoms.(35) Knoth and co-workers have prepared halometal derivatives of W12PO40 3ÿ.(36) Pope and co-workers have reported on the attachment of organic groups to heteropoly anions,(37) Knoth has prepared organic derivatives(38) and metal±metal bonded derivatives(39) of PW12O40 3ÿ, SiW12O40 4ÿ, and SiMo12O40 4ÿ, and Knoth and Harlow have reported the O-alkylation of PW12O40 3ÿ and PMo12O40 3ÿ.(40) Was® and co-workers have prepared (NH4)7[Co3‡ ZnW11O40H2]
20H2O with a Keggin structure in which the Co(III) and Zn2‡ ions are in the peripheral metal and central atom positions.(41) Pope and co-workers have reported the synthesis of heteropoly anions of BiIII and SbIII of stoichiometry [Bi2W22O74(OH)2]12ÿ and [X2W20M2O70(H2O)6]…14ÿ2n†ÿ .(42) 4.2. CATIONS AND CRYSTALLOGRAPHIC STRUCTURE The work of Brown and co-workers has shown that the hexahydrate of HPW has two water molecules hydrogen-bonded to each proton through their oxygen atoms in an approximately planar orientation with a twist from coplanarity between the two H2O entities of 8 (Fig. 4.3).(10) Brown's results show that in CsPW, each cesium ion replaces one H5O2 ‡ ion so that the structurally similar salts of monovalent cations with the heteropoly anions are anhydrous.(10) The acid HPW has a cubic Pn3m structure with the Ê. lattice parameter (a) equal to 12.1506 A Brown and co-workers contend that the ``extra'' protons found in HSiW, for example, must be chemically bonded to and randomly distributed on the 12 terminal oxygen atoms of each anion.(10) The geometry of the H5O2 ‡ ion in (H5O2 ‡ )3(PW12O40 3ÿ ) has been studied by inelastic neutron scattering vibrational spectroscopy at 5 K.(43) A comparison of the observed and calculated spectra shows more bands observed than predicted in the 1200 to 1600 cmÿ1 range which may be due to the presence of pyramidal H2O


H components, frozen in at the lower temperatures of the measurement, but which undergo dynamic inversion at ambient temperatures. A combination of quasielastic neutron scattering and NMR techniques applied to HPW
14H2O identi®es 180 ¯ips of a portion of the water molecules and diffusion of H‡ by chemical exchange (Grotthus mechanism).(44) Because purely vibrational motions do not result in either appreciable line narrowing in broad-line NMR spectra or quasielastic broadening in neutron scattering studies, detectable motions must be rotations and=or self-diffusion.(45)

STRUCTURE AND BULK PROPERTIES

35

Figure 4.3. Anion±cation con®guration in H3PW12O40 showing the dioxonium ion (H5O2 ‡ ) con®guration.(10) Reproduced by permission of IUCr.

Considerable efforts by a number of workers have been applied in an attempt to establish vibrational assignments for the H5O2 ‡ ion.(43,46±51) Because the observed vibrational frequencies can be reproduced with a variety of equally plausible force ®elds, the unambiguous assignment of the former is dif®cult.(51) A combination of IR, Raman, inelastic neutron scattering, and thermal parameters has provided a set of vibrational assignments for HSiW
6H2O and HPW
6H2O which appears to be more probable than the others,(51) and is in agreement with an earlier proposal.(52) The HSiW structure is assumed to contain a H5O2 ‡ ion crystallographically similar to that found in HPW but with an extra hydrogen ion

36

CHAPTER 4

bonded to the Keggin anion which can be represented stoichiometrically as [HSiW12O40]3ÿ . The phase transformations of HPW
29H2O have been extensively studied by application of a variety of spectroscopic and thermal techniques.(53) The 21H2O and 14H2O hydrates of HPW form at 30 and 40  C, respectively, and the phase transformation from 14H2O to 6H2O was fast and complete at approximately 60  C. The 6H2O hydrate remained stable up to 170  C but lost 5.8±6.0 molecules of water in two steps between 170 and 240  C at which temperatures a small decrease in the unit-cell dimensions is evident (Fig. 4.4). The anhydrous phase is stable up to approximately 410  C at which temperature a water molecule, formed from protons and anionic oxygen, is lost. These latter observations are consistent with the results from temperatureprogrammed desorption described elsewhere in this volume.(54) The cubic symmetry and crystal forms of the anhydrous acid are preserved at 400  C. At 450  C the anhydrous phase is converted to a phase identi®ed as PW12O38 which retained cubic symmetry and was stable up to 550  C although showing a substantial decrease in unit-cell dimensions (Fig. 4.4). The Keggin structure was retained in spite of the loss of one molecule of water from the anhydrous phase between 410 and 440  C and indeed to near 550  C. Table 4.4 summarizes the ®ndings of these authors.(53) Raman and IR spectroscopy have also been applied to provide further information on the structural rearrangements and proton species.(53) The dioxonium (H5O2 ‡ ) and oxonium (H3O‡ ) ions are shown to exist in equilibrium in

Figure 4.4. Unit-cell dimensions of the HPW
6H2O phase from 42 to 500  C.(53)

STRUCTURE AND BULK PROPERTIES

37

TABLE 4.4. Phase Transformations of HPW
29H2O(53) Compositiona

Transformation temperature ( C)

20 21 14=6 6 0 PW12O38 Bronze a

28±31 35±42 60 175±230 410±440 580±620 620±1150

Number of water molecules.

the 6H2O hydrate. Evidence for the substitution of terminal oxygen atoms by Ê was found. H2O with a long W±OH2 bond of 2.3 A (55) by the same authors employing inelastic neutron scatterFurther studies ing applied to the parent acid as well as the sodium salts Na2HPW and Na3PW produced conclusions that appear to be consistent with those from earlier photoacoustic FTIR spectroscopic studies of HPW and various of its salts.(56) The latter is discussed in more detail elsewhere in this volume. The phosphorus doped molybdenum oxide compound (P0.18Mo4.00O12.66), isostructural with molybdite, was synthesized in the process of thermally induced phase transformations of HPMo.(57) Colomban and Tomkinson point out that HPW with 29 or 21 water molecules exhibits superprotonic conductivity whereas the conductivity of HPW with 6 down to 0.5 water molecules is only modest (10ÿ4 S cmÿ1 ).(58) After dehydration of HPW at 300  C the inelastic neutron scattering spectrum shows that the protons should not be mobile although they are not covalently bonded to the oxygen atoms. 1 H NMR studies of a solid catalyst prepared from aqueous solutions of Pd(NO3)2 and HPW showed that the addition of H2 results in the formation of acidic protons as well as hydrogen atoms adsorbed on metallic palladium.(59) On heating at or above 60  C the hydrogen atoms are converted to protons whose mobility exceeds that of those in the parent heteropoly acids. 17 O MAS NMR spectra show that the terminal oxygen atoms of HPW are the predominant protonation sites.(60,61) 4.3. LANGMUIR FILMS The formation and properties of monolayer ®lms on liquid substrates have been studied for over 60 years since Langmuir ®rst reported the results for stearic

38

CHAPTER 4

acid on an aqueous HCl solution. Since that time the results of a wide variety of experiments have been reported including the classical surface pressure=area and surface potential measurements and techniques for the removal of the monolayers from the liquid systems and their loading, often as multilayers, on solid substrates. More recently, multilayers with a speci®c organic=inorganic structure have been formed by adsorption of ions on so-called Langmuir ®lms in what is referred to as self-assembled monolayers. The organic self-assembled structure which is formed ®rst acts as a template for the controlled adsorption of the inorganic ions. A Langmuir monolayer of eicosylamine C20H41NH2 was formed in the classical manner by adding droplets of a dilute chloroform solution to ultrapure water in a Langmuir trough.(62) HPW was dissolved in the substrate before the monolayer was formed. X-ray re¯ectivity, surface pressure, and surface potential measurements were employed to study the ®lms. The authors state that no adsorption occurs at basic pH. However, it would appear unlikely that the Keggin anion would be structurally intact under these conditions. At acidic pH the heteropoly anions are adsorbed onto the charged amine lead groups as either a monolayer or a bilayer. The quantity of HPW adsorbed and the thickness of the inorganic layer can be controlled by changing the surface density of the amine. A monolayer of dioctadecylammonium cations (obtained as the bromide) has been used as the template for the adsorption of the anions of H3PW12O40, H4SiW12O40, K5HCoW12O40, and K5BW12O40.(63) The ®lms were transferred by the vertical lifting method onto optically polished calcium ¯uoride or zinc selenide for infrared measurements or onto glass for low-angle X-ray experiments. With all of the acids, regardless of the charge and the central atom, the anions were able to be organized in monolayers. Monolayers of dimethyldioctadecyl-ammonium (as the bromide) have also been constructed as templates for adsorption of the heteropoly anions with results similar to the aforementioned.(64) Brewster angle microscopy has been employed to study the adsorption of the heteropoly anions in the Langmuir ®lms.(65) The anions were found to modify the morphology of the ®lm so that a coexistence of three phases (gaseous, liquid expanded, and liquid condensed) was observed at approximately 20  C. A dense self-assembled monolayer of 1,12-diaminododecane (DD) was formed on an indium tin oxide (ITO) surface followed by the adsorption of HPMo.(66) In the absence of DD the surface concentration of HPMo was calculated as 2:0  10ÿ11 mol cmÿ2 . In contrast, after DD=ITO was dipped into an acidi®ed solution of HPMo the surface concentration of the latter increased to 2:0  10ÿ10 mol cmÿ2 which these authors conclude to be indicative of complete monolayer coverage of HPMo on DD=ITO. Although the aforementioned techniques for the formation of ®lms of heteropoly oxometalates have not yet been employed in catalysis, the potentialities are clearly evident.

STRUCTURE AND BULK PROPERTIES

39

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40

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42. I. Loose, E. Drostl, M. BoÈsing, H. Pohlman, M. H. Dickman, C. Rosu, M. T. Pope, and B. Krebs, Inorg. Chem. 38, 2688 (1999). 43. G. J. Kearley, H. A. Pressman, and R. C. T. Slade, J. Chem. Soc. Chem. Commun. 1801 (1986). 44. R. C. T. Slade, I. M. Thompson, R. C. Ward, and C. Poinsignon, J. Chem. Soc. Chem. Commun. 726 (1987). 45. H. A. Pressman and R. C. T. Slade, Chem. Phys. Lett. 151, 354 (1988). 46. R. D. Gillard and G. Wilkinson, J. Chem. Soc. 1640 (1964). 47. E. Chemouni, M. Fournier, J. Roziere, and J. Potier, J. Chem. Phys. 67, 517 (1970). 48. A. C. Pavia and P. A. Giguere, J. Chem. Phys. 52, 3551 (1970). 49. J. B. Bates and L. M. Toth, J. Chem. Phys. 61, 129 (1974). 50. G. J. Kearley, A. N. Fitch, and B. E. F. Fender, J. Mol. Struct. 125, 229 (1984). 51. G. J. Kearley, R. P. White, C. Forano, and R. C. T. Slade, Spectrochim. Acta 46A, 419 (1990). 52. D. J. Jones, J. Penfold, J. Tomkinson, and J. Roziere, J. Mol. Struct. 195, 283 (1989). 53. U. B. Mioc, R. Z. Dimitrijevic, M. Davidovic, Z. Nedic, M. M. Mitrovic, and P. Colomban, J. Mater. Sci. 29, 3705 (1994). 54. B. K. Hodnett and J. B. Moffat, J. Catal. 88, 253 (1984). 55. U. B. Mioc, P. Colomban, M. Davidovic, and J. Tomkinson, J. Mol. Struct. 326, 99 (1994). 56. J. G. High®eld and J. B. Moffat, J. Catal. 88, 177 (1984). 57. R. Z. Dimitrijevic, P. Colomban, U. B. Mioc, Z. Nedic, M. R. Todorovic, N. Tjapkin, and M. Davidovic, Solid State Ionics 77, 250 (1995). 58. P. Colomban and J. Tomkinson, Solid State Ionics 97, 123 (1997). 59. T. Baba, Y. Hasada, M. Nomura, Y. Ohno, and Y. Ono, J. Mol. Catal. A 114, 247 (1996). 60. I. V. Kozhevnikov, A. Sinnema, R. J. J. Jansen, and H. van Bekkum, Catal. Lett. 27, 187 (1994). 61. I. V. Kozhevnikov, A. Sinnema, and H. van Bekkum, Catal. Lett. 34, 213 (1995). 62. N. Cuvillier, M. Bonnier, F. Rondelez, D. Paranjape, M. Sastry, and P. Ganguly, Prog. Colloid Polym. Sci. 105, 118 (1997). 63. M. Clemente-LeoÂn, C. Mingotaud, B. Agricole, C. J. GoÂmez-Garcia, E. Coronado, and P. DelhaeÁs, Angew. Chem. Int. Ed. Engl. 36, 1114 (1997). 64. M. Clemente-LoÂen, B. Agricole, C. Mingotaud, C. J. GoÂmez-Garcia, E. Coronado, and P. DelhaeÁs, Langmuir 13, 2340 (1997). 65. N. Cuvillier, R. Bernon, J.-C. Doux, P. Merzeau, C. Mingotaud, and P. DelhaeÁs, Langmuir 14, 5573 (1998). 66. S.-Y. Oh, Y.-J. Yun, D.-Y. Kim, and S.-H. Han, Langmuir 15, 4690 (1999).

5 STABILITY

5.1. THERMAL STABILITY A large number of publications have reported studies of the thermal stability of heteropoly oxometalates as part of investigations of other aspects of these materials. Although these reports may have been discussed elsewhere in this monograph where relevant, they have also been included here. The present discussion of thermal stability is, for convenience, restricted to unsupported heteropoly oxometalates with that pertaining to the supported materials included in the section on the properties of the latter. A cautionary note should be added here. The results from studies of stability are dependent not only on the method employed for such evaluations, but in addition are frequently dependent on the conditions under which the experiment was performed. One of the earliest studies of the stability of the heteropoly acids employed differential thermal analysis and X-ray diffraction.(1) With each of the four acids examined, an endotherm and an exotherm are observed at 300  C or less and 340  C or more, respectively (Table 5.1). The low-temperature endotherm is attributed to the removal of water whereas the high-temperature exotherm signals the decomposition of the anion. The authors note that metatungstic acid H8[W12O40]
xH2O decomposes at 50  C. When the void in the center of the [W12O40]8ÿ anion is occupied by Si4‡ or P5‡, the decomposition temperature increases by approximately 400 or 500  C, respectively. The presence of a positively charged species reduces the effect of thermal vibrations and thus of decomposition. Thermogravimetric analyses of H3XW12O40 with X being zinc, iron, cobalt, boron, silicon, or phosphorus, their potassium salts, and their ether-addition compounds have been reported.(2) With the ®rst three acids, three peaks were observed whereas only two were obtained with the latter three acids (Table 5.2). The third peak was attributed to decomposition of the anion. 41

42

CHAPTER 5

TABLE 5.1. DTA of Heteropoly Acids(1) Transition ( C) Acid H3PW12O40 H4SiW12O40 H3PMo12O40 H4SiMo12O40

Pretreatmenta

Endothermal

Exothermal

(a) (b) (a) (b) (a) (b) (a) (b)

175±296

580±595 573±591 470±500 487±508 390±408 397±412 336±355 340±362

140±278 63±159 64±180

(a) Dried in vacuo over concentrated sulfuric acid; (b) Dried to constant weight at temperature greater than the endothermal transition temperature.

The potassium salts showed maximum rates of water loss in the range 180± 200  C with only one peak observed for each salt. The salts appeared to retain their crystalline structure after all of the water of hydration had been removed. With the ether-addition compounds three peaks were observed for the ®rst three acids but the ®rst peak was due to the loss of ether. As with the acids the three ether-addition compounds formed from the latter three acids generated only two peaks representing weight loss. The weight loss and phosphorus content of H3PW12O40
24H2O were measured after heating in air for 3 h (static) and in helium for 2 h (¯ow) at a given temperature.(3) Weight loss was observed in three stages, the ®rst showing the removal of water to yield the 5-hydrate, the second corresponding to the formation of the anhydrous acid at 350±450  C, and ®nally a further weight loss and a substantial change in the XRD pattern above 525  C (Fig. 5.1). The weight

TABLE 5.2. Thermal Decomposition of H3XW12O40(2)

X

Temp ( C) of maximum rate of water loss

Zn Fe Co P B Si

120 120 120 120 120 120

240 240 240 270 270 270

350 390 350 Ð Ð Ð

No. of molecules of water lost 21 19.5 21 22.5 19.5 22

5 5 5 6.5 7.5 7

3 2.5 3 Ð Ð Ð

Energies of activation (kcal molÿ1 ) E1

E2

E3

6 7 7 6 6 7

21 23 21 36 36 38

56 60 58 Ð Ð Ð

STABILITY

43

Figure 5.1. Decrease in weight and phosphorus content in calcination of H3PW12O40
24H2O. s, calcined in air for 3 h (static); d, calcined in helium for 2 h (¯ow); ,, phosphorus content (3); dashed line, DTA from Ref. 1.

loss in the high-temperature region was attributed to loss of phosphorus (Fig. 5.1), which presumably occurs through the sublimation of P2O5. Temperature-programmed desorption (TPD) has been applied to the heteropoly acids.(4) The TPD pro®les of 12-tungstophosphoric acid (abbreviated as HPW) obtained after various pretreatments are shown in Fig. 5.2. The sample pretreated at 298 K exhibited two peaks both due to the desorption of water: one centered around 473 K (peak 1), with an unresolved shoulder on the hightemperature side and a second very broad asymmetrical peak centered around 773 K (peak 2). Outgassing at 463 or 593 K essentially removed peak 1 but had no effect on peak 2. No peaks were observed after outgassing at 723 K. Extremely large quantities of water were desorbed from the unpretreated acid at room temperature under the in¯uence of the helium ¯ow. Water desorbed during outgassing at 463 or 593 K could be replaced by contacting the acid with water vapor at 298 K but outgassing at 723 K destroyed the catalyst's ability to restore its original TPD behavior. The TPD pro®les for HPMo and HSiW are similar to those for HPW except for the presence of a third very sharp peak (peak 3) on the high temperature side

44

CHAPTER 5

Figure 5.2. TPD pro®les for HPW, HSiW, and HPMo after pretreatment at 298 K for 16 h.(4)

of peak 2 (Fig. 5.2). As with HPW, all peaks result from the desorption of water. Peaks 1 and 2 for HSiW appear at temperatures similar to those for HPW, whereas with HPMo, peaks 1, 2, and 3 are centered at lower temperatures of 100, 400, and 450  C, respectively. The relatively low temperatures at which peak 1 appears and the values for the activation energies (Table 5.3) which are approximately twice those usually associated with the hydrogen-bond energy suggest that peak 1 may be attributed to the desorption of water existing on the solid in molecular form held by multiple hydrogen bonds to the acid. In view of the dominance of the peripheral metal element in determining the position of peak 1, such interactions appear to be taking place on the exterior surfaces of the Keggin anion. Further, the quantity of water responsible for peak 1 in the case of HPW is consistent with the elimination of all molecular water from the solid including that contained within the bulk. It is

STABILITY

45

TABLE 5.3. Temperature-Programmed Desorption of Heteropoly Acids(4) Activation energies (kcal molÿ1 ) Acid HPW HPMo HSiW

1a

2a

3a

Water molecules desorbed per anion

15±20 15±20 15±20

47 35 44

Ð 95 60

1.4b 1.5c 1.9c

a

Peak numbers. From area under peak 2. c From areas under peaks 2 ‡ 3. b

apparently possible to remove most, if not all, of the water molecules hydrogenbonded to the protons with the secondary structures of the heteropoly acids. The values for activation energy for peaks 2 and 3 suggest that bond scission and formation are involved. The number of water molecules desorbed per anion (1.4±1.5 with HPW and HPMo and 1.9 with HSiW) are consistent with the formation of water as a result of the extraction of oxygen from the anion by the protons, apparently without collapse of the Keggin structure of the anion but with some slight rearrangement of its secondary structure, as indicated by the XRD data (see below). Thus, H3 PW12 O40 ! PW12 O38:5 ‡ 1:5 H2 O

…5:1†

H3 PMo12 O40 ! PMo12 O38:5 ‡ 1:5 H2 O

…5:2†

H4 SiW12 O40 ! SiW12 O38 ‡ 2:0 H2 O

…5:3†

can be postulated to rationalize the existence of peak 2 with HPW and peaks 2 and 3 for HPMo and HSiW. The XRD patterns for HPW show that pronounced broadening of the lines occurs on heating between 463 and 593 K with some additional lines appearing at 2y values of 19, 21.5, and 25 as indicated by the asterisk (Fig. 5.3). After heating to 723 K these lines became more pronounced. The peak at a 2y value of 25 seemed to shift to 24.5 and a new peak appeared at a 2y value of 26.2 . Reexposure of the acid, pretreated at 593 K, to water vapor at 298 K reproduced the XRD pattern observed after outgassing at 463 K but the lines were noticeably less intense. Decomposition of HPW to its constituent oxides WO3 and P2O5 (only the former was detected by XRD) occurred after heating to 773 K: H3 PW12 O40 ! 12WO3 ‡ 12 P2 O5 ‡ 32 H2 O

…5:4†

Differential thermal analysis yielded two general features: endotherms below 673 K and exotherms above 673 K (Table 5.4). The exotherms probably corre-

46

CHAPTER 5

Figure 5.3. XRD patterns for HPW pretreated in air for 90 min at (a) 463, (b) 593, (c) 723, (d) 773 K.(4)

spond to decompositions to the constituent oxides (see equation above) and the endotherms to loss of water of crystallization. Photoacoustic FTIR (PAS) has been applied to HPW and a variety of its salts.(5) At room temperature the PAS spectrum of HPW contains broad bands assignable to water (3200 cmÿ1 ) and, more speci®cally, the oxonium ion, H3O‡

STABILITY

47

TABLE 5.4. DTA of Heteropoly Acids and Their Salts(4) Sample

Endotherms (K)

HPW HPMo HSiW NaPW MgPW HPW ‡ pyridineaa

383 393 398 Ð Ð 448

a

553 435 553 593 Ð

Exotherms (K) 888 713, 733 803 848 893 830

Sample pretreated at 463 K and then exposed to pyridine vapor at 298 K.

(1710 cmÿ1 ), together with a group of ®ve or six bands below 1100 cmÿ1 characteristic of the Keggin unit, that is, the PW12O40 3ÿ ion (Fig. 5.4). Assuming an initial molecular formula of H3PW12O40
29H2O, the weight loss measured at this temperature (13%) indicates the retention of approximately four to ®ve molecules of H2O KUÿ1. At 200  C (Fig. 5.4) the band at 1710 cmÿ1 disappears, leaving a weak residual band at 1640 cmÿ1 , attributed to the bending vibration of lattice water. Weak bands assignable to the corresponding asymmetric and symmetric stretching vibrations appear at 3060 and 2980 cmÿ1 and a third band at 3160 cmÿ1 which may be the ®rst overtone of the bending fundamental. The development of structure in the KU bands at 1080 and 980 cmÿ1 indicates that some distortion of the anion has occurred. The former band is attributed to the triply degenerate asymmetric stretching vibration of the central PO4 tetrahedron. The appearance of two new bands at 1120 and 1069 cmÿ1 indicates complete resolution of the degeneracy and implies a lowering of symmetry from Td to C2v . The band at 980 cmÿ1 has been assigned to a stretching vibration involving the central W atom and the isolated terminal O(W±Ot) on the periphery of each of the 12 WO6 octahedra. The XRD pattern shows that the secondary Ê . After heating at structure remains cubic with a lattice parameter ao of 12.11 A 350 and 450  C the band at 2240 cmÿ1 appears ®rst to intensify (350  C) and then disappears, indicating complete loss of lattice water by 450  C. The progressive attenuation of the background absorption is believed to result from the loss of water and=or protons. After heating at 450  C the anion bands broaden but remain recognizable indicating partial decomposition. The PAS spectrum of the ammonium salt of HPW remains relatively unchanged with heating up to 350  C (Fig. 5.5).(5) The background absorption is much weaker than that observed with the parent acid and strong new bands are evident at 3200 and 1420 cmÿ1 , attributed to the triply degenerate asymmetric stretching (n3 ) and bending (n4 ) fundamentals of the NH4 ‡ ion, respectively. The good resolution and absence of structure in the anion bands indicate an undistorted anion, that is, Td symmetry. The XRD pattern is sharp and intense and con®rms that the ammonium salt is isostructural with HPW (cubic) but with a

48

CHAPTER 5

Figure 5.4. Effect of heating H3PW12O40
nH2O in vacuo (temp.  C).(5)

STABILITY

49

Figure 5.5. Effect of heating (NH4)3PW12O40 in vacuo (temp.  C).(5)

Ê . The P±O and W±Ot stretching bands are smaller lattice parameter, ao ˆ 11:71 A at 1083 and 987 cmÿ1 , respectively. The bands at 891, 815, and 596 cmÿ1 are assigned to stretching of W±O±W bridges between corner-sharing WO6 octahedra, stretching of W±O±W bridges between edge-sharing octahedra and P±O bend, respectively. Because no bands are observed at 2240 and 1640 cmÿ1 , lattice water is largely absent. With heating above 450  C, bands due to the NH4 ‡

50

CHAPTER 5

ion begin to diminish and broadening and shifting of the KU bands are observed. At 500  C some retention of the anion structure is evident. PAS spectra of the aluminum salt show the presence of residual lattice water as indicated by the broad bands at 3200, 2240, and 1640 cmÿ1 (Fig. 5.6). Distinct shifts and splitting in the high-frequency W±O bands indicate some anion distortion. XRD patterns show that the solid has a cubic structure with an

Figure 5.6. Effect of heating AlPW12O40
nH2O in vacuo (temp.  C).(5)

STABILITY

51

Ê . Most of the lattice water is lost by 450  C but the KU ao value of 12.135 A appears to be stable above 500  C. At 690  C decomposition has evidently occurred. The sodium salt of HPW shows strong sharp bands at 3590, 3537, and 1640 cmÿ1 assigned to the stretching [asymmetric (n1 ) and symmetric (n3 )] and bending (n2 ) fundamentals of water, respectively. The XRD pattern shows that the Ê . The salt appears to be stable up to salt is cubic with an ao value of 11.94 A  500 C or higher. The thermal stabilities of alkylammonium salts of HPW and HPMo have been investigated by DTA.(6) The tungsten-containing salts are more stable than those containing molybdenum and the unsubstituted ammonium compounds are more stable than those that are substituted (Table 5.5). The thermal stability follows the order NH4PM > R4NPM > Rx NH4ÿx PM. The R4NPW salts decompose by elimination of a gaseous base concomitant with decomposition of the secondary structure. However, the mono-, di-, and trisubstituted alkylammonium compounds show no evidence of the production of a gaseous base on thermal decomposition. An additional exothermal peak (usually as a shoulder on the higher temperature peak), found with the Rx NH4ÿx PW compounds, shifts toward lower temperatures in the order Me3NH > Me2NH2 > MeNH3. Apparently, an additional decomposition stage occurs that is facilitated by a lower degree of TABLE 5.5. DTAs of Alkylammonium 12-Heteropoly Salts(6) DTA peaks (K)a

RPW salts R

Exo

PR4N Me4N Me3NH Me2NH2 MeNH3NH NH4 c

765m, br 808m, br 763m, br; 788m, br 738m, br; 790m, br 728w, br; 788m, br Stable to at least 823 K

511w, br; 698w, br

511w

496w; 504w; 539w; 556w; 663s 523w, br; 661m; 691s 583w; 595w; 668m 606m; 663m, br 631m

RPMo salts Pr4N Me4N Me3NH Me2NH2 MeNH3 NH4 d a

b b

578w; 596m 598m; 615m, br 773s

w, weak; m, medium; s, strong; br, broad. No peak observed below 873 K. See Ref. 3. d From Ref. 41. b c

Endo b

485w; 549w b b

b

52

CHAPTER 5

Figure 5.7. Differential thermal analysis of (a) Me4NPW and (b) Me4NPMo.(6)

substitution in the cation. The tungsten-containing salts thermally dissociate by predominantly exothermal processes whereas the alkylammonium molybdophosphates decompose endothermally (Fig. 5.7). 29 Si and 51V MAS NMR have been applied to 10-molybdo-2-vanadosilicic acid (H6SiMo10V2O40) to show that the decomposition begins at 300  C or less and is complete at 300±400  C.(7) In situ powder XRD analyses of H3‡x PM12ÿx VxO40
(13±14)H2O (M ˆ Mo, W; x ˆ 0±1) have been combined with thermal analyses to investigate the thermal stability of these solids.(8) Table 5.6 summarizes the results obtained. All of the triclinic 13±14 H2O hydrates of heteropoly acids containing either tungsten or molybdenum are isomorphous. The authors note that tungstophosphoric acid may be stabilized through the formation of an anhydrous VO2 ‡ salt. This may be related to the relatively high thermal stability of the alkylammonium salts of HPW and HPMo. The H4PMo11ÿx Wx VO40 (x ˆ 0±6) acids and their pyridinium salts have been shown by 31P NMR to be mixtures of species with different values of x.(9) Three stoichiometric ratios of W=Mo (1=10, 2=9, and 3=8, labeled as y ˆ 1, 2, and 3, respectively) were employed in the preparations, with resulting mixtures of 4, 6, and 7 of the MoVW acids, respectively. Table 5.7 summarizes the temperatures at which the weight losses were observed for the acids as obtained from thermogravimetry. Two distinct transitions (weight losses) are observed, the ®rst due to the desorption of hydrogen-bonded water and the second to the loss of

STABILITY

53

TABLE 5.6. In Situ Powder XRD Analyses of H3‡x PM12ÿx VxO40 (M ˆ Mo, W; x ˆ 0, 1)(8) Crystallographic structuresa and transformation temperaturesb Acid

Cubic

HPW

29

HPMo

29

HPW11V

29

HPMo11V

20

a b c

Triclinic 25 ! 25 ! 25 ! 25 !

13 13 13 13

Cubic 40±60 ! 60±80 ! 40±60 ! 60±80 !

6 7±8 6c 7±8c

Tetragonal 180±350 ! 100±350 ! 150±410 ! 100±350 !

Anhy Anhy Anhy Anhy

550 ! 450 ! 550 ! 450 !

Oxides Oxides Oxides Oxides

Numbers beneath each structure represent water of hydration. Numbers above the arrows represent transformation temperature. Phase not isolated.

water formed from the extraction of oxygen atoms from the anions by the protons. The results of DTA are shown in Table 5.8. The DTA curves show that the decomposition of the anhydride requires a relatively small input of energy. The decomposition temperatures of the three mixtures are signi®cantly lower than those of either H4PMo11VO40 or H4PW11VO40 whereas that of the latter is substantially higher than any of the decomposition temperatures. The thermogravimetric (TG) and DTA data for the ®ve pyridinium compounds are shown in Tables 5.9 and 5.10, respectively. As the ratio W=Mo becomes larger, the temperatures at which pyridine is desorbed increase. As these authors note, the peak (Table 5.10) attributed to the crystallization of the oxides also shifts to higher temperatures as the W=Mo ratio increases and for all values of this ratio is higher than that observed for the parent acids.

TABLE 5.7. Characteristic Temperatures from Thermogravimetry(9) Transition temperature ( C) Acid H4PMo11VO40 yˆ1 yˆ2 yˆ3 H4PW11VO40 a

Loss of water.

1a

2a

20±180 20±175 20±160 20±160 20±70, 70±200

235±420 235±420 225±420 230±410 230±410

54

CHAPTER 5

TABLE 5.8. Characteristic Temperatures from Differential Thermal Analysisa…9† Transition temperature ( C) 1b

Acid H4PMo11VO40 yˆ1 yˆ2 yˆ3 H4PW11VO40

105, 105, 110, 115, 70,

135, 140, 145, 150, 105,

160 160 160 165 215

2c

3d

340±420 355±430 355±440 345±440 350±475

485 460 460 460 540

a

Flow of N2 and Ar. Maxima for loss of water of hydration. Endothermal peak. d Onset of crystallization of oxides. b c

TABLE 5.9. Characteristic Temperatures from Thermogravimetry(9) Transition temperature ( C) Acid (PyH)4PMo11VO40 yˆ1 yˆ2 yˆ3 (PyH)4PW11VO40 a b

1a

2b

220±310 220±310 220±310 225±340 225±340

310±550 310±600 340±620 340±620 340±680

First pyridine loss. Second to fourth pyridine and water loss.

TABLE 5.10. Characteristic Temperatures from Differential Thermal Analysis(9) Transition temperature ( C) Acid

1a

2a

3a

4a

5a

6b

(PyH)4PMo11VO40 yˆ1 yˆ2 yˆ3 (PyH)4PW11VO40

310 330 340 340 385

400 430 440 Ð Ð

Ð Ð Ð 480 Ð

430 465 460 Ð 535

510 560 585 590 630

560 630 630 610 >630

a b

Maxima for loss of pyridine. Onset of crystallization of oxides.

STABILITY

55

Compounds of composition Kx (NH4)3ÿx PMo12O40 were prepared from aqueous solutions of the salts by precipitation with HNO3 and were heated at various temperatures to test their thermal stabilities.(10) A single-phase cubic secondary structure with anions of Keggin structure was obtained after heating at 640 K. These authors note that the rate at which the temperature is increased to that ultimately employed is important in retaining the Keggin structure. The surface areas of the four compounds with x equal to 0, 0.92, 2.01, and 2.92 decreased with increasing potassium content (Fig. 5.8). Earlier work has shown that the ammonium salts have higher surface areas than the potassium salts.(11,12) The surface areas decreased with increasing calcination temperature, as previously reported,(13) and, with the exception of K3PMo12O40 which had a surface area of 30 m2 gÿ1 after heating at 753 K, reached negligible values at 693 K. After heating of the solids at 693 K the IR spectra of those initially containing ammonium ions contained bands assignable to MoO3, indicating the onset of the decomposition of the Keggin anion. The apparent degree of decomposition decreased with increasing potassium content, with K3PMo12O40 stable to 753 K.

Figure 5.8. Surface area of Kx (NH4)3ÿx PMo12O40 catalysts as a function of calcination temperature.(10) Reproduced by permission of Academic Press.

56

CHAPTER 5

TABLE 5.11. Fraction of Decomposed Catalysts as a Function of the Calcination Temperature Calculated by X-ray Quantitative Analysis(10) Sample

Temp calcd (K)

MoO3 (wt %)

(NH4)3PMo12O40

640 693 753

0 40 100

K1(NH4)3PMo12O40

640 693 753 843

0 47 73 >90

K2(NH4)2PMo12O40

640 693

0 <20a

K3PMo12O40

640 693 773

0 0 tr

a Approximate evaluation made by comparison of the FTIR spectra of the samples.

The content of MoO3 for each of the samples after heating at various temperatures is summarized in Table 5.11. EPR spectra show that the partial decomposition of (NH4)3PMo12O40 at 693 K leads to an increase in the degree of reduction of the solid whereas at 773 K the signal vanishes probably as a result of the reoxidation of Mo5‡ . A detailed in situ powder XRD analysis of H4PVMo11O40
32H2O at various temperatures has been published.(14) The different crystal structures and the temperatures at which they exist are shown in Fig. 5.9 and their structural parameters are collected in Table 5.12. Infrared spectra and thermal analyses showed that the thermal stabilities of HPW and its K‡ , NH4 ‡ , and Al3‡ salts follow the order KPW* > (NH4PW ˆ NH4PW*) > HPW* > HPW >AlPW* where the asterisk indicates a sample dosed with NH3.(15) These authors propose a three-step mechanism for the dehydration and decomposition of HPW (Fig. 5.10). 31 P NMR spectroscopy has been employed to examine the thermal transformations of bulk and supported HPW.(16) Data for HPW are summarized in Table 5.13. The same chemical shift (d) was found for the 29H2O solid sample as for aqueous solutions of HPW. The minimum value for d was observed after heating at 300±400  C. HPW is completely destroyed at 500±550  C with the formation of an oxide phase that consists of WO3 and P2O5. The authors ®nd that supporting of HPW on g-Al2O3 leads to destruction of the acid. On SiO2 alkaline

Figure 5.9. Conversion versus temperature pro®le for a catalytic run with a sample of H4PVMo11O40 ca. 40 mg within the in situ X-ray diffraction cell. The characters denote the different crystal structures of the catalyst (see also Table 5.1). The empty cell with sample holder was prepassivated for about 48 h at 660 K with the organic feed. After this highly corrosive treatment the conversion of the cell fell to zero. Each data point was reproduced in two independent runs and the conversion was averaged over 2 h under isothermal operation.(14) Reproduced by permission of Academic Press.

STABILITY 57

58

CHAPTER 5

TABLE 5.12. Structural Parameters of Catalyst Phases (Cell Parameters in pm)(14)

Hydration Symmetry a a b b c g Initial T (K) a

A1

A2

B

C

D

E

32 P4=mnc 1289 90 Ð Ð 1844 90 300

14 P1 1410 112.1 1413 109.8 1355 60.7 300

6 Cubic 1216 90 Ð Ð Ð Ð 345

2 Tetragonala 1375.9 90 Ð Ð 1590.6 90 420

0 Pn3m 1160 90 Ð Ð Ð Ð 520

MoO3 Pb n m 369.2 90 1385.5 Ð 369.7 90 570

Data taken from the literature. Better indexing is possible with lower symmetry.

Figure 5.10. Proposed mechanism for the thermal decomposition of HPW.(15) Reproduced by permission of Academic Press.

TABLE 5.13. Ta 20 20 20 150 200 300±400 450 500 550±850 a

31

P NMR Data for HPW(16)

Pb (H2O)

nc

21 13 <6 Ð Ð Ð Ð Ð Ð

29 13 6 3 0 0 0 0 0

ÿd31 Pd 15.0 15.2 15.5 14.6; 12.6; 11.0; 11.0; 11.0; 12.7

15.2 14.6 12.0 12.5; 13.0 12.5; 13.0

Temperature at which samples were heated for 2±4 h. Partial pressure of water vapor (Torr). Water of hydration. d ppm. b c

STABILITY

59

hydrolysis apparently occurs and consequently the loading should be at least 10% and the calcination temperature should not exceed 350  C. The thermal stabilities of HPMo and HSiMo have been investigated with IR and Raman spectroscopies, polarography, and XRD.(17) Infrared spectra show that heating of HPMo up to 380  C has little or no effect but with temperatures of 400  C or higher orthorhombic MoO3 (a-MoO3) begins to form. Similar observations are made with Raman spectroscopy but evidence for the formation of monoclinic MoO3 (b-MoO3) is found at 380  C. XRD shows that the tetragonal form of HPMo exists at 380  C with no evidence for MoO3. However, samples heated at 400 and 440  C show XRD lines for the aforementioned two oxides. Polarographic measurements suggest that the PMo12O40 3ÿ anion begins to decompose at 280  C although the extent of this decomposition is apparently relatively small. The authors note that the decomposition products of HPMo including b-MoO3 can, in the presence of water vapor, reconvert to the acid. In contrast to the observations with HPMo, above 350  C HSiMo is completely converted to a-MoO3. HSiMo remains stable up to 300  C and converts to a-MoO3 at 320  C. XRD, DTA, TG, and 31P MAS NMR have been employed to show that at temperatures greater than 400  C, PMo12O38.5 is formed from HPMo which converts to MoO3 and molybdenyl pyrophosphate (MoO2)2P2O7 at 470  C and higher.(18) This is in agreement with earlier investigations [see Eq. (5.3)].(4) Increases in pretreatment temperature up to 400  C have been shown to reduce the surface area of NH4PMo from values greater than 200 m2=gÿ1 to less than 10 m2=gÿ1 .(13) The micropore volumes decrease to negligible values by 400± 500  C. In contrast, the surface areas of NH4PW, KPW, and CsPW, although decreasing with temperature, retain signi®cant values up to 400±500  C. In particular, the surface area and micropore volume of KPW are relatively unchanged up to 400  C, whereas the micropore volumes of NH4PW and CsPW vanish by 500  C. Diffuse re¯ectance infrared Fourier-transform (DRIFT) and Raman spectroscopy together with powder XRD has shown that the Keggin structure of H4PVMo11O40 is preserved at 520±570 K.(19) The DRIFT spectrum of a sample previously employed in the dehydrogenation of isobutyric acid shows that with partial decomposition the P±O vibration band has split into several new bands and the frequencies of the M±O±M bands shift to lower values while the M ˆ Ot band remains unaffected. The separation of vanadium from the anion also occurs during the reaction. DSC and TG have been employed to examine the thermal stability of four heteropoly acids.(20) On the basis of the DSC results, the authors conclude that the thermal stability follows the order HPW > HSiW > HPMo > HSiMo (Table 5.14). The enthalpies of formation of the hydrated and anhydrous acids were also measured and are summarized in Table 5.15. The DSC and TG results are

60

CHAPTER 5

TABLE 5.14. DSCa and TGAb Results for Heteropoly Acids Peaks(20) 1d

2d

3d

Acid

nc

Ta

ÿH2Ob

Ta

Tp a

ÿH2Ob

Ta

ÿH2Ob

HPW

20 6 8.6 12

40±130 40±125 40±100 40±100

14.0 0.5 2.8 6.1

165±310 160±310 100±220 100±220

280 235 150±170 125±165

6.0 5.5 5.8 5.9

560±600 390±550 435±455 365±385

1.5 2.0 1.5 2.0

HSiW HPMo HSiMo

a

Transition temperatures from DSC. Numbers of water molecules lost are obtained from TGA. Water of hydration. d Endothermal, endothermal, and exothermal, respectively. b c

interpreted as suggesting that the transitions involve the loss of various quantities of water, the loss of 6 molecules of water and the collapse of the Keggin structure with loss of 1.5 or 2.0 molecules of water from P-containing or Si-containing anions, respectively. Cesium salts of molybdovanadophosphoric acids (Csn H3‡xÿn PMo12ÿx VxO40, n ˆ 0 to 3 ‡ x, x ˆ 0 to 2) have been investigated.(21) With the parent acid H4PMo11VO40 heated in situ with O2, a new band, attributed to vanadophosphate, appeared at 1040 cmÿ1 at 473 K, whose intensity increased with temperature up to 623 K, indicating the decomposition of the Keggin structure. On addition of Cs to form Cs2.75Mo11V (Csn Mo12ÿx Vx ) the band at 1040 cmÿ1 appeared at 523 K and remained relatively unchanged up to 623 K, thus indicating that the cesium salts are more stable than the parent acid. Another DSC study was reported in 1998.(22) Two endothermal transitions were observed for the tungsten-containing acids at temperatures of 200  C or below whereas four were found with HPMo and three with HSiMo. Exothermal

TABLE 5.15. Enthalpies of Formation of Hydrated and Anhydrous Acids(20) Hydrated n

DHf a

DHf a

25 20 6 8.6 12

ÿ 18.58 ÿ 17.20 ÿ 13.14 ÿ 12.72 ÿ 13.92

ÿ 11.36 Ð ÿ 11.39 ÿ 10.23 ÿ 10.44

Acid HPW HPW HSiW HPMo HSiMo a

Anhydrous

103 kJ molÿ1 .

STABILITY

61

transitions for HPW, HSiW, HPMo, and HSiMo were found at 612, 545, 444, and 385  C similar to those reported in Ref. 20. Cupric salts of molybdovanadophosphoric acid (Cux H4ÿ2x PVMo11O40), and in particular Cu0.5H3PVMo11O40, are active and selective catalysts for the oxidative dehydrogenation of isobutyric acid.(23) Both the latter salt and the parent acid have vanadium atoms in octahedral symmetry with a slight distortion inside the Keggin anion. After calcination of the parent acid at 593 K for 1 h, the vanadium ions are believed to be in the ‡5 oxidation state and approximately 60% are in a tetrahedral environment outside the Keggin anion. After heating for 12 h a portion of the V5‡ ions appear to be in complex interaction with electrons which is enhanced by the presence of copper. Copper appears to enhance the formation of O2ÿ vacancies in the vicinity of vanadium atoms and may increase the mobility of lattice oxygen atoms. The previously reported evidence for the enhanced thermal stability produced by the introduction of certain cations has been augmented by a study of the tetrabutylammonium salts of HPW and of its singly peripheral metalsubstituted analogue.(24) The onset of the decomposition of the cation is dependent on the composition of the anion (Table 5.16). The decomposition of the cation of (TBA)3PW begins at a higher temperature than that of the substituted tungstophosphate anion and the anion of the former begins to decompose at approximately 450  C and is completely decomposed at 600  C. As (TBA)3PW11M begins to decompose the metal M is expelled from the anion and the anion PW12 is formed as an intermediary. The decomposition temperatures of the potassium salts of the aforementioned anions were at least 150  C higher than those of the tetrabutylammonium salts but the metal M was not expelled with the former cation. TABLE 5.16. Decomposition of Tetrabutylammonium Salts of 12-Tungstophosphoric Acid (HPW) and Its Peripheral Metal Substituted Analogue (HPW11M)(24) Anion (M) Mn Co Ni Cu Fe

Decomposition temperaturea

Formation of (TBA)3PW12 b

150±592 163±603 190±644 186±508 193±562

310 310 320 280 290

a Minimum value corresponds to onset of decomposition of the cation ( C). b Minimum temperature at which the metal M is expelled from the anion ( C).

62

CHAPTER 5

5.2. pH STABILITY The stability of heteropoly anions of Keggin structure has been investigated with a technique based on 31P NMR(25) which permitted the study of the exchange of tungsten and molybdenum atoms between discrete Keggin anions in aqueous solution. The nature of the formation and destruction process of the Keggin anion in aqueous solution over the pH from 0 to 12 has been studied by application of 31P NMR and IR spectroscopies to the heteropoly acids HPW and HPMo.(26) The products formed as HPMo and HPW were mixed with sodium hydroxide were also investigated. Four additional systems HPO4 2ÿ =CrO4 2ÿ, HPO4 2ÿ =VO4 3ÿ, P2O7 4ÿ =MoO4 2ÿ, and HPO3 2ÿ =MoO4 2ÿ were also studied with 31P NMR spectroscopy at various values of pH. 5.2.1. Molybdate=Phosphate System The 31P NMR spectra for acidi®ed phosphate=molybdate (P=Mo ˆ 1=12) solutions, for HPMo, and for the species generated on addition of sodium hydroxide are shown in Fig. 5.11 together with the chemical shifts in Table 5.17. HPMo exists as the predominant species in solutions of pH 41.5 but other species can also be seen in this pH range. Infrared spectra have provided complementary data for the assignments. The spectra for samples at low pH possess the vibrational bands of the Keggin anion at 1065, 960, 870, and 785 cmÿ1 attributed to the P±O symmetric stretch, Mo±O asymmetric stretch, and Mo±O±Mo inter- and intraoctahedral stretches, respectively.(30,31) These bands exist up to a pH of 2, but the 1065 cmÿ1 band splits into two bands at approximately 1063 and 1035 cmÿ1 . The bands present in this spectrum are consistent with those of the PMo11O39 7ÿ anion.(32) At a pH of 5, bands at 965, 895, and 565 cmÿ1 have been attributed to the P2Mo5O23 6ÿ anion.(33) 5.2.2. Tungstate=Phosphate System 31

P NMR spectra for the 12=1 tungstate=phosphate solutions with pH from 0.4 to 11.2 and for solutions containing HPW and sodium hydroxide are shown in Fig. 5.12. The most complex spectra are observed at values of pH between 1.0 and 4.0. At a value of pH greater than 8 a single peak is observed at 3:5 ppm which is attributed to the free phosphate ion. For values of pH of 2.5 to 7.0 a single peak at ÿ10:1 ppm is found and assigned to PW11O39 7ÿ.(27) The up®eld shift of approximately 13 ppm is indicative of a large perturbation in the local phosphorus environment due to the condensation of a heteropoly species. The species PW11O39 7ÿ is stable over a large range of pH values, is a known

STABILITY

63

Figure 5.11. 31P NMR spectra of molybdophosphate and phosphate=molybdate (P=Mo ˆ 1=12) aqueous solutions at various pH.(26) Reprinted from the Journal of Molecular Catalysis, 69, McGarvey and Moffat, p. 137, copyright 1991, with permission from Elsevier Science.

decomposition product of and precursor to the PW12O40 3ÿ anion,(34) and is the predominant species to pH  2:0. Further acidi®cation of PW11O39 7ÿ results in the formation of PW12O40 3ÿ and=or P2W21O71 6ÿ (34) with peaks at ÿ14:4 and  ÿ 12:8 ppm, respectively.(27) Below a pH of 1 the peak assigned to P2W21O71 6ÿ disappears whereas below a pH of 1.5 the PW12O40 3ÿ resonance becomes the major peak, in agreement with the proposed mechanism for the reversible decomposition of the PW12O40 3ÿ anion at values of pH from 1.5 to 2.(34) It should be noted that a number of peaks in both the molybdate=phosphate and tungstate=phosphate systems remain unidenti®ed, particularly with the latter

64

CHAPTER 5

TABLE 5.17. 31P Chemical Shifts and Species Assignments for Molybdate=Phosphate Solutions(26) pH 0.9 1.5 2.0 3.0 4.0 5.0 6.1 7.0 8.0 10.0

Chemical shift (ppm) ÿ 3.05 ÿ 1.14 ÿ 1.17 ÿ 1.06 0.53 ÿ 0.78 1.93 0.48 ÿ 0.69 2.1 2.71 3.26 3.46

Species PMo12O40

Refs. 3ÿ

27, 28

PMo11O39 7ÿ

28

PMo6O25 9ÿ

28

P2Mo5O23 6ÿ

28, 29

combination. However, with both systems the predominant species have been identi®ed. The dependences of the major species present in the tungstate=phosphate and molybdate=phosphate systems on the pH are shown in Figs. 5.13 and 5.14, respectively. The formation or decomposition in each system is evidently strongly dependent on the pH and proceeds in three reasonably well de®ned steps. In the tungstate=phosphate combination, PW11O39 7ÿ is apparently readily produced from WO4 2ÿ and HPO4 2ÿ between pH values of 8 and 7.5 and exists alone to a pH of approximately 2.5. Several NMR peaks are observed between values of pH from 2.5 to 1.0 as well as those attributed to PW11O39 7ÿ and PW12O40 3ÿ. The peak at approximately ÿ12:8 ppm is assigned to the P2W21O71 6ÿ anion which is known to be stable in acidic solution.(34) The species is present down to a pH of 1. The existence of the PW11O39 7ÿ anion from pH 7.5 to 2.5 is con®rmed from 183 W NMR. In view of the demonstrated reversibility of the formation= decomposition process, this species can be prepared from either the parent 12tungstophosphate or the phosphate=tungstate precursors. Such observations may be of relevance in the preparation of Keggin anions with a single substituted peripheral metal atom which evidently can be synthesized from the parent acid or the precursor combinations. The formation and degradation of the PMo12O40 3ÿ anion also appear to proceed through three predominant species (Fig. 5.14). In solutions with values of pH greater than 5 the NMR peak at approximately 2 ppm is identi®ed as P2Mo5O23 6ÿ. Below a pH of 5 this species disappears and a peak at approxi-

STABILITY

65

Figure 5.12. 31P NMR spectra for 12-tungstophosphoric acid aqueous solutions and the species generated from the addition of sodium hydroxide.(26) Reprinted from the Journal of Molecular Catalysis, 69, McGarvey and Moffat, p. 137, copyright 1991, with permission from Elsevier Science.

66

CHAPTER 5

Figure 5.13. 31 P NMR chemical shifts of the predominant species in tungstophosphate solutions at various pH.(26) Reprinted from the Journal of Molecular Catalysis, 69, McGarvey and Moffat, p. 137, copyright 1991, with permission from Elsevier Science.

mately ÿ2 ppm, assigned to PMo11O39 7ÿ, the precursor to PMo12O40 3ÿ,(28) becomes dominant. The two species coexist until the pH reaches 1.5. The vanadate and chromate systems do not show the presence of plateaus in the chemical shift=pH plots, suggesting that the number of heteropoly oxometalates that can be formed with these peripheral metals is small.

STABILITY

67

Figure 5.14. 31P NMR chemical shifts of the predominant species in molybdophosphate solutions at various pH.(26) Reprinted from the Journal of Molecular Catalysis, 69, McGarvey and Moffat, p. 137, copyright 1991, with permission from Elsevier Science.

The stability of HSiMo, HSiW, HPMo, and HPW has been measured in aqueous solutions at various values of pH by use of ion chromatography.(35) The stabilities in aqueous solution with respect to pH follow the order HSiW > HPW > HSiMo > HPMo.

5.3. REGENERATION IN THE PRESENCE OF WATER Studies of SiO2-supported molybdena catalysts prepared from ammonium paramolybdate or the Mo(Z3 -C3H5)4 complex with oxygen chemisorption, ESCA, laser Raman spectroscopy, and ion-scattering spectrometry have indicated the presence of a silicomolybdic anion species at the surface of Mo=SiO2 catalysts at loadings as low as 1% Mo.(36,37) At low loadings a larger quantity of the silicomolybdic species was present in the catalysts prepared from the ammonium paramolybdate. ESR and NMR have been employed to suggest that a rapid reconstruction of Keggin anions from the ®nal decomposition of H3‡n PVn Mo12ÿn O40 supported on silica occurs in the presence of moist air.(38) Raman, XPS, and 31P NMR spectroscopic studies on HPMo=SiO2 have also shown that features of the acid are recovered after exposure of the sample calcined at temperatures lower than 550  C to moist air provided that well-de®ned

68

CHAPTER 5

molybdenum oxide is not formed during the heat treatment.(39) The authors consider that the sample calcined at a temperature lower than 550  C contains a lacunary form of the Keggin structure and not a decomposition product. Further investigations of both HPMo=SiO2 and cesium-doped HPMo supported on SiO2 with laser Raman and X-ray photoelectron spectroscopy con®rm that the spectrum of the Keggin anion is retrieved after heating to 550  C followed by exposure to the atmosphere.(40) In contrast, exposure of the sample previously heated to 560  C to the atmosphere does not lead to the recovery of the Raman features of the unheated supported HPMo.

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26.

S. F. West and L. F. Audrieth, J. Phys. Chem. 59, 1069 (1955). D. H. Brown, J. Chem. Soc. 3189 (1962). H. Hayashi and J. B. Moffat, J. Catal. 77, 473 (1982); J. Catal. 83, 192 (1983). B. K. Hodnett and J. B. Moffat, J. Catal. 88, 253 (1984). J. G. High®eld and J. B. Moffat, J. Catal. 88, 177 (1984). J. B. McMonagle and J. B. Moffat, J. Catal. 91, 132 (1985). R. I. Maksinovskaya, N. N. Chomachenko, and D. V. Tarasova, React. Kinet. Catal. Lett. 28, 111 (1985). M. Fournier, C. Feumi-Jantou, C. Rabia, G. HerveÂ, and S. Launay, J. Mater. Chem. 2, 971 (1992). C. Marchal, A. Davidson, R. Thouvenot, and G. HerveÂ, J. Chem. Soc. Faraday Trans. 89, 3301 (1993). S. Albonetti, F. Cavani, F. Tri®ro, M. Gazzono, M. Koutyrev, F. C. Aissi, A. Aboukis, and M. Guelton, J. Catal. 146, 491 (1994). J. B. McMonagle and J. B. Moffat, J. Colloid Interface Sci. 101, 479 (1984). G. B. McGarvey and J. B. Moffat, J. Catal. 128, 69 (1991). D. Lapham and J. B. Moffat, Langmuir 7, 2273 (1991). T. Ilkenhans, B. Herzog, T. Braun, and R. SchloÈgl, J. Catal. 153, 275 (1995). B. W. L. Southward, J. S. Vaughan, and C. T. O'Connor, J. Catal. 153, 293 (1995). R. I. Maksimovskaya, Kinet. Katal. 36, 836 (1995). C. Rocchiccioli-Deltcheff, A. Aouissi, M. M. Bettahar, S. Launay, and M. Fournier, J. Catal. 164, 16 (1996). T. V. Andrushkevich, V. M. Bondareva, R. I. Maksimovskaya, G. Y. Popova, L. M. Plyasova, G. S. Litvak, and A. V. Ziborov, Stud. Surf. Sci. Catal. 82, 837 (1994). B. Herzog, M. Wohlers, and R. SchloÈgl, Mikrochim. Acta 14, 703 (1997). C. Xian-e, D. Daichun, N. Jianping, J. Yourning, Z. Jing, and Q. Yixiang, Thermochim. Acta 292, 45 (1997). K. Y. Lee, S. Oishi, H. Igarashi, and M. Misono, Catal. Today 33, 183 (1997). M. Varga, B. ToroÈk, and A. Molnar, J. Therm. Anal. 53, 207 (1998). E. Blouet-Crusson, M. Rigole, M. Fournier, A. Aboukais, F. Daubrege, G. Hecquet, and M. Guelton, Appl. Catal. A Gen. 178, 69 (1999). J. A. Gamelas, F. A. S. Couto, M. C. N. Trovao, A. M. V. Cavaleiro, J. A. S. Cavaleiro, and J. D. P. DeJesus, Thermochim. Acta 326, 165 (1999). M. A. Schwegler, J. A. Peters, and H. van Bekkum, J. Mol. Catal. 63, 343 (1990). G. B. McGarvey, and J. B. Moffat, J. Mol. Catal. 69, 137 (1991).

STABILITY

69

27. R. Mossart, R. Contant, J.-M. Fruchart, J.-D Ciabrin, and M. Fournier, Inorg. Chem. 16, 3916 (1977). 28. J. A. Rob Van Veen, O. Sadmeijer, C. A. Emeis, and H. deWit, J. Chem. Soc. Dalton Trans. 1825 (1986). 29. L. Pettersson, L. Andersson, and L.-O. Ohman, Inorg. Chem. 25, 4726 (1986). 30. N. E. Sharpless and J. S. Munday, Anal. Chem. 29, 1619 (1957). 31. C. Rocchiccioli-Deltcheff, R. Thouvenot, and R. Franck, Spectrochim. Acta 32A, 587 (1976). 32. C. Rocchiccioli-Deltcheff and R. Thouvenot, J. Chem. Res. (S), 46 (1977); J. Chem. Res. (M), 546 (1977). 33. L. Lyhamn, Chem. Scr. 12, 153 (1977). 34. M. T. Pope, Heteropoly and Isopoly Oxometalates, Springer, Berlin (1983). 35. A. JuÈrgensen and J. B. Moffat, Catal. Lett. 34, 237 (1995). 36. L. Rodrigo, K. Marcinkowska, A. Adnot, P. C. Roberge, S. Kaliaguine, J. M. Stencel, L. E. Makowsky, and J. R. Diehl, J. Phys. Chem. 90, 2690 (1986). 37. J. M. Stencel, J. R. Diehl, J. R. D'Este, L. E. Makowsky, K. Marcinkowska, A. Adnot, P. C. Roberge, and S. Kaliaguine, J. Phys. Chem. 90, 4739 (1986). 38. E. M. Serwicka and C. P. Grey, Colloids Surfaces 45, 69 (1990). 39. S. Kasztelan, E. Payen, and J. B. Moffat, J. Catal. 125, 45 (1990). 40. E. Payen, S. Kasztelan, and J. B. Moffat, J. Chem. Soc. Faraday Trans. 88, 2263 (1992). 41. J. A. Rashkin, E. D. Pierron, and D. L. Parker, J. Phys. Chem. 71, 1265 (1967).

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6 SUPPORTED HETEROPOLY ACIDS AND THEIR DERIVATIVES

As noted elsewhere in this volume, the surface areas of the heteropoly acids are relatively low (less than 10 m2 gÿ1 ). Although polar molecules are capable of penetration into the bulk of these solids, that is, between the cations and anions, reactions involving such molecules will almost inevitably be diffusion-hindered. To enhance the surface of the solids available to the reactant molecules, two possibilities are available: (1) employ high-surface-area solids as supports or (2) develop higher-surface-area microporous forms of the heteropoly oxometalates. The present chapter will focus on the ®rst method, the second being considered elsewhere in this volume. Silica has probably been the most frequently employed as a support, although silica±alumina, alumina, various polymers, titania, carbon and ion-exchange materials have also been used.

6.1. SUPPORTS 6.1.1. Carbon In view of their solubility in aqueous solutions, it is not surprising that leaching of the acids from supported systems has been reported. Although the complete elimination of such losses is unlikely to be achieved, activated carbon (Calgon F-300) has been suggested as a support from which leaching can be reduced.(1) The quantity of the acid extracted from the carbon by hot water or hot methanol was found to depend on the solvent as well as the initial loading on the support. For example, loadings of 12-tungstophosphoric acid of 9, 17, and 18 wt % were extracted up to approximately 2, 20, and 60%, respectively. 71

72

CHAPTER 6

Repeated reactions of butanol and tert-butyl alcohol on a sample of HPW supported on carbon, recovered after each reaction, produced similar results after each exposure, suggesting that little or no leaching of HPW was occurring(2). Although these authors did not characterize the dispersions nor were surface areas provided, the fraction of the loadings remaining after extraction may be related to the monolayer of acid bound directly to the surface of the support. The acidity of HPW supported on activated carbon has been evaluated with microcalorimetry.(3) The differential heats of adsorption of ammonia on HPW=C are smaller than those of the unsupported acid by approximately a factor of two and decrease with loading (Fig. 6.1). The authors argue that the interaction between the Keggin anion and the support does not result from cation exchange nor the protonation of basic groups of the surface but from an encapsulation of the anion by the carbon. The esteri®cation of phthalic anhydride with a number of alcohols,(4) of acrylic acid by butanol,(5) and of propanoic acid by butanol and 2-ethylhexanol(6) have been investigated in the liquid phase with heteropoly acids supported on carbon. The supported acids were found to be superior to sulfuric acid or resins and under continuous use the low deactivation of the catalyst was ascribed to dissolution of the acid.(5,6) The activity of the supported catalyst in the propanoic acid esteri®cation was found to be lower than that of the unsupported acid, as expected, but these authors contend that one proton per anion does not participate

Figure 6.1. Differential heats of adsorption of ammonia on HPW and HPW=C.(3) Reproduced by permission of Akademiai Kiado.

SUPPORTED HETEROPOLY ACIDS

73

in the esteri®cation process but rather is involved in the interaction between the acid and the support.(6) The results of studies of the adsorption of heteropoly acids from aqueous and organic solutions onto activated carbons have been reported.(7±9,11) In contrast to the earlier work,(3) some relatively strong adsorption involves proton transfer.(7) Adsorption isotherms on active carbon as well as other supports were obtained and a model for the interaction between the heteropoly acid and the hydroxyl groups of the support was proposed.(8) Isotherms for the adsorption of HPW and HSiW from ethanol±water solutions onto carbon have also been measured and the data ®tted to the Langmuir isotherm.(9±11) The limiting values for HPW and HSiW were at loadings of approximately 30 and 15%, respectively, although with the latter adsorbate the saturation plateau was somewhat ambiguous. Evidence for the existence of highly dispersed noncrystalline adsorbed species was provided. This appears to be consistent with the results found earlier for heteropoly acids on SiO2 , as described later in this chapter.(12) HPMo was found to be stable, when supported on carbon, up to 400  C. Contrary to the earlier reports, the tungsten-containing heteropoly acid (H3 PW12 O40 ), when supported on carbon, is found to be thermally stable up to 425  C and has apparently higher acidity than the bulk acid as evidenced from the activity in the dehydration of isopropanol.(11) It is not clear, however, to what extent this observation can be attributed to the increased accessibility of protons in the supported material. 6.1.2. TITANIUM DIOXIDE Recently TiO2 in the form of titania has been considered as a support for the heteropoly acids.(10,11,13±17) Infrared spectroscopy has been employed to investigate the nondissociative adsorption of NO on titania-supported HPMo and its cobalt and nickel salts.(13,14) HPMo loaded on titania by adsorption to equilibrium from a water±ethanol solution was stable up to 400  C.(10) As noted with carbon as a support the TiO2 -supported tungstophosphoric acid (HPW) was thermally stable up to 425  C and possessed higher activity than the bulk acid in the dehydration of isopropanol.(11) HPW appeared to be present in at least ®ve forms on the titania surface: a bulk salt phase, two weakly bound intact Keggin species, a range of partially fragmented clusters such as the 11-``defect'' Keggin ion, and a range of species formed from the fragmentation of the Keggin anion.(15) The relative amounts of these species were found to be dependent on the form of the support. With molybdophosphoric acid (HPMo) supported on titania the heteropoly anion is converted to a- and b-MoO3 at a temperature higher than that of the bulk material due to the interaction between anion and support.(16) Methanol oxidation at 523 K on HPMo=TiO2 produced a high selectivity to dimethyl ether,

74

CHAPTER 6

suggesting an acid catalyst, whereas replacement of the protons in HPMo by nickel and cobalt yielded a catalyst with redox properties as evidenced by the formation of formaldehyde.(17) IR and XPS data obtained after reaction showed that the Keggin structure was preserved up to 623 K. 6.1.3. Silica Silica is probably the most frequently used support for the heteropoly acids. As early as 1971 tungstosilicic acid (HSiW) was employed as a catalyst for the alkylation of benzene with dodecene-1.(18) Silica was found to be superior as a support in comparison with alumina and silica-alumina. Selectivities of 90±94% to phenyldodecane were obtained with silica-supported HSiW. HPMo and its metal salts supported on silica were also employed in the reduction of NO with hydrogen or ammonia.(19) The alkylation of benzene with ethylene, esteri®cation of acetic acid with ethanol and dehydration of 2-propanol were studied on silica gel and activated carbon.(20). Electron spin resonance studies have shown that dehydration of HPMo and HSiMo in air at low temperatures up to 423 K involves the dissolution of the acid in its water of crystallization and is associated with a single narrow line in the spectra characteristic of strong delocalization of the electrons in the Keggin ion.(21) The electron jumps between the 12 Mo atoms at 108 sÿ1 . In the temperature range from 400 to 650 K the odd electron is localized at the molybdenum atoms. The dehydration process appears to be related to a loss of oxygen and a consequent reduction of molybdenum atoms. The authors conclude that the irreversible destruction of the anion occurs at 673 and 623 K with HPMo and HPMo=SiO2 , respectively. However, immediate destruction of the structure occurs at low temperatures on loading the acid on alumina. Oxygen radicals were not observed at either room temperature or 77 K after the adsorption of oxygen on unsupported HSiMo. However, the supported acid previously reduced at 723 K stabilizes O2 ÿ species in high concentrations.(21) Apparently the anion with distorted but intact Keggin structure cannot form or stabilize oxygen radicals. The authors attribute the inability of the Keggin structure to stabilize O2 ÿ to an incorporation into the lattice as O2 …gas† ! O2 ÿ …ads† ! Oÿ …ads† ! O2ÿ …lattice† in contrast to M…nÿ1†‡ ‡ O2 …ads† ! Mn‡ ‡ O2 ÿ …ads† the latter requiring M to be situated in a tetrahedral environment. The oxidation of methane has been investigated on heteropoly acids and their salts supported on silica.(12,22±31) Additional details can be found in the chapter on oxidation processes. The effects of the replacement of protons by a

SUPPORTED HETEROPOLY ACIDS

75

monovalent cation, cesium, as well as a variety of mono-, di-, and trivalent cations have also been considered.(25,28) In the oxidation of methane on HPMo=SiO2 both the acid and the support are participating catalytically, with the support having a detrimental effect on the selectivities to the more desirable products, formaldehyde and methanol, which are easily converted to CO and CO2 on the support itself, while little or no reaction of CO is observed.(22,23) The conversion and selectivities with nitrous oxide on HPMo=SiO2 remained relatively constant for calcination temperatures up to 773 K (Fig. 6.2).(24) The conversion of methane at 843 K decreases sharply at higher calcination temperatures while the production of CO and CO2 remains essentially ®xed up to approximately 900 K. At temperatures greater than 900 K the generation of H2 CO and CO decreases, while that of CO2 increases, apparently approaching the selectivities found with the support itself. A slight loss in loaded molybdenum occurs at temperatures higher than 800 K. At a calcination temperature of 823 K [at which temperature thermal degradation is occurring (Fig. 6.2a,b)], the conversion of methane at the same

Figure 6.2. Effect of the temperature of calcination during 16 h (left) and of the time of calcination at 823 K under air (right) on the CH4 conversion, selectivity, and Mo loading of the 23-HPMo catalyst. Reaction conditions CH4 (67%), N2 O (33%), TR ˆ 843 K, W ˆ 0:5 g, F ˆ 30 ml minÿ1 . Symbols: (n) CO, (s) CO2 , (u) CH2 O, (X) CH4 conversion, () Mo loading.(24)

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temperature decreases slowly with the duration of calcination, a small loss of molybdenum occurs, and the production of CO increases slightly while the selectivities to CO2 and H2 CO decrease slightly (Fig. 6.2c,d).(24) It seems clear that the catalytic activity in the methane oxidation process can be ascribed to either HPMo or the effect of supporting HPMo on SiO2 which, on heating to a suf®ciently high temperature, generates decomposition products that are signi®cantly less catalytically active. The rates of formation of the various products from the oxidation of CH4 with nitrous oxide at 843 K increase approximately linearly with loading of HPMo=SiO2 , extrapolate, at low loading values, to the results expected for the support itself, reach maxima at a loading of approximately 20 wt % or 120 mmol of KU per gram of support, and for further increase in loading decrease markedly.(24) This optimum loading corresponds to a coverage of approximately Ê 2 of the Ê 2 KUÿ1 to be compared with the cross-sectional area of 100 A 1000 A Keggin unit. Not surprisingly, the selectivities to the various products are also dependent on the loading. At 843 K and at low loadings the selectivities to CO and to CO2 increase and decrease sharply ultimately reaching constant values at approximately 30 wt %.(24) Concomitantly the selectivity to formaldehyde increases, passes through a maximum, and decreases to negligible values at 30 wt %. Of some interest, methanol does not appear until 80 mmol KU per gram of support and vanishes for loadings higher than 150 mmol KU per gram of support. Infrared spectra of acetonitrile solutions after their use in washing of various supported HPMo samples show that the bands at 1080 and the doublet at 969± 960 cmÿ1 characteristic of the Keggin structure and attributed to the triply degenerate asymmetric stretch of the central PO4 tetrahedron and that of the outer Mo±O bond, respectively, are intact after heating 20% HPMo=SiO2 at 640  C for 16 h but are reduced to very small intensities after heating this sample at 730  C for 16 h.(30) Laser Raman spectra of bulk HPMo and of HPMo=SiO2 , the latter of various loadings, after calcination at 350  C show the main Raman bands at 998, 970, 610, and 247 cmÿ1 assigned to v s Mo±Ot, v as Mo Ot , v s Mo Oc , and Mo Oa (t, terminal oxygen atom; c, oxygen bridging two Mo atoms; a, oxygen joining Mo and P atoms). A 23 wt % sample calcined at temperatures up to 550  C for 90 h shows the same bands whereas after 700  C for 3 h the intensities of these bands had signi®cantly diminished.(30) 31 P NMR spectra of HPMo=SiO2 contained one relatively broad peak and the peak positions and FWHM are consistent with the 31P NMR spectra of heteropolymolybdate for samples of 31 wt % and lower (Table 6.1) (30) Although, after calcination of a 23 wt % sample at 550  C for 16 h or use in the oxidation of methane at 570  C for 4 h the NMR data are consistent with the retention of the HPMo structure, use of 600  C evidently introduces a perturbation.

SUPPORTED HETEROPOLY ACIDS

77

TABLE 6.1. 31P NMR Peak Parameters of SilicaSupported 12-Molybdophosphoric Acid Catalyst: (a) 23-HPMo Calcined at Various Temperatures and for Various Lengths of Time; and (b) with Different Loadings and Calcined at 350 C, 2 h(30) NMR 31P peak position (ppm)

FWHM (Hz)

a. Treatment Bulk HPMo 200  C, 2 h 350  C, 16 h 450  C, 16 h 550  C, 16 h AT (570  C, 4 h)* 600  C, 16 h 730  C, 16 h

ÿ8 ÿ8 ÿ8 ÿ8 ÿ8 ÿ8 ÿ5 ÿ5

1.5 1.0 1.2 1.1 1.2 1.1 1.8 9.4

b. Loading (wt %) Bulk HPMo 1.2 8.3 16 31 39

ÿ8 ÿ8 ÿ8 ÿ8 ÿ8 ÿ8

1.5 1.5 1.2 1.0 1.1 4.1

*After use in methane oxidation at temperature and duration of time indicated.

XPS has provided further information on the dispersion of the supported HPMo on SiO2 .(30) A plot of the Mo3d =Si2p XPS intensity ratio versus the loading contains three identi®able portions, the ®rst linear up to a loading of 0.04 KU nmÿ2 (10 wt % HPMo) with a slope close to the position of the calculated curve and probably the result of single or small aggregates dispersed over the surface (Fig. 6.3) A change in slope can be attributed to the formation of larger aggregates up to 23 wt % HPMo at which large aggregates begin to form. The aforementioned results show that the highly dispersed HPMo=SiO2 is stable up to approximately 580±600  C whereas the particles of HPMo on the highly loaded catalysts are converted to MoO3 under more severe conditions as expressed in the stoichiometric equation H3 PMo12 O40 ! 12 P2 O5 ‡ 12MoO3 ‡ 32 H2 O The increase in the stability of the supported HPMo relative to that of bulk HPMo which degrades at 350  C(12) is tentatively attributed to its relatively strong

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CHAPTER 6

Figure 6.3. Evolution of the XPS Mo3d =Si2p intensity ratio versus the 12-molybdophosphoric acid loading for the series of catalysts calcined at 350  C (s), 2 h, and after catalytic testing at 570  C, 3 h (‡), as calculated for a monolayer dispersion (± ± ±).(30)

interaction with the silica, although the surface sites for such interaction and its nature are as yet unknown. Further insight into the interactions between heteropoly acids and SiO2 as a support have been obtained from studies of MoO3 =SiO2 catalysts.(26) IR and Raman spectra together with methane oxidation experiments have provided evidence for the formation of 12-tungstosilicic acid on SiO2 when the impregnation of MoO3 on the support occurs at pH 2 and 7. This suggests that a portion of the silica is dissolved during the impregnation in an acidic medium, forming HSiMo. In basic solution the heptamolybdate and a small amount of molybdosilicic acid are formed. At least a portion of the HSiMo remains stable up to 873 K. However, at high calcination temperatures MoO3 crystallites result either from the decomposition of HSiMo or from heptamolybdate at relatively low temperatures. The proportion of the aforementioned three species existing on the support is evidently dependent on both the pH and temperature with that of HSiMo increasing as the pH decreases (Table 6.2) The conversion of methane on MoO3 =SiO2 appears to be primarily due to the formation of molybdosilicic acid, although, not surprisingly, the selectivities to desired products are not favored at

SUPPORTED HETEROPOLY ACIDS

79

TABLE 6.2. Catalytic Properties and Characteristics of the Samples(26)

Samples c

SiO2 support MoO3 =SiO2 (pH 11) MoO3 =SiO2 (pH 7) MoO3 =SiO2 (pH 2) H4 SiMo12 O40 =SiO2

Conversiona

Selectivitiesa

Loading wt % MoO3

CH4

N2 O

N2 O TONb

CO

CO2

CH2 O

CH3 OH

Ð 3.0 3.2 3.2 5.7

0.8 1.1 2.6 3.3 5.0

3.0 4.0 8.4 10.0 19.6

Ð 9.8 19.0 22.6 25.2

75 72 67 61 46

25 19 24 37 54

Ð 9 9 2 tr

Ð tr tr tr Ð

T ˆ 773 K, mass of catalyst 2 g, ¯ow rate 15 ml minÿ1 , ¯ow composition CH4 (33%), N2 O (67%). Apparent turnover number (10ÿ4 mol at. Moÿ1 sÿ1 ). c SiO2 Davison-Grace, Grade 400, 750 m2 gÿ1 . a b

the higher conversions. Subsequent studies of MoO3 =SiO2 at various loadings have provided further evidence relating to the aforementioned.(33) Impregnation of silica with aqueous paratungstate solutions produces a Si± O±W crystalline species with Keggin-like structures whose thermal stability, although dependent on the loading, is higher than that of HSiW.(34) These results are analogous to those for Mo=SiO2 in which evidence for the formation of HSiMo was found.(26) Solid-state NMR and IR studies of both HPMo and HPW supported on SiO2 , g-Al2 O3 , and SiO2 Al2 O3 have provided further information on the thermal stabilities of these supported acids.(35) With 14 wt % HPMo=SiO2 the linewidth of the 31P NMR signal (75 Hz) nearly coincides with that of pure HPMo (70 Hz) whereas the signal on alumina and silica±alumina is much broader. The authors conclude that the catalyst primary structure is only weakly perturbed on silica whereas on alumina and silica±alumina it is more strongly modi®ed or even partially destroyed. In situ IR spectra show that the Keggin unit of HPMo=SiO2 and of HPW=SiO2 is thermally stable in the range 273±473 K. Silica-supported heteropoly acids have also been employed as catalysts in liquid-phase Friedel±Crafts reactions.(36) Both the pure heteropoly acids and those loaded on silica up to 30 wt % were employed in the work. The supported catalysts were generally superior to the pure materials. Further 1H and 31P MAS NMR have been reported with the chemical shifts, referenced to TMS and phosphoric acid, respectively.(37) The 1H chemical shifts for the pure acids are shown in Table 6.3. The intensity of the most intense line, at d ˆ 1:8 ppm, of the 1H NMR spectra of HPW=SiO2 , attributed to the silica surface hydroxyl groups, decreases as the loading increases. The line at 9.3 ppm, as found for pure HPW, while not evident in the spectra of HPW=SiO2 for loadings less than 50 wt %, appears in the spectrum of the latter. A line with d ˆ 5 ppm forms for loadings greater than or equal to 20 wt % and is assigned to the proton of supported HPW.

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TABLE 6.3. 1H Chemical Shifts of Pure Heteropoly Acids(37) Acid HPW HSiW HSiMo HPMo

d (ppm) 9.3 9.7 8.5 7.4

The 31P MAS spectra of HPW=SiO2 are shifted from those found for pure HPW, indicative of the interaction of the acid with the support.(37) The evidence from both 1H and 31P NMR spectra suggests that at loadings up to 20 wt % the HPW species are isolated on the surface but that at higher loadings clusters exist while crystals of HPW form at 50 wt %. These results appear to be in good agreement with those obtained earlier for HPMo=SiO2 .(30) These authors argue that the interaction of HPW with SiO2 results in the formation of new protonic sites on the silica surface, resulting in a decrease in acid strength: H3 PW12 O40 ‡ m… SiÿOH† ! …Si†m ‡ …H3ÿm PW12 O40 †mÿ ‡ mH2 O The effect of supporting HSiMo on SiO2 has also been investigated by application of IR spectroscopy and the methanol oxidation reaction.(38) The thermal stability of the supported. (9 and 17.6 wt %) was found to be 20±30  C lower than that of the pure acid. The thermal stability of the supported H3‡n PVn Mo12ÿn O40 acids has been examined by application of ESR and MAS NMR spectroscopy.(39) These authors conclude that deposition on silica lowers the thermal stability but that a reformation of the Keggin structure from the decomposition products occurs in the presence of water vapor. The properties of H5 PV2 Mo10 O40 =SiO2 strongly depend on the surface coverage for 0.05±1.0 monolayer.(40) Interactions between the hydroxyl groups of silica and the protons of the heteropoly acid result, at low loadings, in the hydrolysis of the Keggin anions into triads of edge-linked octahedra strongly bound to the surface (Fig. 6.4) Consequently, the acidity is markedly diminished and in the oxidation of methanol only formaldehyde and methyl formate are formed. At high acid loadings, anions of unperturbed Keggin structure exist on the surface of silica leading to a slight increase in thermal stability and forming dimethyl ether from methanol. Reconstruction of the thermally fully decomposed heteropoly anion occurs on exposure to water vapor. The effect of supports on the activity of heteropoly acids has been evaluated.(41) HPW=SiO2 and HPW=Al2 O3 catalyze cumene cracking at 150  C

Figure 6.4. At low loadings of H5 PV2 Mo10 O40 =SiO2 the interaction between the acid and the support results in the opening of the Keggin structure to form triads of edge-linked MoO6 octahedral.(40) Reproduced by permission of Baltzer Science Publishing.

SUPPORTED HETEROPOLY ACIDS 81

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with conversions of approximately 90% and toluene disproportionation at 250  C, with maximum conversions for loadings of 38±50 wt %. Low-temperature isomerization of hexane was also observed. However, magnesia-supported HPW showed no activity for the aforementioned reactions. The application of 29Si and 31P MAS NMR to unsupported and silicasupported HSiMo has provided further information on the properties of these materials.(42) The T1 relaxation time for pure HSiMo was found to be approximately 16 s whereas that for the loaded samples was considerably smaller (Table 6.4), the latter attributed to the interaction of (SiMo12 O40 †4ÿ anions with the support. The HSiMo=SiO2 samples of loadings from 2 to 40 wt % have a resonance line at ÿ74:5 to ÿ74:2  0:1 ppm, similar to that observed (ÿ74:5  0:1 ppm) for the unsupported acid, suggesting that the Keggin unit is only slightly perturbed by the SiO2 support. Interestingly, application of 29Si NMR shows that HSiMo=SiO2 is stable up to 500  C. After heating HPMo=SiO2 to 500  C, only orthorhombic a-MoO3 is evident on the surface and methanol=oxygen produced the redox products, formaldehyde and methyl formate.(43,44) For temperatures of 250±450  C, a mixture of species was found one of which was the monoclinic b-MoO3 , which is the predominant species after heating at 350  C. The calcium and magnesium salts of HPMo have been loaded on SiO2 by ®rst impregnating the support with the cation (in the form of the nitrate) followed by exposure to an aqueous solution of the acid.(45) A comparison of the results obtained in the oxidation of methanol is given in Table 6.5. HPMo=SiO2 and MgPMo=SiO2 were obtained by impregnating the support with an aqueous solution of the acid and HMgPMo12 O40, respectively, whereas HPMo=Mg-SiO2 and HPMo=Ca-SiO2 were prepared by ®rst impregnating with an aqueous solution of the nitrate followed by that of the acid. The results show that the concentration of protons has been decreased in forming the salts, as expected, but no discrimination between the two cations is evident.

TABLE 6.4. T1 Relaxation Times for HSiMo and HSiMo=SiO2 (42) Loading (wt %) 17 25 33 40

T1 (s) 0.3 0.5 1.5 2 16

SUPPORTED HETEROPOLY ACIDS

83

TABLE 6.5. Catalytic Behavior during Methanol Oxidation(45) Sample

HPMo=SiO2

HPMo=Mg-SiO2

MgPMo=SiO2

HPMo=Ca-SiO2

Coverage (mnl)

0.3

0.5

0.3

0.5

0.5

0.3

0.5

S CH3 OCH3 …%† S CH2 O …%† S HCOOCH3 …%† S …CH3 O†2 CH2 …%† S CO; CO2 …%† Conversion (%)

73 13 8 0 6 12.3

74 13 7 0 6 13.0

35 40 17 6 2 9.4

47 30 16 4 3 8.9

44 31 21 2 2 9.7

39 30 16 7 9 8.0

43 28 14 14 1 7.3

Note. T ˆ 250  C; CH3 OH=O2 =He ˆ 7=16=77 mol %; mnl ˆ monolayer

129

Xe NMR has also been applied to HPW=SiO2 .(46) Micropores of Ê are detected in the heteropolyacid overlayers. Because approximately 7 A heteropoly acids are nonporous, the presence of a pore structure in the supported HPW due to the acid itself seems rather puzzling. Two species, H3 PW12 O40 , and H6 P2 W18 O62, were believed to exist on HPW supported on either SiO2 or MCM-41 with the former dominating on loadings from 30 to 50 wt %.(47) On lower loadings the fraction of the latter species increases as the loading decreases. In contrast, use of a methanol solution for the impregnation solvent produced catalysts containing the former species exclusively. Studies of the isomerization of 1-butene on HPW=SiO2 have shown that the selectivity to isobutene at temperatures from 250 to 400  C reaches a maximum at a loading of approximately 23 wt %,(48) similar to the observations noted earlier in this chapter.(24) The 1H MAS NMR chemical shifts increase with the loading (Fig. 6.5).(48) 6.1.4. Alumina The adsorption of HPMo, ammonium 9-molybdophosphate ‰…NH4 †6 P2 Mo18 O62 ], and 9-molybdophosphoric acid (H6 P2 Mo18 O62 ) on galumina during incipient wetness impregnation was studied by application of 31 P and 95Mo NMR.(49) The 95Mo chemical shifts were referenced to sodium molybdate. The 95Mo NMR spectrum of HPMo shows a single peak with a chemical shift of ÿ59 ppm, the former as expected because the 12 molybdenum atoms in HPMo are equivalent. The dimeric 9-molybdophosphoric acid has two types of molybdenum atoms, 6 in the capping position and 12 in the equatorial position, and has a spectrum with two broad peaks. The 31P chemical shifts are referenced to trimethyl phosphite. The 31P spectrum of HPMo shows a single peak at ÿ142:1 ppm. These authors noted that

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Figure 6.5. 1H MAS NMR chemical shifts for HPW=SiO2 .(48) Reproduced by permission of Baltzer Science Publishing.

the pH of the impregnating HPMo solution was initially 1.48 but ultimately reached 4.62 probably due to the release of hydroxyl species from the alumina surface during adsorption of the molybdophosphate anions. HPMo adsorbs on the alumina without decomposition because interaction between the phosphate groups and alumina is prevented by the shell of MoO6 octahedra surrounding the former species. Methanol has been shown to be converted to dimethyl ether and methane on HPW=a-Al2 O3 and HPMo=a-Al2 O3 with the former being more active than the latter for the production of DME.(50) The authors believe that the a-Al2 O3 decreases the deactivation rate of the catalyst. In contrast to the aforementioned, IR spectroscopy studies have suggested that molybdenum-containing heteropoly acids are stable on SiO2 and TiO2 surfaces but at low concentrations decompose on Al2 O3 .(51) The results of further studies of methanol conversion on HPMo and HPW supported on a-Al2 O3 have been reported.(52) 6.1.5. MgF2 1

Magnesium ¯uoride (MgF2 ) has been employed as a support for HPW.(53) A H NMR resonance line at d ˆ 5:1 ppm appears for HPW=MgF2 samples of

SUPPORTED HETEROPOLY ACIDS

85

loadings greater than 3 wt % and increases in intensity with loading. This line is attributed to protons in isolated heteropoly acids bonded with the surface of the support. Crystals of HPW begin to form on the MgF2 for loadings larger than 20 wt %, as evidenced by the appearance of a resonance line at d ˆ 9:3 ppm typical of the pure HPW. The HPW=MgF2 samples of various loadings were employed as catalysts in the oligomerization of isobutene (Fig. 6.6) A maximum in the activity was observed at a loading of approximately 20 wt %. It is of some interest to note that the rates of the oxidation of methane and the formation of products were also found to reach a maximum on HPMo=SiO2 at a loading of approximately 20 wt %.(24)

Figure 6.6. The dependence of the second-order rate constant k of isobutylene oligomerization on the H3 PW12 O40 content on the MgF2 surface.(53) Reprinted from the Journal of Molecular Catalysis, A95, Mastikhin et al., p. 135, copyright 1995, with permission from Elsevier Science.

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6.1.6. SiC The properties of the heteropoly acid H4 PMo11 VO40 supported on silicon carbide have been evaluated and compared with those of SiO2 (aerosil).(54) The thermal stability is higher on the supported samples at low loadings. With SiO2 the decomposition of the acid occurs at a temperature 100  C higher than that for the unsupported acid for an acid content of 0.1 of a theoretical monolayer, decreasing with increasing loading until at 1 theoretical monolayer the decomposition temperature is lower than that of the unsupported acid. With SiC as the support the stabilizing effect is observed up to 2.5 monolayers. The coverage of the support is not complete, even at high acid content. The deposited phase forms blocks of different height covering only 20% of the SiO2 surface and 50±60% of the SiC surface, even when the acid content is equal (SiO2 ) or much higher (SiC) than 1 theoretical monolayer. 6.1.7. ZrO2 Tungstophosphoric acid loaded on silica and on zirconia±silica have been characterized by 31P and 1H MAS NMR, XPS, and FTIR spectroscopies of pyridine adsorption.(55,56) The 31P NMR spectra show that there are two species in the fresh supported catalystsÐcrystalline and interacting HPMoÐa conclusion similar to that for HPMo=SiO2 reported earlier.(30) In agreement with earlier reports,(24,30) the authors ®nd that silica signi®cantly enhances the thermal stability of HPMo to 773 K relative to that of the bulk (673 K) whereas ZrO2 ÿSiO2 leads to a destruction of HPMo at lower temperatures, which is attributed to its stronger interaction with zirconia. With HPMo supported on silica loaded with 1.6 wt % Zr, the thermal stability of HPMo is similar to that of HPMo=SiO2 . However, for higher loadings of Zr, decomposition of HPMo occurs at lower temperatures. As found with HPW=SiO2 ,(48) the 1H NMR lines for HPMo=SiO2 are sharp, indicative of a similar environment for each of the protons. Loading of HPMo on either of the supports results in an increase in acidity. Ethanol solutions of HPW have been employed to impregnate zirconia.(57) This is a technique similar to that employed with carbon(9±11) and with titania(10) by other workers. DTA, XRD, surface area, infrared, Raman, and 31P NMR were employed for characterization purposes as well as the decomposition of isopropanol. Below 673 K HPW exists, on ZrO2 , as distorted intact Keggin species interacting with Zr OH or Zr‡ groups, whereas at 773 K species similar to those in unsupported HPW were found. Above 773 K HPW on ZrO2 decomposed to the constituent oxides. The heteropoly acids have been supported on a variety of oxides, principally ZrO2 , in the presence of a small concentration of a mineral acid such as H2 SO4 and have been employed in the alkylation of isobutane with butene.(58)

SUPPORTED HETEROPOLY ACIDS

87

HPW=ZrO2 =H2 SO4 produced the highest conversion and best selectivity but suffered from reversible deactivation, easily reversed by heating in air at 250 350  C. The deactivation of HPMo=ZrO2 H2 SO4 was due, at least partly, to the degradation of the heteropoly acid. 6.1.8. MCM-41 MCM-41 has been investigated as a support for HPW and HSiW.(47,59,60). With HPW=MCM-41 two supported species were identi®ed, one with Keggin structure (A) and the other with a different structure (B) which the authors believe to be H6 P2 W18 O62 and=or H6 P2 W21 O71, the relative amounts of A and B being dependent on loading, with A dominant.(47) At loadings of 30±50 wt %, the supported material is practically pure structure A. In contrast, 10±40 wt % HPW=MCM-41 prepared in MeOH contains only structure A but these are largely located inside the MCM-41 pores. The B species were found to be approximately eight times more active than the A species in the liquid-phase trans-de-tert-butylation of 2,6-di-tert-butyl-4-methylphenol. In the liquid-phase esteri®cation of 1-propanol and hexanoic acid and in the gas-phase esteri®cation of acetic acid and 1-butanol, HPW=MCM-41 and HSiW=MCM-41 were found to be active catalysts.(59) However, after the reactions large clusters of the acids were found on the external surface of MCM-41, although to a lesser extent with the vapor-phase reactions. Recent studies on HPW=MCM-41 with 1H and 31P MAS NMR, FTIR, and the conversion of 1,3,5-triisopropylbenzene in the vapor phase at 300  C and atmospheric pressure, show that a maximum in acidity and catalytic activity in the aforementioned reaction exists at a loading of approximately 23 wt %.(60) Distortions of the Keggin structure occur at higher loadings. Of interest is the similarity of the loading at the maximum with those reported for HPMo and HPW on a variety of supports.(24,30,37,53) However, because the surface areas of these supports are not identical the rationalization of these similarities is not straightforward. 6.1.9. Zeolite Y Dealuminated zeolite Y has also been investigated as a support for HPW.(61) XRD, DTA=DTG, IR, argon adsorption, and MAS NMR spectroscopy were employed for characterization purposes. The conversion of m-xylene was used to assess the catalytic properties of the supported material. The structure of dealuminated zeolite is retained after introduction of HPW. Dispersion of HPW on the zeolite increases the thermal stability to 650  C in contrast to pure HPW, which decomposes at 610  C. Two types of Keggin anions were identi®ed on the surface. One form, which is predominant at low coverages, results from the

88

CHAPTER 6

interaction of the anions with the OH groups of the support, whereas at high coverages a type resulting from the weak interaction between the Keggin anion and the support is predominant. The selectivity to disproportionation from m-xylene is increased with loading of HPW on the zeolite. 6.1.10. Heteropoly Salts Haber and coworkers have published a series of papers concerned with the properties of H3‡n PVn Mo12ÿn O40 …n ˆ 0 to 3) unsupported and supported on the potassium salt K3 PMo12 O40 .(62±65) As discussed in the chapter on microporosity, it was shown in 1984 that salts of the heteropoly acids formed from monovalent cations have relatively high surface areas and microporous structures.(66) The loadings employed in the aforementioned work correspond to one monolayer coverage.(62±65) As vanadium is added, the ratio of formaldehyde to dimethyl ether from methanol oxidation at 533 K increases but is similar for both the unsupported and supported materials. However, with increase in vanadium the activity of the supported acid increases whereas that of the unsupported acid decreases (Fig. 6.7) The authors argue that the increased thermal stability of the supported acids results from an epitaxial relationship with the support.(62±65) Grinding of mechanical mixtures of HPW and its cesium salt has been found to disperse the acid on the salt, producing ef®cient catalysts for the isomerization of n-butane.(67) As noted earlier in this chapter and in the chapter on microporosity, work by the author of this volume has shown that the monovalent salts of HPW possess relatively high surface areas and microporous structures.(66) The effectiveness of the grinding technique as compared to the separate constituents of the mixture is illustrated in Fig. 6.8, which also shows results for the catalyst formed from wet impregnation of the cesium salt with its parent acid. 6.1.11. Clays HPW supported (20 wt %) on a montmorillonite clay (K-10) produced a conversion of 71% in the reaction of tert-butyl alcohol and methanol at 85  C with a selectivity to methyl tert-butyl ether of 99%.(68) On the support itself the conversion and selectivity are 56 and 70%, respectively. The same catalyst produced a conversion of 69% in the reaction of phenethyl alcohol with methanol at 70  C and phenethyl methyl ether as the only product.(69) Heteropoly acids supported on silica have been compared with clays for the hydrolysis of ethyl acetate in the liquid phase.(70) Three lamellar silicatesÐa natural montmorillonite with high alkaline metal content, an activated montmorillonite strongly delaminated with low alkaline metal content, a kenyaite (a lamellar sodium silicate) together with a mesoporous hexagonal silicaÐwere investigated as supports for HPW.(71) With the lamellar

SUPPORTED HETEROPOLY ACIDS

89

Figure 6.7. Oxidation of methanol at 533 K on H3‡n PVn Mo12ÿn O40 …n ˆ 0 to 3). s, pure acid; d, acid supported on K3 PMo12 O40 .(62) Reproduced by permission of The Royal Society of Chemistry.

structures, which contain relatively large quantities of alkaline cations, the Keggin anions are decomposed, as seen from IR spectroscopy. The authors speculate that the pH at the interface with the impregnating solution is increased by the alkaline cations leading to the formation of lacunary structures of the acid. In contrast, on activated clay or hexagonal mesoporous silica (HMS) with high surface areas and low concentrations of alkaline cations, the Keggin structure is retained and well dispersed where the quantity of supported material is less than the monolayer coverage. On the activated montmorillonite and for low loadings the charge compensating cations of the clay were found to exchange with the protons of HPW, thus decreasing the catalytic activities.

90

CHAPTER 6

Figure 6.8. Catalytic activity of Csx H3ÿx PW in butane isomerization at 473 K. I, activities of the sample prepared by wet impregnation; P, activities of the sample prepared by precipitation.(67) Reproduced by permission of Baltzer Science Publishing.

6.1.12. Polymers The entrapment of heteropoly acids in polymeric matrices, while strictly speaking not analogous to the use of a high-area support, is nevertheless of interest from a catalytic viewpoint and as such is discussed here. Heteropoly acids have been supported on polymers such as poly (4vinylpyridine).(72) The anchoring of H6 ‰Co…II†W12 O40 Š produced an increased stability in the oxidation of 1,4-butanediol. Other workers have introduced heteropoly oxometalates into the polymer matrix for electrochemical purposes.(73±76) Various polymeric materials, including polypyrrole,(77) polyacetylene,(78) polysulfone,(79±81) poly(ethersulfone),(80,81) and poly(phenylene oxide),(81) have been doped with HPMo and tested for the conversion of ethanol. At 503 K the yield of acetaldehyde from ethanol with HPMo-doped (20%) polyacetylene was approximately 40 times greater than that obtained with the acid itself.(78) While HPMo was found to be evenly distributed over the cross section of the polymer ®lm, its concentration at the surface of the polymer was higher than that in the bulk.(78) In the conversion of 2-propanol, HPMo blended with polysulfone produced a higher yield for acetone but a lower yield for propylene than the parent acid.(82) Dimethylformamide, used as a solvent in the preparation of the blended composition, strongly adsorbed on the acidic sites of the acid, decreasing the number of acidic sites. The authors suggest that the increase in the oxidation

SUPPORTED HETEROPOLY ACIDS

91

activity resulted from the uniformly and ®nely distributed acid in the ®lm. The latter appears to be in contrast to that observed with polyacetylene.(77) HPW and HPMo combined with polyaniline, has been prepared by chemical polymerization of aniline in the presence of HPW and protonation of polyemeraldine base with HPW.(83±86) The acid is incorporated in the polyaniline which is simultaneously protonated by what the authors refer to as the most acidic proton of the acid. Not surprisingly the two preparative methods yield solids with different catalytic properties. The former method distributes the acid evenly throughout the surface and bulk whereas the latter preparation produces a solid with the acid bound to the surface only and consequently more accessible to the reactant molecules in alcohol conversion processes. A membrane preparation technique has been employed to form HPMopolymer composite ®lm catalysts by blending these materials using dimethylformamide or methanol±chloroform mixtures.(87,88) These were tested in the conversion of ethanol(86) and the liquid-phase synthesis of tert-butanol from isobutene and water.(87) A polyaniline-supported HSiW after pretreatment in helium at 373 K was found to have a low activity (less than 2%) in the generation of MTBE from methanol and isobutene.(89) However, after activation of the catalyst in air at 473 K the activity increased to 15±20%. Selectivities to MTBE were approximately 60 and 85%, respectively. In the latter treatment, protons apparently migrate from the bulk of the polymer matrix to the surface. Electrically conducting polyaniline doped with H3 PMo12 O40 , H4 PMo11 VO40 , H5 PMo10 V2 O40 , and H6 PMo9 V3 O40 were synthesized by electrochemical polymerization and reversible redox systems were found for both the polymer and the acids.(90) Catalytic membranes with HPW entrapped in poly(vinyl alcohol) on porous ceramic plates were examined for their separative and catalytic properties.(91) The method employed to affect the cross-linking was found to be important; glutaraldehyde as the cross-linking agent proved to be superior. Unsubstituted and dimethoxy-substituted aromatic poly(azomethines) doped with HPW and HPMo have been examined for the oxidation of propylene.(92) The acids were shown to retain their structure and be molecularly dispersed. The predominant product from propylene was hexadiene, in contrast to acrolein with crystalline HPMo. 6.2. FORMATION OF HETEROPOLY ACIDS FROM OXIDES SUPPORTED ON SILICA Although, strictly speaking, this topic is not concerned with the spreading of a heteropoly acid on a support, nevertheless the numerous observations that the

92

CHAPTER 6

interactions between MoO3 and SiO2 are suf®ciently strong to produce HSiMo, among other species, bound to the support, are of suf®cient importance and relevancy to warrant the inclusion of these reports within this chapter. HSiMo was identi®ed on the surface of MoO3 =SiO2 at low loadings with polymolybdates and MoO3 whiskers at higher loadings.(93) The calcination of Mo(VI)=SiO2 catalysts at 500  C in air produced changes in the UV±visible diffuse re¯ectance spectra which were attributed to the destruction of a molybdosilicate structure.(94) The bands of octahedral molybdenum disappear but reappear when the sample is exposed to water vapor at 25  C. Laser Raman spectroscopy of catalysts prepared by the impregnation of SiO2 with aqueous solutions of ammonium paramolybdate followed by drying at 130  C for 12 h and calcination at 500  C in air for 15 h indicated that silicomolybdic anion species were present on the surface at loadings as low as 1% Mo.(95,96) Subsequent SIMS investigations con®rmed these results.(97) H4 SiMo12 O40 =SiO2 catalysts and MoO3 =SiO2 catalysts prepared at various values of pH were examined by IR and laser Raman spectroscopy and were employed in the oxidation of methane (Table 6.2).(26) Both the conversion of CH4 and the product distribution approach those of the molybdosilicic acid as the pH of the preparative solution decreases. The laser Raman spectra of the MoO3 =SiO2 samples prepared at pH values of 2 and 7 clearly show the bands of HSiMo. At a pH of 11, bands of the heptamolybdate species can be seen and bands at 1000 and 820 cmÿ1 observed with samples prepared at pH 7 and 11 are attributed to MoO3 crystallites. The existence of molybdosilicic acid is also shown by IR and laser Raman analyses of the ®ltered solution from the washing of a calcined sample with acetonitrile to extract the supported HSiMo species selectively. HSiMo is apparently synthesized during the formation of the MoO3 =SiO2 catalyst at pH 2 and 7. However, at least three speciesÐHSiMo, the heptamolybdate and MoO3 crystallitesÐexist on MoO3 =SiO2 depending on the pH of the solution, Mo loading and calcination conditions. Mo=SiO2 catalysts prepared by the impregnation of aerosil silica with ammonium heptamolybdate were characterized by XPS, laser Raman spectroscopy, surface potential measurements, and acidimetric titration.(98) The activity in methane oxidation was also evaluated. Results similar to the aforementioned were obtained. For low loadings of Mo, HSiMo was the principal Mo species. At Mo loadings of 5±10 wt % polymolybdate species on the HSiMo layer were observed whereas at higher loadings crystalline MoO3 appeared. In the oxidation of methane by N2 O the yield and selectivity to formaldehyde increased with the quantity of HSiMo. At higher Mo content the predominant reaction was total oxidation. Regardless of the Mo content with oxygen as the oxidant total oxidation was observed. These observations are similar to those reported earlier.(22,24,26)

SUPPORTED HETEROPOLY ACIDS

93

Dimethylformamide solutions of tetrabutylammonium hexamolybdate (TBA2 Mo6 O19 ) were used to impregnate silica.(99) Exposure to water produces catalysts with acidic properties whereas those not exposed show a redox catalytic behavior. The activity increases with decreasing molybdenum content. Evidence for trimolybidic units bonded to silica through Si±O±Mo bridges as well as HSiMo is obtained. The authors argue that the latter play only a minor role in the catalytic oxidation of methanol. A catalyst with HPMo on silica, at a loading equivalent to or slightly larger than that for a monolayer, on heating to temperatures higher than 300  C, formed species believed to be the dehydrated or defect KU, formed on removal of an oxygen ion in the process of the desorption of water, and=or MoO3, the former of which was reconverted to Keggin species on exposure to water vapor.(31) Studies of the partial oxidation of methane on MoO3 =SiO2 catalysts have also provided evidence for a semiquantitative correspondence between results obtained earlier on HPMo=SiO2 and those for the aforementioned catalyst.(22,24,100). REFERENCES 1. Y. Izumi and K. Urabe, Chem. Lett. 663 (1981). 2. Y. Izumi, Res. Chem. Intermed. 12, 461 (1998). 3. F. Lefebvre, P. Dupont, and A. Auroux, React. Kinet. Catal. Lett. 55, 3 (1995); F. X. Liu-Cai, B. Sahut, E. Faydi, A. Auroux, and G. HerveÂ, Appl. Catal. A 185, 75 (1999). 4. M. A. Schwegler, H. van Bekkum, and N. A. de Munck, Appl. Catal. 74, 191 (1991). 5. P. Dupont, J. C. VeÂdrine, E. Paumard, G. Hecquet, and F. Lefebvre, Appl. Catal. A 129, 217 (1995). 6. P. Dupont and F. Lefebvre, J. Mol. Catal. A 114, 299 (1996). 7. M. A. Schwegler, P. Vinke, M. van der Eijk, and H. van Bekkum, Appl. Catal. A Gen. 80, 41 (1992). 8. Y. Wu, X. Ye, Y. Yang, X. Wang, W. Chu, and Y. Hu, Ind. End. Chem. Res., 35, 2546 (1996). 9. L. R. Pizzio, C. V. Caceres and M. N. Blanco, J. Colloid Interface Sci. 190, 318 (1997). 10. G. Vazquez, M. N. Blanco and C. V. Caceres, Catal. Lett. 60, 205 (1999). 11. L. R. Pizzio, C. V. Caceres, and M. N. Blanco, Appl. Catal. A Gen. 167, 283 (1998). 12. S. Kasztelan, E. Payen, and J. B. Moffat, J. Catal. 128, 479 (1991). 13. A. Spojakina, S. Damyanova, D. Shopov, T. Shokhireva, and T. Yurieva, React. Kinet. Catal. Lett. 77, 333 (1985). 14. S. Damyanova, J. L. G. Fierro, and A. Spojakina, React. Kinet. Catal. Lett. 56, 321 (1995). 15. J. C. Edwards, C. Y. Thiel, B. Benac, and J. F. Knifton, Catal. Lett. 51, 77 (1998). 16. S. Damyanova and J. L. G. Fierro, Chem. Mater. 10, 871 (1998). 17. S. Damyanova, M. L. Cubeiro and J. L. G. Fierro, J. Mol. Catal. 142, 85 (1999). 18. R. T. Sebulsky and A. M. Henke, Ind. Eng. Chem. Process Des. Dev. 10, 272 (1971). 19. I. Mochida, T. Nakashima, and H. Fujitsu, Bull. Chem. Soc. Jpn. 57, 1449 (1984). 20. Y. Izumi, R. Hasebe, and K. Urabe, J. Catal. 84, 402 (1983). 21. R. Fricke and G. Ohlmann, J. Chem. Soc. Faraday Trans. 82, 263 (1986).

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59. M. J. Verhoef, P. J. Kooyman, J. A. Peters, and H. van Bekkum, Microporous Mesoporous Mater., 27, 365 (1999). 60. A. Ghanbari-Siahkali, A. Philippou, J. Dwyer, and M. W. Anderson, Appl. Catal. A 192, 57 (2000). 61. K. Pamin, A. Kubacka, Z. Olejniczak, J. Haber, and B. Sulikowski, Appl. Catal. A 194±195, 137 (2000). 62. K. Bruckman, J. Haber and E. M. Serwicka, Faraday Discuss. Chem. Soc. 87, 173 (1989). 63. K. Bruckman, J. Haber, E. M. Serwicka, E. N. Yurchenko, and T. P. Lazarenko, Catal. Lett. 4, 181 (1990). 64. E. M. Serwicka, K. Bruckman, J. Haber, E. A. Paukshtis, and E. N. Yurchenko, Appl. Catal. 73, 153 (1991). 65. K. Bruckman, J.-M. Tatibouet, M. Che, E. Serwicka, and J. Haber, J. Catal. 139, 455 (1993). 66. J. B. McMonagle and J. B. Moffat, J. Colloid Interface Sci. 101, 479 (1984). 67. P.-Y. Gayrand, N. Essayem, and J. C. Vedrine, Catal. Lett. 56, 35 (1998). 68. G. D. Yadav and N. Kirthivasan, J. Chem. Soc. Chem. Commun. 203 (1995). 69. G. D. Yadav and V. V. Bokade, Appl. Catal. A 147, 299 (1996). 70. Y. Izumi, K. Urabe, and M. Onaka, Microporous Mesoporous Mater. 21, 227 (1998). 71. F. Marme, G. Coudurier, and J. C. Vedrine, Microporous Mesoporous Mater. 22, 151 (1998). 72. K. Nomiya, H. Murasaki, and M. Miwa, Polyhedron 5, 1031 (1986). 73. B. Keita and L. Nadjo, J. Electroanal. Chem. 240, 325 (1988). 74. B. Keita, L. Nadjo, and J. P. Haeussler, J. Electroanal. Chem. 243, 481 (1988). 75. B. Keita, K. Essaadi, and L. Nadjo, J. Electroanal. Chem. 259, 127 (1989). 76. A. Mahmoud, B. Keita, L. Nadjo, O. Oung, R. Constant, S. Brown, and Y. de Kouchkovsky, J. Electroanal. Chem. 463, 129 (1999) and references therein. 77. J. Pozniczek, A. Bielanski, I. Kulszewicz-Bajer, M. Zagorska, K. Kruczala, K. Dyrek, and A. Pron, J. Mol. Catal. 69, 223 (1991). 78. J. Pozniczek, I. Kulszewicz-Bajer, M. Zagorska, K. Kruczala, K. Dyrek, K. Bielanski, and A. Pron, J. Catal. 132, 311 (1991). 79. I. K. Song, S. K. Shin, and W. Y. Lee, J. Catal. 144, 348 (1993). 80. I. K. Song, J. K. Lee, and W. Y. Lee, Appl. Catal. A 119, 107 (1994). 81. J. K. Lee, I. K. Song, and W. Y. Lee, J. Mol. Catal. A 120, 207 (1997). 82. J. K. Lee, I. K. Song, W. Y. Lee, and J. J. Kim, J. Mol. Catal. A 104, 311 (1996). 83. M. Hasik, W. Turek, E. Stochmol, M. Lapkowski, and A. Pron, J. Catal. 147, 544 (1994). 84. M. Hasik, A. Pron, J. Pozniczek, A. Bielanski, Z. Piwowarska, K. Kruczala, and R. Dziembaj, J. Chem. Soc. Faraday Trans. 90, 2099 (1994). 85. M. Hasik, J. Pozniczek, Z. P. Wowarska, R. Dziembaj, A. Bielanski, and A. Pron, J. Mol. Catal. 89, 329 (1994). 86. R. Dziembaj, A. Malecka, Z. Piwowarska, and A. Bielanski, J. Mol. Catal. A 112, 423 (1996). 87. G. I. Park, S. S. Lim, J. S. Choi, I. K. Song, and W. Y. Lee, J. Catal. 178, 378 (1998). 88. S. S. Lim, Y. H. Kim, G. I. Park, W. Y. Lee, I. K. Song, and H. K. Youn, Catal. Lett. 60, 199 (1999). 89. A. Bielanski, R. Dziembaj, A. Malecka-Lubanska, J. Pozniczek, M. Hasik, and M. Drozdek, J. Catal. 185, 363 (1999). 90. M. Barth, M. Lapkowski, and S. Lefrant, Electrochim. Acta 44, 2117 (1999). 91. Q. Liu, P. Jia, and H. Chen, J. Membr. Sci. 159, 233 (1999). 92. W. Turek, E. Stochmal-Pomarzanska, A. Pron, and J. Haber, J. Catal. 189, 297 (2000). 93. A. Castellan, J. C. J. Bart, A. Vaghi, and N. Giordano, J. Catal. 42, 162 (1976). 94. K. Marcinkowska, L. Rodrigo, P. C. Roberge, and S. Kaliaguine, J. Mol. Catal. 33, 189 (1985). 95. L. Rodrigo, K. Marcinkowska, A. Adnot, P. C. Roberge, S. Kaliaguine, J. M. Stencel, L. E. Makovsky, and J. R. Diehl, J. Phys. Chem. 90, 2690 (1986).

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7 MICROPOROSITY

7.1. INTRODUCTION As noted elsewhere in this volume, the heteropoly acids have been found to have low surface areas (10 m2 gÿ1 ) and no porous structure, although polar molecules such as ammonia are capable of penetrating into the bulk structure, that is, between the cations and anions.(1,2) Not surprisingly, the protons can be replaced by other cations either through the slow addition of an aqueous solution of a compound containing the desired cation to an aqueous solution of the heteropoly acid or by an appropriate ion exchange process. The latter will be discussed in detail later in this chapter. Thus, for example, the ammonium salt of 12-tungstophosphoric acid can be prepared by the slow addition of an aqueous solution of ammonium nitrate to an aqueous solution of 12-tungstophosphoric acid. As will be discussed later in this chapter, the use of stoichiometric quantities of the preparative reagents does not, in general, yield a stoichiometric product, but rather one that contains residual quantities of protons.(2) It is also worth noting that addition of the acidic solution to the solution containing the desired cation is not recommended because the Keggin structure may be partially (or completely) destabilized at values of pH even as low as 4.(3) Further, rapid addition of the cationic solution to the acidic solution is similarly disadvantageous. Earlier work has shown that 12-tungstophosphoric acid (HPW) is active and selective in the conversion of methanol to hydrocarbons, predominantly ole®nic in nature.(4,5) However, surprisingly, the ammonium salt of HPW prepared by precipitation is more active than the parent acid and, in contrast, yields largely saturated hydrocarbons.(5,6) Although a number of factors undoubtedly play a role in producing this result, the ammonium 12-tungstophosphate was found to have a surface area 10±15 times larger than that of the parent acid. 97

98

CHAPTER 7

As a consequence of the aforementioned observations, the salts of monovalent cations of the group 1 alkali metals have been prepared from a series of heteropoly acids, including 12-tungstophosphoric, 12-molybdophosphoric, 12tungstosilicic, and 12-tungstoarsenic acids and examined for their morphological, surface, and catalytic properties. 7.2. MICROPOROSITY OF SALTS PREPARED FROM THE MONOVALENT CATIONS OF THE GROUP 1 ALKALI METALS AND AMMONIUM AND RELATED CATIONS As noted above, these have been prepared by precipitation from aqueous solution, taking advantage of the very low solubilities in water. The morphologies have been evaluated by a variety of methods, although the appropriate analyses of adsorption±desorption isotherms may generate the most useful data. 7.2.1. Surface Areas and Pore Structures from the Analysis of Nitrogen Adsorption±Desorption Isotherms Typical nitrogen adsorption±desorption isotherms measured on the cesium, ammonium, and potassium salts held at 77 K are characterized by a sharp increase in the quantities of nitrogen adsorbed at low relative pressures (Figs. 7.1 and Ê ). The 7.2),(1,7±13) which can be attributed to the ®lling of micropores (d < 20 A Ê > d > 20 A Ê ) in cesium 12-tungstophosphate can existence of mesopores 500 A be inferred from the presence of hysteresis loops in their isotherms. The surface areas, as calculated from the adsorption branch of the nitrogen adsorption-desorption isotherms by application of the Brunauer±Emmett±Teller in®nite layer multilayer theory, generally increase with increase in the diameter of the monovalent cations employed in the preparation of the salts (Table 7.1, Figs. 7.3 and 7.4). With all of the heteropoly anions examined, the surface areas of the sodium salts are very small. The surface areas of all of the ammonium and cesium salts are high as are those of the potassium salts with the exception of those for 12-tungstosilicic and 12-molybdosilicic acid. Although the surface areas for all of the rubidium salts are not available, those prepared from 12-tungstosilicic and 12tungstoarsenic acids are relatively high in comparison with that for the salt of molybdosilicic acid. Increases in the diameters of the cations lead to a reversion to small surface areas although it must be noted that the organic cations are signi®cantly larger than the alkali metal cations. It should be recalled that the walls of pores smaller than the size of the adsorbate molecule, in the present case, N2 , will not, of course, contribute to the surface area and additionally the nitrogen molecules, being nonpolar, will not penetrate into the bulk structure of the MOCC.

MICROPOROSITY

99

Figure 7.1. Nitrogen adsorption±desorption isotherms (78 K) for the high-surface-area cesium, ammonium, and potassium salts of HPW.(7)

While it is clear from Table 7.1 that the surface area is dependent on the size of the monovalent cation (Fig. 7.3), it is evident that the nature of the heteropoly anion, and in particular the size of the central atom, is also a factor (Fig. 7.4). The values of the BET C parameters for the salts of the monovalent cations are high for those solids with high surface areas (Table 7.2), which implies high

100

CHAPTER 7

Figure 7.2. Nitrogen adsorption±desorption isotherms (78 K) for the high-surface-area cesium, ammonium, and potassium salts of HPMo.(7)

values for the energy of adsorption in the ®rst adsorbed layer. This is as would be expected for pores whose sizes are of the same order of magnitude as the adsorbate and therefore is indicative of the presence of micropores. Micropore size distributions, as obtained by application of the t-plot method, are shown for the ammonium salts of the various acids (Figs. 7.5±7.8). Because

MICROPOROSITY

101

TABLE 7.1. Surface Areasa of Metal±Oxygen Cluster Compounds(1,7±13) Anion Cation ‡

H Na‡ K‡ NH4 ‡ Rb‡ Cs‡ MeNH3 ‡ Me4 N‡ a b

diamb

PW

PMo

SiW

SiMo

AsW

Ð 1.90 2.66 2.86 2.96 3.30 4 Ð

8 3.7 90.0 128.2 Ð 162.9 3.0 4.5

8 3.5 39.9 193.4 Ð 145.5 1.3 Ð

3.2 1.3 3.3 116.9 116.3 153.4 3.0 Ð

8 Ð 2 102 13 95 Ð Ð

6.5 Ð 46.0 82.1 101.4 65.4 2.1 Ð

BET N2 77.4 K; m2 gÿ1 . Ê. A

Ê in these solids, the the radii of the pores range from approximately 7 to 15 A pores can be considered as large micropores or small mesopores. The distribution Ê , which is not seen in for NH4 PMo shows an increase in Dv=Dr at a radius of 20 A the remaining distributions and is indicative of the existence of pores that are suf®ciently large so as to ®ll by capillary condensation and thus produce a hysteresis loop as seen in the adsorption±desorption isotherm for this solid (Fig. 7.2). Values for r, the mean micropore radius, for the high-surface-area microporous salts (Table 7.3) correlate linearly with the number of adsorbed layers (n), the latter calculated from the ®nite-layer BET equation (Fig. 7.9) As expected for a microporous solid, the values of n fall in the range from 1 to 3, where nitrogen is employed as the absorbate. The slope of the line in Fig. 7.9 is approximately Ê , in good agreement with the diameter of the nitrogen molecule. Because the 3.5 A values for r and n are obtained by independent calculational methods, although these are based on the experimentally determined adsorption±desorption isotherm, the linear correlation and the realistic value for the slope provide evidence for the internal self-consistency of the two procedures. 7.2.2. The Source of the Microporosity The total micropore volumes, as obtained from the appropriate analysis of the nitrogen adsorption isotherms using the t-plot method, are, as expected, and regardless of the elemental composition of the anion, negligible for the parent acids and the sodium salts but signi®cant for those of potassium, ammonium, rubidium, and cesium (Figs. 7.10±7.13). With further increase in the size of the cation to methylammonium, the micropore volumes again become insigni®cant. The intensities of the (110) XRD re¯ection relative to that of the (222) plane are

102

CHAPTER 7

Figure 7.3. Surface areas as a function of the diameter of the cations for 12-heteropoly salts.(7±13)

MICROPOROSITY

103

Figure 7.4. Surface area as a function of the diameter of the anionic central atom for 12-heteropoly salts.(13)

high for the parent acids and the sodium and methylammonium salts but have considerably smaller values for those salts with signi®cant micropore volumes. Interstitial voids, separated from each other by the terminal oxygen atoms of the heteropoly anions, are present in the parent acids of cubic structure, but passage of molecules, such as those of nitrogen, between the interstitial voids is effectively prevented by the aforementioned oxygen atoms. Because these oxygen atoms lie in the (110) plane of the crystal, the smaller values of the (110)=(222) XRD intensity ratio found where porous structures are present may be attributed to a translation and=or rotation of the anions so that the terminal oxygen atoms are removed, at least partially, from the (110) plane. This permits the interstitial voids to become interconnected thereby permitting molecules, such as those of nitrogen, to pass between the voids. Although the pores thus formed

104

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TABLE 7.2. BET C Parametersa for Metal±Oxygen Cluster Compounds(1,7±13) Anion Cation ‡

H Na‡ K‡ NH4 ‡ Rb‡ Cs‡ MeNH3 ‡ Me4 N‡ a

PW

PMo

SiW

SiMo

AsW

Ð 18 2430 760 Ð 1720 Ð Ð

Ð 142 1070 877 Ð 568 Ð Ð

100 473 127 1919 489 Ð Ð Ð

290 380 900 8000 1500 6500

58 Ð 22 1292 149 1359 292 Ð

Calculated from BET in®nite-layer equation.

Figure 7.5. Micropore size distribution for NH4 PW.(1,7±13) Reprinted from the Journal of Molecular Catalysis, 52, Moffat, p. 169, copyright 1989, with permission from Elsevier Science.

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Figure 7.6. Micropore size distribution for NH4 PMo.(1,7±13) Reprinted from the Journal of Molecular Catalysis, 52, Moffat, p. 169, copyright 1989, with permission from Elsevier Science.

are assumed to be cylindrical, it is likely that the channels are somewhat constricted at regular intervals. 7.2.3. The Dependence of Morphological Properties of Salts Prepared from Group 1 Monovalent Cations on the Source of the Cations and the Stoichiometry of the Salt The surface areas and pore structures of the salts of HPW and HPMo prepared from group 1 monovalent cations have been shown to be dependent on the nature of the cation, the source of the cation, the preparative stoichiometry, the peripheral metal atom of the anion, and the temperature to which the salt has been exposed.(1,7±9,11±14) For example, NH4 PW, prepared as stoichiometric, shows

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Figure 7.7. Micropore size distribution for NH4 SiW.(1,7±13) Reprinted from the Journal of Molecular Catalysis, 52, Moffat, p. 169, copyright 1989, with permission from Elsevier Science.

signi®cantly different BET total surface areas, those due to meso=macro and micropores as well as the volume and radius of the micropores for different cation sources (Table 7.4).(14) The largest BET areas and the largest areas due to micropores are obtained with ammonium nitrate and chloride, with which the aforementioned properties are virtually identical, whereas the sulfate possesses considerably smaller values of the morphological parameters. Although the carbonate may not be as advantageous from a morphological viewpoint, the ability to minimize the anion content from the ®nal preparation may be suf®cient justi®cation for its use. The effects of pretreatment temperature and stoichiometry with NH4 PW and NH4 PMo are summarized in Tables 7.5 and 7.6, respectively.(14) In general, as the cation=H‡ preparative ratio is increased, the surface areas increase, as would be expected because the numbers of residual protons in the salts are decreasing. Several observations are worth noting in these latter two tables. The most substantial changes in total surface area and that due to micropores as well as the micropore volume occur with the increase in preparative cation=H‡ ratio from de®cit (ÿ15%) to stoichiometric values. In contrast, the average micropore radius changes relatively little. These observations appear to be consistent with the

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107

Figure 7.8. Micropore size distribution for NH4 AsW.(1,7±13) Reprinted from the Journal of Molecular Catalysis, 52, Moffat, p. 169, copyright 1989, with permission from Elsevier Science.

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Figure 7.9. Average radius (r) versus value of n from ®nite-layer BET equation for microporous heteropoly oxometalates.(1,7±13) Reprinted from the Journal of Molecular Catalysis, 52, Moffat, p. 169, copyright 1989, with permission from Elsevier Science.

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109

Ê) TABLE 7.3. Mean Micropore Radii (A of High-Surface-Areas MOCC(1,7±13) Anion Cation ‡

K NH4 ‡ Rb‡ Cs‡

PW

PMo

SiW

AsW

8.8 10.3 Ð 13.9

9.3 13.0 Ð 14.3

Ð 9.5 10.3 10.5

11.2 11.3 9.4 9.9

tentative explanations for the existence of the micropores. The protons and the orientation of the terminal oxygen atoms separate the interstitial voids from each other, preventing the formation of the pore channels. As the protons are replaced by larger cations and the terminal oxygen atoms rotate from their blocking positions, the pore sizes are expected to change relatively little while the micropore volumes increase, as observed.

Figure 7.10. Correlation of estimated micropore volume and XRD [110]=[222] intensity ratio with monovalent cation diameter for the 12-tungstophosphates.(7)

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Figure 7.11. Correlation of estimated micropore volume and XRD [110]=[222] intensity ratio with monovalent cation diameter for the 12-molybdophosphates.(7)

With increase in pretreatment temperature the BET and micropore surface areas and volumes decrease while the meso=macropore surface areas generally increase. Although the source of this phenomenon is not clear, it is possible that the higher temperature alters the mobility, both translational and rotational, of both the cations and anions.

7.3. MICROPOROSITY OF SALTS OF OTHER CATIONS 7.3.1. Microporosity of Salts Prepared from the Monovalent Cations of the Group 11 and 13 Elements The synthesis of microporous solids is not restricted to salts of the cations of the elements of the 1A group of the periodic table.(15,16,18,19) The silver (group 11) and thallium (group 13) salts of HPW and HSiW have been prepared with microporous structures and high surface areas (Table 7.7). For reasons not yet understood, the silver salt of HPMo, prepared in a similar manner to the

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111

Figure 7.12. Correlation of estimated micropore volume and XRD [110]=[222] intensity ratio with monovalent cation diameter for the 12-tungstosilicates.(12)

aforementioned salts, was found to have a low surface area and no evidence of microporosity. The micropore volumes of the silver and thallium salts of HPW and HSiW also show the approximately reciprocal relationship with the I[110]=I[222] XRD ratios noted earlier with the salts prepared from the cations of the group 1 elements (Fig. 7.14). The analysis of adsorption±desorption isotherms of the thallium salts of HSiW and HPMo has been illustrated by application of a variety of calculational methods.(15) 7.3.2. Stoichiometric and Nonstoichiometric Salts of Groups 11 and 13 and Their Morphological Properties PAS FTIR studies have shown that the microporous salts synthesized from stoichiometric quantities of the preparative reagents contain residual protons.(2,17)

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Figure 7.13. Correlation of estimated micropore volume and XRD [110]=[222] intensity ratio with monovalent cation diameter for the 12-tungstoarsenates.(13)

TABLE 7.4. Effect of the Source of the Cation on the Morphological Properties of NH4 PW(14) Source of NH4 ‡ Nitrate Chloride Carbonate Sulfate a

SBET a

Smeso macro a

Smicro a

Vmicro b

rmicro c

165 161 134 63

20 15 24 38

147 145 110 26

61.8 58.9 48.8 11.0

9.6 8.9 9.9 11.6

SBET , Smeso macro , and Smicro refer to the BET surface area, that due to meso- and macrophores, and that due to micropores, respectively …m2 gÿ1 †. Vmicro refers to the micropore volume (103 cm3 liquid gÿ1 . c Ê ). rmicro is the average micropore radius (A b

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113

TABLE 7.5. The Effect of the Preparative Stoichiometry and Pretreatment Temperature of NH4 PW on Its Morphological Properties(14) Pretreat. temp … C† 200 300 400

Stoichiometrya

SBET b

SMM b

SMP b

VMP c

rMP d

E S D E S D E S D

140 120 35 130 104 14 111 93 29

29 30 21 42 46 14 59 63 18

110 90 15 87 57 0 52 34 13

48.9 40.0 5.8 39.7 27.2 0.0 25.9 15.0 5.2

10.0 10.7 14.9 10.4 11.6 11.8 12.3 8.5

a

E, D, and S refer to stoichiometric excess of cation (15%), stoichiometric de®cit of cation …ÿ15%†, and stoichiometric. SBET , BET surface area; SMM , area due to meso- and macropores, SMP , area due to micropores, m2 gÿ1 . c VMP , micropore volume (103 cm3 liquid gÿ1 ). d Ê ). rMP, average micropore radius (A b

Salts prepared from nonstoichiometric quantities of the preparative reagents have microporous properties and in particular, their micropore volumes, suggest that the excess cations or anions, depending on the nature of the departure from stoichiometry, may partially block the channels which produce the microporosity.(14) Additional studies of the effect of deliberate variations in the stoichiometries of the preparative reagents for the thallium and silver salts show that, for example, the maximum in the pore size distribution for the thallium salt of HPW TABLE 7.6. The Effect of the Preparative Stoichiometry and Pretreatment Temperature of NH4 PMo on Its Morphological Properties(14) Pretreat. temp … C† 200 300

a

Stoichiometrya

SBET b

SMM b

SMP b

VMP c

rMP d

E S D E S D

190 206 66 161 156 21

38 59 34 66 93 22

157 139 33 94 62 0

65.8 63.3 14.5 43.5 29.1 0

10.2 11.1 12.0 11.8 13.4

E, D, and S refer to stoichiometric excess of cation (15%), stoichiometric de®cit of cation …ÿ15%†, and stoichiometric. SBET , BET surface area; SMM , area due to meso- and macropores, SMP , area due to micropores, m2 gÿ1 . c VMP , micropore volume (103 cm3 liquid gÿ1 ). d Ê ). rMP, average micropore radius (A b

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TABLE 7.7. Morphological Properties of Stoichiometric Salts Prepared from Representative Elements of Groups 11 and 13(15,16,18,19) Sample Ag3 PW12 O40 Ag3 PMo12 O40 Ag4 SiW12 O40 Tl3 PW12 O40 Tl3 PMo12 O40 Tl4 SiW12 O40 a

CBET

SBET …m2 gÿ1 †

na

VMP …ml gÿ1 †

rMP Ê) (A

7,400 14,000 35,600 2,700 900 5,000

100.9 1.5 106.0 131.6 157.0 97.5

1.46 Ð 1.38 2.08 2.68 1.50

0.037 Ð 0.040 0.045 0.046 0.036

7.9 Ð 7.6 8.2 8.4 8.0

n refers to the numbers of adsorbed layers calculated from the ®nite BET equation.

Figure 7.14. Micropore volumes and XRD I[110]=I[222] ratios for stoichiometric silver and thallium salts of HPW together with values reported earlier for other cations.(15,16,18,19) Reprinted with permission from Parent and Moffat.(16) Copyright 1996 American Chemical Society.

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115

diminishes in size but not in position, suggesting that changes in the relative quantities of [cations]=[anions] primarily alter the micropore volume rather than the diameters of these pores (Fig. 7.15).(15,16,18,19) The mean micropore radii of the silver and thallium salts of HPW vary relatively little as the preparative cation=proton ratio is varied from 0.5 to 1.5 while the micropore volumes increase signi®cantly (Table 7.8).(15,16,18,19) The surface areas of the aforementioned salts increase with the preparative cation=proton ratio up to a value of 1.15 for the latter and remain constant or decrease, in the case of the thallium salt, for further increase of this ratio (Table 7.9 and Fig. 7.16).(16,18,19) As noted earlier in this chapter, PAS FTIR evidence combined with the results from the analyses of adsorption±desorption isotherms suggests that the changes in porosity that occur as the stoichiometry of the salts is altered may be related to the concentration of residual protons. Further data in support of this contention have been obtained from 1H MAS NMR. With TlPW, for example, the NMR spectra contain only one peak regardless of the stoichiometry, showing that the residual protons in the lattice framework have only one environment

Figure 7.15. Pore size distribution for thallium 12-tungstophosphate for various preparative stoichiometries.(15,16,18,19). Reprinted with permission from Parent and Moffat.(16) Copyright 1996 American Chemical Society.

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TABLE 7.8. Micropore Volumes …VMP † and Mean Micropore Radii …rMP † as a Function of the Cation=Acid Ratio Used in the Synthesis of the Salt(15,16,18,19)

Salt Ag3 PW12 O40 Ag3 PMo12 O40 Ag4 SiW12 O40 Tl3 PW12 O40 Tl3 PMo12 O40 Tl4 SiW12 O40

50% shortage

15% shortage

Stoichiometric

15% excess

VMP …ml gÿ1 †

rMP Ê) (A

VMP …ml gÿ1 †

rMP Ê) (A

VMP …ml gÿ1 †

rMP Ê) (A

0.027

8.0

0.031

7.8

0.037

7.9

0.037

7.8

0.021

7.6

0.040 0.043 0.038 0.039

7.2 8.2 8.2 7.8

0.040 0.045 0.046 0.036

7.6 8.2 8.4 8.0

0.039 0.043 0.041 0.035

7.2 8.0 8.8 7.8

50% excess

VMP rMP VMP rMP Ê ) …ml gÿ1 † (A Ê) …ml gÿ1 † (A 0.035

7.7

0.030

7.9

TABLE 7.9. Surface Areas (SBET † a of the Stoichiometric and Nonstoichiometric Salts(16,18,19) Preparative stoichiometryb Salt

0.50

0.85

1.00

1.15

1.50

Ag3 PW12 O40 Ag3 PMo12 O40 Ag4 SiW12 O40 Tl3 PW12 O40 Tl3 PMo12 O40 Tl3 SiW12 O40

77.6

86.2 1.3 107.1 126.9 149.7 102.5

100.9 1.5 106.0 131.6 157.0 97.5

101.4 0.9 105.9 128.6 147.2 92.7

100.5

a b

62.0

103.6

m2 gÿ1 . Cation=proton ratio.

(Fig. 7.17).(15,16,18,19) As the ratio of cation=anion employed in the synthesis increases, the peak areas in the 1H MAS NMR spectra decrease as a result of the shift in equilibrium toward the products in the precipitation process, consistent with the increase in the micropore volume due to the interconnection of interstitial spaces. 7.4. PORE STRUCTURES FROM

129

Xe NMR

The correlation between the isotropic 129Xe chemical shift and the free space of the Xe trapping site has been employed to obtain further information on the pore structures of the ammonium, cesium, and potassium salts of HPW, HPMo, and HSiW.(20,21) Xenon adsorption isotherms for the salts of HPW (Fig. 7.18) show that, for a given pressure, the quantities adsorbed follow the order

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117

Figure 7.16. Surface areas (SBET ) of the (a) silver and (b) thallium salts for various preparative stoichiometries.(15,16,18,19) Reprinted with permission from Parent and Moffat.(16) Copyright 1996 American Chemical Society.

Figure 7.17. 1H MAS NMR spectrum of Tl3 PW12 O40 for various preparative stoichiometries.(19)

118 CHAPTER 7

MICROPOROSITY

119

NH4 ‡ > Cs‡ > K‡ , similar to that of the micropore volumes calculated from nitrogen adsorption isotherms. The NMR spectra for the salts with adsorbed xenon show one remarkably narrow signal. …40 < DH < 150 Hz† as illustrated in Fig. 7.19 for the salts of HPW. For quantities of xenon adsorbed …nXe † greater than 2  1019 atoms gÿ1 , the measured chemical shifts increase linearly (Fig. 7.20). The 129Xe NMR results provide further evidence for the existence of micropores in the monovalent salts of the heteropoly acid. The narrow shapes of the lines show that the microporosity is homogeneous and organized. The variations of the chemical shifts are similar to those observed with Y zeolites exchanged with the same cations. However, the values of pore diameters obtained from the 129Xe NMR technique are lower than those obtained from analysis of the nitrogen adsorption±desorption isotherms. 7.5. SORPTION AND DIFFUSION IN METAL±OXYGEN CLUSTER COMPOUNDS Studies of the sorption and diffusion of organic molecules of various shapes and sizes in porous solids provide useful information concerning the nature of the channels. The sorption and diffusivities of aromatic hydrocarbons,(22) aliphatic saturated hydrocarbons,(23) alkanes,(24) and alcohols(25) have been measured on various of the MOCC. The sorption capacities of the nonporous solid acids HPW, HPMo, and HSiW at a particular temperature were found to be in the order benzene, toluene > p-xylene > m-diethylbenzene, the approximate order of molecule sizes (Fig. 7.21).(22) The heats of sorption fall in the range 9.41.5 kcal molÿ1 , indicative of physisorption. In contrast, the sorption capacities of the porous ammonium salts of the aforementioned acids with these aromatics are approximately 10 times larger than those of the parent acids and apparently depend on the molecular size, shape, and electron density in the benzene ring of the particular aromatic (Fig. 7.22) The diffusivities of the aromatic hydrocarbons on the ammonium salts of HSiW and HPMo are factors of 30 and 3, respectively, larger than those on the parent acids, whereas somewhat surprisingly, the diffusivities on HPW and its ammonium salt for the aromatics are quite similar. With the aliphatic saturated hydrocarbons, n-hexane, 3-methylpentanes, cyclohexane, n-heptane, and n-octane, the sorption capacities of the porous ammonium salts NH4 PW, NH4 SiW, and NH4 PMo are factors of 15±20 higher than those of the parent acids (Fig. 7.23).(23) The sorption capacities of the latter decrease with increasing temperature, as expected with physisorption, but for a given temperature vary little with the nature of the sorbate. The sorption capacities of the ammonium salts for the various alkanes increase with decreasing boiling point of the alkanes as would be expected where physical interaction forces are involved. Where the boiling points are similar, as with n-heptane and

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Figure 7.18. Xenon adsorption isotherms at 300 K for M3 PW12 O40 : M ˆ NH4 ‡ , m; Cs‡ , d; K‡ , ..(21)

isooctane, the higher sorption capacity is obtained with the sorbate molecule of smaller diameter. Of the three ammonium salts, the highest sorption capacity for all alkanes is found with NH4SiW which, as noted earlier in this chapter, has the smallest mean pore radius and micropore volume. The diffusivities of the alkanes in the heteropoly acids are in general a factor of approximately 5 lower than those observed with NH4 PMo and a factor of approximately 20 lower than those with NH4 SiW.(24) In view of the absence of pore structures in the acids, the differences in diffusivities between these and the ammonium salts are not unexpected. Surprisingly, however, the diffusivities of the alkanes with the ammonium salt of HPW differ little from those observed with the acids. The diffusivities show a marked dependence on the kinetic diameters of the sorbate molecules with the lowest and highest values found with cyclohexane and n-hexane, respectively. However, for a given MOCC the diffusivities are not monotonically dependent on the kinetic diameters of the sorbate molecules, although correlations between the diffusivity values and the product of kinetic diameter and boiling point of the sorbates are evident

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Figure 7.19. Shape of 129Xe resonance adsorbed on M3 PW12 O40 : (a) M ˆ NH4 ‡ ; (b) M ˆ Cs‡ ; (c) mixture 52% (a) ‡48% (b); (0T, uncompressed sample; xT, x tonnes cmÿ2 compressed sample).(21)

(Fig. 7.24). Application of the ®nite-layer BET equation to the isotherms for the alkanes produced values for n similar to those obtained from the nitrogen adsorption data (Table 7.10). The sorption and diffusion of alcohols in the MOCC have also been studied and the results compared with those for ZSM-5.(25)

7.6. ION EXCHANGE AND STRUCTURE RETENTION 7.6.1. Cation Exchange and Microporosity The earliest reports on cation exchange were concerned with salts of the metal±oxygen cluster compounds as inorganic ion exchangers and focused on their potential as nondegradative resins for the separation of radioactive isotopes from nuclear reactor washwater.(26±28) The exchange capacity and exchange kinetics of the ammonium salts were shown to be dependent on the size and charge of the cation under exchange from solution with exchange up to 50% with

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Figure 7.20. d ˆ f …nXe † curves at 300 K for M3 PW12 O40 : M ˆ NH4 ‡ , m; Cs‡ ; d; K‡ , w.(21)

K‡ , Rb‡ , Cs‡ , and Tl‡, but negligible exchange with Li‡ and Na‡ for the ammonium ion.(26,29) Explanations of these results have included correlations between the exchange capacity of NH4 PMo and the ionic radius and thermodynamic data for the exchanging cations.(30) and lattice orientation of the ammonium ion,(31,32) but none of these have provided satisfactory rationalizations. It is of considerable interest to note, however, that the salts that have now been shown to be microporous, namely, those with cations K‡ , Rb‡ , and NH4 ‡, have superior ion-exchange properties relative to those of the parent acids and the alkylammonium, alkaline earth, and Li‡, Na‡ salts. More recently, studies of several microporous monovalent salts (Cs‡ , K‡ , ‡ NH4 ) of HPW and HPMo have been undertaken to investigate the retention of morphological properties under ion exchange(33±35) and the possibility of preparing microporous MOCC catalysts with a variety of cations by this method. The salts employed were nonstoichiometric (cation=anion ratio 6ˆ 3) and contained differing proton concentrations, thus allowing the examination of the exchange of the proton as well as that of the alkali and ammonium cations. For a given anion the most signi®cant exchange occurred following the ®rst exposure whereas subsequent exposures resulted in progressively smaller quantities of cations exchanged into the salts. The initial exchange was found to

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123

Figure 7.21. Sorption capacities of HPW, HPMo, HSiW for benzene, toluene, xylenes, mesitylene, and m-diethylbenzene.(22)

involve the removal of both protons and cations from the solid phase, although diminished exchange of protons was also observed in subsequent steps. Because discontinuities are absent from the isotherms (Figs. 7.25 and 7.26), the exchange of protons from the solid phase apparently follows a similar pathway to that of the cations. Additionally, with a given anion and pair of cations the exchange of cations is strongly dependent on the phase location of each cation. For a given anion the exchange of cations depends signi®cantly on the respective phases in which the cations are found. Thus, the magnitude of the exchange of NH4 ‡ into K3 PW12 O40 exceeds that of K‡ into (NH4 †3 PW12 O40 , although such differences are less apparent with the PMo12 O40 3ÿ anion (Figs. 7.25 and 7.26). With cesium as an exchanging cation and either NH4PW or KPW as the recipients, as much as 50% of the ammonium or potassium ions, respectively, is exchanged, in contrast with 10% where K‡ is displacing Cs‡ from CsPMo (Figs. 7.27 and 7.28). Similar results have been obtained with the K‡=Cs‡=PW, NH4 ‡=Cs‡=PW, and NH4 ‡=Cs‡=PMo systems. Measurements of the concentrations of the two exchanging cations as the exchange process proceeds have shown that a one-to-one correspondence exists between the cations exchanging into and out of the solid. As the uptake of cations by the solid phase could not be attributed to adsorption on the peripheral surfaces,

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Figure 7.22. Sorption capacities of NH4 PW, NH4 PMo, NH4 SiW for benzene, toluene, xylenes, mesitylene, and m-diethylbenzene.(22)

the sites for exchange must be located within the microporous structure of the catalysts. The maximum exchange capacities have been found to depend on the two cations under study and the nature of the cation initially present in the solid phase (Table 7.11) (Fig. 7.29). The maximum exchange capacity decreases as the radius of the cation in the solid phase increases, but is signi®cantly less dependent on the composition of the anion. The largest exchange capacities are achieved where the cations entering into and exiting from the solid have similar radii and where the cation initially present in the solid is smaller than the entering cation. The maximum exchange capacity is considerably smaller where the cation initially present in the solid is larger than that entering into the solid. Although no completely adequate explanation can be proffered at this time, the higher binding energies of the larger cations may play a role in the aforementioned observations.

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Figure 7.23. Sorption capacities of NH4 PW, NH4 PMo, and NH4 SiW for aliphatic saturated hydrocarbons.(23)

7.6.2. Crystal Structure and Morphology Retention on Ion Exchange As noted earlier in this chapter, the ability of a microporous solid to retain its morphology and structure on exchange of its cation is of both fundamental and practical importance. Powder XRD and nitrogen adsorption±desorption measurements have been employed to provide evidence in support of such retention.(34) The measured XRD patterns obtained for the ion-exchanged microporous solids with pairs of the cations, K‡ , Cs‡ , NH4 ‡ , were consistent with the cubic Pn3m space group and a single crystallographic phase, in contrast to earlier reports of two phases in K3ÿx Hx PM12 O40 .(36) These observations suggest that biphasic systems can be expected where substantial concentrations of protons exist. Single-crystal XRD measurements on ion-exchanged K‡=NH4 ‡=PMo12 O40 3ÿ

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Figure 7.24. Correlation between diffusivity and the product of kinetic diameter and boiling point of the sorbates.(23)

have also been interpreted as indicative of the presence of single crystallographic phases for these dicationic systems.(32) Nearly one half-century ago it was proposed that the isomorphous 12heteropoly salts possess ¯exible lattices that are capable of expansion and contraction.(26) As noted earlier in this chapter the lattice parameters (a) of the precipitated microporous monovalent salts prepared with cations from the group 1 elements as well as NH4 ‡ were found to increase with the size of the cation.(1,7±19) A similar observation was made with the ion-exchanged materials

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127

TABLE 7.10. Values of n from Finite-Layer BET Equation(24) Sorbent Sorbate N2 n-Hexane Isooctane

NHPW

NHSiW

NHPMo

2.0 2.6 2.3

1.6 1.7 1.3

2.6 2.4 2.0

(Figs. 7.30±7.35). As the larger cation is exchanged by the smaller species, the lattice parameter decreases, generally linearly. The values of a0 obtained for the dicationic exchanged systems are also consistent with the values of the monocationic systems containing the same cations. Evidently, the secondary structure is retained during the ion-exchange process.

Figure 7.25. Ion-exchange isotherm for the exchange of K‡ into …NH4 †3 PW12 O40 (d) and NH4 ‡ into K3 PW12 O40 (m).(33±35)

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Figure 7.26. Ion-exchange isotherm for the exchange K‡ into …NH4 †3 PMo12 O40 (d) and NH4 ‡ into K3 PMo12 O40 (m).(33±35)

Figure 7.27. Ion-exchange isotherm for Cs‡ exchanging into …NH4 †3 PW12 O40 (d) and K3 PW12 O40 (m).(33±35)

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129

Figure 7.28. Ion-exchange isotherms showing the effect of ions entering or exiting the solid ion exchanger. d, Cs‡ exchanging from the liquid phase into the solid K3 PMo12 O40 phase; m, K‡ exchanging from the liquid phase into the solid Cs3 PMo12 O40 phase.(33±35)

Figure 7.29. Maximum exchange capacities of the heteropoly ion exchangers as a function of the ratio of the radius of the cation exiting the solid to that of the cation entering the solid. , PMo12 O40 3ÿ salts; m, PW12 O40 3ÿ salts.(33±35)

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TABLE 7.11. Maximum Exchange Capacities for Monovalent 12-Heteropoly Salts(33) Measured exchange capacity % anion Cation in solid phase K‡ K‡ NH4 ‡ NH4 ‡ Cs‡ Cs‡

Cation in liquid phase NH4 ‡ Cs‡ K‡ Cs‡ K‡ NH4 ‡

PW12 O40 3ÿ

PMo12 O40 3ÿ

67.2 45.2 43.7 42.3 14.6 12.2

45.9 51.8 33.6 50.7 9.8 4.8

The retention of microporosity during ion exchange has also been demonstrated.(34) As the larger cation is replaced by the smaller species in the ionexchanged solid, the surface area decreases and approaches the value obtained for the monocationic material (Figs. 7.36±7.38 and Table 7.12). Pore size analyses also demonstrate that the morphological properties are retained under ion exchange. The mean micropore radii of the ion-exchanged material correlate

Figure 7.30. Lattice parameter as a function of cation composition for the NH4 ‡ =K‡ =PW12 O40 3ÿ ion-exchange system.(33±35)

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131

Figure 7.31. Lattice parameter as a function of cation composition for the NH4 ‡ =Cs‡ =PW12 O40 3ÿ ion-exchange system.(33±35)

Figure 7.32. Lattice parameter as a function of cation composition for the K‡ =Cs‡ =PW12 O40 3ÿ ionexchange system.(33±35)

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Figure 7.33. Lattice parameter as a function of cation composition for the NH4 ‡ =K‡ =PMo12 O40 3ÿ ion-exchange system.(33±35)

Figure 7.34. Lattice parameter as a function of cation composition for the NH4 ‡ =Cs‡ =PMo12 O40 3ÿ ion-exchange system.(33±35)

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133

Figure 7.35. Lattice parameter as a function of cation composition for the K‡ =Cs‡ =PMo12 O40 3ÿ ionexchange system.(33±35)

Figure 7.36. BET surface area as a function of the cation composition for the ion-exchanged K‡ =NH4 ‡ =PW12 O40 3ÿ system.(33±35)

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Figure 7.37. BET surface area as a function of the cation composition for the ion-exchanged K=NH4 ‡ =PMo12 O40 3ÿ system.(33±35)

Figure 7.38. BET surface area as a function of the cation composition for the ion-exchanged Cs‡ =NH4 ‡ =PW12 O40 3i system.(33±35)

MICROPOROSITY

135

TABLE 7.12. BET Surface Areas and Micropore Volumes of Ion-Exchanged 12-Heteropoly Salts(34) M ˆ Mo

MˆW

Ion-exchanged salt

SBET a

Vmp b

SBET a

Vmp b

NH4 PMNH4 1 NH4 PMNH4 2 NH4 PMNH4 3 NH4 PMK1 NH4 PMK2 NH4 PMK3 NH4 PMCs1 NH4 PMCs2 NH4 PMCs3 KPMNH4 1 KPMNH4 2 KPMNH4 3 KPMK1 KPMK2 KPMK3 KPMCs1 KPMCs2 KPMCs3 CsPMNH4 1 CsPMNH4 2 CsPMNH4 3 CsPMK1 CsPMK2 CsPMK3 CsPMCs1 CsPMCs2 CsPMCs3

88.0 80.5 75.7 72.2 67.9 56.8 137.5 113.6 85.7 66.0 60.8 44.9 56.8 45.8 43.6 93.9 21.7 7.7 110.1 71.9 60.2 98.0 44.3 80.9 88.5 50.1 41.3

1.58 1.58 1.50 1.62 1.55 1.61 1.62 1.16 0.77 1.52 1.47 1.08 1.41 1.24 0.54 2.70 0.42 Ð 0.23 Ð Ð 0.54 Ð 0.45 0.34 Ð Ð

119.9 137.7 134.2 131.2 119.5 107.1 121.9 119.1 116.3 122.7 123.1 126.4 110.9 104.2 100.6 118.1 121.0 119.4 131.1 124.3 114.6 127.2 117.3 104.0 127.6 117.3 108.4

2.90 2.75 2.70 2.90 3.01 2.82 2.89 2.78 2.98 1.86 3.00 3.13 2.66 2.78 2.55 3.45 3.33 3.60 1.78 1.16 1.93 2.17 0.77 0.62 1.47 0.85 0.93

a b

m2 gÿ1 . cm3 gÿ1  102 .

with the number of adsorbed layers of nitrogen, as was observed for the monocationic heteropoly salts (Fig. 7.39 and Table 7.13). The surface areas of the NH4 =K‡ exchange systems of the PW and PMo anions have been shown to increase as the lattice parameter increases (Figs. 7.40 and 7.41), consistent with the hypotheses advanced for the existence of pore structures in the heteropoly salts. 7.7. ARGON ADSORPTION Nonstoichiometric cesium salts of HPW were subjected to the adsorption of argon(37) and the pore size distributions were calculated with a mean-®eld

136

CHAPTER 7

Figure 7.39. Mean micropore radius as a function of the number of adsorbed layers for ionexchanged 12-heteropoly salts: r, KPMoBn; ., NH4 PMoBn; m, KPWBn; j, NH4 PWBn; w, CsPMoBn.(33±35)

method(38,39) extended to pores of cylindrical symmetry(40) and independently evaluated(15) Salts prepared with 2.1 mol Cs per mole of the salt and those Ê whereas promoted with 0.5 wt % Pt were found to have micropores of 5.1±5.3 A those prepared with 2.5 mol Cs per mole of the salt and the Pt-promoted analogue Ê. were calculated to have bimodal distributions of 5.0±14 and 45 A

7.8. DIVALENT SALTS IR spectroscopy, powder and single-crystal XRD, DTA, and nitrogen adsorption±desorption measurements have been performed on a series of solids prepared from HPW and HPMo and the alkaline earth hydroxides.(41,42) Although the infrared spectra provide no evidence for the presence of decomposition species of the 12-heteropoly anions, the DTA analyses suggest that complex

MICROPOROSITY

137

TABLE 7.13. CBET Parameters and Mean Micropore Radii of Ion-Exchanged 12-Heteropoly Salts(34) M ˆ Mo

MˆW

Ion-exchanged salt

CBET

Ê) rmp (A

CBET

Ê) rmp (A

NH4 PMNH4 1 NH4 PMNH4 2 NH4 PMNH4 3 NH4 PMK1 NH4 PMK2 NH4 PMK3 NH4 PMCs1 NH4 PMCs2 NH4 PMCs3 KPMNH4 1 KPMNH4 2 KPMNH4 3 KPMK1 KPMK2 KPMK3 KPMCs1 KPMCs2 KPMCs3 CsPMNH4 1 CsPMNH4 2 CsPMNH4 3 CsPMK1 CsPMK2 CsPMK3 CsPMCs1 CsPMCs2 CsPMCs3

1409 877 1136 1252 1098 1797 452 460 376 1896 2705 1739 2203 2131 2674 806 93 501 472 194 177 385 150 252 205 153 181

10.8 12.0 10.9 12.1 10.4 9.5 13.4 14.2 14.6 8.9 8.5 9.1 8.8 8.8 8.6 7.9 9.5 Ð 13.8 Ð Ð 13.6 Ð 15.1 14.9 Ð Ð

1643 1207 643 996 1649 2795 1301 561 651 1958 1880 2147 1580 3704 2861 1539 1219 3638 524 493 611 449 467 361 636 452 1749

11.6 12.3 11.7 15.6 10.1 10.6 11.3 12.0 11.7 7.6 10.5 10.9 11.4 9.4 10.4 10.3 10.1 10.2 Ð Ð Ð Ð Ð Ð Ð Ð Ð

interactions between the divalent cations and the heteropoly anions are occurring. The powder XRD studies and the single-crystal study suggest that the alkaline earth cations are not incorporated into the lattice. With the Ba2‡ =PMo12 O40 system the measured lattice parameters are almost identical to those of H3 PMo12 O40 . The BET surface areas of the preparations calculated from the nitrogen adsorption isotherms are less than 20 m2 gÿ1 and the CBET values are less than 100 in all cases. Further, the isotherms show little or no evidence for the formation of porous structures. The acid apparently coexists in a two phase system with the cation present as another salt. The formation of divalent salts of the heteropoly acids of Keggin structure appears to be dif®cult if not impossible under conditions of retention of the structure of the parent acid.

138

CHAPTER 7

Figure 7.40. The relationship between surface area and lattice parameter for the K‡ =NH4 ‡ =PW12 O40 3ÿ ion-exchange system.(33±35)

Figure 7.41. The relationship between surface K‡ =NH4 ‡ =PMo12 O40 3ÿ ion-exchange system.(33±35)

area

and

lattice

parameter

for

the

MICROPOROSITY

139

7.9. REACTIONS ON THE MICROPOROUS SALTS AND SHAPE SELECTIVITY The ®rst evidence, albeit indirect, of the impact of the porous structure on catalytic processes with heteropoly oxometalates may be that found in the conversion of methanol to higher hydrocarbons on the heteropoly acids and the ammonium salt of 12-tungstophosphoric acid, referred to at the beginning of this chapter.(4±6) In contrast to the ole®ns observed as products with HPW, paraf®ns were the predominant products with NH4 PW. Bimolecular hydrogen transfer reactions are believed to be enhanced in zeolites as a result of the high density of acidic sites and the concentration of the hydrocarbon reactants in the pores as well as the polarization of the reacting species within the channels. Such a phenomenon can be occurring in the pores of NH4 PW. The isomerization of C6 ±C8 alkenes has been investigated in the liquid phase at 303±343 K on HPW, HPMo, and HSiW and their ammonium salts.(43a) Double-bond and cis±trans isomerizations, but no skeleton isomerizations, are observed. HPW, HSiW, and NH4 PW are active at room temperature in the liquid phase but HPMo and NH4 PMo are relatively inactive as is also, surprisingly, NH4 SiW. The conversion of C6 ±C8 alkenes in the gas phase on NH4 PW at temperatures from 600±700 K yields products ranging from C3 hydrocarbons to aromatics, the latter appearing in selectivities as high as 30%.(43b) The primary products provide strong evidence that condensation processes are strongly favored on NH4 PW, rather than the usual cracking process involving b-scission of carbocations. As with hydrogen transfer processes the presence of micropores is undoubtedly advantageous in condensation processes. The conversion of methylethylbenzenes and, in particular, 1-methyl-2ethylbenzene has been studied on NH4 PW and NH4 SiW.(44) In view of the presence of micropores in these catalysts, a monomolecular, as opposed to a bimolecular, transalkylation process would be expected with these dialkylbenzenes. Indeed, toluene is observed as a primary product from 1-methyl-2ethylbenzene with NH4 PW reaction temperatures up to 200  C and after pretreatment at 350 or 450  C and with NH4 SiW after pretreatment at 450  C. As noted with the isomerization of C6 ±C8 alkenes, the conversion obtained with the siliconbased catalyst is signi®cantly lower than that with the phosphorus analogue. The cracking activity for n-hexane at 648±698 K is found to be a factor of 10 higher for NH4 PW relative to that on NH4 PMo and NH4 SiW although the sorption capacities and the diffusivities at 20  C for n-hexane follow the order NH4 SiW > NH4 PMo > NH4 PW.(23,45) The reaction of toluene with methanol on NH4 PW and NH4 SiW occurs predominantly through the alkylation of the benzene ring to produce xylenes and the trimethyl- and tetramethylbenzenes.(46) Experiments in which methanol was

140

CHAPTER 7

introduced to the catalysts followed by the later addition of toluene showed that the initial step in the mechanism involved the methylation of the terminal oxygen atom as observed in photoacoustic FTIR studies of the conversion of methanol to higher hydrocarbons.(47,48) A comparison of the selectivities to the isomers of xylene obtained on NH4PW and NH4SiW with those on the zeolites H-theta-1 and H-ZSM-5(49) at similar conversions shows that the NH4 salts of HPW and HSiW compare favorably with the highly selective H-theta-1 for the production of p-xylene, demonstrating the existence of shape selectivity with the microporous heteropoly salts (Table 7.14) Several of the cesium salts of HPW together with the parent acid itself have been synthesized and characterized by chemical analysis, thermogravimetric analysis, NH3 adsorption±desorption, and 31P NAS NMR techniques.(50) Nitrogen adsorption coupled with t-plot analysis con®rm the earlier work showing that porous structures are present in the cesium salts. The most active catalysts for nbutane isomerization which requires strong acidic sites are those with a cesium content of approximately 2 Cs per KU while for methanol dehydration to dimethylether samples with Cs content of 2±2.7 Cs per KU are the most active. The liquid-phase isobutane=butene alkylation process was shown to require larger pores to permit the reactants to access the acidic sites and the desorption of reactant products.(51) In con®rmation of the earlier work on the effect of the stoichiometry on the porosity of salts of the heteropoly acids,(14) further studies of the nonstoichiometric cesium salts have been undertaken.(52,53) These authors contend that cesium salts of HPW up to and including 2 Cs per KU have negligible surface areas in contrast to the work reported by others.(50,51)

TABLE 7.14. Illustration of Shape Selectivity in the Selectivity to the Isomers of Xylene Formed from the Reaction of Methanol and Toluene(46) Catalysta Xylene composition (%) p-Xylene m-Xylene o-Xylene a b c

NH4 SiW 250  C

NH4 PWb 160  C

Theta-1c 550  C

ZSM-5c 600  C

45.1 36.0 18.9

49.9 31.3 18.8

44.9 28.6 26.5

41.4 41.4 18.2

Temperatures are those for reaction. 200 ml NH3 per 0.50 mg catalyst added to catalyst at 160  C prior to reaction. Reference 49.

MICROPOROSITY

141

Acidic cesium salts of various stoichiometries have been prepared by grinding mixtures of HPW and CsPW in various relative amounts to produce a dispersion of the acid on the high-surface-area cesium salt.(54) The authors show that this technique produces high-surface-area solids with signi®cantly higher catalytic activity for the isomerization of n-butane at 473 K. The activity, in the liquid-phase alkylation of isobutane with 2-butene at 80  C, of the cesium, ammonium, and potassium salts of HPW prepared with different stoichiometries, was shown to correlate with the surface acidity of the solid.(55) Evidence was provided to show that the nonstoichiometric ammonium salt with 2:5 NH4 ‡ per KU contains sites of a higher acid strength relative to the cesium and potassium analogues. This is of particular interest in view of the earlier observations obtained with NH4PW in the conversion of methanol to higher hydrocarbons.(5,6) Powder X-ray and electron diffraction techniques together with computer simulations have been employed to investigate the secondary structures of Csx H4ÿx PVMo11 O40
nH2 O…0  x  4†.(56) As reported earlier for CsPW from X-ray and neutron diffraction studies(57) and discussed in detail in the chapter on structure, the structures of the cesium salts with 2  x  4 are cubic with space group Pn3m, whereas the acidic form has a triclinic lattice.

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9.

10. 11. 12. 13.

J. B. Moffat, J. Mol. Catal. 52, 169 (1989). J. G. High®eld and J. B. Moffat, J. Catal. 88, 177 (1984). G. B. McGarvey and J. B. Moffat, J. Mol. Catal. 69, 137 (1991). H. Hayashi and J. B. Moffat, J. Catal. 77, 473 (1982). J. B. Moffat in: Catalytic Conversions of Synthesis Gas and Alcohols to Chemicals (R. G. Herman, ed.), Plenum Press, New York (1984). H. Hayashi and J. B. Moffat, J. Catal. 83, 1982 (1983). J. B. McMonagle and J. B. Moffat, J. Colloid Interface Sci. 101, 479 (1984). J. B. Moffat, J. B. McMonagle, and D. Taylor, Solid State Ionics 26, 101 (1988). (a) J. B. Moffat, in: Studies in Surface Science and Catalysis, Vol. 30 (B. Delmon, P. Grange, P. A. Jacobs, and G. Poucelet, eds.), Elsevier, Amsterdam (1987); (b) J. B. Moffat, in: Studies in Surface Science and Catalysis, Vol. 38 (J. Ward, ed.), Elsevier, Amsterdam (1988); (c) J. B. Moffat, G. B. McGarvey, J. B. McMonagle, V. Nayak and H. Nishi, in: Guidelines for Mastering the Properties of Molecular Sieves, (D. Barthomeuf, ed.), Plenum Press, New York (1990); (d) J. L. Bonardet, K. Carr, J. Fraissard, G. B. McGarvey, J. B. McMonagle, M. Seay and J. B. Moffat in: Advanced Catalysis and Nanostructured Materials (W. Moser, ed.), Academic Press, San Diego (1996). (a) S. J. Gregg and R. Stock, Trans. Faraday Soc. 53, 1355.(1957).; (b) S. J. Gregg and M. M. Tayyab, Trans. Faraday Soc. 74, 348 (1977). J. B. McMonagle, V. S. Nayak, D. Taylor and J. B. Moffat, in: Proc. 9th Int. Congr. on Catal. (M. J. Phillips and M. Ternan, eds.), Chemical Institute of Canada, Ottawa (1988). D. B. Taylor, J. B. McMonagle, and J. B. Moffat, J. Colloid Interf. Sci. 108, 278 (1985). G. B. McGarvey and J. B. Moffat, J. Colloid Interf. Sci. 125, 51 (1988).

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14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49.

D. Lapham and J. B. Moffat, Langmuir 7, 2273 (1991). M. A. Parent and J. B. Moffat, Langmuir 11, 4474 (1995). M.A. Parent and J. B. Moffat, Langmuir 12, 3733 (1996). J. G. High®eld and J. B. Moffat, J. Catal. 89, 185 (1984). M. A. Parent and J. B. Moffat, Catal. Lett. 48, 125 (1997). M. A. Parent and J. B. Moffat, J. Catal. 177, 335 (1998); Catal. Lett. 60, 191 (1999). J. L. Bonardet, J. Fraissard, G. B. McGarvey, and J. B. Moffat, Colloids Surfaces A72, 191 (1993). J. L. Bonardet, J. Fraissard, G. B. McGarvey, and J. B. Moffat, J. Catal. 151, 147 (1995). V. S. Nayak and J. B. Moffat, J. Colloid Interface Sci. 120, 301 (1987). V. S. Nayak and J. B. Moffat, J. Colloid Interface Sci. 122, 475 (1988). V. S. Nayak and J. B. Moffat, J. Phys. Chem. 92, 2256 (1988). V. S. Nayak and J. B. Moffat, J. Phys. Chem. 92, 7097 (1988). H. Buchwald and W. P. Thistlewaite, J. Inorg. Nucl. Chem. 5, 341 (1958). J. V. R. Smit, W. Robb, and J. J. Jacobs, Nucleonics, Sept., 116 (1959). J. V. R. Smit and J. J. Jacobs, Ind. Eng. Chem. Process Des. Dev. 5, 117 (1966). J. V. R. Smit and W. Robb, J. Inorg. Nucl. Chem. 26, 509 (1964). C. J. Coetzee and E. F. C. H. Rohwer, J. Inorg. Nucl. Chem. 32, 1711 (1970). T. V. Healy, Radiochim. Acta 3, 106 (1964). J. C. A. Boeyens, G. J. McDougall, and J. V. R. Smit, J. Solid State Chem. 18, 191 (1976). G. B. McGarvey and J. B. Moffat, J. Catal. 128, 69 (1991). G. B. McGarvey and J. B. Moffat, J. Catal. 130, 483 (1991). G. B. McGarvey and J. B. Moffat, J. Catal. 132, 100 (1991). T. Haeberle and G. Emig, Chem. Eng. Technol. 11, 392 (1988). Y. Yamada, K.-I. Johkan, and T. Okuhara, Microporous Mesoporous Mater. 26, 109 (1998). G. Horvath and K. J. Kawazoe, J. Chem. Eng. Jpn. 16, 470 (1983). R. D. Kaminsky, E. Maglara, and W. C. Conner, Langmuir 10, 1556 (1994). A. Saito and H. C. Foley, AIChE J. 37, 429 (1991). G. B. McGarvey, N. J. Taylor, and J. B. Moffat, J. Mol. Catal. 80, 59 (1993). G. B. McGarvey and J. B. Moffat, Catal. Lett. 16, 173 (1992). (a) V. S. Nayak and J. B. Moffat, Appl. Catal. 36, 127 (1988); (b) 77, 251 (1991). H. Nishi and J. B. Moffat, J. Mol. Catal. 51, 193 (1989). V. S. Nayak and J. B. Moffat, Appl. Catal. 47, 97 (1989). H. Nishi, K. Nowinka, and J. B. Moffat, J. Catal. 116, 480 (1989). J. G. High®eld and J. B. Moffat, J. Catal. 95, 108 (1985). J. G. High®eld and J. B. Moffat, J. Catal. 98, 245 (1986). A. G. Ashton, S. A. I. Barri, S. Cartlidge, and J. Dwyer, in: Chemical Reactions in Organic and Inorganic Constrained Systems (R. Setton, ed.), Reidel, Dordrecht (1980). N. Essayem, G. Coudurier, M. Fournier, and J. C. VeÂdrine, Catal. Lett. 34, 223 (1995). N. Essayem, S. Kieger, G. Coudurier and J. C. VeÂdrine in: Studies in Surface Science and Catalysis, Vol. 101 (J. W. Hightower et al., eds.), p. 591, Elsevier, Amsterdam (1996). T. Okuhara, T. Nishimura, and M. Misono, in: Studies Surface Science and Catalysis, Vol. 101 (J. W. Hightower et al., eds.), p. 581, Elsevier, Amsterdam (1996). T. Okuhara, T. Nishimura, and M. Misono, Chem. Lett. 55 (1995). P.-Y. Gayraud, N. Essayem and J. C. VeÂdrine, Catal. Lett. 56, 35 (1998). A. Corma, A. Martinez, and C. Martinez, J. Catal. 164, 422 (1996). S. Berndt, D. Herein, F. Zemben, E. Beckmann, G. Weinberg, J. SchuÈltze, G. Mestl, and R. SchloÈgl, Ber. Bunsenges, Phys. Chem. 102, 763 (1998). G. M. Brown, M. R. Noe-Spirlet, W. E. Busing, and H. A. Levy, Acta Crystallogr. B33, 1038 (1977).

50. 51. 52. 53. 54. 55. 56. 57.

8 THE TWO FUNCTIONS Acidity and Oxidation±Reduction

8.1. ACIDITY Whereas acidity in single-phase aqueous systems can be rigorously de®ned and the measurement techniques can be unambiguously related to the de®nition, the application of these concepts to solids and their surfaces is considerably more complex.(1) The usual interpretation of acid strength as related to the equilibrium constant(s) for the dissociation of the protonic acid is at best dif®cult to apply to solid surfaces and the data obtained from techniques that are currently available for the estimation of acidic strengths on surfaces are not easily correlated with each other.(2) Further, the often encountered energetic and geometric heterogeneity of the surfaces of catalysts inevitably leads to a distribution of acid strengths, an additional complexity to the problem. 8.1.1. Electrical Conductivity Because the acid strength of a catalyst is related to the mobilities of the protons, some measure of acidity in the heteropoly oxometalates can be gained from studies of the protonic conductivity,(3±15) that have employed conductivity, nuclear magnetic resonance and neutron scattering techniques. The ®rst evidence for the high protonic conductivity was reported in 1979(3) followed by a second publication from the same authors.(4) The electrical conductivities at 25  C of H3 PMo12 O40 , H4 PMo11 VI Mo V O40 , H4 PMo11 VO40 , H4 SiMo12 O40 , and H4 SiW12 O40 in ethanol and acetone.(5) and in acetic acid(6) have been measured and the dissociation constants have been calculated. The ambient temperature conductivities for the 21H2 O hydrates of HPW and HPMo were reported as 0.5 and 3 Smÿ1, respectively, with activation energies of 43 and 40 kJ molÿ1, 143

144

CHAPTER 8

respectively,(7) to be compared with 17 Smÿ1 and 17 kJ molÿ1 for the 29 hydrate.(3) Studies of the n ˆ 6, 14, 21, and 29 hydrates of HPW
nH2 O show that the conductivities increase with n, and self-diffusion and conduction occur through different mechanisms in these hydrates.(8,9) The proton conductivities of a number of molybdovanadophosphoric(10) and molybdotungstovanadogermanic acid(11) have been measured. Measurements with pelletized HPW
6H2 O and HSiW
6H2 O suggest that conduction occurs via crystallite surfaces.(12) Any conductivity within the crystallites is less than 210ÿ7 Smÿ1 . The conductivity of pellets of HPW exposed to water at 350  C increased with the duration of exposure.(13) Although beyond the scope of the present report, it should be noted that the high proton conductivity of the heteropoly acids has been utilized in multicomponent systems. For example, polymeric materials such as polyacetylene(14) and polyaniline(15) have been protonated with these acids. The decrease in electrical conductivity in conducting polyaniline and polypyrrole due to thermal aging at 120  C for periods up to 11 h has been reported.(16) The electrical conductivities of supported heteropoly acids have also been measured. The proton conductivities of HPW, HPMo, and HSiW adsorbed on porous glass have been shown to depend on the quantity sorbed and the relative humidity.(17) Silica-gel ®lms containing HPW prepared by the sol±gel method have also been studied(18±21) and a tin-mordenite containing HPW, prepared by wet impregnation has been evaluated.(22) The conductivities of salts of the heteropoly acids have also been measured.(23) Nonstoichiometric ammonium 12-tungstophosphate prepared from an excess of the parent acid relative to the ammonium salt has a higher conductivity than the corresponding stoichiometric composition, providing evidence for protonic conductivity. Although the heteropoly acids have high conductivities, because as noted these are dependent on the degree of hydration and thus on the humidity and temperature, they are not as suitable for applications. The salts of the heteropoly acids are less susceptible to such conditions and consequently have been further investigated for applications where both high conductivities and relatively high stabilities are required.(24,25) A recent publication has investigated fuel cells with the 29 hydrate of HPW as the solid electrolyte.(26) 8.1.2. Calorimetry Differential heats of ammonia sorption onto and into heteropoly acids have been measured by microcalorimetry and have been employed to characterize the acid strengths of these solids.(27±29) After pretreatment of the catalysts at 523 K and NH3 sorption at 323 K, the differential heats for HPW and HSiW for the initial sorption are approximately 150 kJ molÿ1 (Figs. 8.1 and 8.2) and remain at this value until approximately 3 and 4 molecules per anion [Keggin Unit (KU)],

ACIDITY AND OXIDATION±REDUCTION

145

Figure 8.1. Differential heats of the sorption of NH3 on HPW at 323 K. Pretreatment: (1) 423 K=2 h=UHV, (2) 523K=1 h=UHV.(28) Reprinted from Microporous Materials, 1, Lazefowicz et al., p. 313, copyright 1993, with permission from Elsevier Science.

respectively, had been sorbed at which a precipitous decrease in the differential heat is evident. It should be noted that the latter values correspond to the total number of protons per anions contained in the solid acids, as represented by the stoichiometric formulas H3 PW12 O40 and H4 SiW12 O40 , respectively. Because the surface areas of these acids are small (< 10 m2 gÿ1 ), the preponderance of the protons are contained within the solid and evidently the NH3 molecules are

Figure 8.2. Differential heats of the sorption of NH3 on HSiW at 323 K. Pretreatment: (1) 423 K=2 h=UHV, (2) 523 K=1 h=UHV(28). Reprinted from Microporous Materials, 1, Lazefowicz et al., p. 313, copyright 1993, with permission from Elsevier Science.

146

CHAPTER 8

penetrating into the bulk structure of the acids as shown by photoacoustic FTIR studies.(30) This property will be discussed in more detail subsequently. A lower pretreatment temperature (423 K) results in initially higher values for the differential heat but the aforementioned sharp drop in these values occurs at lower values of the quantity of NH3 sorbed. The latter observations suggest that at the lower pretreatment temperature, suf®cient water remains, hydrogen-bonded to the protons, to prevent the interaction of ammonia with the totality of protons present. The higher adsorption energies found for NH3 with HPW pretreated at the lower temperature (423 K) may result from an inductive effect produced by the water molecules hydrogen-bonded to the protons or crystallographic structural changes due to the presence of the water and=or ammonia. The value for the differential heat of ammonia adsorption obtained with 12molybdophosphoric acid is signi®cantly smaller than that observed with the tungsten-containing acids, consistent with the expectations of lower acidic strength of the former acid as found from both theoretical(31±33) and experimental(34±36) studies. A relationship between the acid strength (Ho ) and the differential heat of adsorption of ammonia on silica±alumina was employed to show that differential heats of 137.0 and 76.1 kJ molÿ1 correspond to Ho values of ÿ14:5 and ÿ5:6, respectively.(37,38) Because a superacid is conventionally de®ned as any acid stronger than 100% H2 SO4 ,(39) which has an Ho value of ÿ12, the aforementioned two tungsten-containing acids can be classed as superacids. However, as noted in a subsequent section of this chapter, contrary opinions have been advanced. Microcalorimetric measurements on HPW and HSiW supported on silica have also shown the effect of pretreatment temperature as well as the dependence of acid strength on the composition of the anion.(40) HPW=SiO2 , pretreated at 423 K, has acidic sites …qd ˆ 220 kJ molÿ1 † that are stronger than those on HSiW=SiO2 pretreated at 623 K (200 kJ molÿ1 ) and both are stronger than those of the unsupported acids. In contrast, the acidic sites on HPW supported on activated carbon are weaker than those of the bulk material.(41) Heat capacities of the heteropoly acids have also been measured.(29a) The acidities of various carbon supported and unsupported heteropoly acids have been measured by ammonia sorption microcalorimetry.(29b) The acidity decreases in the order H3 PW12 O40 > H4 SiW12 O40 > H5 PW12 O40  H21 B3 W39 O132. When supported on carbon, the acidity of the ®rst three acids is slightly reduced whereas that of the latter is increased. Most recently the acidity of solid HPW has been measured using slurry calorimetry in cyclohexane with pyridine as the probe and analysis by the Cal-ad method.(42) The Cal-ad method combines the measurements from a calorimetric titration and adsorption isotherm, both of which, in this work, employed pyridine. The protons were shown to have different strengths (Table 8.1) which were assigned to the face (n1 ) and edge (n2 ) positions in the unit cell of HPW. The total

ACIDITY AND OXIDATION±REDUCTION

147

TABLE 8.1. Solid Acidity of HPW(42) Type of proton

Population na

Enthalpyb of interaction

1 2

0.08 0.16

ÿ32:7 ÿ19:6

a b

mmol gÿ1 . kcal molÿ1 .

number of protons titrated (0.24 mmol gÿ1 ) shows that not all of the available protons. (1.04 mmol gÿ1 ) react with pyridine. This appears to be in agreement with earlier results from the photoacoustic FTIR of pyridine on HPW.(43) However, the number of protons titrated with pyridine (0.24 mmol gÿ1 ) is larger than the number of surface protons (0.008 mmol gÿ1 ), showing that penetration into the lattice occurs. As found in earlier work,(43) XRD shows that the lattice of HPW expands on exposure to pyridine. The enthalpies of the most acidic protons in solid HPW (ÿ32:7 kcal molÿ1 ) and for that dissolved in acetonitrile (ÿ30:4 cal molÿ1 ) are in good agreement. 8.1.3. Theoretical Semiempirical extended HuÈckel calculations on PW12 O40 3ÿ , PMo12 O40 3ÿ , and SiW12 O40 4ÿ have shown that the magnitude of the negative charges on the terminal oxygen atoms in the anions containing molybdenum are higher than those in which tungsten is the peripheral metal element (Fig. 8.3).(31±33) Under the assumption that the protons are coulombically bound to the terminal oxygen atoms, it may be concluded that the protons in the compounds containing tungsten are more mobile and hence more acidic than those in the corresponding materials containing molybdenum. These conclusions are consistent with the results from microcalorimetry as well as those from various catalyzed reactions, as noted elsewhere in this volume. NMR experiments with 17O-enriched HPW and HSiW have found that the terminal oxygen atoms are the predominant protonation sites.(44) 8.1.4. Solid±Liquid Phase Titrations Both the numbers and strengths of surface acidic sites can be estimated by titration with a weak base such as butylamine dissolved in an inert solvent and a series of basic Hammett indicators with a range of pKa values which are sorbed on the solid surface.(1,2,45) As with any method, de®ciencies exist (see, e.g., Refs. 1 and 2 and references therein). In particular, Bronsted and Lewis acid sites cannot be distinguished. The method has been applied to a number of heteropoly

148

CHAPTER 8

Figure 8.3. Net atomic charge on bridging (b) and outer (o) oxygen atoms and bond energy of peripheral metal±oxygen bonds of the Keggin anion as calculated from extended HuÈckel calculations.31 Reprinted from the Journal of Molecular Catalysis, 26, Moffat, p. 385, copyright 1984, with permission from Elsevier Science.

acids and their salts.(36) As with the results from microcalorimetry, the effects of pretreatment temperature are again evident (Fig. 8.4). The most strongly acidic sites (Ho  ÿ5:6) on HPW increase in number with increase in pretreatment temperature up to approximately 673 K, followed by a precipitous decrease with further increase in temperature. It is noteworthy that the onset of the second peak observed in TPD experiments with HPW occurs at approximately this temperature.(32) This TPD peak has been attributed to water produced by the extraction of oxygen atoms by protons which would decrease the concentration of acidic sites.

ACIDITY AND OXIDATION±REDUCTION

149

f

Figure 8.4. Effect of pretreatment temperatures on the acid strength distribution of (A) HPW and (B) HSiW.(36) ( ) Ho  ÿ5:6; …D† ÿ 5:6 < Ho < ÿ3:0; (d) ÿ3:0 < Ho < 1:5; (s) 1:5 < Ho < 3:3; (,) 3:3 < Ho < 5:0; (u) 5:0 < Ho < 6:8; (e) Ho  6:8.

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The acid strengths of a series of HXW acids, where X is Co, B, Si, Ge, or P, have been evaluated using dicinnamylideneacetone (pKa is ÿ3:0) as an indicator and following the color changes by UV spectrophotometry.(46) The acid strength was found to increase with the oxidation state of the central atom so that Co < B < Si, Ge < P. A previous investigation found that HPW > HPMo > HSiW > HSiMo > HGeMo.(47) 8.1.5. Adsorption±Desorption The adsorption and desorption of basic molecules, such as pyridine, have been employed to provide information on the acidic properties of the heteropoly oxometalates.(48) The number of acidic sites can be estimated from the number of pyridine molecules irreversibly sorbed on and in the solid while the acid strength increases with the temperature at which the pyridine molecules remain irreversibly sorbed (Fig. 8.5). The numbers of acid sites on HPW increase up to a pretreatment temperature of approximately 700 K and decrease with further increase in temperature. 8.1.6. Spectroscopy 8.1.6.1. Photoacoustic FTIR Although the spectroscopic studies will be discussed in more detail elsewhere in this volume, it is relevant to note at this juncture that photoacoustic FTIR experiments with ammonia and pyridine have shown that polar molecules such as these are capable of penetrating into the bulk of the heteropoly acids.(30,43) In addition, bands attributed to the ammonium and pyridinium ions indicate the presence of BroÈnsted acidic sites whereas the absence of those due to coordinated ammonia and pyridine show that Lewis acidity is not present. With ammonia, quantities up to 5 molecules anionÿ1 were sorbed at room temperature in 5 min.(30) A signi®cant fraction of these desorbed on heating in vacuum at 150  C. When smaller quantities of NH3 were dosed stepwise at 150  C after 3 molecules NH3 anionÿ1 were taken up no further sorption occurred, implying the penetration of NH3 into the bulk structure and the formation of a stoichiometric salt. Bands at  3200 and 1420 cmÿ1, characteristic of the NH4 ‡ ion, together with sharpening of the KU bands, with the spectrum ultimately approaching that of ammonium 12-tungstophosphate, provide supporting evidence for the aforementioned interpretation (Fig. 8.6). In view of its progressive loss, the background continuum can be attributed to the presence of mobile protons. The PAS spectra of HPW heated to 450  C and that after dosing with NH3 to form the stoichiometric salt show nearly identical background absorption suggesting that heating to 450  C results in the removal of protons as heat

! H3ÿx PW12 O40ÿx=2 ‡ x=2 H2 O H3 PW12 O40 ÿ vac

ACIDITY AND OXIDATION±REDUCTION

151

Figure 8.5. Dependence of the acid strength distribution of (A) NH4 PW and (B) HPW on pretreatment temperature. s, d, n, m and u represent the number of acidic sites that can sorb pyridine irreversibly at 323, 423, 523, 623, and 723 K, respectively.(48) Reprinted from the Journal of Molecular Catalysis, 80, Nayak and Moffat, p. 75, copyright 1993, with permission from Elsevier Science.

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Figure 8.6. Photoacoustic FTIR spectra of (a) HPW preevacuated at 473 K, (b)±(e) after stepwise dosing with NH3 at 423 K (f) The ammonium salt prepared in aqueous solution is included for comparison.(30)

This is consistent with results from temperature-programmed desorption of HPW.(49±50) PAS in the infrared region has also been employed to study the interactions of pyridine with HPW.(43) The results show interesting differences from those

ACIDITY AND OXIDATION±REDUCTION

153

obtained with NH3 . With pyridine, as with NH3 , after evacuation of HPW at 250  C (Fig. 8.7a), exposure of HPW at 25  C to excess pyridine produced a rapid initial sorption followed by a slow continuous uptake attaining a total limiting value of approximately 6 pyridine molecules anionÿ1 in 1 h. After evacuation at 25  C the PAS spectrum (Fig. 8.7b) contains new bands attributed to sorbed pyridine (1700±1100 cmÿ1 ) along with the expected bands in the range 1100± 600 cmÿ1 characteristic of the Keggin structure, superimposed on a background whose intensity increases with decreasing wave number. Because the background continuum is attributed to the presence of mobile protons(30) and the band (1540 cmÿ1 ) characteristic of the pyridinium ion is very weak, the formation of the pyridinium ion appears to be inhibited. However, the observation of major bands at 1605, 1489, 1443, and 1425 cmÿ1 suggests that pyridine is associated in hydrogen-bonded forms. When the pyridine is dosed at 25  C, the spectrum (Fig. 8.7c) shows strong bands at 1640, 1610, 1537, and 1485 cmÿ1 , characteristic of protonated pyridine. Comparison of the spectrum with those of pyridinium salts and the absence of bands characteristic of other types of bound pyridine indicates that the sorbed pyridine has been converted to the pyridinium ion. After dosing of the stoichiometric amount of pyridine (3 py anionÿ1 ) at  25 C, most of the pyridine (2.7 py anionÿ1 ) was sorbed in approximately 2.5 h. Pyridinium ions are evidently present as the major species (Fig. 8.7d) but the 1540 cmÿ1 band is suppressed and a band characteristic of H-bonded pyridine is also present at 1443 cmÿ1 . When this sample was heated under vacuum at 100  C, bands (Fig. 8.7e) due to the pyridinium ion were enhanced and the band at 1443 cmÿ1 and the background continuum were suppressed, indicating that a fraction of pyridine in H-bonded form had been converted to the pyridinium ion. The spectrum is now similar to that of the pyridinium salt (Fig. 8.7f). At stoichiometric loading (or greater), pyridine apparently interacts with the pyridinium ion to form an H-bonded complex (py2 H†‡ . The aforementioned background continuum is now attributed to the pyridinium=pyridine complex. Experiments with deuterated pyridine have con®rmed that a dimer in which the two pyridine molecules are equivalent is formed. The aforementioned results can be taken as suggesting that pyridinium ions in the stoichiometric salt occupy two crystallographically distinct sites in roughly equal numbers. The dimer is apparently unstable above 100  C. 8.1.6.2. Nuclear Magnetic Resonance The acidities of the heteropoly acids have been investigated in both solution and solid state by application of NMR. It is relevant to note in this context that very dilute solutions of heteropoly acids in water and in organic solvents have been investigated by conductimetry(51±53) and by potentiometric titrations.(54) Although only one dissociation constant was measured in these experiments, the others were estimated in the ratio K1 =K2 =K3 as 1=0.04=0.004 for HPW.(55) The

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Figure 8.7. Photoacoustic FTIR spectra of (a) HPW preevacuated at 250  C; (b) after exposure to excess pyridine at 25  C and evacuation; (c) as (a) exposed to a controlled dose of pyridine (0.94 py KUÿ 1), (d) as (a) exposed to a larger dose of pyridine (2.7 py KUÿ 1), (e) as (d) after heating under static vacuum at 100  C, (f) pyridinium salt […PyH†3 PW12 O40 preevacuated at 200  C] prepared by adding pyridine to an aqueous solution of HPW in stoichiometric proportions.(43)

ACIDITY AND OXIDATION±REDUCTION

155

acidities of concentrated solutions in water and in acetic acid have also been measured(56,57) by the Hammett indicator method.(58±60) More recently, careful work on the acidities of HPW dissolved in acetic acid has been reported.(61) In this work the authors note the ¯aws in the aforementioned work and attempt to rectify these in their studies using their 13C NMR method(62±68) for the measurement of acidity with mesityl oxide (MO) as an indicator. The reaction between MO and HPW a

b

a

b

ÿ H3 CÿCOÿ CH ˆC…CH†3 †2 ‡ HPW! H2 CÿC…OH†ˆCH ÿ C…CH3 †‡ 2 ‡ PW

is fast so the C-13 chemical shifts for the mixture of MO and MOH‡ are the weighted average of the values for these two species. The difference in the chemical shifts associated with the a and b carbon positions, the latter of which carries the largest positive charge, is taken as indicative of the degree of protonation: Dd ˆ dC…b† ÿ dC…a† Earlier work by these authors showed that Dd is linearly dependent on the concentration of the indicator and that in protonated form. The value of Dd at in®nite dilution (Dd ) of the indicator can be employed to develop an acidity function Dd ˆ s…‰MOŠ ‡ ‰MOH‡ Š† ‡ Dd where the slope s is related to the acidity. The authors ®nd that the three dissociation steps of HPW occur independently so that a molecule of the acid is equivalent to three molecules of strong acid in solution. The authors relate these observations to the size of the Keggin anion. HPW is found to be stronger than perchloric acid but the authors are careful to caution against concluding on this basis that HPW is a superacid. The 1H MAS NMR chemical shift has been proposed as a measure of acidity with an increase in this quantity corresponding to an increase in the Bronsted acid strength.(69±72) The results from the application of this suggestion to zeolites have been found to be in agreement with IR spectroscopic data.(73) The contention has also been examined and applied to HPW supported on silica and various salts of the heteropoly acids.(74±77) Four characterization techniquesÐthermal gravimetric analysis, powder XRD, MAS NMR, and inelastic neutron scatteringÐhave been employed to investigate the changes in HPW, Csx H3ÿx PW12 O40 and, in particular, Cs1:9 H1:1 PW12 O40 on heating.(78) Evidence for the presence of hydrated protons H‡ …H2 O†n is found after heating HPW and CsPW up to 373 and 323 K, respectively. This agrees with the results obtained earlier from photoacoustic FTIR and discussed earlier in this chapter.(30, 43) Not surprisingly, the strongest

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acidity was found in those samples in which the water had been removed. This has been discussed earlier.(79) As noted in an earlier section of this chapter, evidence for the terminal oxygen atoms of the KU as the predominant protonation sites has been obtained from 17O-enriched HPW and HSiW.(44) 8.1.7. Temperature-Programmed Desorption TPD can generate information on both the acid strength distribution of catalytically active sites and the thermal stability of adsorbed species.(80,81) The TPD pro®les for the acids containing phosphorus or silicon as central atoms and molybdenum or tungsten as peripheral metal atoms have been discussed in the section on stability. The TPD of heteropoly acids exposed to pyridine provides useful data on the acidities.(49) TPD pro®les for HPW pretreated in helium at 463 K followed by exposure to various quantities of pyridine at 298 K are presented in Fig. 8.8. Traces (a) to (d) of Fig. 8.8 represent desorption from aliquots of HPW containing previously sorbed 0.15, 0.38, 0.75, and 1.5 molecules of pyridine per anion, respectively. Trace (e) corresponds to HPW previously exposed to an excess of pyridine. The desorbing species, as measured by mass spectrometry, are shown under the peaks. Two cases are evident; that as in trace (e) in which pyridine was the major species desorbed and that in traces (a)±(c) where products CO2 , N2 and H2 arising from the decomposition of pyridine were desorbed. Trace (d) is evidently intermediate between the aforementioned. The effect of pretreatment temperatures is illustrated in Fig. 8.9. For these experiments the samples after pretreatment at the indicated temperature were exposed to pyridine at 298 K and either 0.38 pyridine molecule per anion (Fig. 8.9A) or an excess (Fig. 8.9B) were taken up. Where the quantity of desorbable material was relatively small (Fig. 8.9A), peaks similar to those in Fig. 8.8 were obtained with some shifting of peak maxima to lower temperatures as the outgassing temperature of HPW was increased. Where the samples had been exposed to excess pyridine, large quantities of pyridine were desorbed from HPW outgassed at 463 or 593 K (Fig. 8.9B). However, at pretreatment temperatures of 298 and 723 K some or all of the peaks associated with decomposition products were observed. Evidently the latter pretreatment temperature thermally modi®es the heteropoly acid. After pretreatment at 463 K and on exposure to pulses of pyridine at 298 K, relatively small quantities were sorbed on the Na and Mg salts of HPW and their TPD pro®les showed only residual traces of water at approximately 473 K and no peaks attributable to pyridine or its decomposition products. Ammonia at 593 or 693 K passed over HPW presorbed with pyridine to saturation levels was found to desorb pyridine into the gas phase. When followed

ACIDITY AND OXIDATION±REDUCTION

157

Figure 8.8. TPD of HPW after various pulses (5.9 mol per pulseÿ1 ) of pyridine sorbed at 298 K; number of pulses: (a) 2, (b) 5, (c) 10, (d) 20, and (e) sample saturated. Pretreatment at 463 K.(49)

by temperature programming, neither pyridine nor its decomposition products were detected but the TPD pro®les resembled those of NH4 PW. The quantities of pyridine irreversibly sorbed at 463 K after sorption of pyridine under static conditions at 298 K on HPW, NaPW, and MgPW are summarized in Table 8.2. At a temperature between 593 and 723 K, the quantity of pyridine sorbed on HPW decreased markedly.

Figure 8.9. Effect of pretreatment temperatures on subsequent TPD of pyridine from HPW. Pretreatment temperatures: (a) 298, (b) 463, (c) 593, (d) 723 K.(49) (A) Exposure to 0.38 pyridine molecule KUÿ1 at 298 K; (B) saturated with pyridine at 298 K.

158 CHAPTER 8

ACIDITY AND OXIDATION±REDUCTION

159

TABLE 8.2. Sorption of Pyridinea(49) Sample

Pretreatment temp (K)

HPW HPW HPW HPW NaPW MgPW

298 463 593 723 463 463

Pyridineb irreversibly sorbed 5.7 6.2 6.5 1.4 13.7 0.4

Ð (2.9) (2.4) (0.7) (0.2)c Ð

a

At 298 K. Amount irreversibly sorbed at 463 K. c Amount irreversibly sorbed after outgassing at 598 K. b

Because pyridine desorbs from HPW without decomposition to CO2 , H2 O, and N2 only where more than one pyridine molecule per anion is sorbed, comparisons of the acidic properties should be carried out under the latter conditions. The TPD of NH3 from HPW and its cesium and ammonium salts has been investigated and the results compared with sulfated zirconia.(82) Gaseous NH3 was found to react with HPW to form the ammonium salt as shown earlier in photoacoustic FTIR experiments.(30) For all of the heteropoly samples examined, the similarity of the TPD pro®les of NH3 desorption suggests that these are a re¯ection of the stabilities of the samples and not the acid strengths.(82) Some authors claim that hydrogen species (labeled as H*) that are more reactive than anticipated can exist on the heteropoly oxometalates.(83) The H* concentrations are dependent on the hydrogen pretreatment (TH2) temperature with little or no ``reactive hydrogen'' measured for TH2 less than 523 K (Table 8.3) For higher temperatures, anionic vacancies are created as a result of the elimination of water. This appears to be consistent with the results found earlier in studies of the partial oxidation of methane in which anionic vacancies resulting from the loss of water were related to the necessary presence of protons in the TABLE 8.3. Dependence of H* on the H2 Pretreatment Temperature(83) Pretreatment tempa Catalyst …NH4 †2:5 Cs1:6 P1:7 Mo11 V1:1 O40 …NH4 †3:5 Cs0:16 P1:3 Mo11 V1:2 O40 a

H* …10ÿ3 mol gÿ1 †:

523 K

593 K

623 K

0.04 0

1.2 1.4

1.3 1.3

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active catalyst.(84±87) In this earlier work, pretreatment in a reducing atmosphere was found to increase the activity of the catalyst while having little or no in¯uence on the selectivity.(84) Earlier TPD experiments showed that the protons contained within the heteropoly acids will extract oxygen atoms from the heteropoly anion and be desorbed as water.(49) This process occurs at approximately 773 K with HPW whereas with HPMo the maximum of the corresponding peak is found at approximately 673 K. This difference was attributed to the greater lability of the thermal oxygen in the anions of HPMo as contrasted with that in HPW as shown from extended HuÈckel calculations.(31±33) Photoacoustic FTIR spectra of these two acids demonstrated that the structure of the heteropoly anion was preserved at these temperatures in spite of the loss of up to two oxygen atoms per anion.(30,43) Temperature-programmed reduction experiments with hydrogen generated peaks of similar shape and locations to those obtained in the TPD experiments.(50) However, in the TPR experiments eight water molecules were desorbed at the higher temperature, showing that the added hydrogen performs similarly to the protons in stripping oxygen atoms from the anion. The authors argue that H2 splits heterolytically with one half of the hydrogen reservoir (H*) consisting of hydride ions (Hÿ ) and the remaining one-half as protons (H‡ ).(83) The numbers of H* species were determined from the hydrogenation of 2methylbuta-1,3-diene after pretreatment with H2 , producing anionic vacancies from the elimination of H2 O.(83) No hydrogenation activity was observed on the untreated solid. 8.1.8. Probe Reactions 8.1.8.1. Methanol The earlier studies of the conversion of methanol to hydrocarbons on various heteropoly acids and their salts together with the assessment of thermal stabilities and acidic strengths have shown that the tungsten-containing acids are more acidic than those containing molybdenum.(34,88±90) The conversion of methanol to dimethyl ether at 180  C and the isomerization of n-butane have been employed to estimate the acidity of Csx H3ÿx PW12 O40 salts.(78) On the basis of conversion of methanol the more acidic salts have values of x between 2 and 2.7. Those with very strong acidity as shown from the isomerization of butane have values of x from 2 to 2.1. 8.1.8.2. Isobutane The exchange and conversion of isobutane with deuterated HPW and Cs1:9 H1:1 PW12 O40 has shown that these compounds, while strongly acidic, are

ACIDITY AND OXIDATION±REDUCTION

161

not superacidic.(78b) Both surface and bulk structure deuterons in the D3 PW12 O40 sample exchange at 473 K with isobutane, but only with the hydrons in the a position to the branched carbon, in contrast with the observations with liquid superacids with which all the hydrons of the isobutane are exchanged. As noted elsewhere in this chapter, earlier TPD experiments(49) have shown that water is desorbed at two separate ranges of temperature, the lower resulting from the loss of hydrogen-bonded water, the higher due to water formed from the extraction of anionic oxygen atoms by the acidic protons as H3 PW12 O40 ! 1:5H2 O ‡ PW12 O38:5 Temperature-programmed exchange between D2 and HPW was found to begin at 623 K, a temperature approximately the same as that at which the aforementioned high-temperature peak emerged.(50) Under the conditions of the experiment, essentially complete exchange took place. The authors of the aforementioned isobutane exchange report employed a temperature of 673 K to pretreat HPW followed by deuteration with D2 O at 473 K. Further discussions of the conversion of isobutane are included in the chapter on acid-catalyzed processes. 8.1.8.3. Butanol The conversions of 1-, 2- and tert-butanols have been employed as a means of assessing the acidic strengths of various salts of HPW prepared from monovalent cations of groups 1, 11, and 13 of the periodic table.(75) Earlier work has shown that these salts have high surface areas and micro=mesoporous structures in contrast to their parent acids.(91±102) In addition, as the protons in the heteropoly acids are replaced by larger cations there is evidence that the distribution of the Bronsted acidic strengths is altered.(34,88±90) These effects appear to be dependent on both the nature of the substituting cations and the number of these relative to the remaining protons. Further evidence for these conclusions can be seen in Table 8.4. To eliminate variables other than the effect of the cations on the protons the conversions of the butanols have been divided by the surface areas and numbers of protons. The conversions of tert-butanol are generally higher than those for 2-butanol and those for the latter are higher than those for 1-butanol, for a given cation and stoichiometry, as expected for a carbocation mechanism. In addition, the conversion of a given alcohol increases, in general with the cation in the order Ag < Tl < Cs. Signi®cant changes are also observed for a given cation, as the preparative stoichiometry is altered. As expected the most signi®cant differences are found with 1-butanol. However, in general, the conversion decreases for a given cation and alcohol as the preparative stoichiometry increases.

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TABLE 8.4. Conversions of the 1-, 2- and tert-Butanols(75) Conversiona,b,c=[(m2 †…H‡ †Š

Preparative stoichiometry

1-Butanold

2-Butanold

tert-Butanole

Ag3 PW12 O40

0.85 1.00 1.15

2.1 2.1 0.9

8.7 8.4 11.6

10.8 12.6 10.0

Tl3 PW12 O40

0.85 1.00 1.15

1.4 0 0

21.1 42.2 1.4

25.8 28.1 31.9

Cs3 PW12 O40

0.85 1.00 1.15

4.2 0.8 0.6

28.9 38.2 11.6

33.2 35.5 51.9

Salt

a

For measurements taken at 5 min. Moles converted multiplied by 1028. c Relative number of protons from 1H MAS NMR data. d Reaction temperature of 108  C. e Reaction temperature of 46  C. b

8.1.8.4. Butene The effect of the nature of the cations and the stoichiometry of the salt has also been investigated with the conversion of 1-butene (Table 8.5).(77) cis- and trans-2-butene were the only products formed at 100, 200, and 300  C. The thallium salts generally produced higher conversions than the corresponding silver salts at each of the three reaction temperatures and the areal conversions per proton decreased with increasing cation=proton preparative ratio. TABLE 8.5. Conversions of the 1-Butene at 10 min with Salts of H3 PW12 O40 (77) Conversiona=[(m2 †…H‡ †Š

Preparative ratio

100  C

200  C

300  C

AgPW

0.50 1.00 1.50

3.4 1.8 1.9

3.7 2.9 0.2

3.2 2.0 1.6

TlPW

0.50 1.00 1.50

8.9 9.0 0.0

7.4 3.6 0.0

5.9 1.4 0.0

Salt

a

Moles of products multiplied by 1026 ; relative number of protons from 1H MAS NMR data.

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163

The variation in the acidic strengths of the stoichiometric and nonstoichiometric thallium (I) salts of HPW, HSiW and HPMo has been investigated with the isomerization of 1-butene (Table 8.6).(76) The only products at 100±300  C were cis- and trans-2-butene. The speci®c conversions at 100, 200, and 300  C on the TlPW salts are generally higher than those found with the corresponding salts of HPMo and HSiW. However, with the tungsten-containing salts the speci®c conversions at a given temperature decrease to insigni®cant values whereas with TlPMo the speci®c conversion increases as the cation=proton ratio increases. It should be recalled that, although the number of protons decreases as the preparative cation=proton ratio increases, signi®cant numbers remain at the highest values of this ratio (Table 8.6). As noted earlier in this chapter, semiempirical quantum-mechanical calculations showed that the magnitude of the negative charges on the terminal oxygen atoms of the tungsten-containing anions is lower than that of the molybdenum based materials. Because the mobility of the protons and, hence, their acidic strengths is expected to be inversely related to the aforementioned charges, the tungsten-containing heteropoly acids should be more strongly acidic than those with molybdenum. The observations from the studies of butene isomerization and dehydration of butanols suggest that the introduction of larger cations into the lattice structure perturbs the mobility of the protons. This may result from a perturbation of the electron densities and consequently the magnitude of the charge on the terminal

TABLE 8.6. Conversion of 1-Butene at 10 min with Thallium Salts(76) Conversiona,b,c=[(m2 †…H‡ †Š

Cation= proton ratio

100  C

200  C

300  C

TlPW

0.50 1.00 1.50

8.9 9.0 0.0

7.4 3.6 0.0

5.9 1.4 0.0

TlPMo

0.85 1.00 1.15

0.8 0.6 0.0

1.3 2.3 3.0

1.4 2.9 3.6

TlSiW

0.85 1.00 1.15

0.5 0.1 0.0

1.9 1.8 0.0

0.8 1.9 0.0

Salt

a

For measurements taken at 10 min. Moles converted multiplied by 1026. Relative number of protons from 1H MAS NMR data. d Reaction temperature. b c

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oxygen atoms by the nonprotonic cations. The dissimilar effects of the larger cation, thallium, on the tungsten- and molybdenum-containing salts appear to be dependent on the magnitude of the negative charge on the terminal oxygen atoms, producing a bene®cial effect where the magnitude is high and a disadvantageous effect with low magnitudes. Where the acidic strength of the parent acid is relatively high, the introduction of larger cations depresses the acidic strength whereas with the weaker parent acids the acidity is apparently enhanced.

8.1.8.5. 2-Methylpent-2-ene As noted elsewhere in this volume, a wide variety of techniques are available to provide information on the nature, numbers, and strengths of acidic sites on heterogeneous catalysts and solids in general. While each of these has its advantages and de®ciencies, the use of a probe reaction has much to commend it, particularly as a consequence of the evaluation being conducted under reactor conditions. Kramer and McVicker have recommended the isomerization of 2-methylpent-2-ene. (2M2PE) as an appropriate reaction for probing the relative acidity of solids.(103,104) The alkene is converted to a variety of products each re¯ecting the acidic strength required for their production. The reaction has been applied to three heteropoly acids HPW, HPMo, and HSiW each separately supported on silica at a loading of 20%.(105) This loading has been shown to correspond at least approximately to a monolayer.(106) Further, examination of the conversion of 2M2PE and the products formed at various loadings indicated that both the activity and the selectivities reached approximately constant values at 20% loading. In general, the results re¯ect the predictions that showed that the protons in the tungsten-containing acids would be expected to be more mobile and hence more acidic than those in the tungstencontaining solids.(31±33) As is evident from Fig. 8.10, the conversions and selectivities are similar for the tungsten-containing acids but quite dissimilar for the molybdenum-containing solid. In particular, the conversion of 2M2PE at 300  C for the former two acids is approximately twice that obtained with HPMo=SiO2 . The selectivities to products re¯ecting methyl group shiftsÐ3methylpent-2-ene, 2,3-dimethylbut-2-ene, and hexeneÐshow that the acidic strengths in HPW and HSiW are greater than those found in HPMo, whereas those involving double-bond shifts, for example, 4-methylpent-2-ene, show little or no dependence on the elemental compositions of the acids. These results suggest that the isomerization of 2M2PE on the heteropoly acids proceeds via a carbocation mechanism(103) with the latter being formed from the protons contained within the acids.

ACIDITY AND OXIDATION±REDUCTION

165

Figure 8.10. Conversions of 2-methylpent-2-ene and selectivities to products on silica-supported HPW, HPMo, and HSiW.(105)

Some exceptions to the aforementioned comments should be noted. The selectivity to 2-methylpent-1-ene. (2M1PE) is markedly higher with HPMo than with either HPW or HSiW, presumably as a consequence of the predominance of weaker acidic sites on HPMo and the probability that the 2,1 double bond shift would occur more readily than the 2,3 shift on HPMo. The isomer 2M1PE is the predominant product on HPMo at this temperature. The formation of the saturated 3-methylpentane is also of interest, although the selectivities to this compound on both HPW and HSiW are small, none was found with HPMo. Kramer and McVicker employed the ratio of the rates of formation of trans3-methylpent-2-ene to that of cis- and trans-4M2PE as a convenient parameter for assessing the acidity.(104) If the selectivities to these compounds are employed as convenient substitutes for the rates, values of approximately 2.7 and 0.7 are obtained for the tungsten- and molybdenum-containing acids, respectively. The former corresponds to an acidic strength greater than that of 1.5 wt % ¯uorided galumina.(104) At 250  C the equilibrium value of the aforementioned ratio is 2.70.

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8.2. OXIDATION±REDUCTION AND THE PROPERTIES OF THE ANIONIC OXYGEN ATOMS The relative binding strengths of the terminal and bridging anionic oxygen atoms in the Keggin structure and their dependence on the elemental composition of the anion are of considerable importance in understanding the oxidation± reduction process in the heteropoly oxometalates. IR spectroscopy of Ag3 PMo12 O40 subjected to H2 reduction and 18O2 oxidation cycles has shown that the oxygen of the central tetrahedron was not involved in the reduction process.(107) However, rapid exchange between the terminal and bridging oxygen takes place. Because, in the H2 ‡ 18 O2 reaction of AgPMo, 18O predominantly replaced the bridging oxygen, the reaction apparently occurs via redox cycles of these oxygen atoms.(108) The reduction of HPMo and its potassium salt has been investigated by following the changes in intensity of IR bands as the reduction proceeds.(109) Two bridging oxygen atoms per anion are ®rst lost followed by the further removal of terminal oxygen atoms. Lattice oxygen of HPW and HPMo was found to exchange rapidly with H2 18O.(110) However only 20±30% of the lattice oxygen of CsPMo exchanged rapidly while the remainder exchanged very slowly. In view of the relatively large surface area of the cesium salt, it appears that the rapid exchange can be assigned to the surface oxygen. Because HPMo is soluble in water whereas CsPMo is not, the slow exchange of oxygen from water within the interior of the cesium salt is not unexpected. The rate of exchange of the bridging oxygen is similar to that of the terminal oxygen. Reduction of HPMo with H2 is proposed to proceed via the reduction of the terminal oxygen atoms from the 6‡ to 5‡ state and the formation of protons which extract a bridging oxygen to form water. Information on the reactivity and lability of lattice oxygen, of relevance in selective oxidation processes, can be obtained from temperature-programmed reduction (TPR) with hydrogen and the mobility of protons can be assessed from temperature-programmed exchange (TPE) with deuterium.(50) All of the peaks observed in TPR of HPW, HSiW, and HPMo are due to water and bear a strong similarity to those observed in analogous TPD experiments that were performed under essentially identical conditions except that helium was employed as the carrier gas (Figs. 8.11±8.13).(50) Because peak 1 was previously attributed to the desorption of water molecules held on the catalyst in hydrogen-bonded form, the water molecules desorbed as peak 2 (together with peak 3 for HPMo) are summarized in Table 8.7.(50) As the values obtained from TPD correlate to the numbers of protons present initially in these acids, peak 2 in TPD was attributed to deprotonation of the acids with concurrent nonreductive loss of lattice oxygen.(50) TPR peak 1 was signi®cantly smaller than the corresponding peak

ACIDITY AND OXIDATION±REDUCTION

167

Figure 8.11. (a) TPR pro®le for HPW (b±d) Mole fractions D2 …g† (g), HD(g) (s), and H2 (g) (d) during TPE of HPW.(50)

in TPD experiments, possibly due to the relatively higher pretreatment temperature, in the TPR experiments, than those at which peak 1 appears. TPE experiments showed that exchange between D2 and HPW and HSiW occurred with the loss of D2 and the production of HD and H2 (Figs. 8.11b±d and 8.12b±d). No exchange of D2 was observed with HPMo but consumption of D2 and reduction of the acid were noted as with the corresponding TPR experiment. The inception of the exchange occurs at approximately 623 K, which is similar to the temperature at which peak 2 began to appear in the TPR experiments (Figs. 8.11a and 8.12a) and a maximum was attained at 675±700 K for both HPW and HSiW. HD was the principal product in the exchange, with the HD=H2 ratio never less than 7=1. Essentially complete exchange was found for both acids under the conditions of the experiment, showing that all protons were exchanged. Because the exchange was complete before the major portion of peak 2 had evolved, exchange was not occurring between H2 O and D2 . TPR experiments on NaPW and MgPW produced peaks due to the desorption of water but only the magnesium salt, which contained the only

168

CHAPTER 8

Figure 8.12. (a) TPR pro®le for HSiW. (b±d) Mole fractions D2 …g† (g), HD(g) (s), and H2 (g) (d) during TPE of HSiW.(50)

divalent cation, desorbed signi®cant quantities of water near the temperature at which peak 2 for HPW was observed (Fig. 8.13). As with the acids, continuous reduction was observed, but with the Na and Mg salts the onset of this occurred at considerably higher temperatures. With these salts strong peaks were seen at 923 K, apparently due to decomposition of the salt. Little or no exchange of D2 with these salts was observed. The results of TPR and TPE experiments on the pyridinium and ammonium salts of HPW are shown in Figs. 8.14 and 8.15.(50) With the former the peak at approximately 823 K resulted from pyridine and water. The maximum exchange rate for the pyridinium salt occurred at approximately the same temperature as that due to the desorption of water and pyridine in the TPR experiments. Evidently pyridine interacts strongly with the protons and exchange can only occur after removal of the pyridine. The exchange peak at 623 K may be due to the presence of some unreacted free acid.

ACIDITY AND OXIDATION±REDUCTION

169

Figure 8.13. TPR pro®les for HPMo, NaPW, and MgPW.(50)

With the ammonium salt no signi®cant peaks emerged under TPR although continuous reduction began at approximately 673 K. TPE of D2 began at approximately 673 K and reached a maximum at approximately 873 K, considerably higher than that observed with the parent acid.

TABLE 8.7. Desorption of Water from TPD and TPR(49,50) Acid HPW HSiW HPMo a b c

TPDa

TPRa

1.4 1.9 1.5

8.6b 9.4b 8.0c

Water molecules per anion (KU) From peak 2. From peaks 2 ‡ 3.

170

CHAPTER 8

Figure 8.14. (a) TPR pro®le for PyPW. (b±d) Mole fractions D2 …g† (g), HD(g) (s), and H2 …g† (d) during TPE of PyPW.(50)

Because a ¯ow system was employed in this work, the results can be taken to re¯ect the initial product distribution the equations for which have been suggested as(111) D2 …g† ‡ H…s† ! HD…g† ‡ D…s† D2 …g† ‡ 2H…s† ! H2 …g† ‡ 2C…s† In view of the substantially greater quantities of HD than H2 that were produced, the former stoichiometry appears to be the more probable. TPR peaks 1 and 2 in the tungsten-containing compounds are believed to result from dissociation of H2 to form protons which then strip oxygen atoms from the anions as O2ÿ

H2 …g† ! 2H‡ …ads† ! H2 O

ACIDITY AND OXIDATION±REDUCTION

171

Figure 8.15. (a) TPR pro®le for NHPW (b±d) Mole fractions D2 …g† (g), HD(g) (s), and H2 …g† (d) during TPE of NHPW.(50)

with peak 1 due to the reaction of protons with extremely labile lattice oxygen. The water so formed then remained in the structure until the temperature of decomposition of (H5 O2 †‡ ions occurred as shown by TPD.(50) The exchange process could occur in a similar fashion as D2 …g† ‡ H‡ …s† ! 2D‡ …ads† ‡ H‡ …s† ! HD…g† ‡ D…s†

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4. 5. 6. 7. 8. 9.

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9 ACID-CATALYZED PROCESSES

9.1. METHANOL CONVERSION TO HYDROCARBONS 9.1.1. Heteropoly Acids The process in which methanol is converted to hydrocarbons has been well studied from both practical and fundamental viewpoints although some details of the mechanism remain to be established.…1† While the preponderance of the earlier work was concerned with zeolites as catalysts and particularly ZSM-5, the reaction has been employed as a useful probe of catalytic properties. The conversion of methanol has been studied on 12-molybdophosphoric, 12-tungstophosphoric, and 12-tungstosilicic acids under a variety of conditions.…2;3† At temperatures between 325 and 403  C the conversion of methanol is 90% or higher (Fig. 9.1). At all three temperatures dimethyl ether (DME) forms at the lowest residence times (W=F) and decreases as W=F increases. Hydrocarbons form at a W=F equal to approximately 100 mg catalyst min mlÿ1 He at the lowest temperature and their yields pass through maxima as W=F increases. At a reaction temperature of 352  C hydrocarbons begin to appear at a W=F of 20 mg catalyst min mlÿ1 He and the dominant hydrocarbon is ethylene although methane is produced in almost equal amounts. The yields of hydrocarbons at 352  C are generally higher than those found at 325  C whereas at 403  C the yields of methane are considerably higher than those at the two lower temperatures. As noted elsewhere in this volume, thermogram and powder XRD data indicated that H3 PW12 O40
24H2 O forms anhydrous acid at 350±450  C without decomposition.…2† The pretreatment conditions were, however, found to have considerable in¯uence on the results (Fig. 9.2). The yields of hydrocarbons obtained at 352  C with HPW pretreated in helium were found to vary considerably with calcination temperature. After pretreatment in helium at 350 or 400  C 175

176

CHAPTER 9

Figure 9.1. Effect of residence time on the conversion of methanol over H3 PW12 O40 at various temperatures. W=F, apparent residence time in mg catalyst min mlÿ1 He; dashed lines, equilibrium for DME formation.…2;3†

Figure 9.2. Effect of calcination conditions on the conversion of methanol over H3 PW12 O40 at 352  C. W=F, apparent residence time in mg catalyst min mlÿ1 He.…2;3†

ACID-CATALYZED PROCESSES 177

178

CHAPTER 9

TABLE 9.1. Product Distributions from 18-Tungstodiphosphoric Acid and 12-Molybdophosphoric Acid…3† H6 P2 W18 O62 a Calcined in Methanol converted (%) Dimethyl ether produced (%) CO CO2 CH4 C2 C3 C4 C5 a b

He 99.0 56.0 5.5 0.9 1.2 Ð Ð Ð Ð

H2 94.2 41.8 0.1 Ð 4.1 6.1 9.1 13.7 8.7

H3 PMo12 O40 b He 98.0 25.0 61.7 3.9 Ð Ð Ð Ð Ð

H2 100.0 26.1 30.4 2.9 0.9 0.6 Ð Ð Ð

Calcined at 400  C, reaction at 350  C, W=F ˆ 99. Calcined at 350  C, reaction at 350  C, W=F ˆ 99.

the C4 hydrocarbons appeared in larger quantities than any of the other measured hydrocarbons; with the 450  C pretreatment, methane is produced in the largest quantities. In contrast, pretreatment in hydrogen at 450  C produced the highest yield of C4 hydrocarbons. The maxima in yields of hydrocarbons observed after pretreatment in helium are also observed after pretreatment in hydrogen. Pretreatment in air evidently has a strongly deleterious effect. The importance of the elemental composition of the anion in the methanol conversion process is illustrated in Table 9.1. Whereas the conversion of methanol is high on 12-molybdophosphoric acid, the products are predominantly those expected from oxidation.

9.1.2. Metallic Salts of 12-Tungstophosphoric Acid The sodium, boron, calcium, magnesium, zinc, aluminum, and zirconium salts of HPW have been studied as catalysts for the conversion of methanol.…3;4† The conversion of methanol and the yields of hydrocarbons vary widely with the nature of the cation (Fig. 9.3). However, the yields of hydrocarbons with two or more carbon atoms can be seen to correlate, at least approximately, with the highest acidic strengths (Fig. 9.4), with aluminum generating the highest yield of C4 hydrocarbons. As noted elsewhere in this volume, the Bronsted acidity is expected to decrease with the increase in the magnitude of the negative charge on the terminal oxygen atoms of the anion (ÿdo ). The maximum yield of C4 hydrocarbons decreases from that obtained by the Al salt to that for Na as the magnitude of the charge on the terminal oxygen increases. Salts with values of ÿdo less than that for Al generate C4 yields that are less than those for Al

Figure 9.3. Methanol conversion over metal salts of 12-tungstophosphoric acid. Reaction temperature: 350  C. Catalyst calcined at 400  C in He for 2 h.…3;4†

ACID-CATALYZED PROCESSES 179

180

CHAPTER 9

Figure 9.4. Maximum yield for C4 hydrocarbons in methanol conversion and the partial charge on oxygen (do ) in the metal salts of 12-tungstophosphoric acid.…3;4†

possibly due to irreversible chemisorption of products and=or their precursors leading to deactivation of the surface. 9.1.3. Ammonium 12-Tungstophosphate Methanol conversion has also been studied with ammonium 12tungstophosphate.…3;5† As discussed in the section on microporous compounds, ammonium 12-tungstophosphate (abbreviated as NH4 PW) can be prepared with high surface area and microporous structure.…6ÿ14† With pretreatment the surface area increases to approximately 163 m2 gÿ1 at 300  C and thereafter decreases with further heating to 600  C. At 350  C the yield of C1 ±C5 hydrocarbons was 76.9% for NH4 PW, in contrast to 32.4% obtained with HPW (Fig. 9.5). The major products were C4 hydrocarbons with iso content of 75% at 400  C and 86% at 325  C. Whereas with HPW the products are largely ole®ns, with NH4 PW these are predominantly paraf®nic, although a small amount of ethylene is formed (Fig. 9.5). As with the parent acid, pretreatment conditions alter the yield of hydrocarbons, but with the ammonium salt the effect is less substantial. At a reaction temperature of 350  C the yield of C4 hydrocarbons reaches 40±50% at pretreatment temperatures of 400±500  C regardless of the environment (air, H2 , He) in which the pretreatment occurs. With pretreatment in air or H2 at 400  C or in He

ACID-CATALYZED PROCESSES

181

Figure 9.5. Effect of reaction temperature on the methanol conversion over H3 PW12 O40 (a) and …NH4 †3 PW12 O40 (b). Catalysts: calcined at 400  C in helium for 2 h; W=F ˆ 246 mg catalyst min mlÿ1 He.…3;5†

182

CHAPTER 9

Figure 9.6. Typical desorption curves for CH3 OH on 12-tungstophosphoric acid (preevacuated at 330±350  C), dosed in excess at 25  C and evacuated at 25  C.…17†

at 500  C a maximum in the C4 yield occurs at values of W=F between 100 and 300 mg catalyst min mlÿ1 He, whereas no maximum is apparent up to W=F equal to 600 after pretreatment in He at 400  C. Although stoichiometric quantities of the preparative reagents were employed in the synthesis of the NH4 PW samples used in the methanol studies, as has been discussed elsewhere in this volume, residual protons have been shown to remain.…15;16† The distribution of acidic strengths of these protons is, not surprisingly, different from that of the protons in the parent acid. This undoubtedly accounts, at least in part, for the dissimilar observations in methanol conversion. The ammonium salt produces less carbon and consequently high yields. Although the higher surface area for NH4 PW could contribute to the superior results, this has been shown not to be so.…5†

9.1.4. Mechanistic Studies with Photoacoustic FTIR Spectroscopy Information on the methanol conversion process can be obtained from PAS FTIR spectroscopy.…17;18† On exposure of HPW (preevacuated at 330±350  C) to an excess of methanol at 25  C, up to 8 molecules of methanol per anion are rapidly taken up (Fig. 9.6). On evacuation the number of methanol molecules sorbed per anion decreases to 3. This indicates that a 1 : 1 correspondence exists

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between the methanol sorbed and the protons contained in the catalyst, as indicated by the stoichiometry H3 PW12 O40 . Because the surface area of HPW is small (< 10 m2 gÿ1 ), most of the protons must reside within the bulk of the structure and consequently the methanol is evidently interacting with both the surface and bulk protons. The PAS FTIR spectrum of HPW with 3 molecules of methanol chemisorbed per anion (Fig. 9.7b) is signi®cantly different from that of the acid itself (Fig. 9.7a). The characteristic bands of the anion (below 1200 cmÿ1 ) remain along with new bands in the region 1750±1300 cmÿ1 superimposed on a background continuum similar to that previously observed in the hydrated acid (see elsewhere in this volume). The large bandwidths of the methanol deformation vibrations at 1535 and 1405 cmÿ1 (Fig. 9.7c) show that the sorbate is subject to

Figure 9.7. (a) 12-Tungstophosphoric acid, preevacuated at 350  C, then (b) exposed to excess CH3 OH at 25  C for 1 h, and reevacuated at 25  C for 2 h; (c) detail of (b); (d) corresponding detail of spectrum derived from equivalent treatment of the acid with CD3 OH shown for comparison.…17†

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strong H-bonding interaction. This contrasts with the relatively weak interaction observed with methanol and the sodium salt of HPW (not shown). The stoichiometry of the sorbed methanol (3 molecules per anion) suggests the formation of protonated methanol, CH3 OH2 ‡ . For the structure

H3C O

H H

+

two active IR deformation modes (asymmetric and symmetric) involving the COH group would be expected. The spectrum of sorbed CD3 OH (Fig. 9.7d) shows a broad band centered at  1460 cmÿ1 which is attributed to the asymmetrical mode. The band at 1535 cmÿ1 (Fig. 9.7c) is assigned to the asymmetric COH deformation in CH3 OH2 ‡ while the symmetric mode is believed to lie at approximately 1430 cmÿ1 . Stepwise heating of HPW containing 3 molecules of methanol per anion resulted in the partial desorption of the methanol as well as related dissociation processes (Fig. 9.8). While the broad bands at 1535 and 1405 cmÿ1 diminish in intensity, a sharp band develops at 1453 cmÿ1 attributed to the CH3 symmetric deformation in the CH3 O group. The progressive development of a band at 1022 cmÿ1 (inset, Fig. 9.8) provides evidence for the increasing formation of a metal alkoxide structure. The C±O stretch in transition metal methoxides is found in the region 1100±1000 cmÿ1…19† and, in particular, at 1070 cmÿ1 in W(OCH3 †6 .…20† After exposure of the sample (150  C pretreatment, Fig. 9.8d) to excess CH3 OH at 25  C for 5 min, a signi®cant amount of DME (approximately 0.6 KUÿ1 ) was recovered, demonstrating that the active intermediate in the dehydration of CH3 OH has already been produced in the thermal pretreatment step and that it reacts rapidly with CH3 OH in a subsequent step. The PAS FTIR spectrum after exposing the sample whose spectrum was shown in Fig. 9.8d to excess CH3 OH at 25  C and evacuation at 25  C shows that the band at 1453 cmÿ1 remains, although probably diminished in intensity, while the broad band centered at approximately 1600 cmÿ1 which is characteristic of CH3 OH2 ‡ is regenerated. Separate experiments showed that at 25  C ammonia displaces the sorbed methanol easily with the accompanying formation of the ammonium ion, providing further evidence for the interaction of alcohol with and hydrogenbonding to the proton of the catalyst while remaining in undissociated form. When HPW with presorbed CH3 OH is heated to 150  C followed by exposure to ammonia at 100  C, CH3 NH3 ‡ is formed, as evidenced from the PAS spectra. Approximately stoichiometric (3.00:1 NH3 KUÿ1 ) quantities of ammonia were sorbed. The aforementioned observations provide further support for the conten-

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Figure 9.8. Effect of stepwise heating in vacuo on spectrum of ``irreversibly sorbed'' CH3 OH on 12tungstophosphoric acid. (a) 50  C, (b) 70  C, (c) 110  C, (d) 150  C [inset peak obtained by subtraction of spectrum of preevacuated acid (Fig. 9.7a), normalized at 1080 cmÿ1 ], (e) effect of dosing (d) with excess CH3 OH at 25  C and evacuation at 25  C.…17†

tion that CH3 OH2 ‡ dissociates on heating to produce chemisorbed methyl and=or methoxy groups. The initial steps in the methanol process may thus be summarized in the following manner. Methanol penetrates into the bulk structure of the solid where it is protonated at the oxygen atom: 25  C

! CH3 OH2 ‡ …OKU † CH3 OH ‡ H‡ …OKU † ÿ fast

…1†

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Figure 9.9. Assisted methanolic C±O bond cleavage.…17†

The electron density is suf®ciently perturbed to weaken the C±O bond which dissociates at slightly elevated temperatures to produce water and a methyl cation 50ÿ150  C

ÿ! CH3 ‡ ‡ H2 O CH3 OH2 ‡ ÿÿÿ slow

…2†

the latter of which associates with the negatively charged terminal oxygen atom of the anion (Keggin Unit): ! CH3 Oÿ KU CH3 ‡ ‡ O2ÿ KU ÿ fast

…3†

DME is formed from the interaction of methanol and the surface CH3 group 25  C

! CH3 OCH3 ‡ H‡


O2ÿ KU CH3 Oÿ KU ‡ CH3 OH ÿ fast

…4†

The methanolic C±O bond cleavage and the methylation of the surface may occur as a concerted process rather than in the separate steps (2) and (3), with a neighboring oxygen ion in the Keggin anion assisting in the scission by electron donation to the carbon atom as shown (Fig. 9.9). The conversion of C2 ±C4 alcohols to hydrocarbons on HPW has also been studied by PAS FTIR.…18† Dehydration of C2 ±C4 alcohols occurs on Bronsted acid sites by a mechanism similar to that for CH3 OH. The sorbed alcohol is rapidly protonated after which thermally induced C±O bond cleavage occurs. The carbene (:CH2 ) mechanism of initial C±C bond formation from sorbed CH3 ‡ is tentatively favored over the onium ylide route. 9.1.5. Alkylammonium Oxometalates The ammonium and alkylammonium salts of heteropoly acids with either phosphorus or arsenic as central atoms and molybdenum or tungsten as peripheral metal atoms have been studied as catalysts for the conversion of methanol.…21† As noted elsewhere in this volume, the ammonium salts of HPW and HPMo have been shown to have high surface areas and micro=mesoporous structures whereas

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the alkylammonium salts have low surface areas (< 10 m2 gÿ1 ) and no porosity.…6ÿ14† It should be noted that the presence of cations of size greater than CH3 NH3 ‡ resulted in crystal structures other than the cubic (Pn3m) form reported for the Keggin structures.…22† Results obtained for methanol conversion at 573  K in a continuous ¯ow reactor with NH4 PW pretreated at 673  K in helium show that whereas the activity decreases by only 15% between initial and steady-state conditions, the product distributions change markedly with the selectivities to hydrocarbons decreasing while that to DME increases (Fig. 9.10).

f

Figure 9.10. Methanol conversion over NH4 PW; continuous ¯ow reactor, 573 K, 0.249 g catalyst min mlÿ1 MeOH, pretreated in helium at 673 K; (d) conversion, (s) dimethyl ether, (l) C1 , ( ) C2 , (u) C3 , (j) C4 , and (g) C5 .…21†

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Under pulse reactor conditions the conversion and C1 ±C3 yields obtained for Me4 NPW are similar to those found with HPW but the yields of C4 and C5 are smaller than those for HPW and NH4 PW (Table 9.2). The remaining alkylammonium 12-tungstophosphates produced results similar to those with Me4 NPW. A comparison of the results from methanol conversion over a number of catalysts shows the dependencies on both the cation and anion compositions (Table 9.3). The presence of Bronsted acid sites is a necessary condition for the catalysis to occur. PAS FTIR measurements have shown that a one-to-one interaction between methanol molecules and all protons, both surface and bulk, occurs with the formation of CH3 OH2 ‡ .…15ÿ18† However, the selectivity to hydrocarbons decreases with time on stream either as a result of blocking of the acidic sites by irreversibly chemisorbed species and=or coking or the irreversible consumption of the protons. The data suggest that the more acidic sites are preferentially and irreversibly consumed while the less acidic are retained. As has been shown directly by PAS FTIR measurements, the methanol molecules interact with the protons both on the surface and in the bulk of the catalyst to methylate the terminal oxygen atoms of the Keggin anion.…17ÿ18† Migration of carbocations from the bulk to the surface where protons are released may lead to a surplus of protons relative to the availability of surface sites and a restricted ability of the surplus to diffuse to receptor sites in the bulk structure. Another possibility exists. Although the number of protons in the catalyst may remain unchanged during the catalysis, the distribution of acidic strengths may change as a result of protons which formerly were mobile and hence acidic TABLE 9.2. Catalytic Dataa for Methanol Conversion over Tetramethylammonium 12-Tungstophosphate…21† Yield (as % Cl) Product C1 C2 C3 C4 C5 DME Conversion (%)g a

Me4 NPW

HPWb

NH4 PWc

4.9 5.9d 5.9d 5.0 f 0.8 f 3.8 95.7

4.0 5.8d 6.2d 12.0f 1.8 12.0 95.0

2.5 3.1d 18.0e 45.2f 6.7 0 100

Pulse reactor, reaction temp ˆ 623 K. From Ref. 2, W=F ˆ 400 g catalyst min mlÿ1 . From Ref. 5, W=F ˆ 246 g catalyst min mlÿ1 . d Ole®n predominates. e Alkane predominates. f Iso-alkane predominates. g Including carbonization. b c

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TABLE 9.3. Summary of Catalytic Selectivities of 12-Heteropoly and Related Compounds for Methanol Conversiona…21† Major product Hydrocarbons (C1 ±C6 ) Compound NH4 PWb NH4 AsWb AgPWb AgPWc HPWb HPMob H6 P2 W18 Ob62

Initial

Prolonged

Oxidation (CO, CO2 )

References

‡ ÿ ÿ ‡ ‡ ÿ ÿ

ÿ ÿ ‡ ‡ ‡ ÿ ÿ

ÿ ‡ ÿ ÿ ÿ ‡ ‡

21, 5 21 23, 24 23, 24 2, 25 2, 3 2, 3

a

Excluding dehydration or esteri®cation products. Pretreated in helium. c Pretreated in hydrogen. b

becoming more strongly attached to their receptor sites (negatively charged terminal oxygen atoms) and thus less mobile and less acidic.

9.2. ETHANOL The most stable of the salts of HPMo, NH4 PMo, and Me4 NPMo were examined for ethanol conversion under various pretreatment and reaction conditions.…21† The predominant products were ethylene (C02 ), ethane (C2 ), and acetaldehyde. Only trace amounts of diethyl ether (DEE) were found under the conditions employed, in contrast to results from independent authors in which DEE was reported as a major product from ethanol over HPW.…26† Pretreatment conditions, as with methanol, have a substantial effect on the selectivities from ethanol over NH4 PMo (Fig. 9.11). Although pretreatment at 623 K in helium resulted in relatively poor yields, activation in air or hydrogen produced signi®cantly increased selectivities to acetaldehyde and ethylene, respectively. Catalyst samples activated in hydrogen retained their activities with time on stream. With increasing pretreatment temperature and a hydrogen atmosphere the conversion of ethanol and selectivity to ethylene increased whereas those to ethane and acetaldehyde decreased (Fig. 9.12). The overall activity of Me4 NPMo was approximately 25±50% of that of NH4 PMo although the effects of pretreatment temperature were similar (Table 9.4). The effect of hydrogen pretreatment may be attributed to the reduction of molybdenum from 6‡ to 5‡. Acetaldehyde and water are produced from ethanol with the concomitant removal of oxygen from the anion, a process that may result in

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Figure 9.11. Activation environment effects on ethanol conversion over NH4 PMo; continuous ¯ow reactor, 523 K, 0.46 g catalyst min mlÿ1 EtOH.…21†

f

Figure 9.12. Activation temperature effects (in hydrogen) on ethanol conversion over NH4 PMo; continuous ¯ow reactor, 498 K, 1.61 g catalyst min mlÿ1 EtOH; (u) conversion (omitting carbonization), (d) C02 , ( ) C2 , and (s) acetaldehyde.…21†

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TABLE 9.4. Comparisona of Activities and Selectivities of NH4 PMo and Me4 NPMo for Ethanol Conversion…21†

Conv.b C02 c C2 c CH3 CHOc

NH4 PMo (623 K, H2 )

Me4 NPMo (623 K, H2 )

NH4 PMo (693 K, H2 )

Me4 NPMo (693 K, H2 )

38.3 15.3 32.9 51.8

9.7 27.6 30.6 41.8

68.2 56.6 13.0 30.0

35.9 47.0 15.4 37.6

a

Continuous ¯ow reactor, reaction tempˆ 498 K, W=Fˆ1.61 g catalyst min mlÿ1 EtOH. Excluding carbonization products. c Selectivities (%). b

the maintenance of the catalyst in a reduced state. Ethane results from the secondary hydrogenation of ethylene, the expected product from ethanol. Photoacoustic FTIR spectra of the interaction of ethanol with HPW show some resemblances to those with methanol as well as some differences.…18† The spectra show that, as with methanol, the decomposition of ethanol begins with its protonation, followed by dehydration as a result of C±O bond cleavage with concomitant ethylation of the solid, which can be represented as 30  C

2ÿ C2 H5 OH ‡ H‡


O2ÿ KU ÿ! C2 H5 OH‡ 2


OKU D

ÿÿ! ‰ÿ OC2 H5 ŠKU ‡ H2 O O2ÿ KU


C2 H5 OH2 ‡ ÿÿ 50ÿ110  C

…5† …6†

where, as before, O2ÿ KU represents an oxygen ion in the Keggin unit (anion). The aforementioned process is evidently analogous to that for methanol and suggests that alkylation of the catalyst is a vital intermediate step in the dehydration of all alcohols. The catalytic cycle may be completed by transfer of a proton to form the alkene D

‰ÿ OC2 H5 ŠKU ÿ! C2 H4 ‡ H‡


O2ÿ KU

…7†

and=or interaction of gas-phase ethanol with the ethylated anion to form the corresponding ether ‰ÿ OC2 H5 ŠKU ‡ C2 H5 OH ! …C2 H5 †2 O ‡ H‡


O2ÿ KU

…8†

With ethanol the formation of ethylene is favored probably due to the facile nature of step (7). The spectra obtained by heating the HPW, previously exposed to ethanol, to 150  C show evidence for both chain growth and chain branching in a surprisingly low temperature range, 110±150  C. This implies that a process not

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observed with methanol is operative with ethanol, probably resulting from the interaction of the product ole®n with the ethylated anion as C2 H4 ‡ C2 H5 Oÿ KU ! CH3 ÿ …CH2 †3 ÿ Oÿ KU CH3 ÿ …CH2 †3 ÿ Oÿ KU ! …CH3 †2 ÿ CH ÿ CH2 Oÿ KU

…9† …10†

Supporting evidence for chain growth and rearrangement was obtained by analyses of the products desorbed at temperatures greater than 110  C as well as experiments in which the ethylated acid was exposed to excess C2 H4 . The PAS spectra from the latter clearly show the development of bands attributable to CH2 and on further heating up to 150  C the characteristic spectrum of the iso species is obtained. Additional PAS FTIR studies provided further evidence for the tertbutylation of the anion. After desorption at 205  C, the quantities of the products recovered were in the order C2 H4 > n-C4 H10  C3 H8 > i-C4 H10 > C2 H6. For temperatures higher than 250  C the order was C3 H8 > n-C4 H10  C2 H6 > C2 H4 .

9.3. PROPENE The oligomerization of propene to distillate fuels has been shown to be catalyzed by HSiW.…27ÿ29† As found in the alcohol-to-hydrocarbons process the catalytic activity was shown to increase with calcination temperature up to 240  C. The activity was also found to be maintained for longer periods of time with HSiW supported on bauxite.…28† At 170±180  C and a pressure of 12 MPa conversions of propene as high as 95% were obtained with HSiW. The oligomerization of propene on the ammonium, potassium, nickel, copper, cobalt, iron, cerium, and aluminum salts of HPW and the ammonium and aluminum salts of HSiW has been studied at 5 MPa.…30† The trimer of propene was the principal product. The salts of the monovalent cations NH4 ‡ and K‡ had surface areas of 37 and 140 m2 gÿ1 , respectively, whereas the remainder were less than 15 m2 gÿ1 , in semiquantitative agreement with earlier work.…6ÿ14† After calcination at 325  C, HPMo, HSiW, and NH4 SiW had little or no activity whereas AlSiW showed activity at 220±230  C. The activity of the HPW salts decreased in the order Al  Co > Ni, NiH, NH4 > H, Cu > Fe, Ce > K. The results obtained after the addition of sand to dilute the catalyst bed are of particular interest (Fig. 9.13). The conversion of propene can be seen to be maintained for longer periods of time and that with AlPW, as prepared from the nitrate, remains virtually complete for approximately 35 h. The conversions followed similar trends to those obtained with the undiluted acids: AlPW > CuPW > FePW.

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Figure 9.13. The effect of dilution with sand on the activity of AlPW, CuPW, and FePW in the oligomerization of propene.…30† Reproduced by permission of Academic Press.

HPW and HSiW supported on a-alumina (25 wt % loading) produced higher conversions of propene for longer periods of time than the unsupported acids. The conversion on HSiW supported on silicotungsten remained constant at the initial values. The C12 ‡ yield was found to be a function of conversion and not the cation contained in the catalyst (Fig. 9.14). It is of interest to note that the highest yield was obtained with AlPW, in agreement with the results obtained for the various salts of HPW in the methanol conversion process.…4†

9.4. PROPANOL PAS FTIR studies, similar to those with methanol and ethanol, provided similar but not identical conclusions.…18† As with the lower alcohols after sorption of n-C3 H7 OH and i-C3 H7 OH at 25  C on HPW followed by evacuation at the same temperature, approximately 2.5 molecules KUÿ1 remained on the catalyst.

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Figure 9.14. Yield of C12 ‡ is a function of conversion of propene but not of the nature of the cation.…30† Reproduced by permission of Academic Press.

Spectral evidence for protonation of the alcohols and C±O bond cleavage followed by alkylation of the anion was again obtained. 9.5. BUTANE As noted elsewhere in this volume, the cesium salts of the common heteropoly acids have high surface areas and microporous structures, both of which are dependent on the Cs‡ =H‡ stoichiometry.…6ÿ14† The nonstoichiometric cesium salts of HPW, Csx H3ÿx PW12 O40 , were investigated for the skeletal isomerization of n-butane.…31† The activity was found to be dependent on the cesium content, increasing to a maximum at x equal to 2.5. The conversions of nbutane at 300  C on the 2.5 catalyst and SO4 2ÿ =ZrO2 were reported as approximately 15 and 25%, respectively, in the initial stages of the reaction but after 3 h on stream the values had decreased to 7.5 and 4%, respectively. Selectivities to isobutane on the two catalysts were given as 83 and 61%, respectively. The conversion of butane was employed to test the surface acidity of a number of salts of HPW and HSiW, as referred to in the section on propene.…30†

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The authors classi®ed their salts as type A or B, the former being those with low surface area and a number of endothermic peaks associated with loss of mass as measured from TG-DTA, while the latter have high surface area and only one endothermic mass loss. HPW and AlPW, the latter type A prepared from aluminum nitrate, did not crack butane between 350 and 450  C whereas NH4 PW (type B) converted 1.9% of butane at 400  C, largely to isobutane. A catalyst was prepared by the addition of an aqueous solution of Pt(NH3 )4 Cl2 to an aqueous solution of HPW followed by the addition of an aqueous solution of Cs2 CO3 at 50  C.…32;33† The resulting suspension was reduced to a solid by evaporation at 50  C. The authors state that the molar ratio was 0.25=2.5=1 for Pt2‡ , Cs‡ , PW, respectively. Although IR spectra show that the Keggin structure is present, the nature of the material is not clear. The authors claim that the presence of Pt suppressed the catalyst deactivation in the conversion of n-butane but unfortunately time-on-stream data were not supplied. Another study of the properties of the cesium salts of HPW also investigated the isomerization of n-butane.…34† These authors carefully investigated the properties of the samples by chemical analysis, TGA, NH3 adsorption±desorption, powder XRD, and 31 P MAS NMR spectroscopy. The Cs2 HPW12 O40 preparation appeared to be the more ef®cient catalyst as a consequence of its high acidic strength and high surface area. When expressed per unit mass this sample produced the highest rate of formation of isobutane, although on an areal basis, the parent acid occupied this position. An interesting paper examined the effect of mechanical mixtures and grinding of HPW and its cesium salt.…35† Because rapid deactivation was observed for all samples, catalytic activities after 4 min on stream were reported. The initial catalytic activity in the isomerization of n-butane was shown to be higher after grinding in comparison with the aforementioned constituents of the mechanical mixture. The authors concluded that the acid is well dispersed on the higher area cesium salt. It should be noted that the rate of formation of isobutane increased with the duration of grinding up to 6 min at which time a plateau apparently was reached. It is also worth noting that mild mixing produced a catalyst with higher activity than that of the constituents of the mixture. The isomerization of n-butane has also been studied on HPW supported on a variety of porous silica-based materials.…36† Three lamellar silicatesÐa natural montmorillonite with high alkaline metal content, an activated montmorillonite strongly delaminated with low alkaline metal content, a kenyaite (a lamellar sodium silicate) together with a mesoporous hexagonal silicaÐwere used as supports. In H=D exchange studies which are discussed elsewhere in this volume, HPW and Cs1:9 H1:1 PW12 O40 in their anhydrous forms were found to be very active for the isomerization of n-butane at 473 K, with a conversion of 12.8% and selectivity to isobutane of 84%.…37† In the presence of 0.67 kPa of water in the

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feedstream, the conversion of butane decreased to 4.7% while the selectivity to isobutane increased to 92.5%.

9.6. BUTENE The surface acidity of a variety of salts of HPW and HSiW has been studied with the isomerization of butene.…30† NH4 PW and KPW initially produced cracking, double-bond and skeletal isomerization products but with increased time on stream the double-bond isomerization became dominant. The HPW salts of Ca, Co, Ni, Ce, Fe, and Al, although similar in activity to the aforementioned, produced only double-bond shift products. Unsupported 12-tungstophosphoric acid and its ammonium salt produced high conversions (80±90%) of 1-butene but cis- and trans-2-butene were the principal products.…38† Silica-supported HPW (HPW=SiO2 ) yielded appreciable quantities of isobutene. The conversion of butene changes relatively little with the loading of HPW=SiO2 but the conversions at the smallest loading, at any of the four reaction temperatures, are smaller than those found at the higher two loadings (Fig. 9.15). In contrast, the selectivities to isobutene are strongly dependent on the loading, at all temperatures, increasing to a maximum of approximately 33% on 23% HPW=SiO2 at 300±350  C. A similar trend with catalyst loading was found for the partial oxidation of methane…39† and the oxidative dehydrogenation of ethane…40† on HPMo and HPW supported on silica. Laser Raman, X-ray photoelectron, and 31 P NMR spectroscopies have been employed to show that 12-molybdophosphoric acid can be deposited uniformly on the surface of silica in a highly dispersed form up to a loading of approximately 10 wt %.…41† Aggregates begin to form for loadings from 10 to 29 wt % and at higher loadings particles of the acid are present. Laser Raman and 31 P NMR spectroscopies have provided evidence for the enhancement of the thermal stability of the silica-supported catalysts resulting from the strong interaction between the supported material and the support.…41† With increase in the reaction temperature and with 17% HPW=SiO2 the selectivity to isobutene reaches a maximum at 300  C while the selectivities to the cis-2- and trans-2-butene are at their minimum values (Fig. 9.16). The selectivities to cis-2- and trans-2-butene decreased while that to isobutene increased with increase in the residence time (Fig. 9.17). Exposure of 17% HPW=SiO2 to aliquots of NH3 drastically reduced the selectivity to isobutene while that to the 2-butenes as well as the conversion remained relatively unchanged. With increases in the time on stream the selectivities to the 2-butenes increased while those to isobutene and the oligomerization products decreased.

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Figure 9.15. Conversion of 1-butene and selectivity to isobutene for four loadings of HPW on SiO2.…38† Reproduced by permission of Baltzer Science Publishing.

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Figure 9.16. Conversion of 1-butene on 17% HPW=SiO2 and selectivities to products at various reaction temperatures.…38† Reproduced by permission of Baltzer Science Publishing.

IR spectra of HPW=SiO2 before and after exposure to 1-butene, cis-2butene, and isobutene show the predominance of saturated C±C bonds and little or no evidence of the ole®nic structure (Fig. 9.18).…38;42† The 2962 cmÿ1 band, attributed to asymmetric CH3 stretching, is the most prominent band with all of

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Figure 9.17. Effect of residence time on conversion of 1-butene on 17% HPW=SiO2 and selectivities to products.…38† Reproduced by permission of Baltzer Science Publishing.

the aforementioned. Evidence for the polymerization of butene on the surface of the catalyst can also be seen from the spectra. The 2962, 2933, and 2875 cmÿ1 bands observed essentially vanished on evacuation at 300  C of HPW=SiO2 previously exposed to 1-butene, but the intensity of the band at 1590 cmÿ1 , attributed to conventional coking, remained relatively unchanged (Fig. 9.19). It is clear that the most acidic sites are required to catalyze the formation of isobutene while isomerization of 1-butene to form the 2-butenes is more facile. Because the selectivities to the 2-butenes decrease with increase in contact time while that to isobutene increases, the former species are primary products while the latter is a secondary product. This is consistent with a carbenium ion mechanism. A common secondary carbenium ion may then be formed on the Bronsted acid sites from 1-butene and the 2-butenes and an equilibrium established between these four species (Fig. 9.20).…38† As the temperature is increased in the presence of appropriate surface properties, the secondary carbenium ion can be converted to a primary species which on loss of a proton forms isobutene. As the secondary carbenium ion is converted to the primary ion, the 2-butenes will be depleted in number. The isomerization of 1-butene has also been studied on HPW supported on zirconia.…43† Loadings up to 25.0 wt % and a reactor temperature of 673 K were employed. The maximum conversion was obtained with a 16±20% loading with

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Figure 9.18. Infrared spectra of 15% HPW=SiO2 after exposure to 3 Torr of 1-butene at 25  C followed by evacuation at various temperatures. Trace 1, 3 mm Hg 1-butene, 25  C, 8 min; trace 2, 3 mm Hg 1-butene, 25  C, 24 h; trace 3, evacuated, 25  C, 30 min; trace 4, evacuated, 115  C, 25 min; trace 5, evacuated 200  C, 30 min; trace 6, evacuated, 300  C, 25 min.…42† Reprinted from Colloids and Surfaces, A105, Gao and Moffat, p. 133, copyright 1995, with permission from Elsevier Science.

selectivity to isobutene of 17%. As discussed in more detail in the section on supported catalysts, the authors ®nd predominantly Lewis acid sites with HPW supported on hydrated ZnO2 calcinated at 673 K but both Bronsted and Lewis sites with a previously calcined ZnO2 as support. The isomerization of 1-butene has been employed as a probe reaction for the evaluation of acid strengths in stoichiometric and nonstoichiometric silver and thallium salts of HPW, HPMo, and HSiW.…44;45† This is discussed in more detail elsewhere in this volume.

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Figure 9.19. Absorbance of the bands at 1590, 1460, 2933, and 2962 cmÿ1 after exposure of HPW=SiO2 to 70 Torr of 1-butene at 150  C for various periods of time.…42† Reprinted from Colloids and Surfaces, A105, Gao and Moffat, p. 133, copyright 1995, with permission from Elsevier Science.

9.7. ISOBUTANE The conversion of isobutane has been proposed as a suitable reaction for the assessment of acid strengths of catalysts and the elucidation of reaction mechanisms.…46;47† On HPW=SiO2 at 375  C the maximum conversion of isobutane was obtained at loadings of 30±70%, although values of only 2% or less were measured (Fig. 9.21).…48† The principal product was butane, although the selectivities to butenes, including isobutene, exceeded that to butane. Kramer and McVicker…47† noted that silica±alumina produced little or no n-butane at 650  C whereas ultrastable Y zeolite at 450  C formed n-butane as well as a number of products expected to result from cracking involving carbonium ions. With 50% HPW=SiO2 the conversion reaches a maximum of approximately 10% at 300  C followed by a precipitous decrease in conversion for further increases in temperature (Fig. 9.22), the latter of which may be attributed to irreversible loss of protons and the concomitant partial decomposition of the heteropoly anion. At 200  C the selectivity to butane is as high as 75%, decreasing to 45% at 350  C. With further increase in temperature the selectivity

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Figure 9.20. Mechanistic scheme for 1-butene isomerization.…38† Reproduced by permission of Baltzer Science Publishing.

to n-butane decreases precipitously while that to isobutene increases to approximately 80% at 600  C. Radical-initiated cracking processes such as those producing isobutene from carbon±hydrogen bond scission are expected to become important at and above approximately 400  C, whereas skeletal isomerization is favored by lower temperatures.…46† Isopentane is the only signi®cant C5 product formed on HPW=SiO2 and at temperatures below 400  C, the selectivities to isopentane and propane are similar as also found with faujasite.…46† Because at temperatures higher than 400  C the butenes are the dominant products, the hydride transfer processes are less signi®cant. Hall and co-workers have studied the cracking of isobutane over amorphous and crystalline aluminosilicates and have reported values for the n-butane= butenes ratio at various temperatures.…49† This ratio was less than 1 with SiO2 ± Al2 O3, H ˆ ZSM-5, and HY at temperatures from 400 to 500  C whereas with

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203

Figure 9.21. Conversions of isobutane and selectivities to products as a function of loadings of HPW=SiO2 (375  C).…48†

slightly dealuminated mordenite values of 1; 1, 74, 39, 22, and 7 were found at 200, 250, 300, 350, 400, and 450  C, respectively, similar to those found for HPW=SiO2. On the HY zeolites n-butane is believed to form from the isomerization of the tert-butyl to the sec-butyl cation followed by H-transfer from the parent isobutane, thus establishing a carbenium ion chain reaction.…50† The formation of C3 results from the transfer of the hydride ion (Hÿ ) to the isopropyl cation formed either in the primary step or from the cracking of a C8 oligomer. As observed with HPW=SiO2, C3 and isopentane were formed with the Y zeolites, probably through oligomerization, skeletal isomerization, and cracking. Because the metastable carbenium ion that is formed must be suf®ciently long-lived so that bimolecular processes can occur, the strongest solid acids are required.…50†

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Figure 9.22. Conversion of isobutane and selectivities to products as a function of temperature (50 wt % HPW=SiO2 ).…48†

As noted in detail elsewhere in this volume, PAS FTIR investigations have shown that carbenium ions produced from methanol, ethanol, and tert-butanol form alkoxy groups with the terminal oxygen atoms of the anions of HPW.…17;18† Butenes chemisorbed on the acid are converted to saturated species by the protons which are coulombically attached to the terminal oxygen atoms.…42† Higher temperatures are required to desorb these species from the surface of strong acids such as HPW. With increase in the time on stream the conversion of isobutane at 300  C on 50% HPW=SiO2 decreases from 10% to 4% in 1 h (Fig. 9.23).…48† However, the selectivity to butane remains at approximately 60% with almost equal selectivities to C3 and isopentane. Deactivation of 10% HPW=SiO2 was insigni®cant but increased with the loading while the selectivities remained relatively unchanged.

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Figure 9.23. Conversion of isobutane and selectivities to products on 50% HPW=SiO2 at 300  C as a function of time on stream.…48†

As discussed in detail in the chapter on microporosity, salts of the heteropoly acids with monovalent cations have been shown to be microporous with surface areas at least one order of magnitude higher than those of the parent acid.…6ÿ14† With increasing preparative Ag‡ =H‡ ratios the surface areas and pore volumes increase to a maximum at an Ag‡ =H‡ ratio of 1 and thereafter remain relatively unchanged while the micropore radius remains relatively constant (within 2%) for all compositions (Fig. 9.24).…44;45† PAS FTIR…15;16† and 1 H MAS NMR…44;45;51† have shown that residual protons are contained within the salts of the heteropoly acids for virtually all preparative stoichiometric ratios. The number of residual protons per anion in AgPW as estimated from 1 H MAS NMR is found to be signi®cant at preparative ratios as high as 1.5 (Table 9.5). The effects of the aforementioned morphological and acidity properties and their changes with preparative stoichiometric ratio (Ag‡ =H‡ ) have been assessed with isobutane.…51†

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TABLE 9.5. Protons in Silver 12-Tungstophosphate…44;45† PreparativeAg‡ =H‡ ratio 0.50 0.85 1.00 1.15 1.50 a

From 1 H MAS NMRa 0.94 1.78 1.27 1.26 1.14

From integration of peaks.

With AgPW the rate of conversion of isobutane increases with preparative Ag‡ =H‡ ratio, reaching a maximum in the 1.00±1.20 range, depending on the reaction temperature (Fig. 9.25). The conversion on AgPW reaches a maximum at approximately 500  C, probably indicative of the onset of the decomposition of the anion (Figs. 9.25 and 9.26). The number of protons, while signi®cant at all Ag‡ =H‡ examined, not surprisingly, decreases with the preparative ratio, possibly accounting for the decrease in the conversion for values of Ag‡ =H‡ greater than 1.0 (Table 9.5 and Fig. 9.25). The increase in conversion with the lower preparative ratios may be related to the increase in surface area up to the stoichiometric value (Figs. 9.24 and 9.25). In contrast to the results with HPW=SiO2 the products from AgPW are primarily butene although small selectivities to butane are observed for the lower Ag‡ =H‡ values and lower temperatures (Figs. 9.26 and 9.27). The high selectivities of isobutene are indicative of a radical-initiated process together with a carbenium ion mechanism in which the linear butenes are formed. Hydride transfer processes are occurring but to a considerably smaller extent than observed with HPW=SiO2. No resistance of the molecules diffusing into or out of the pores is expected here because the diameters of the micropores are large relative to the sizes of the molecules involved.

9.8. ISOBUTENE The gas-phase hydration of isobutene on heteropoly acids has been investigated by several authors.…52;53† HSiW and HPW on various supports have been employed at 343 K with HSiW=Amberlyst-15 sulfonic resin showing the highest activity.…52† The aluminum salt of HSiW was found to have activity approximately 30% higher than the latter. Unsupported HSiW has been studied for the same process at 313±353 K.…53† In addition to tert-butyl alcohol, the predominant product with selectivity of 40±80%, the dimer of isobutene was also

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Figure 9.24. Morphological properties of silver 12-tungstophosphate.…44;45†

207

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Figure 9.25. Rate of conversion of isobutane on AgPW of various preparative stoichiometries at various temperatures.…48†

detected. The dehydration of HSiW
nH2 O was also studied by the latter research group.…54† 9.9. BUTANOL PAS FTIR investigations of the interaction of n-, iso-, and tert-butanol again provided evidence for dehydration at 60  C although the steady-state concentrations of sorbed intermediates with n- and t-C4 H9 OH were too small to permit characterization by PAS FTIR.…18† The PAS spectra show that the rate of C±O bond cleavage and alkylation of the Keggin anion is dependent on the alcohol as t-C4 H9 OH > i-C3 H7 OH  n-C3 H7 OH > C2 H5 OH > CH3 OH which corre-

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209

Figure 9.26. Conversion of isobutane and selectivities to products on AgPW (1.00) at various temperatures.…48†

sponds closely to the order of decreasing stability of the corresponding alkyl carbenium ion.…55† The dehydration of butanols has also been employed as a reaction to probe the acidic strengths of stoichiometric and nonstoichiometric silver and thallium salts of HPW, HPMo, and HSiW.…51† This is discussed in detail elsewhere in this volume.

9.10. ALKYLATION PROCESSES 9.10.1. Isobutene and Methanol The elimination of tetraethyl lead from gasoline necessitated the introduction of a new octane enhancer.…56† MTBE, although not as effective as tetraethyl

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Figure 9.27. Conversion of isobutane and selectivities to products at 375  C on AgPW of various preparative stoichiometries.…48†

lead, has found wide use for this purpose.…57† MTBE is produced commercially in the liquid phase from methanol and isobutene with strongly acidic ion-exchange resins.…58† The latter suffer from several disadvantages, in particular their low thermal stability, and consequently a number of potentially useful acidic catalysts have been examined.…59† However, the function and importance of acidity remain in dispute.…60;61† Various heteropoly MTBE oxometalates have been investigated for MTBE synthesis. The silver salt of HPW supported on carbon was found to be active and selective for this reaction.…62† Studies of the synthesis of MTBE from methanol and tert-butanol on HPW at 90±140  C and various of its salts at 90  C found that the manganese and lead salts produced the highest conversions of tert-butanol (95 and 91%, respectively) whereas the cobalt and strontium salts yielded the highest selectivities to MTBE (67 and 64%, respectively).…63†

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211

MTBE was prepared in the liquid phase from methanol and isobutene using a Dawson heteropoly acid (H6 P2 W18 O62 ) which had a higher activity than the corresponding Keggin acid.…64† HPW supported on a modi®ed clay was also active.…65† A more recent report has con®rmed that the Dawson acid (H6 P2 W18 O62 ) is more active than the Keggin acids with P, Si, Ge, B, and Co as central atom and methanol and isobutene as reactants.…66† Because the acidic strength of the aforementioned Dawson acid was found, by the Hammett indicator method, to be lower than that of any of the Keggin acids studied, the authors concluded that other factors play a more important role in the catalytic process. The rates of absorption of methanol and MTBE on the Dawson acid were found to be substantially higher than those measured for the Keggin acids although the quantities taken up were similar to those for HPW, which may be partly responsible for the superior performance of the Dawson acid in the MTBE synthesis. With methanol and isobutene as reactants the activity of the silica-supported Dawson acid was found to be similar to that of a sulfonated polymer resin, Amberlyst 15, while the selectivity to MTBE was higher with the former.…67† HPW=SiO2 was also active and selective. Although HPW is more acidic than H6 P2 W18 O62 , when both are supported on SiO2 the difference diminishes. The selectivity to MTBE calculated from methanol increased on supporting the Dawson acid on SiO2 but the selectivity calculated from isobutene decreased. The high activity of the Dawson acid has been attributed to a ¯exible crystalline structure that facilitates rapid absorption and desorption of molecules.…68† Gas-phase synthesis from methanol and isobutene at 373 K on the Dawson acid H6 P2 W18 O62 produced a selectivity nearly 100% to MTBE.…69† The conversion of isobutene at this temperature decreased from 17% to an insigni®cant value as the calcination temperature increased from 373 K to 673 K. The Dawson structure is claimed to be retained up to 873 K. The reaction was stable for 3 h. Tungstosilicic acid (H4 SiW12 O40 ) supported on polyaniline has a low activity for MTBE synthesis after pretreatment in helium but increases an order of magnitude after activation in air at 473 K.…70a† Experimental evidence suggests that protons diffuse from the polymer to the surface, resulting in an increase in the catalytic activity. Further studies at the same laboratory employed the pure heteropoly acid.…70b† The synthesis of MTBE from tert-butanol and methanol on HPMo, HPW, and HSiW supported on titania, silica, and alumina has been compared with HFtreated and mineral acid-activated montmorillonite clays at temperatures of 100± 180  C.…71† The authors ®nd that HF=clay catalysts generate selectivities to MTBE as high as 94%. Extended use of the heteropoly acids is limited by loss of these species from the support, changes in the acidity, and the formation of organophosphate species, although coking was not signi®cant.

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9.10.2. Isobutene and 2-Butene The Clean Air Act places a variety of requirements on re®ners in North America.…72† The reformulation of gasoline to meet these requirements will encourage an increase in the alkylation capacity in view of the high motor octane number of alkylate which is constituted from low-reactive branched saturated hydrocarbons. The current dependency on liquid acid catalysts (H2 SO4 and HF) for alkylation processes constitutes an impediment to capacity increases which may be alleviated by the introduction of new solid, stable catalysts which are the focus of current research. The chemistry, catalysts, and processes for isoparaf®n±ole®n alkylation have been reviewed recently by Corma and Martinez.…72† Although the mechanism is not fully understood, the ®rst step in the simpli®ed mechanism involves the proton addition to the ole®n to form a tert-butyl cation followed by the addition of the latter to the ole®n to produce C8 carbocations, which may isomerize through hydride and methyl shifts to produce more stable carbenium ions. Isobutane transfers a hydride ion to the carbenium ions forming isomers of octane and reproducing the tert-butyl cation, which continues the chain process. Secondary reactions that reduce the quality of the alkylate, such as polymerization, disproportionation, cracking, and self-alkylation, also occur. The liquid-phase alkylation of isobutane with 2-butene at 80  C has been investigated with the cesium, ammonium, and potassium salts of HPW of various stoichiometries.…73† As discussed in the chapter on microporosity, the salts of the heteropoly acids prepared from monovalent cations have high surface areas and microporous structures.…6ÿ14† The surface areas and pore volumes generally increase with the cation=H‡ ratios up to the stoichiometric values. The activities in isobutane alkylation were found to correlate with the surface acidity where the latter was calculated from the product of the bulk acidity, 3±x (x is equal to the number of cations, other than H‡ , per anion) and the speci®c surface area, assuming a uniform composition in the solid (Figs. 9.28 and 9.29). The data in Figs. 9.28 and 9.29 have similar shapes. The conversions were also dependent on the nature of the cation. With the ammonium and potassium salts a maximum in the conversion of 2-butene at 1 min time on stream was observed for a cation concentration of x equal to 2.5 whereas for the cesium compounds the maximum was found in the range 2 < x < 2:5. The authors concluded that the high activity of Cs2 PW was due to its relatively high surface area. The parent acid HPW was believed to have protons of higher acidic strength than those of the salts, although the numbers of the former are smaller. However, the …NH4 †2:5 PW salt appears to have a higher concentration of strong surface acid sites than the corresponding Cs2 PW and K2:5 PW salts all of which have similar surface areas.

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213

Figure 9.28. Initial conversion of 2-butene as a function of the number of cations KUÿ1 in the monovalent salts of HPW: (d) Csx PW, (m) (NH4 )x PW, and (j) Kx PW (1 min time on stream.…73† Reproduced by permission of Academic Press.

Figure 9.29. Relative surface acidity as a function of number of cations KUÿ1. Symbols as in Fig. 9.28.…73† Reproduced by permission of Academic Press.

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The alkylation of isobutane with 1-butene has been shown to be catalyzed by Cs2:5 H0:5 PW12 O40 at room temperature.…74† Total yields of C8 alkylates as high as 79% were obtained. 12-Tungstophosphoric acid supported on silica, an aluminosilicate (MSA), and an all-silica mesoporous MCM-41 have been compared for the alkylation of 2-butene with isobutane at 33  C and 2.5 MPa.…75† A higher alkylation activity is found with amorphous silica in comparison with the remaining two supports. The silica-supported catalyst with a loading of 40 wt % produced the highest conversion, selectivity to trimethylpentanes, and catalyst stability. Although all-silica MCM-41 possesses a high surface area, a partial blockage of the channels of MCM-41 by HPW decreased the availability of acidic sites. The authors point out that HPW=MCM-41 may be improved by increasing the pore diameters of the support. Decreases in conversion with time on stream were not due to losses of HPW during the alkylation reaction but were the result of deposition of heavy hydrocarbons on the surface of the catalyst. HPW and HSiW supported on various silicas as well as MCM-41 and HMS and also prepared by a sol±gel technique were found to have good catalytic properties for the alkylation of isobutane with butene.…76† The activity and selectivity of the supported catalysts correlated with the surface acidity and their activities decreased relatively little over periods of 8±12 h. Heteropoly acids (HPW, HPMo, and H6 P2 W18 O62 ) supported on ZrO2, TiO2, Al2 O3, or Fe2 O3 in the presence of a small amount of a mineral acid have been shown to be active and selective in the alkylation of isobutane with 1-butene at 70  C and 20 bars.…77† On HPW supported on ZrO2 in the presence of sulfuric acid the conversion of butene was 79% with a selectivity to saturated octanes of 97% of which 30% is trimethylpentane. The deactivation of the latter catalyst is attributed to the deposition of heavy organics which are removed by in situ heating, under air ¯ow, at 250±350  C. 9.10.3. Benzene What may be the ®rst study of alkylation processes catalyzed by heteropoly acids was published in 1971…78† although it could be argued that the oligomerization of propene should be considered a self-alkylation process.…27ÿ29† The alkylation of benzene with dodecene-1, which can be represented stoichiometrically by the following equations, benzene ‡ dodecene ! phenyldodecane phenyldodecane ‡ dodecene ! heavy alkylate dodecene ! dodecene dimer benzene ‡ dodecene dimer ! heavy alkylate

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215

was studied on various supports with silica gel proving superior to the others tested.…78† The authors found that the selectivity was strongly dependent on the benzene=dodecene-1 ratio in the feed with selectivities of 90±94% of phenyldodecane attainable. The alkylation of benzene with propylene to form cumene has been investigated on HPMo and HPW supported on ZrO2 containing small quantities of SO4 2ÿ .…79† With the former catalyst the propylene conversion and selectivity to aromatics were 71 and 95%, respectively, whereas with the latter catalyst these were 99 and 93%, respectively, at 100  C and 4 MPa. Relatively small quantities of n-propylbenzene were formed. HPW and HSiW supported on SiO2, Al2 O3, and diatomaceous earth have been employed in the alkylation of octene and nonene with phenol at 353± 393 K.…80† The catalytic activity was highest with SiO2 as the support. The selectivities for p- and o-alkylphenol on HSiW=SiO2 were 90 and 10%, respectively. The authors conclude that HSiW=SiO2 and HPW=SiO2 are both excellent catalysts for the aforementioned processes. 9.11. FRIEDEL±CRAFTS REACTIONS Alkylation ArH ‡ RCl ! ArR ‡ HCl Acylation ArH ‡ RCOCl ! ArCOR ‡ HCl The interest in replacing homogeneous by heterogeneous catalysts in industrial processes for both practical and environmental considerations has, not surprisingly, extended to Friedel±Crafts reactions.…81† It is also not surprising to ®nd that heteropoly oxometalates have been examined as possible replacements for their homogeneous counterparts. The earliest work was reported in the 1980s…82ÿ86† with Olah taking up the challenge in the late 1990s.…87;88† HPMo was found to be an effective catalyst for the alkylation by benzyl chloride and tert-butyl chloride, acylation by acetyl chloride, and sulfonylation by tosyl chloride.…82;83† However, heteropoly acids with tungsten as the peripheral metal were found to be more effective than those containing molybdenum for the latter two reactions.…84† Decomposition products, such as MoO3, the latter of which may itself be transformed, appeared to be the most active in the acylation of chlorobenzene with o-chlorobenzoyl chloride.…85† HPW, HPMo, and HSiW supported on SiO2 ef®ciently catalyzed Friedel± Crafts alkylation and acylation reactions of aromatic hydrocarbons with octene, benzyl chloride, and benzoyl chloride in the liquid phase.…86† For the alkylation of

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benzene with 1-octene, HPW=SiO2 was the preferred catalyst at 35  C. The maximum conversion of 1-octene was obtained after pretreatment of the catalyst at 150  C. In the acylation process with HPMo=SiO2 the heteropoly acid was leached out of the support and its decomposition products or those formed from the interaction with the reactants were apparently responsible for the catalysis. In contrast, however, the silica-supported HSiW and HPW acids after pretreatment at 200±500  C remained intact and were catalytically active during the acylation process. Olah examined the adamantylation of toluene, anisole, and ¯uoro- and bromobenzene with various heteropoly acids.…87;88† Only the m- and p-isomers were produced together with adamantane as a by-product. With H4 SiMo12 O40 , the anisole=1-bromoadamantane reaction at 152  C, the re¯ux temperature, produced no adamantane and selectivities to m- and p-isomers of 0.8 and 99.2%, respectively, with conversion of 100% after a reaction time of 2 h. The regioselectivity of the substitution was found to have a strong dependence on the acidity with increasing acidic strength leading to the formation of the m-isomer. Although the process is kinetically controlled, the distribution of isomers is altered by isomerization of the primary products. 9.12. C5 ±C8 Alkanes The isomerization of alkanes usually employs dual-function catalysts with noble metal(s) supported on an acidic solid. Ono and co-workers appear to be the ®rst to consider the incorporation of a metal, as the cation, into a heteropoly acid (HPW), supported on silica gel, for the isomerization of pentane and hexane.…89† The palladium salt Pd3 …PW12 O40 †2 , supported on silica gel with a loading of 30 wt %, heated in a hydrogen stream at 573 K for 1 h, produced the highest conversion of hexane at a reaction temperature of 523 K and hexane and hydrogen pressures of 30 and 71 kPa, respectively, as well as the highest selectivity to isomers (84.4%). The Ag, Ni, Al, and Cu salts of HPW produced conversions a factor of 5 or more smaller than that found with the palladium salt. The authors believe that the Pd(II) cations are reduced to Pd(0) by the hydrogen, generating protons as Pd…II† ‡ H2 ! Pd…0† ‡ 2H‡ Evidence for the location of the Pd(0) atoms was not provided. Heteropoly oxometalates and ZSM-5 have been compared as catalysts for the cracking of hexane at 648±698 K.…90† The activity of ZSM-5 for this process was found to be a factor of 20±100 times higher than that of the ammonium salts of HPW, HSiW, and HPMo, and particularly with NH4 PW, the activity decreased precipitously with time on stream whereas that of ZSM-5 remained relatively

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217

stable. Activation energies for ZSM-5 and the three ammonium salts were 10  1 kcal molÿ1 , indicative of the existence of diffusion resistance in the catalysts. Whereas pretreatment temperature had relatively little effect on ZSM5, NH4 PMo, and NH4 SiW, with NH4 PW the activity dropped precipitously after pretreatment of the catalyst at 750 K, suggesting the onset of anion decomposition. With ZSM-5 and NH4 PW, propane was the principal product with relatively little change in its selectivity for conversions of hexane up to 80%. In general the products obtained with the two catalysts are similar although the selectivities and their variations with conversion are catalyst dependent. The isomerization=cracking of C6 ±C8 alkanes has been investigated with the ammonium salt of HPW as catalyst.…91† The cracking of n-hexane at 373±723 K has been studied on HPW with particular attention being given to the deactivation of the catalyst.…92† The catalytic activity and thermal stability of HSiW, H6 P2 W18 O62 , and H6 P2 W21 O71 …H2 O†3 were also investigated.…93† None of these acids was active for n-hexane cracking at 623 K, apparently as a result of decomposition of the latter two acids. However, HSiW is active at lower temperatures. Supported catalysts were also investigated. Pentane is isomerized on HPW and its nonstoichiometric cesium salt at 35  C:…93† The conversions increase to maximum values of approximately 50 and 25% at pretreatment temperatures of 200 and 250  C on HPW and the cesium salt, respectively. The selectivities to isobutane vary widely with cesium content, the highest value (73.6%) being obtained with the lowest cesium content and a pretreatment temperature of 200  C. A mechanical mixture of Pt=Al2 O3 and a nonstoichiometric cesium salt of HPW produced, in the presence of hydrogen and 453 K, a higher activity in the isomerization of either n-pentane or n-hexane than the salt promoted with Pt by direct impregnation.…94† Addition of Pt reduced the deactivation observed with the salt and increased the selectivity to isopentane from n-pentane. Conversions of 35% and selectivities of 97% were obtained with the mixture. Cumene cracking on HSiW at 150±244  C produced predominantly benzene and propylene.…95† After 10±20 min rapid deactivation of the catalyst was observed. Below 170  C the surface reaction is rate controlling whereas above this temperature the reaction is diffusion controlled. 9.13. C6 ±C8 ALKENES The isomerization of 1-hexene, 1-heptene, and 1-octene has been investigated in the liquid phase at temperatures of 303±343 K on the ammonium salts of HPW, HPMo, and HSiW and their parent acids.…96† The acids HPW, HSiW, and NH4 PW are active in the double-bond isomerization of long-chain alkenes at room temperature in the liquid phase. In contrast, HPMo, NH4 PMo, and NH4 SiW

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are relatively inactive. In all cases only double-bond and E±Z, but no skeletal isomerization, are observed. With only one exception, the cis-2-alkene is formed most rapidly initially, with its concentration reaching a maximum, followed by a diminution as the remaining isomers are formed. The isomerization is believed to occur through the formation of carbocations, as a result of the presence of protons on both the acids and the ammonium salts, followed by subsequent 1,2-hydride shifts.

9.14. RING EXPANSION OF METHYLCYCLOPENTANE AND RING CONTRACTION OF CYCLOHEXANE Methylcyclopentane (MCP) is ring-expanded to cyclohexane (CY) and CY is ring-contracted to MCP on 12-tungstophosphoric acid supported on silica and on microporous silver 12-tungstophosphate with Ag‡ =H‡ preparative ratios of 0.8±1.8.…97† The conversions of CY and MCP increase with the loading of HPW on SiO2 (Figs. 9.30 and 9.31) and reach maximum values at reaction temperatures of approximately 200 and 300  C (Figs. 9.32 and 9.33), respectively. The conversions of CY and MCP are signi®cantly higher on AgPW, increase with

Figure 9.30. Conversion of cyclohexane and selectivity to methylcyclopentane on HPW=SiO2 at 300  C, F ˆ 11 ml minÿ1 , W ˆ 200 mg.…97† Reproduced by permission of Baltzer Science Publishing.

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219

Figure 9.31. Conversion of methylcyclopentane and selectivity to cyclohexane on HPW=SiO2 at 300  C. Reaction conditions as in Fig. 9.30.…97† Reproduced by permission of Baltzer Science Publishing.

Figure 9.32. Conversion of cyclohexane and selectivity to methylcyclopentane on 20% HPW=SiO2. Reaction conditions as in Fig. 9.30.…97† Reproduced by permission of Baltzer Science Publishing.

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Figure 9.33. Conversion of methylcyclopentane and selectivity to cyclohexane on 20% HPW=SiO2. Reaction conditions as in Fig. 9.30.…97† Reproduced by permission of Baltzer Science Publishing.

Figure 9.34. Conversion of cyclohexane and selectivity to methylcyclopentane versus preparative Ag=H ratio at 250  C.…97† Reproduced by permission of Baltzer Science Publishing.

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preparative Ag‡ =H‡ ratio, and pass through maxima at approximately 1.3 and 1.5, respectively (Figs. 9.34 and 9.35). On 1.00 AgPW the conversion of CY reaches a maximum at approximately 250  C whereas on 1.30 AgPW the maximum in the conversion of MCP appears at 225  C. As noted earlier in this chapter, the selectivity to isobutene reached a maximum for 23 wt % HPW=SiO2 and at a reaction temperature of approximately 300  C, the latter with a 17 wt % loading.…38† HPW has strong acidic sites…98† which have been shown to be predominantly Bronsted in nature.…15;16† Although the surface area of HPW=SiO2 decreases with increasing loading, the total number of protons contained within the loaded HPW increases (Table 9.6). However, the number of protons in the surface will reach a plateau after the surface of the support is completely covered. As can be seen in Figs. 9.30 and 9.31, after the initial sharp increases in conversions of CY and MCP as the loading increases, changes in these parameters become considerably reduced, possibly re¯ecting the quasiconstancy of available protons. The monovalent salts of HPW, synthesized from stoichiometric quantities of the preparative reactants, have been shown to possess residual protons.…15;16† With AgPW, as the preparative Ag‡ =H‡ ratio increases, the number of protons TABLE 9.6. Surface Areas of HPW=SiO2 a…97† Loading (wt %)

As (m2 gÿ1 )

0.0 4.7 9.1 23 50 70 100 a

195.5 179.4 154.1 122.1 94.3 55.0 5.7

N2 adsorption, 77 K, BET.

TABLE 9.7. Morphological Properties of the Stoichiometric and Nonstoichiometric Silver Salts of 12-Tungstophosphoric acid…97† Preparative stoichiometrya Property As (m2 gÿ1 ) VMP (ml gÿ1 ) Ê) rMP (A a

0.50

0.85

1.0

1.15

1.50

77.6 0.027 8.0

86.2 0.031 7.8

100.9 0.037 7.9

101.4 0.037 7.8

100.5 0.035 7.7

Preparative cation=proton ratio.

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Figure 9.35. Conversion of methylcyclopentane and selectivity to cyclohexane versus preparative Ag=H ratio at 225  C.…97† Reproduced by permission of Baltzer Science Publishing.

Figure 9.36. Mechanistic scheme for the ring contraction of CY and the ring expansion of MCP.…97† Reproduced by permission of Baltzer Science Publishing.

ACID-CATALYZED PROCESSES

223

decreases, remaining small but ®nite at values of this ratio exceeding 1 (Table 9.5). Concomitantly the surface areas and micropore volumes increase and become essentially constant at the stoichiometric AgPW composition (Table 9.7). In contrast, the micropore radius remains unchanged with the Ag‡ =H‡ ratio. Evidently the pores develop in length and=or number as the Ag‡ =H‡ ratio increases but change relatively little, if at all, in cross section. Consequently, protons are available in both HPW=SiO2 and AgPW and their numbers change with the composition of these catalysts. A suggested mechanism for the ring contraction of CY and the ring expansion of MCP is shown in Fig. 9.36.

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10 OXIDATION PROCESSES

Oxidation reactions have been studied with a variety of heteropoly oxometalates, both supported and unsupported. The oxidation of molecules ranging in size from methane to cyclohexane has been investigated although, somewhat surprisingly, relatively little research on propane has been reported.

10.1. METHANE 10.1.1. Methane with N2 O and O2 as Oxidants The ®rst studies of the oxidation of methane on heteropoly oxometalates were reported in 1987.…1† In this work heteropoly acids supported on silica were employed to increase the contact between reactants and catalyst because the surface areas of the heteropoly acids are low (< 10 m2 gÿ1 ). Both O2 and N2 O were initially tested as oxidants but as the selectivities to partial oxidation products were vanishingly small on the former oxidant, the present review will focus on N2 O. The results of the work show that not all heteropoly oxometalates with anions possessing the Keggin structure are equally effective as catalysts in the partial oxidation of methane because their catalytic capabilities are dependent on the nature of the central and peripheral metal elements of the anions. The effect of changes in the peripheral metal atom and in the central atom of the anion is illustrated in Table 10.1. With tungsten as the peripheral metal atom the conversions are less than 0.5% and little or none of the partial oxidation products, formaldehyde and methanol, are observed. Changes in the central atom with the tungsten-containing acids produce only minor differences. In contrast, with the acids containing molybdenum in the peripheral metal position the conversions of methane are a factor of 10 higher than with the tungsten227

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TABLE 10.1. Comparison of Heteropoly Oxometalates for the Oxidation of Methane…1† Selectivity (%) Catalysta HPW (26.2) HSiW (26.2) HPMo (20.0) HSiMo (19.9)

Conversion of CH4 (%)

CO

CO2

CH2 O

CH3 OH

0.40 0.35 5.1 2.5

56.0 44.0 65.0 56.6

44.0 56.0 22.5 32.3

tr tr 12.0 8.7

nd nd 0.5 0.4

Note. Reaction conditions: TR 843 K, W ˆ 0.35 g, F ˆ 30 ml minÿ1 , CH4 (67%), N2 O (33%). a Numbers in parentheses represent the loading (wt %) of the acid on the SiO2 support.

containing acids and signi®cant selectivities to the partial oxidation products, particularly H2 CO, are now evident. With phosphorus as the central atom the conversion of methane is twice as large as that with silicon as the central atom and the selectivities to partial oxidation products are higher with the former than the latter solid. In addition, the silicon central atom favors the production of CO2 at the expense of CO. These observations appear to be consistent with the results of earlier extended HuÈckel (EXH) calculations…2;3† which showed that the strength of the bond between the peripheral metal atoms and the outer or terminal oxygen atoms of the heteropoly anion was much higher with tungsten than with molybdenum, implying that the terminal oxygen atoms with the latter acids are suf®ciently labile to permit their participation in oxidation processes whereas those in the tungsten-containing species are relatively more tightly bound and hence less able to participate in such a reaction. In view of the superior results obtained with HPMo=SiO2, the present review will focus on this catalyst. Pretreatment of this solid in a reducing atmosphere resulted in an increase of the activity of the catalyst with little or no in¯uence on the selectivity. TPD experiments…4† have shown that, at suf®ciently high temperatures, protons contained within the heteropoly acids will extract oxygen atoms from the heteropoly anion and be desorbed as water. With HPW this process occurs at approximately 773 K, some 100 degrees higher than found with HPMo. It may be noted that this provided further support for the results of the calculations referred to earlier. That the heteropoly anion remains intact at these temperatures in spite of the loss of up to two oxygen atoms per anion has been demonstrated from photoacoustic FTIR studies of HPW and HPMo.…5† Although TPR experiments…6† with hydrogen produced peaks of similar overall shape and positions to those found in the analogous TPD experiments, whereas in the TPD experiments 1.5 water molecules per KU were desorbed with HPMo, 8.0 water molecules were measured with the high-temperature TPR peak. The added

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hydrogen in the TPR experiments is evidently capable of supplementing the protons in stripping oxygen atoms from the heteropoly anion. TPR experiments with deuterium…6† showed that HPW exchange begins at approximately 623 K, a temperature similar to that at which the high-temperature peak began to appear in the TPR experiments. In contrast, no exchange was observed with HPMo but consumption of deuterium and reduction of the acid were observed. From the aforementioned observations it may be concluded that pretreatment of HPMo=SiO2 in a reducing atmosphere may remove a portion of the terminal oxygen atoms from each heteropoly anion, thus establishing vacancies for the replacement of oxygen by the oxidant and the ultimate consumption of this oxygen in the oxidation process. At the temperatures employed for the partial oxidation of methane in this work, no reaction occurs in the absence of the catalyst with either oxidant, N2 O or O2. Thus, the oxidant is apparently functioning as a source of regenerative oxygen for the catalyst with N2 O better suited for this purpose than O2. With N2 O as oxidant the selectivity to H2 CO increases with decreasing contact time (W=F) until a maximum of approximately 18% is reached (Fig. 10.1). A maximum is also found for the selectivity to CO whereas that for CO2 passes through a minimum. Because an increase in reaction temperature displaces these extrema to lower residence times, higher temperatures favor further oxidation of the partial oxidation products as well as CO. The conversion of CH4 increases linearly with increase in contact time up to approximately 5%. With increase in reaction temperature the selectivities to CO and formaldehyde decrease whereas that to CO2 increases (Fig. 10.2). At 843 K low values of the N2 O=CH4 molar ratio favor the production of partial oxidation products with a maximum selectivity to H2 CO of approximately 22% reached at N2 O=CH4 ˆ 0:1 (Fig. 10.3). The introduction of various partial pressures of water vapor into the reactant stream with HPMo=SiO2 results in a decrease in the conversion with either N2 O or O2 as the oxidant. The initial conversion is recovered after elimination of the water. With N2 O as the oxidant the selectivity to formaldehyde increases for small partial pressures of water but decreases with increases in the water content. With oxygen little change in the selectivity is observed. With either oxidant the original selectivity is recovered on removal of the added water. The results from the reaction of the products from the oxidation of CH4 with and without N2 O or O2 on SiO2 under one set of experimental conditions show that the support is not inert in the present studies and that both hydrogenation and oxidation processes are occurring with H2 CO and CH3 OH, the latter processes being detrimental to the generation of partial oxidation products in the oxidation of CH4 (Table 10.2). Addition of HPMo to the support enhances the catalysis of the hydrogenation process, when CH3 OH and CH2 O react in the absence of N2 O, as more CH4

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Figure 10.1. Effect of the contact time on the conversion and selectivity of 29.4% HPMo=SiO2 for the CH4 ‡ O2 reactions. Reaction conditions: (A) TR ˆ 793 K, W ˆ 1.0 g, CH4 (67%), N2 O (33%); (B) TR ˆ 843 K, open symbols W ˆ 1.0 g, solid symbols W ˆ 0.17 g, CH4 (67%), N2 O (33%); (C) TR ˆ 843 K, W ˆ 1.0 g, CH4 (67%), O2 (33%). Symbols: (n) CO, (s) CO2 , (u) CH2 O, (X) CH4 conversion.…1†

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Figure 10.2. Effect of the reaction temperature on the selectivity of the CH4 ‡ N2 O or CH4 ‡ O2 reactions on 29.4% HPMo=SiO2 and the SiO2 support. Reaction conditions: (A) SiO2 , F ˆ 15 ml minÿ1 , W ˆ 2.0 g, open symbols CH4 (67%), N2 O (33%), solid symbols CH4 (67%), O2 (33%); (B) HPMo open symbols F ˆ 7.5 ml minÿ1 , solid symbols F ˆ 90 ml minÿ1 , W ˆ 1.0 g, CH4 (67%), N2 O (33%); (C) HPMo open symbols F ˆ 15 ml minÿ1 , W ˆ 1.0 g, CH4 (67%), O2 (33%). Symbols: (n) CO, (s) CO2 , (u) CH2 O).…1†

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Figure 10.3. Effect of the CH4 or N2 O concentration on conversion and selectivity of 29.4% HPMo=SiO2. Reaction conditions: F ˆ 60 ml minÿ1 , W ˆ 1.0 g, TR ˆ 843 K. (A) CH4 (50%), He (50N2 O%); (B) N2 O (25%), He (75-CH4 %). Symbols: (n) CO, (s) CO2, (u) CH2 O, (d) CH4 conversion.…1†

is produced than with SiO2 as catalyst. In the presence of N2 O the conversion of CH3 OH or CH2 O and the selectivity to CO and CO2 are increased. The hydrogenation function remains effective as evidenced by the presence of CH4 and a trace of H2 . The addition of HPMo increases the conversion of both CH3 OH and H2 CO and with the former increases the production of CH4 , consistent with observations on HPMo studied in unsupported form.…7;8†

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TABLE 10.2. Conversion and Selectivity of Different ReactantsÐCO, CH2 O, CH3 OHÐ Involved in the Partial Oxidation of Methane…1† Catalyst

SiO2 20.0 HPMo

Flow CO conversion Selectivity CO2 Flow CH2 O conversion Selectivitya CO CH4 CO2 CH3 O H H2 Otherb Flow CH3 OH conversion Selectivitya CO CH4 CO2 CH2 O H2 Otherb

SiO2 20.0 HPMo

SiO2 20.0 HPMo

CO (50%) 0.75 8 100 100

CO (50%) N2 O (50%) 16 16 100 100

CO (50%) O2 (50%) 34 34 100 100

CH2 O (10%) 52 87

CH2 O (10%) N2 O (17%) 48 68

CH2 O (10%) O2 (17%) 55 83

83 5 9 3 lt tr

51 13.5 34 0.5 lt l

CH3 OH (10%) 49 94 30 32 2 16 lt tr

22 58 7 7 lt 6

86 2 12 1 lt tr

82 0.5 16.5 0.5 tr 0.2

CH3 OH (10% N2 O (28%) 65 80 50 18 11 16 lt 5

47 24 11 17 tr 1

67 nd 33 0.5 nd nd

18 nd 82 nd nd nd

CH3 OH (10%) O2 (28%) 96 100 46 1.5 39.5 9 nd nd

61 nd 30 9.3 nd nd

Note. Reaction conditions: TR ˆ 843 K, W ˆ 0.35 g, F ˆ 60 ml minÿ1 . Complementary gas: helium. tr, trace; lt, large trace; nd, not detected. a C deposits are not accounted for in the selectivity. b C2 , C3 , C4 hydrocarbons and dimethylether.

With O2 as oxidant the conversion of the oxidation products as well as the selectivities to complete oxidation products are increased. With HPMo present the conversion of CH3 OH and CH2 O is increased. Oxygen is apparently a more powerful oxidant for CH3 OH and CH2 O than N2 O whereas the latter is more ef®cient in the oxidation of CH4 than O2. The absence of partial oxidation products in the CH4 ‡ O2 reaction may thus be the result of their complete oxidation in subsequent steps rather than their absence in the initial product composition. In view of the importance of supporting the heteropoly acids where the reactants are nonpolar species, additional factors in the methane oxidation process were examined, in particular the effects of conditions on the resulting activity and selectivity…9† (Fig. 10.4). The conversion and selectivity remain relatively constant for calcination temperatures up to 773 K (Fig. 10.4A,B) but for higher temperatures the conversion decreases sharply while the selectivities to

Figure 10.4. Effect of the temperature of calcination during 16 h (left) and of the time of calcination at 823 K under air (right) on the CH4 conversion, selectivity, and Mo loading of the 23%-HPMo catalyst. Reaction conditions: CH4 (67%), N2 O (33%), TR ˆ 843 K, W ˆ 0.5 g, F ˆ 30 ml minÿ1 . Symbols: (n) CO, (s) CO2, (u) CH2 O, (X) CH4 conversion, (d) Mo loading.…9†

234 CHAPTER 10

OXIDATION PROCESSES

235

CO and CO2 remain constant up to approximately 900 K. Above 900 K the selectivities to H2 CO and CO decrease while that to CO2 increases, with all three approaching that found with the SiO2 support itself. A small but signi®cant decrease in the quantity of supported molybdenum is seen for temperatures higher than 700 K while the conversion vanishes at 1000 K. At a calcination temperature of 823 K, at which temperature thermal degradation is occurring (Fig. 10.4A,B), the conversion of methane decreases relatively slowly with the duration of calcination with a concomitant small loss of HPMo (Fig. 10.4D). The rates of formation of the various products from the reaction of CH4 and N2 O at 843 K increase with the loading of HPMo on the support, and at low values of loading extrapolate to those results found for the support itself …9† (Fig. 10.5). The rates of formation of the various products reach maxima at a loading of approximately 120 mmol of HPMo ( 25 wt %) per gram of support. The product distributions also vary with the loading (Fig. 10.6). At low loadings the selectivities to CO and CO2 increase and decrease sharply, respectively, with the loading and, at least with CO, reach a constant value at relatively small loadings. In contrast, the selectivity to CO2 passes through a minimum while that to H2 CO increases to a maximum. IR spectra of solutions obtained by washing the HPMo=SiO2 with acetonitrile show that the Keggin structure is retained on samples of various loadings…9† (Fig. 10.7). The bands at 1080 and the doublet at 969±960 cmÿ1 characteristic of the Keggin structure and attributed to the triply degenerate asymmetric stretch of the central PO4 tetrahedron and that of the outer Mo±O bond, respectively, are clearly evident, even after heating of the sample at 623 K for 2 h. Some diminution of the intensities occurs in 20.1% HPMo=SiO2 samples heated up to 923 K for 16 h in air and after use in an oxidation experiment at 843 K, and even after heating to 1003 K, some vestiges of the Keggin structure remain. The linear increase in the rate of formation of various products as the loading is increased demonstrates that the active species in the oxidation process are associated with the supported materials. Because, at temperatures of approximately 923 K, the activity and selectivity of the supported catalyst change to those expected for the support, the active species is thermally sensitive. Evidently an optimum coverage of the support by HPMo exists. Heavier loadings result in the formation of particles which may block access of the reactant to the HPMo attached to the support. However, the polar products H2 CO and CH3 OH can penetrate into these particles where the production of CO and=or CO2 occurs. The effect, on the oxidation of methane with nitrous oxide, of the addition of cesium to silica-supported HPMo is illustrated in Fig. 10.8.…10† The turnover rates of methane and the oxidant, N2 O, decrease approximately linearly as the quantity of cesium is increased up to approximately two cations per KU. With continued

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Figure 10.5. Effect of the HPMo loading of the support on the production rate of the different products of the CH4 ‡ N2 O reaction at 843 K. Reaction conditions: CH4 (67%), N2 O (33%), Wˆ0.5 g, Fˆ30 ml minÿ1 . Symbols: (X) N2 , (‡) total carbon detected, (d) H2 O, (,) CH3 OH, (n) CO, (s) CO2 , (u) CH2 O.…9†

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237

Figure 10.6. Effect of the HPMo loading of the support on the selectivity of the CH4 ‡ N2 O reaction at 843 K. Same experimental conditions as Fig. 10.5. Symbols: (n) CO, (s) CO2 , (u) CH2 O, (,) CH3 OH.…9†

addition of Cs the rates fall more precipitously and approach the values found for the silica support when 3.5±4 Cs‡ KUÿ1 are present. The changes observed in the selectivities as the cesium is added are also of interest. As the cesium is added, the selectivities to CO and CO2 decrease and increase, respectively, become equal at approximately 3±4 Cs‡ KUÿ1, and then remain constant, for further increases in the cesium content, at values virtually the same as those found for silica. Concomitantly the selectivity to H2 CO decreases to zero after approximately 4 Cs‡ KUÿ1 have been added. The catalytic properties of HPMo have been completely eliminated by the addition of numbers of Cs‡ similar to the numbers of protons in the stoichiometric HPMo. As with HPMo, the activities of HSiMo and HPV2 Mo10 and HPW supported on SiO2 are almost completely removed by the addition of cesium although the shapes of the curves for turnover rates show some differences (Fig. 10.9). The selectivities to the products also are altered in similar, but not identical, fashion.

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Figure 10.7. IR spectra of acetonitrile solution after washing of the following supported HPMo samples calcined in different conditions. (a) Bulk H3 PMo12 O40 ; (b) 1.16 HPMo, 350  C, 2 h; (c) 11.1% HPMo, 350  C, 2 h, then 20.1% HPMo sample; (d) 350  C, 16 h; (e) 450  C, 16 h; (f) 550  C, 16 h; (g) 640  C, 16 h; (h) 640  C, 16 h, followed by a test at 570  C, 10 h; (i) 730  C, 16 h.…9†

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Figure 10.8. Turnover rate and selectivity of the CH4 oxidation by N2 O versus the amount of cesium on the 16 wt% HPMo catalyst. Turnover rates are expressed as 10ÿ2 molecule KUÿ1 sÿ1 . As each KU occupies 14.7 nm of the SiO2 support surface the equivalent turnover rate for SiO2 is 10ÿ2 molecules (15 nm2 support)ÿ1 sÿ1 . Reaction conditions: TR ˆ 843 K, W ˆ 0.5 g, F ˆ 30 ml minÿ1 , CH4 (67 mol%), N2 O (33 mol%). Symbols: (n) CO, (s) CO2 , (u) CH2 O, (X) N2 O turnover rate, (d) CH4 turnover rate  2. Dashed lines are calculated for substitution of H‡ by Cs‡ .…10†

The poisoning by cesium was also investigated with experiments on the reoxidation of samples prereduced by hydrogen at 570  C for 1 h. The quantity of oxygen necessary to reoxidize the sample was determined from the total amount of nitrogen produced after the injection of 15 successive pulses of nitrous oxide (Fig. 10.10B). The amount of oxygen decreases linearly as the Cs content

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Figure 10.9. Turnover rate and selectivity of the CH4 oxidation by N2 O versus the amount of cesium on the 16.2 wt % HPV2 Mo10 , 16 wt % HSiMo, and 21.2 wt % HPW. Same reaction conditions and symbols as those in Fig. 10.8. Symbols: (n) CO, (s) CO2, (u) CH2 O; turnover rate: (X) HPV2 Mo10 , (d) HSiMo, (m) HPW.…10†

increases to 3 Cs‡ KUÿ1 and for further increases remains constant. In the absence of Cs, 19 oxygen atoms KUÿ1 are required to reoxidize the reduced HPMo, but this decreases to approximately 4 oxygen atoms KUÿ1 on addition of cesium. The activation energy for the methane oxidation process is also reduced on addition of cesium (Fig. 10.10A). Up to 2 Cs‡ KUÿ1 the activation energy

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Figure 10.10. Variation of (A) the apparent activation energy and (B) the amount of oxygen consumed for reoxidation after H2 reduction at 843 K for 1 h versus the amount of cesium on the 16 wt % HPMo. Reaction conditions for (A): W ˆ 0.5 g, F ˆ 30 ml minÿ1 , CH4 (67 mol %), N2 O (33 mol %). Symbols: (d) ®rst measure, (s) second measure.…10†

remains at 32  2 kcal molÿ1 but is reduced to 12  2 kcal molÿ1 , a value similar to that for SiO2, on further increase in the cesium content. XP, laser Raman, 31 P NMR, and ion scattering spectroscopies have been employed to provide further information on the nature of the cesium-doped

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HPMo=SiO2.…11† The Raman spectra of the HPMo=SiO2 (16 wt %) and various CsPMo=SiO2 samples (denoted as x-CsPMo where x is the number of Cs added KUÿ1 ) show no evidence of the characteristic bands of MoO3 (strong bands at 995 and 820 cmÿ1 ) and molybdate MoO4 species (main bands at 890 cmÿ1 and a characteristic band at 320 cmÿ1 ) which would be expected if decomposition of the Keggin structure had occurred (Fig. 10.11). At Cs loadings

Figure 10.11. Laser Raman spectra of cesium-doped silica-supported 12-molybdophosphoric acid (16 wt %) catalysts calcined at 350  C, 2 h: (a) 16 wt % HPMo, (b) 0.63-CsPMo, (c) 1.45-CsPMo, (d) 2.85-CsPMo, (e) 4.3-CsPMo, (f) 5.6-CsPMo, (g) 6.9-CsPMo, (h) 8.3-CsPMo. Plasma lines of the laser are indicated by p.…11†

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up to 5.6-CsPMo (spectra b±f ), small shifts of the Raman bands as compared with HPMo can be seen. At higher loadings (spectra g and h), bands attributed to isopolymolybdate species can be seen. 31 P NMR of the CsPMo=SiO samples doped with cesium are characteristic of the supported HPMo,…11† the results for which are discussed in more detail elsewhere in this volume, up to the 2.15-CsPMo sample (Fig. 10.12). For loadings from 2.85 to 4.3 Cs KUÿ1, a splitting of the peak is observed and a broadening of the 31 P NMR band is seen for the 8.3 Cs sample, the latter of which may be ascribed to a partial degradation of the KU with formation of adsorbed phosphate ions. The doublet found with the 2.85-CsPMo sample may result from a mixture of HPMo, substituted HPMo, and phosphate.

Figure 10.12. 31 P NMR spectra of the cesium-doped silica-supported 12-molybdophosphoric acid catalysts calcined at 350  C, 2 h. The widths of the peaks are shown in brackets. (a) 0.63-CsPMo, (b) 1.45-CsPMo, (c) 2.15-CsPMo, (d) 2.85-CsPMo, (e) 4.3-CsPMo, (f) 8.3-CsPMo.…11†

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XPS intensity ratios for Cs3d5=2 =Si2p suggest that the Cs is well dispersed up to a loading of 3.85 Cs KUÿ1 whereas for larger quantities of Cs a line of decreased slope demonstrates that a portion of the added Cs is not detected by XPS and therefore is contained within aggregates (Fig. 10.13). Because the Mo3d =Si2p intensity ratio for the CsPMo samples is constant but higher than that for the HPMo sample, the addition of Cs apparently induces a perturbation on the HPMo, possibly in its dispersion. The changes in the Cs3d5=2 binding energy and the full width at half-maximum, which occur for Cs KUÿ1 loadings greater than 2.85, can be interpreted as indicative of the existence of two different cesium

Figure 10.13. Evolution of XPS data versus cesium loading of the 12-molybdophosphoric acid catalysts. Mo3d5=2 binding energies (j); Cs3d5=2 binding energies (.); Cs3d5=2 FWHM (m); Mo3d =Si2p intensity ratio (s); Cs3d5=2 =Si2p intensity ratio (d).…11†

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species at the higher loadings. The aforementioned changes together with those observed for Mo3d show that the substituted KU is strongly perturbed relative to the free KU when Cs is added. Ion scattering spectra (Fig. 10.14) provide evidence for a strong shielding effect, the large Cs ions protecting the neighboring Mo ions, which demonstrates the proximity of the two elements and therefore the direct interaction between the KU and the Cs ion. The aforementioned spectroscopic data show that the structure of the heteropoly anion remains intact although a direct interaction of the Cs ions with the KU occurs. The Cs ions are well dispersed up to a loading of 3.85 Cs KUÿ1, corresponding to the exchange of all of the protons with possible formation of aggregates of cesium isopolymolybdate at higher loading. The methane oxidation results taken together with the information from TPD experiments and various spectroscopies show that at suf®ciently high temperatures, protons in the heteropoly acids extract terminal oxygen atoms from the

Figure 10.14. 20 Ne‡ ion scattering spectra of the 0.63-CsPMo and 5.6-CsPMo samples taken at different 20 Ne‡ ion beam energies. All of the spectra have been recorded at the same detector sensitivity. Ion energies: (a) 500, (b) 1000, (c) 1500, (d) 2000 eV.…11†

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anions to form water which desorbs from the solid, leaving oxygen vacancies (u) in the anion, but with retention of the Keggin structure: 2KUOH ! KUO ‡ KUu ‡ H2 O The oxygen vacancies can be ®lled by oxygen from an oxidant such as N2 O: N2 O ‡ KUu ! N2 ‡ KUO The replacement of the protons, or a portion thereof, by species such as Cs‡ , reduces the numbers of vacancies that can be generated and hence the capacity of the catalyst to oxidize methane. The substitution of the protons by lithium, barium, magnesium, bismuth, aluminum, iron, silver, cobalt, chromium, copper, cadmium, mercury, nickel, lead, or vanadyl cations on HPMo=SiO2 produced effects semiquantitatively similar to those obtained with cesium.…13† Three effects have been discerned. One is attributable to the removal of protons by substitution, leading to the blocking of the oxygen extraction step which generates vacancies in the Keggin structure. Apparently this blocking effect does not depend on the nature or charge of the cation. The second effect appears to have an electronic basis, and is strongly dependent on the nature of the cation, particularly the electronegativity (Fig. 10.15). Although somewhat tangentially related to the present discussion, it is not inappropriate to comment on Mo=SiO2 catalysts for methane oxidation and their relationship to silica-supported molybdosilicic acid.…14† The pH of the Mo solution used in the preparation has considerable effect on the methane conversion properties (Table 10.3). Both the conversion and the product distribution approach those of the molybdosilicic acid as the pH of the solution decreases. The laser Raman spectra of the samples shown in Table 10.3 contain bands characteristic of HSiMo (Fig 10.16). The spectra of MoO3 =SiO2 prepared at pH values of 2 and 7 contain bands at 995, 969, 636 and 252 cmÿ1 , characteristic of HSiMo, whereas the sample prepared at pH 11 can be identi®ed from the characteristic band at 965 cmÿ1 as due to a polymolybdate phase, or more speci®cally a heptamolybdate species. The existence of molybdosilicic acid is also shown by IR and laser Raman spectra analyses of the ®ltered solution resulting from the washing of a calcined sample with acetonitrile to extract the supported HSiMo species selectively. Depending on the pH of the solution, the Mo loading, and the calcination conditions, at least three species are generated on MoO3 =SiO2, namely, HSiMo, the heptamolybdate, and MoO3 crystallites.

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Figure 10.15. Promotion factor of the promoted HPMo catalyst at a loading of two cations per KU versus the electronegativity (Pauling's scale). In parentheses, samples likely to be incorrectly exchanged.…13†

TABLE 10.3. Catalytic Properties and Characteristics of the Samples…14† Conversiona Samples SiO2 supportb MoO3 =SiO2 (pH 11) MoO3 =SiO2 (pH 7) MoO3 =SiO2 (pH 2) H4 SiMo12 O40 =SiO2 a

Loading wt % MoO3 CH4 Ð 3.0 3.2 3.2 5.7

0.8 1.1 2.6 3.3 5.0

Selectivitiesa

N2 O N2 O TONc CO CO2 CH2 O CH3 OH 3.0 4.0 8.4 10.0 19.6

Ð 9.8 19.0 22.6 25.2

75 72 67 61 46

25 19 24 37 54

Ð 9 9 2 tr

T ˆ 773 K, mass of catalyst 2 g, ¯ow rate 15 ml minÿ1 , ¯ow composition CH4 (33%), N2 O (67%). SiO2 Davison±Grace, Grade 400, 750 m2 gÿ1 . c Apparent turnover number [10ÿ4 molecule (atom Mo)ÿ1 sÿ1 ]. b

Ð tr tr tr Ð

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Figure 10.16. Raman spectra of (a) bulk H4 SiMo12 O40 calcined at 623 K, 2 h, and of supported catalysts calcined at 773 K, 2 h; (b) H4 SiMo12 O40 =SiO2 ; (c) MoO3 =SiO2 , pH ˆ 2; (d) MoO3 =SiO2 , pH ˆ 7; (e) MoO3 =SiO2 , pH ˆ 11.…14†

10.1.2. The Effect of Gas-Phase Additives in the Conversion of Methane on Heteropoly Oxometalates The addition of small quantities of chlorinated methanes has been found to have both interesting and advantageous effects on the methane conversion process catalyzed by the heteropoly acids.…15ÿ18† The effects of the addition of tetrachloromethane (TCM) to the methane feedstream with HPMo=SiO2 and HPW=SiO2 as catalysts are strikingly dissimilar (Fig. 10.17). On each of the

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Figure 10.17. Conversion and selectivity as a function of concentration of TCM in the feed (T ˆ 450  C). (a) HPMo=SiO2 W ˆ 1.05 g, F ˆ 60 ml minÿ1 , CH4 =N2 O  1; (b) HPW=SiO2 W ˆ 2.0 g, F ˆ 60 ml minÿ1 , CH4 =N2 O  4. Selectivities: s H2 CO, m CO2, u CO2, . CH3 Cl, e: conversion [(products=methane fed)(100)].…15ÿ18† Reproduced by permission of Baltzer Science Publishing.

catalysts, methyl chloride and the usual oxidation products are formed. However, on HPMo=SiO2 oxidation products account for the major portion of methane reacted, whereas on HPW=SiO2 the selectivity to methyl chloride is as high as 80%. With either of the catalysts only small amounts of TCM are required to produce signi®cant changes in both the conversion and selectivities and for larger

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Figure 10.18. Conversion and selectivity as a function of residence time (W=F) in the presence and absence of CCl4 in the feed (T ˆ 450  C, W ˆ 0.25 g±4.0 g). (a) HPMo=SiO2, [CCl4 ]fed ˆ 0.17% (by mole), F ˆ 60 ml minÿ1 , CH4 =N2 O  4; (b) HPW=SiO2, [CCl4 Šfed ˆ 0.17% (by mole), F ˆ 60 ml minÿ1 , CH4 =N2 O  1. Filled symbols: CCl4 absent; e, r conversion [(products=methane fed)(100)]; m CH3 Cl; s, d H2 CO (balance CO ‡ CO2 omitted for clarity).…15ÿ18† Reproduced by permission of Baltzer Science Publishing.

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Figure 10.19. Rates of production of CH3 Cl and CO with time on stream (a), (b) 20 and 10% HPW=SiO2, respectively. ,, ., CH3 Cl; s, d, CO. Open symbols: pretreatment, CCl4 ; reactants, CH4 , N2 O. Solid symbols: pretreatment, CCl4 ; reactants, CH4 , N2 O, CCl4 . W ˆ 2.0, F ˆ 11 ml minÿ1 , [CCl4 Š ˆ 0:38 mol %, CH4 =N2 O  1, Tˆ 450  C (pretreatment conditions as above except read He for CH4 ).…15ÿ18†

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quantities (greater than approximately 0.3 mol % of TCM) relatively small changes are observed. On HPMo=SiO2 the overall yield of H2 CO is increased on addition of TCM to the feedstream. The time during which the reactant stream is in contact with the catalyst has a considerable effect on both the conversion and selectivities in both the presence and absence of TCM (Fig. 10.17). The conversion of methane is substantially higher with TCM in the feedstream over the entire range of residence time examined. In the absence of TCM the conversion increases with W=F on both catalysts whereas in the presence of TCM the rate of increase with W=F is higher (Fig. 10.18). On HPMo=SiO2 the selectivities to H2 CO are similar, regardless of W=F, whereas on HPW=SiO2 the selectivity to H2 CO is reduced on addition of TCM but the selectivity to CH3 Cl is substantially higher than that found with the molybdenum-containing catalyst. Evidently HPMo=SiO2, in the presence of TCM, acts primarily as an oxidation catalyst, whereas HPW=SiO2, a relatively poor oxidation catalyst, functions as an effective oxychlorination catalyst. Ancillary experiments have been performed to provide information on the role of TCM in the methane conversion process. Two different loadings of HPW=SiO2 were pretreated with a mixture of He, N2 O, and TCM followed by exposure of the catalyst to a stream of CH4 and N2 O (no CCl4 ) and subsequently a mixture of CH4 , N2 O, and TCM (Fig. 10.19). In the early stages of reaction, the rates of production of CH3 Cl in the two methods of operation (that is, with or without TCM in the feed) are similar. Evidently the chlorine appearing in the product CH3 Cl results from that remaining on the catalyst after pretreatment. In the reaction in which TCM was not present in the feed, the rates of production of CH3 Cl decreased rapidly with time on stream as a result of depletion of the chlorine taken up by the catalyst during the pretreatment stage. However, with a reaction stream containing TCM, a high rate of production of CH3 Cl was maintained, indicating a continuous process of consumption (by CH4 ) and buildup (by TCM) of chlorine on the catalyst. Although the possibility of a gas-phase component in the process involving TCM cannot be excluded, nevertheless the aforementioned experiments indicate that the process is largely one in which TCM interacts with and chlorinates the solid catalyst.

10.2. ETHANE The oxidation of ethane has been widely studied on a variety of catalysts, although less extensively than methane. The study of ethane conversion provides a means of comparing and contrasting the mechanisms of the two processes as well as shedding further light on the properties of the catalyst itself.

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The conversion of ethane and selectivities to products of the SiO2 -supported (20 wt %) and unsupported metal±oxygen cluster compounds are shown in Table 10.4.…19† The conversions obtained with the supported molybdenum-containing catalysts, although small, are factors of 3 or 4 higher than those obtained with the catalysts containing tungsten as well as the unsupported Mo catalysts. The predominant product is ethylene with N2 O as oxidant but CO with O2. However, with N2 O signi®cant selectivities to acetaldehyde are obtained at least with the supported heteropoly acids containing only one peripheral metal element. At a reaction temperature of 540  C the conversion of ethane and the yield of acetaldehyde each pass through a maximum at a loading of approximately 20% similar to the observations with methane, while the selectivities to ethylene and acetaldehyde increase and decrease, respectively (Fig. 10.20). On increasing the contact time with N2 O as oxidant, the conversion of ethane decreased and the selectivity to the partial oxidation product, acetaldehyde, and the oxidative dehydrogenation product, ethylene, increased with decreasing contact time, whereas the selectivities to CO and CO2 decreased, suggesting that acetaldehyde and ethylene are primary products. At conversions of approximately 3% the sum of the selectivities to acetaldehyde and ethylene is

TABLE 10.4. Conversion and Selectivity of Supported MOCC and Reference Catalysts…19† Conversion (%) Catalyst

Selectivity (%)

C2 H6

N2 O

O2

CO

CO2

C 2 H4

CH3 CHO

(A)

HPW HSiW HPMo HSiMo HPV3 Mo9 HPMoa V Mo SiO2

0.7 0.7 2.7 3.1 3.2 0.2 9.3 0.5 0.1

3.4 2.7 19.3 23.5 28.0 1.2 74.5 5.6 1.3

Ð Ð Ð Ð Ð Ð Ð Ð Ð

2.6 2.4 13.8 15.5 19.9 16.4 18.1 17.9 10.5

5.7 5.0 5.5 7.2 6.2 2.4 5.3 14.7 4.5

62.5 61.8 47.7 48.7 69.3 81.2 73.6 51.1 72.1

28.8 30.1 30.0 26.1 2.7 ndb nd 12.7 nd

(B)

HPV3 Mo9 SiO2 HPMo V Mo

5.1 0.3 4.8 7.9 1.4

Ð Ð Ð Ð Ð

25.9 6.5 30.2 27.6 10.9

52.0 59.3 53.5 28.4 40.1

13.6 11.4 14.3 15.3 20.1

34.5 26.4 31.1 55.9 37.1

nd nd 1.1 nd 2.7

Note. Reaction conditions: TR ˆ 540  C, (A) W ˆ 0.5 g. F ˆ 25 ml minÿ1 , C2 H6 (80%), N2 O (20%); (B) C2 H6 (80%), O2 (20%). a Unsupported HPMo (parent acid). b nd, not detected.

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Figure 10.20. Effect of loading of HPMo on conversion of ethane and selectivity over supported catalyst; TR ˆ 540  C, W ˆ 0.5 g, F ˆ 25 ml minÿ1 , C2 H6 =N2 O ˆ 4=1, (d) conversion of C2 H6 , (n) selectivity of C2 H4 , ( ) CH3 CHO, (,) CO2, (e) CO, ( ) yield of CH3 CHO.…19† Reprinted from Applied Catalysis, A109, Hong and Moffat, p. 117, copyright 1994, with permission from Elsevier Science.

greater than 70%. At high contact time, the lower selectivities to these two products probably result from their further oxidation. Increase in the reaction temperature leads to an increase in the conversion of ethane and selectivities to CO and CO2 at the expense of acetaldehyde. Increases in the pretreatment temperature produced results similar to those found with methane (Fig. 10.21). With 20 wt % HPMo=SiO2 and a reaction temperature of 540  C the conversion and selectivities in the oxidation of ethane with nitrous oxide on 20 wt % HPMo=SiO2 remain relatively constant for pretreatment temperatures up to 500  C. For higher temperatures, the conversion of ethane and selectivity to acetaldehyde decrease whereas the selectivities to deep oxidation products, CO and CO2, increase gradually. In the absence of an oxidant the products from ethane are initially predominantly ethylene and acetaldehyde on 20 wt % HPMo=SiO2 at 510  C (Fig. 10.22). However, after 2 h on stream, little or no acetaldehyde is produced and ethylene is the predominant product. With the same catalyst and a variety of reaction temperatures, contact times, and C2 H6 =N2 O ratios, the selectivities to ethylene and acetaldehyde decrease

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Figure 10.21. Effect of pretreatment temperature on ethane conversion and selectivity over 20 wt % HPMo=SiO2 ; TR ˆ 540  C, W ˆ 0.5 g, F ˆ 25 ml minÿ1 , C2 H6 =N2 O ˆ 4=1, (d) conversion of C2 H6 , (e) selectivity of CO, (,) CO2, (n) C2 H4 , ( ) CH3 CHO.…19† Reprinted from Applied Catalysis, A109, Hong and Moffat, p. 117, copyright 1994, with permission from Elsevier Science.

Figure 10.22. Effect of time on stream on ethane conversion and selectivity over 20 wt % HPMo=SiO2 ; W ˆ 0.75 g, F ˆ 5 ml minÿ1 , C2 H6 only and no oxidants, TR ˆ 510  C; (d) conversion of C2 H6 ; (e) selectivity of CO, (,) CO2, (n) C2 H4 , ( ) CH3 CHO, (u) CH4 .…19† Reprinted from Applied Catalysis, A109, Hong and Moffat, p. 117, copyright 1994, with permission from Elsevier Science.

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with conversion whereas those to CO and CO2 increase with conversion, indicating that the former two species are primary whereas the latter two are secondary products (Fig. 10.23). After use of the 20 wt % HPMo=SiO2 for 2 h in the absence of an oxidant, followed by exposure to N2 O, the catalytic properties are essentially regenerated whereas with a similar treatment with oxygen such regeneration is not successful (Fig. 10.24). The similarities between the results for the conversion of methane and of ethane are evident. The optimum loadings of HPMo on SiO2 are similar as is the effect of the pretreatment temperature. As with methane the existence of the Keggin structure is important, as well as the presence of the proton, at least initially, in order to generate the oxygen vacancies and presumably to reduce the peripheral metal atoms, Mo6‡ , to Mo5‡ which are subsequently oxidized by the oxidant, N2 O: Mo5‡ ‡ N2 O ! Mo6‡ ÿOÿ ‡ N2 Interaction between ethane and the terminal oxygen atoms can generate ethyl radicals and replenish the protons, Oÿ ‡ C2 H6 ! C2 H5
‡ OHÿ

Figure 10.23. Selectivities to product versus conversion of ethane at different reaction conditions over 20 wt % HPMo=SiO2 : TR ˆ 450±570  C, W=F ˆ 0.01±0.08 g catalyst mlÿ1 minÿ1 , C2 H6 =N2 O ˆ 1=1ÿ4=1, (r) CO, (.) CO2 , (m) C2 H4 , ( ) CH3 CHO.…19† Reprinted from Applied Catalysis, A109, Hong and Moffat, p. 117, copyright 1994, with permission from Elsevier Science.

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Figure 10.24. Effect of time on stream on ethane conversion and selectivity over 20 wt % HPMo=SiO2 ; TR ˆ 510  C, W ˆ 1.00 g, F ˆ 5 ml minÿ1 , C2 H6 only and no oxidants, (d) conversion of C2 H6 , (e) selectivity of CO, (,) CO2 , (n) C2 H4 , ( ) CH3 CHO.…19† Reprinted from Applied Catalysis, A109, Hong and Moffat, p. 117, copyright 1994, with permission from Elsevier Science.

with the ethyl radicals remaining on the catalyst, desorbing into the gas phase, ethylating the anion, or forming an ethoxy species: C2 H5
‡ Oÿ ! C2 H5 ÿOÿ Although, as with the oxidation of methane, the formation, at least initially, of ethanol would be anticipated, none was observed in the present work. Ethanol may result from the interaction between the protons in the acid and the ethoxy group or, alternatively, ethoxy species may react with water to produce ethanol and a hydroxide ion.…20† The oxidation and dehydrogenation pathways may be summarized as %

C2 H4

C2 H6 ! C2 H5 OH ! CH3 CHO & CO; CO2 It is of interest to note that under the conditions employed in the present work the ethylene=acetaldehyde and CO=CO2 ratios produced from ethanol are 1.8 and

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2.9, respectively, whereas under similar conditions those from ethane are 1.6 and 2.5, respectively, providing some support for the contention that ethanol is a short-lived intermediate in the oxidation of ethane. As with methane, the effect of the addition of TCM to the feedstream for ethane conversion has also been studied on HPMo=SiO2 with N2 O as an oxidant.…21† At a reaction temperature of 450  C the predominant products are acetaldehyde and ethylene, whose selectivities vary relatively little with the concentration of TCM in the feed, at least up to 0.6 mol %. The primary products appear to be acetaldehyde, ethylene, and ethyl chloride, although the selectivities to the latter are small. As before, 20 wt % HPMo on SiO2 appears to be an optimum loading at least with respect to the conversion of ethane. As noted elsewhere in this volume, this corresponds, at least approximately, to a monolayer of HPMo. Changes in the composition of the anion have been found to have an advantageous effect in the conversion of ethane on heteropoly oxometalates.…22ÿ24† As noted elsewhere in this volume, replacement of the proton in heteropoly oxometalates by larger cations increases the thermal stability of the resulting material. Thus, HPMo begins to decompose at 300  C, and its ammonium salt has an exothermic peak at 500  C.…25† The replacement of one of the molybdenum atoms in the anion by antimony was found to have an advantageous effect on the thermal stability.…22† The XRD patterns for the potassium ammonium salts of HPMo show that in the absence of antimony the anion is decomposed at 500  C whereas with one molybdenum atom replaced by antimony the decomposition begins at 550  C. However, the addition of antimony results in a decrease in catalytic activity which these authors attribute to a stabilization of the reduced state of molybdenum. The addition of small quantities of iron, chromium, and cerium restored the activity in the oxydehydrogenation of ethane (Table 10.5). Variations in the nature and composition of the cations have also been found to have bene®cial effects on the conversion of ethane as well as other small alkanes.…26;27† At 360  C Cs2:5 Fe0:08 H1:26 PVMo11 O40 produced a conversion TABLE 10.5. Conversions of Ethane and Selectivities to Ethylene on K3 PMo12 O40 and Its Peripheral Metal-Substituted Analogue…22† Catalyst composition K3 PMo12 O40 K3 PMo11 Sb1 On K3 PMo11 Sb1 Fe1 On K3 PMo11 Sb1 Fe1 Ce0:25 On K3 PMo11 Sb1 Fe1 Ce0:25 Cr0:5 On

Ethane conversion (%)

Ethylene selectivity (%)

3.8 2.5 2.3 3.2 5.9

40 44 83 62 60

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of 1.7% and selectivities to ethylene, CO, and CO2 of 65, 14, and 21%, respectively.…26† At 400  C the conversion of ethane and yield of ethylene were 4.0 and 2.9%, respectively. In a subsequent report the effects of the addition of Cu2‡ , Mn2‡ , Co2‡ , and 2‡ Ni to Cs2:5 H1:5 PVMo11 O40 were investigated.…27† At 425  C conversions ranged from 8.6% for the former to 3.4% for the latter with selectivities to ethylene of 52 and 58%, respectively.

10.3. PROPANE Although relatively little work has been reported on the direct oxidation of propane to acrylic acid with heteropoly oxometalates the one-step production of this monomer offers considerable advantages in comparison with the usual industrial process involving the two-step oxidation of propene into acrolein and acrolein into acrylic acid. The reaction of propane with oxygen can produce the deep oxidation products CO and CO2 but a number of partial oxidation and oxidative dehydrogenation products have also been observed: 3=2O2

H3 CÿCH2 ÿCH3 ƒƒƒƒ! ÿH2 O 2O2

ƒƒƒƒ! ÿ2H2 O 3=2O2

ƒƒƒƒ! ÿ2H2 O 1=2O2

ƒƒƒƒ! ÿH2 O

H3 CÿCH2 ÿCOOH Propionic acid H2 CˆCHÿCOOH Acrylic acid H2 CˆCHÿCHO Acrolein H3 CÿCHˆCH2 Propene

as well as acetic acid and acetone. The direct oxidation of propane into acrylic acid has been recently studied on MOCC with two peripheral metal elements, V and Mo, in particular PVx Mo12ÿx O40 4ÿ,…28† and with partial substitution of the cationic protons by Cs and one of Fe, Ni, Co, Cu, Mn. Of the latter, partial substitution (0.08 mol=molÿ1 anion) of Fe produced the highest conversion of propane at 360  C with PMo12 O40 as the anion. Substitution of 1±3 vanadium atoms for Mo demonstrated that the anion PVMo11 O40 together with 2.5 Cs and 0.08 Fe generated a conversion of propane and selectivities and yields to acrylic acid of 17, 27, and 4.5%, respectively, at 360  C. The authors claim that the structure of the catalyst is retained during the reaction.

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The oxidative dehydrogenation of propane has been examined on 12molybdophosphoric acid with three of the molybdenum atoms substituted by vanadium, and cesium and copper replacing a portion of the hydrogen atoms.…29† The material producing the highest yield of propene carried the stoichiometry Cs2:5 Cu0:8 H3:34 PV3 Mo9 O40 and produced a selectivity to propene of 27% and a conversion of propane equal to 40% at 380  C. Carbon monoxide was the predominant product with a selectivity of 47%. Further studies on heteropoly acids with partial substitution of protons by heavy atoms have been done with propane as noted above for ethane.…26;27† With n‡ the stoichiometry listed as Mn‡ as Rh3‡ , 0:08 Cs2:5 H0:05ÿ0:08n PMo12 O40 and M Fe3‡ , Ni2‡ , Co2‡ , Cu2‡ , and Mn2‡ , the former produced the highest conversion of propane at 360  C at 17% and a selectivity and yield to acrylic acid of 8 and 1.4%, respectively. However, the highest yield to acrylic acid (2.4%) was obtained with Fe3‡ . 12-Molybdophosphoric acid (HPMo), after treatment with pyridine in aqueous solution, shows interesting properties in the oxidation of propane.…30ÿ32† The conversion of propane and the selectivity to acrylic acid at 340  C passed through maxima of approximately 9 and 45%, respectively, at approximately 420 and 470  C as the pretreatment temperature was increased. Earlier photoacoustic FTIR studies of the adsorption of pyridine showed that pyridinium ions were formed on and in HPW…33† but the process through which the pyridine diffuses from the surface to the bulk was found to be slower and more complex than that observed with ammonia. The authors of the propane paper attribute the activity and selectivity to acrylic acid to a reduced state of the catalyst. A vanadium-containing heteropoly acid (H4 PMo11 VO40 ) was prepared from MoO3, V2O5, and phosphoric acid.…34† A separate sample containing copper as well employed a similar preparative method but with CuO added. Evidence for the existence of the Keggin structure was found for the ®rst sample but the copper in the second sample was not incorporated in the Keggin anion. A physical mixture of CuO with vanadium-substituted HPMo had a catalytic activity similar to that of the aforementioned copper-containing solid. 10.4. BUTANE The results of several studies of the oxidation of n-butane on MOCC have appeared in the last decade or so.…35;36† As with propane the principal interest in butane relates to the possibility of obtaining economically advantageous yields to a partial oxidation product, in this case maleic anhydride, which is found along with acetic acid, acrylic acid, and the oxides of carbon. In one report…35† the catalyst was prepared from an aqueous solution of 12molybdophosphoric acid to which vanadyl oxalate was added. Cs2 CO3, with the

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261

ratio Cs=HPMo equal to 2.5, was then introduced to the resulting solution yielding a yellow precipitate, to which system pumice was added, followed by drying at temperatures less than 200  C and calcination in a ¯ow of oxygen at 380  C. Although the authors label this material as Cs2:5 H0:05 PMo12 O40 ‡ 2VO2‡ , no results of analysis of the preparation are reported. As noted elsewhere in this text, the ®nal composition may not match the preparative stoichiometries. The selectivities of the principal products decrease with increase in reaction temperature and conversion (Fig. 10.25). This is, as noted earlier, a rather typical result usually interpreted as re¯ecting the further conversion of the products. It is to be noted that the selectivity to acids is considerably more signi®cant than that to maleic anhydride and that at a reaction temperature of 360  C the conversion of n-butane is somewhat smaller than 50%.

Figure 10.25. Selectivities as a function of the conversion of n-butane and reaction temperature. MA, maleic anhydride; AcOH, acetic acid; AA, acrylic acid.…35† Reproduced by permission of Academic Press.

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Studies of the conversion of the products showed that the stabilities of maleic anhydride and acetic acid were high in comparison with those of the expected intermediates such as acetaldehyde, butadiene, butene, furan, and methyl ethyl ketone, none of which were observed. The vanadium-substituted 12-molybdophosphoric acid (H5 PV2 M10 O40
xH2 O) and its Cs salt (Cs3 H2 PV2 Mo10 O40 ) employed in the oxidation of n-butane have been characterized in detail…36† as discussed elsewhere in this volume. The conversion of n-butane and selectivity to maleic anhydride with the aforementioned acid and its salt are shown in Fig. 10.26. Both the conversion and selectivity found with the acid increase signi®cantly between 563 and 613 K, in contrast to the reciprocal behavior of these quantities commonly found in oxidation processes, whereas with the cesium salt both the conversion and selectivity are considerably lower than those found with the parent acid.

Figure 10.26. Catalytic behavior of H5 PV2 Mo10 O40 (u) and of Cs3 H2 PV2 Mo10 O40 (X) samples in the conversion of n-butane and in the selectivity to maleic anhydride as a function of the reaction temperature, (e) and (n), respectively.…36† Reproduced by permission of The Royal Society of Chemistry.

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263

These authors attribute the sharp increase at 573 K to a transformation in the acid which creates the sites active in the conversion of n-butene to maleic anhydride. In particular, this may be related to the presence of bridging inter-Keggin oxygen atoms (V±O±V) and the resulting greater reducibility of V5‡ sites, high oxygen mobility in the vicinity of such bridges, easier delocalization of electrons during the catalytic reaction due to electron transfer between the anions, and the existence of Lewis acid sites from oxygen vacancies. The reactivity of metal substituted heteropoly acids was evaluated from the steady-state and transient response kinetics of the oxidation of butane.…37† H4 PMo11 VO40 , H8 PMo10 MVO40 (MˆMn, Co, Ni, Cu, or Zn), Cs2:5 H1:5 PMo11 VO40 , and CuH2 PMo11 VO40 were examined. The preparations employed MoO3, V2 O5, and phosphoric acid and a dichloride to introduce the metal M. The cesium and copper salts were prepared from either a cesium or copper(II) carbonate. The authors claim that Cu or Zn had stoichiometrically substituted for Mo and were located in the Keggin anion, but did not comment on the location of the remaining elements. No elemental analysis results were provided. Of the acids studied the highest selectivity to maleic anhydride at 40% conversion was found with H8 PMo10 VCuO40 (45%) whereas a commercial (VO)2 P2 O7 catalyst produced a selectivity of 75% at 340  C. The rate of diffusion of oxygen in the lattice of the prepared catalysts during reoxidation correlates with the steady-state activity but the latter does not correlate with the rate constant for surface reoxidation. The authors contend that surface acidity also plays an important role. This appears to be consistent with earlier similar conclusions concerning the role of protons in the partial oxidation of methane.…1ÿ14†

10.5. ISOBUTANE Isobutane contains one tertiary C±H bond which has a bond energy of 390 kJ molÿ1 , less than those of the secondary and primary C±H bonds found in ethane and propane and consequently more readily broken.

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TABLE 10.6. Effect of V5‡ Substitution in Cs2:5 Ni0:08 H0:34‡x PVx Mo12ÿx O40 on the Conversion of Isobutane and the Selectivities to Methacrylic Acid…39† x

Conversion of isobutane (%)

Selectivity to methacrylic acid (%)

0 1 2 3 4

10 15 13 12 10

27 36 28 10 4

Note. Reaction temperature 320  C

Vanadium-substituted HPMo with hydrogen atoms partially substituted by Cs and Ni has been shown to be active in the oxidation of isobutane.…26;27;38;39† On substitution of Cs2:5 Ni0:08 H0:34‡x PVx Mo12ÿx O40 with vanadium the conversions increased, reached a maximum of 15% for x equal 1, and decreased (Table 10.6), whereas the selectivity to MAA behaved similarly reaching a maximum of 36% also for x ˆ 1. Methacrolein, acetic acid, CO, and CO2 were also formed. Variations in the cesium content were also found to in¯uence the oxidation of isobutane (Table 10.7).…38;39† The highest conversion was found for x equal to 2.5 in Csx Ni0:08 H3:84ÿx PVMo11 O40 . The catalytic activity for isobutane oxidation of salts of HPMo prepared with ammonium and potassium as cations has been shown to be enhanced by the addition of iron.…40;41† At a reaction temperature of 350  C with 1.5 iron atoms KUÿ1 the conversion approached 12% in contrast to 4% in the absence of added iron, while the selectivities to products decreased. Although the authors imply that the iron occupies the charge-balancing cation positions in the cubic structure, partially replacing the ammonium ion, they admit that ®nely dispersed iron oxide may be present. The added iron is suggested to affect the molybdenum ion redox properties although the direct participation of iron may also be occurring. TABLE 10.7. Effect of Cs in Csx Ni0:08 H3:84ÿx PVMo11 O40 on the Oxidation of Isobutane…39† x 0 1 2 2.5 3

Conversion of Isobutane (%)

Yield of Methacrylic Acid (%)

6 4 12 31 8

1.3 1.8 5.6 9.0 0.8

Note. Reaction temperature 340  C:

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265

The kinetics of the oxidation of isobutane has been evaluated on a cesium± ammonium mixed salt of 11-molybdo-1-vanadophosphoric acid.…42† The ratelimiting step appears to be the reaction of isobutane on the oxidized surface of the catalyst. The total selectivity to methacrolein (MA) and MAA was found to depend only on the conversion of isobutane and not on the hydrocarbon=oxygen ratio or the reaction temperature in the 300±350  C range. While both the direct route from isobutane to MAA and that through MA are signi®cant, MA appears to be particularly reactive compared with isobutane and the degradation of the former plays an important role in the ®nal yield of MAA. In contrast, MAA is very stable and its degradation is negligible at low conversions. Further work on Csx H3ÿx PMo12 O40 as a catalyst for the oxidation of isobutane has been reported.…43† The same authors had reported previously that the Cs2:5 Ni0:08 H0:34 PMo12 O40 catalyst produced the highest conversion (24%) and total yield (8.0%) of MAA and MA at 340  C. The use of Fe3‡ in lieu of Ni2‡ at the same reaction temperature yielded a conversion of 14% and a selectivity to MAA of 35%, the latter being the highest found in this work. The effect of Cs‡ substitution in HPMo was also examined at 340  C (Table 10.8). With increasing x (Csx H3ÿx PMo12 O40 ) the conversion and selectivity to MAA increased and passed through maxima, the former of 17% at x equal to 2.85 and the latter of 34% at x equal to 2. However, the rate of reaction decreased continuously as x was increased. The oxidation of isobutane has also been investigated on a vanadiumsubstituted HPMo (Cs1:6 H2:4 P1:7 Mo11 V1:1 O40 ).…44† At 349  C a conversion of 10.6% and a selectivity to MAA of 37.6% were obtained. For temperatures higher than 250  C anionic vacancies are created as a result of the elimination of hydroxyl groups as water and it is believed that these vacancies abstract hydride ions from the isobutane to form isobutane which the authors believe to be the ®rst

TABLE 10.8. Conversions and Selectivities in the Oxidationa of Isobutane over Csx H3ÿx PMo12 O40 …43† Selectivitiesb x 0 1 2 2.5 2.85 3 a b

Conversion

MAA

MA

HOAc

COx

7 6 11 16 17 8

4 23 34 24 5 0

18 17 10 7 10 10

8 10 7 7 5 6

70 50 50 62 81 67

Reaction temperature 340  C. MAA, methacrylic acid; MA, methacrolein; HOAc, acetic acid.

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step in the oxidation of isobutane to form MAA. This suggested mechanism is similar to that proposed for the partial oxidation of methane on silica-supported heteropoly acids.…1;9ÿ18† In the section on propane oxidation, 12-molybdophosphoric acid treated with pyridine was reported to catalyze the oxidation of propane to acrylic acid and acetone.…32† The same authors have extended their work to the oxidation of isobutane on the same catalyst.…45† Because as with propane the pretreatment temperature had a substantial effect on the results, the experiments were done with a pyridine-treated HPMo after activation at 420  C in a nitrogen atmosphere for 2 h. Reaction temperatures from 280 to 340  C were employed. The highest selectivity (55.9%) was obtained at a conversion of 8.4% and 280  C reaction temperature. At 300  C a conversion of 22% and selectivities to MAA and acetic acid of 51 and 22%, respectively, were measured. The authors suggest that the pyridine treatment produces highly reduced HPMo with oxygen-de®cient anions which provide sites for activating molecular oxygen and isobutane. This is similar to the proposals put forth in earlier studies of the oxidation of methane.…9ÿ18† Salts of 12-molybdophosphoric acid prepared from cesium and copper or iron of stoichiometry Cs2 Mx y‡ H1ÿxy PMo12 O40 with M as Fe3‡ or Cu2‡ and 0  x  0:43 have been tested as catalysts in the oxidation of isobutane to MAA.…46;47† These salts were found to consist of two phases with the acid covering the particles of the cesium salt and the transition metal doping the supported acid phase. These results appear to be consistent with earlier observations that demonstrated that salts of heteropoly acids prepared from divalent cations could not be prepared as single-phase, one-component systems but rather as mixtures of the parent acid and a salt of the multivalent cation.…48;49† The introduction of copper decreases the conversion (although the rate of conversion is increased) and the selectivity to MAA whereas the presence of iron increases the conversion and the MAA selectivity (Table 10.9). The authors conclude that copper is directly involved in the redox mechanism by catalyzing the reduction of the solids whereas, in contrast, iron has no direct effect on the

TABLE 10.9. Oxidation of Isobutane to Methacrylic Acid on a…47† Cs2 My‡ x H1ÿxy PMo12 O40 Compound

Conversion (%) of isobutane

Selectivity to MAA (%)

Cs2 Cs2 Cu0:30 Cs2 Fe0:30 Cs2:5 Cs2:5 Fe0:10

7.2 6.6 6.3 10.3 10.3

12 6 18 12 19

a

Reaction temperature 613 K.

OXIDATION PROCESSES

267

reduction rate, but decreases it indirectly by decreasing the number of protons, but as a result of the latter the selectivity to MAA is increased. Although the present volume is primarily concerned with the surface and catalytic properties of heteropoly oxometalates with Keggin structure, it is relevant to note that the oxidative dehydrogenation of isobutane to isobutene has been studied on a Wells±Dawson catalyst (K6 P2 W18 O62
10H2 O).…50;51† Calcination of this material at temperatures higher than 850 K produced a compound (K3 PW12 O40 ) with Keggin structure plus a second phase resulting from the decomposition product potassium phosphate and the aforementioned Keggin solid. The Keggin product was found to have a higher activity in the conversion of isobutane to isobutene than the Wells±Dawson and K3 PW12 O40 which was suggested to result from the presence of an amorphous layer of unknown stoichiometry at the surface of the calcined Wells±Dawson compound.

10.6. METHACROLEIN The oxidation of MA to MAA

on alkali molybdophosphate catalysts was initially reported in a 1977 patent.…52† Subsequently a variety of salts of HPMo were tested.…53† The highest oxidation rate per surface area at 320  C was found with the cobalt salt whereas the highest selectivity (46%) to MAA was obtained on the thallium salt.…53† The authors report a correlation between the rate and the ionic potential taken as Z=r. Work with molybdenum phosphate catalysts has shown that the addition of rubidium produces a substantial increase in conversion of MA and selectivity to MAA at 350  C.…54† Values as high as 73 and 65% for the former and latter, respectively, were obtained. The increased activity is believed to result from the formation of the rubidium salt of HPMo which coexists with molybdenum oxide. In later work HPMo and its copper and cesium derivatives along with HSiMo and HPMo10 V2 were found to be active and selective for the MA process.…55ÿ57† Both the conversion and selectivity to MAA at 300  C on HPMo were found to increase on addition of small partial pressures of water. Introduction of a silica support for HPMo increased both the conversion and selectivity with maxima for each being found at a 60 wt % loading. The addition of vanadium has been found to increase the selectivity to MAA…58† whereas the introduction of cesium produced a decrease in catalytic activity.…59;60† As already noted, this appears to be consistent with the results

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reported earlier for the oxidation of methane on HPMo=SiO, as the protons are replaced by cesium.…1;9ÿ18† Kinetic studies of the oxidation of MA on Csx H3ÿx‡y ‰PMo12ÿy Vy O40 Š (x ˆ 0; y ˆ 0; 1; x ˆ 1; y ˆ 0; 1; 2) have shown that the process can be ®tted to a Mars±van Krevelen model.…61;62† Substitution of one or two molybdenum atoms by vanadium results in a substantial decrease in the rate of oxidation of MAA whereas the rate of oxidation of MA is less signi®cantly altered. Apparently with the catalysts containing vanadium the increase in selectivity to MAA can be attributed to a reduction in the subsequent oxidation of MAA. The replacement of hydrogen by cesium produces a decrease in the formation of byproducts via parallel and consecutive oxidation and a reduced blocking of the reoxidation of the catalyst resulting from strong adsorption of MA.

10.7. ISOBUTYRIC ACID The oxidative dehydrogenation (ODH) of isobutyric acid (IBA) to MAA

is an important step in the methyl methacrylate process, offering the possibility of avoiding the conventional acetone cyanohydrin process. In addition to MAA, acetone, propene, and carbon oxides are produced:

The ODH process has been found to be catalyzed by heteropoly oxometalates of various compositions.…63ÿ73† With the salts of 12-molybdophosphoric acid prepared from the alkali metals and alkaline earths, MAA was formed with a selectivity of 44±66% at 300  C.…63† The conversion appears to increase with the decreasing electronegativity (x) of the alkali metal cation (Table 10.10), although as the proton is stepwise substituted by sodium a maximum is seen at Na2 HPMo. No systematic changes are observed in the selectivities. 12-Molybdosilicic acid produced a conversion of 100% but a selectivity to MAA of 35.7%. The authors propose that the reoxidizability of Mo5‡ increases with decreasing electronegativity of the cations.

OXIDATION PROCESSES

269

TABLE 10.10. Conversion of Isobutyric Acid (IBA) and Selectivity to Methacrylic Acid (MAA) on Salts of 12-Molybdophosphoric Acida…63† Cation

xb

Conversion of IBA (%)

Selectivity to MAA (%)

H NaH2 Na2 H Na

2.1 1.7c 1.3c 0.9

32.0 67.8 77.4 54.6

50.9 50.9 49.7 55.9

Li Na K Rb Cs

1.0 0.9 0.8 0.8 0.7

50.5 54.6 75.8 92.0 95.7

50.5 55.9 66.0 60.9 58.0

a

Reaction temperature 300  C. Pauling electronegativity. c Average. b

Further work with the salts of HPMo formed from various elements of groups 3, 8±12 has led the authors to propose that the reoxidizability of Mo5‡ increases with increasing standard electrode potential (SEP) of the cations.…64† Although the strength of the molybdenum±oxygen bond was also expected to increase with the increasing SEP, the authors found that the reducibility of the group 3, 8±12 salts by IBA increased with increasing SEP of the cations and as a consequence the ability of the cations to capture the electrons formed by the reduction process is also important. The cations are thus proposed to alter the strength of the molybdenum±oxygen bond by changing the electrochemical properties of the molybdenum atoms as well as in¯uencing the scission of the molybdenum±oxygen bond through delocalization of the electrons formed by reduction with the former role predominating in the case of alkali metal and alkaline earth cations and the latter role predominating with salts formed from groups 3, 8±12. Additionally, kinetic analysis suggests that in the oxidation of IBA at 300  C the reduction of the catalyst by IBA is rate-controlling. Further work by the aforementioned authors showed that both the conversion of IBA and the selectivity to MAA were decreased by the stepwise substitution of molybdenum atoms by tungsten (Table 10.11).…65† In contrast, substitution by vanadium produced a maximum in conversion and in selectivity at substitutions of one and two vanadium atoms, respectively (Table 10.12). The authors argue that the Mo±O bond in H3 PMo10 W2 O40 is the same. MAA was shown to be produced from IBA on H5 PMo10 V2 O40 at 250  C with approximately 100% conversion and greater than 70% selectivity.…66† The source of the differences between these results and those shown in Table 10.12 for a value of x equal to 2 is not clear. It is to be noted that the reaction temperatures

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TABLE 10.11. The Effect of Substitution of W in HPMo on the ODHa of IBA…65† xb

Conversion (%)c

Selectivity (%)d

0 1 2 3 4 5 6 12

80.6 70.0 54.5 44.0 28.1 23.8 22.3 9.0

60.5 57.8 43.7 32.5 19.5 16.7 10.5 1.5

a

300  C. x in H3 PMo12ÿx Wx O40 . c Of IBA. d To MAA. b

in Table 10.12 are 300  C. Although the latter authors…66† employed a pulse reactor, they compared their results with those obtained with a micro-integralreactor. The latter authors attribute the high selectivity to the lattice oxygen bonded to vanadium, both in the bulk and in the surface. Gaseous oxygen has a detrimental effect on the selectivity to MAA. The catalytic properties of V2 O5 ÿP2 O5 (V±P) and H5 PMo10 V2 O40 have been compared for the ODH of IBA at 200±300  C.…67† With the former the selectivity to MAA passes through a maximum (50±60 mol %) at P=V ratios of 1.0±1.6. The selectivities to acetone, propylene, and MAA are attributed to basic, acidic, and combination sites. The selectivities to MAA with the P=V (1.0±1.6) catalysts are similar to those obtained with H5 PMo10 V2 O40 , but the principal byproducts are propylene and acetone, respectively, at conversions greater than TABLE 10.12. The Effect of Substitution of V in HPMo on the ODHa of IBA…65† xb

Conversion (%)c

Selectivity (%)d

0 1 2 3

48.0 75.0 52.3 33.2

57.1 68.8 71.9 71.6

a

300  C. x in H3‡x PMo12ÿx Vx O40 . c Of IBA. d To MAA. b

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271

50%, which is interpreted as suggesting that the V±P catalyst is more acidic and=or less basic than H5 PMo10 V2 O40 . The oxidation activities of the V±P (1.0± 1.6) and the H5 PMo10 V2 O40 catalysts are found to be of the same order of magnitude. This author related the selectivities to the acidic=basic properties of the catalysts. Kinetic studies have shown that the ODH of IBA may be ®tted to the Hans± van Krevelen mechanism for each primary product.…68† Signi®cantly different activation energies were measured for propylene, acetone, and MAA suggesting the existence of three speci®c sites, acid or cationic for propylene and bridged oxygen for acetone and MAA. A separate kinetic study on H5 PV2 Mo10 O40 and Csx H5ÿx V2 Mo10 PO40 showed that the conversion and surface area passed through a maximum at 225  C at three Cs‡ , whereas the selectivities to MAA, propene, and acetic acid reached maxima at x ˆ 2 (78%), 1 (8%), and 3 (19%), respectively.…69† Catalyst deactivation was attributed to the loss of molybdenum and vanadium. Studies of the structure retention and morphological properties of ionexchanged heteropoly oxometalates were reviewed in the chapter on the microporosity of certain of these solids.…70ÿ72† 12-Molybdophosphoric acid, its cesium and ammonium salts, and a series of catalysts containing both of the latter cations prepared by ion exchange were investigated to determine the contribution of the cation in the IBA-to-MAA process.…73† The nature of the cation associated with the PMo12 O40 ÿ3 anion has a substantial effect on the product selectivity (Table 10.13). At a given temperature the selectivity to MAA is highest for the ammonium salt and lowest for the cesium salt. The differences in the three catalysts are also re¯ected in the selectivities to acetone and propene. The highest and lowest selectivities to acetone are obtained with the cesium and ammonium salt, respectively, whereas those to propene were found with the parent acid and cesium salt, respectively. The selectivities to acetone and propene increase with temperature from 250 to TABLE 10.13. Selectivitiesa to Major Products over Various 12-Heteropoly Compounds…73† 250  C Compound H3 PMo12 O40 …NH4 †3 PMo12 O40 Cs3 PMo12 O40

300  C

350  C

W=Fb MAA Propene Acetone MAA Propene Acetone MAA Propene Acetone 1.45 1.45 1.45

49.8 58.1 17.5

20.3 14.4 11.3

29.3 26.8 71.2

43.8 55.8 40.1

Note: Data for X=F equal to 0.85 can be found in Fig. 10.1. a mol %. b mg min mlÿ1 .

21.2 10.4 9.9

34.6 32.1 49.0

36.4 42.6 24.7

24.6 20.6 11.4

38.7 36.0 63.2

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350  C, whereas those to MAA decrease. The cesium salt is evidently a less effective catalyst for the ODH of IBA than either the parent acid or its ammonium salt. The disadvantageous effect of the replacement of the proton by cesium has also been observed with HPMo=SiO2 in the partial oxidation of methane.…9ÿ18† Previous studies have demonstrated the existence of deviations from stoichiometry and the presence of residual protons in salts prepared from the heteropoly acids.…74;75† The role of protons in the IBA-to-MAA process has been investigated by exchanging the residual protons with the cations already present in the given salt. At 250, 300, and 350  C the selectivity to acetone increases markedly after one exchange of cesium into CsPMo, whereas that to MAA decreases proportionately (Fig. 10.27). The effect of self-exchange with the ammonium salts is less dramatic with exchange only slightly increasing the selectivities to MAA and propene whereas that to acetone decreases, in contrast to the observations with the cesium self-exchanged solids (Fig. 10.28).

Figure 10.27. Selectivities to methacrylic acid, propene, and acetone as a function of the number of exchange reactions for CsPMoCsn (W=F ˆ 0.85 mg min mlÿ1 ) at 300  C: (d) methacrylic acid, (m) propene, (j) acetone.…70;71;73†

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273

Figure 10.28. Selectivities to methacrylic acid, propene, and acetone as a function of the number of exchange reactions for NH4 PMoNH4 n (W=Fˆ0.85 mg min mlÿ1 ) at 300  C: (d) methacrylic acid, (m) propene, (j) acetone.…70;71;73†

To determine the effect of the two cations, Cs‡ and NH4 ‡ , and their concentrations on the catalytic behavior of the salts, the ODH of IBA was investigated on catalyst samples prepared by ion-exchanged modi®cation of the unexchanged Cs‡ and NH4 ‡ salts. With the NH4 PMo salt into which Cs is exchanged up to six successive exposures the changes in selectivities were relatively small in spite of the signi®cant quantities (40±50%) of cesium ions which were shown to be exchanged into the ammonium salt after a single exposure (Fig. 10.29). In contrast, the changes in selectivities brought about by exchanging NH4 ‡ ions into the cesium salt are more signi®cant (Fig. 10.30), with the MAA and propene selectivities increasing and that to acetone decreasing. Evidently the ammonium ion, even in small quantities, has a bene®cial effect in comparison with the cesium ion. Additionally, only small quantities of ammonium ion are

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Figure 10.29. Selectivities to methacrylic acid, propene, and acetone as a function of the number of exchange reactions for NH4 PMoCsn (W=F ˆ 0.85 mg min mlÿ1 ) at 300  C: (d) methacrylic acid, (m) propene, (j) acetone.…70;71;73†

suf®cient to increase the activity of the catalysts (Table 10.14). Conversions with the cesium salts exchanged with ammonium were higher than those selfexchanged with cesium. Additionally the ammonium salts exchanged with cesium ions appeared to be less susceptible to deactivation than the selfexchanged ammonium salts. With only IBA and no oxygen in the feedstream, appreciable conversions of IBA and selectivities to the products were initially noted indicating the participation of lattice oxygen in the process, although after 4 h on stream the production of MAA had decreased to a negligible value apparently resulting from the depletion of lattice oxygen atoms and the destruction of the Keggin anion

OXIDATION PROCESSES

275

Figure 10.30. Selectivities to methacrylic acid, propene, and acetone as a function of the number of exchange reactions for CsPMoNH4 n (W=F ˆ 0.85 mg min mlÿ1 ) at 250  C: (d) methacrylic acid, (m) propene, (j) acetone.…70;71;73†

(Table 10.15). The gaseous oxygen in the feedstream apparently serves, at least in part, as an oxygen source for the reoxidation of the catalysts. Although decomposition of the Keggin anion eventually occurs under reaction conditions in the absence of oxygen, after either the ion-exchange experiments or the reaction studies with oxygen present no evidence for the decomposition of the anion was found. The conversions and selectivities are evidently in¯uenced by the nature and concentrations of the cations as well as the number of protons present in the salts. The effect of the concentrations of the cations can be seen in Figs. 10.31 and 10.32. The MAA selectivity increases slightly with NH4 ‡ content whereas that to acetone decreases and to propene remains almost unchanged (Fig. 10.31).

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TABLE 10.14. Conversions of IBAa for 12-Heteropoly Salts…73† Compound H3 PMo12 O40 …NH4 †3 PMo12 O40 Cs3 PMo12 O40 NH4 PMoNH4 1 NH4 PMoNH4 5 NH4 PMoCs1 NH4 PMoCs6 CsPMoCs1 CsPMoCs5 CsPMoNH4 1 CsPMoNH4 6 a b

W=Fb

250  C

300  C

350  C

0.85 1.45 0.85 1.45 0.85 1.45 0.85 1.45 0.85 1.45 0.85 1.45 0.85 1.45 0.85 1.45 0.85 1.45 0.85 1.45 0.85 1.45

11.7 13.1 48.0 48.5 11.6 5.7 40.9 40.1 18.2 28.8 65.5 67.4 58.5 91.3 3.9 8.5 6.2 7.6 43.4 31.5 61.6 40.2

43.0 60.0 76.0 71.0 20.6 23.5 65.0 74.6 19.6 59.0 89.6 96.0 79.5 84.9 21.4 11.3 24.4 22.7 75.0 60.0 58.3 94.8

88.3 94.0 88.1 83.4 43.1 35.6 74.4 67.9 70.9 79.7 96.8 95.9 92.8 99.0 25.6 33.2 30.6 40.6 99.6 99.3 86.1 98.7

mol %. mg min mlÿ1 .

TABLE 10.15. Selectivitiesa to Products from Isobutyric Acid over 12-Heteropoly Compounds in the Absence of Oxygen…73† Compound

MAA

Propene

Acetone

Nitriles

H3 PMo12 O40 …NH4 †3 PMo12 O40 Cs3 PMo12 O40 NH4 PMoNH4 1 NH4 PMoNH4 5 CsPMoCs1 CsPMoCs5 NH4 PMoCs1 NH4 PMoCs6 CsPMoNH4 1 CsPMoNH4 6

11.0 2.0 11.0 5.5 17.5 23.5 8.8 10.0 1.0 31.5 10.0

74.9 86.0 4.0 81.5 54.0 6.0 5.8 78.0 91.0 19.0 84.5

13.0 3.0 86.5 3.0 5.0 70.0 84.3 4.0 3.0 51.0 5.0

Ð 8.0 Ð 8.5 22.0 Ð Ð 5.5 5.0 Ð 1.0

Note: All experiments at 300  C and W=F 0.85 mg min mlÿ1 . a mol %.

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277

However, the MAA selectivity is strongly dependent on the concentration of Cs‡ in the salts. With the CsPMo salt the selectivity to MAA is approximately 5% at 3 Cs‡ anionÿ1 (Fig. 10.32). However, as NH4 ‡ cations are introduced and the Cs‡ are removed, the MAA selectivity at 300  C increases dramatically whereas the selectivity to acetone decreases proportionally and that to propene decreases relatively little (Fig. 10.32). It is clear that the cations play a signi®cant role in the IBA process. The effect may result from the interaction of the cationic species with the Keggin anion. The cesium cation may be considered as a stable species, apparently stabilizing the CsPMo salt and leading to an enhanced thermal stability, whereas, in contrast, protons and ammonium ions are more reactive species which can interact with both the anion and the IBA. The selectivities to MAA and acetone show a stronger dependence on the Cs‡ content than on that of NH4 ‡ , with which only small quantities are required to maximize the MAA selectivity for the NH4 ‡ =Cs‡ salts. As noted earlier, PAS FTIR studies have shown that residual quantities of protons exist in salts of the heteropoly acids prepared to be stoichiometric.…74;75† The importance of the proton in the partial oxidation of methane has been described earlier in this chapter. With silica supported HPMo and its salts, as the number of cesium cations per Keggin anion increased, up to a cesium loading of approximately 3.5 per anion, concurrently the selectivity to formaldehyde, turn-

Figure 10.31. Selectivity to methacrylic acid (d), propene (j), acetone (m) as a function of the number of NH4 ‡ cations per Keggin anion (W=F ˆ 0.85 mg min mlÿ1 ) at 300  C.…70;71;73†

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Figure 10.32. Selectivity to methacrylic acid (d), propene (j), acetone (m) as a function of the number of Cs‡ cations per Keggin anion (W=F ˆ 0.85 mg min mlÿ1 ) at 300  C.…70;71;73†

over rate of oxidant N2 O, and the consumption of oxygen decreased.…1;9ÿ18† A mechanism in which protons extract (presumably terminal) oxygen atoms from the anion to release water and form anions each de®cient in one proton and a second anion containing a vacancy or empty site (u) was proposed: 2KUOH ! KUO ‡ KUu ‡ H2 O The vacancy could serve as a site for dissociation of the gas-phase oxidant for replacement of the missing oxygen atom in the anion. Alternatively, oxygen molecules may be converted to O2 ÿ at the vacancy. While the self-exchange process increases the number of cesium cations in the structure, this may result from replacement of protons and=or insertion of cations in empty cationic sites, thus altering the distribution of acidic sites and potentially the activity and selectivities of the catalyst in the IBA process. As discussed in an earlier chapter, the surface areas and pore structures of the heteropoly oxometalate salts are strongly in¯uenced by both the nature and concentrations of the cations.…76† Although the morphology of the catalyst may in¯uence its catalytic properties in the IBA process, the observations appear to have little or no dependency on the surface areas (Table 10.16). The nature of the

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279

TABLE 10.16. Surface Areas of Catalysts Used for the Oxidative Dehydrogenation of Isobutyric Acid…73† Compound H3 PMo12 O40 …NH4 †3 PMo12 O40 Cs3 PMo12 O40 NH4 PMoNH4 1 NH4 PMoNH4 5 NH4 PMoCs1 NH4 PMoCs6 CsPMoCs1 CsPMoCs5 CsPMoNH4 1 CsPMoNH4 6

Surface area (m2 gÿ1 )  5.0 88.5 144.1 88.0 75.7 137.5 85.7 88.5 41.3 110.0 60.2

cation and its concentration are evidently more important than the morphology of the catalyst. Few, if any, suggestions have appeared in the literature as to molecular mechanisms and reaction intermediates in the ODH of IBA process. In view of the apparent reciprocal relationship between the formation of MAA and acetone as found in this work, any such mechanistic proposal must take this observation into account. A carbocation mechanism, in which a hydride ion is abstracted from the IBA, may serve such a purpose. An E1 mechanism in which the elimination of a proton produces MAA or a nucleophilic SN 1 assimilation of oxygen forming CO2 and acetone can be postulated:

It is worth noting that at temperatures of 300±350  C isobutyronitrile and methacrylonitrile were obtained, in addition to the usual products, from IBA with the ammonium salts of 12-molybdophosphoric acid.…77† Although the selectivities to nitriles were less than 5% in the presence of oxygen, values as high as 22% were measured at 300  C in the absence of oxygen. The in¯uence of the addition of small amounts of tungsten on the catalytic properties of H4 PMo11 VO40 and their pyridinium salts in the ODH of IBA has

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been evaluated.…78† The synthesized solids were found from 31 P NMR spectra to be mixtures of species with different relative quantities of tungsten. In order to vary the proportion of each species, three stoichiometric ratios of W=Mo (1=10, 2=9, 3=8) were employed with each of the solids thus containing mixtures of the title species. The ratios were labeled as y equal to 1, 2, 3, respectively. The Bronsted acidity increases with increasing tungsten content (increasing y) while the oxidative activity decreases (Table 10.17). The increasing selectivity to propene with increasing tungsten content on both unsupported and supported catalysts appears to be consistent with earlier contentions that the formation of propene is acid-catalyzed. The authors conclude that small amounts of tungsten may modify either the electron transfer or the oxygen lability but as can be seen from Table 10.17 the presence of both vanadium and tungsten is disadvantageous for good oxidative catalytic activity. The oxidative dehydrogenation of IBA to MAA has been further investigated on H4 PVMo11 O40 prepared from V2 O5 , MoO3, and phosphoric acid.…79† The sodium and copper salts were obtained from the parent acid by exposure to aqueous solutions of sodium bicarbonate and copper nitrate. The aforementioned compounds were impregnated on silica and employed as catalysts at 340  C with IBA, air, and water in the reactant stream. The H4 PVMo11 O40 catalyst produced a conversion of 94±95% and a selectivity to MAA of approximately 70% both of which remained relatively unchanged over 15 days. The introduction of copper increased both the conversion and selectivity whereas that of sodium had a detrimental effect. With Cuy H4ÿ2y PVMo11 O40 =silica the conversion reached a maximum of 98% at y ˆ 0:3 whereas the selectivity was the highest (75±76%) for y ˆ 0:5. The conversions and selectivities with the copper-containing catalysts changed relatively little over 40 days. The authors attribute the increases obtained with copper to the presence of Bronsted acidity and oxygen vacancies. Unfortu-

TABLE 10.17. Catalytic Properties in the ODHa of IBA…78† Selectivities Catalyst H4 PMo11 VO40 y ˆ 0:5 yˆ1 yˆ2 yˆ3 H4 PW11 VO40 a b

Conversion 98 (62)b 99 (58) 83 (52) 82 (25) 80 (12) 29

MAA 69 63 58 25 26 3

(72) (51) (48) (41) (25)

Acetone

Propene

16 16 14 17 16 0

14 21 28 64 66 97

(15) (16) (15) (15) (13)

(12) (33) (38) (45) (62)

320  C, 12 h on stream. Numbers in parentheses refer to results for silica-supported catalysts with a loading of 25%.

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281

nately the nature of the copper-containing catalysts and the location of the added copper were not investigated. Potassium=ammonium salts of 12-molybdophosphoric acid have been examined for the IBA reaction.…80† Although additional discussion of this work may be found in the chapter on stability, it is relevant to note here that these workers found only a single phase after calcination at 640 K with patterns characteristic of the cubic structure expected for the Keggin anion with decomposition beginning at approximately 693 K. The potassium salt did not decompose until 753 K. The conversion of IBA at 533 K and selectivities to MAA obtained with (NH4 †3 PMo, K(NH4 †2 PMo, K2 …NH4 )PMo, and K3 PMo (abbreviated as K0 , K1 , K2 , K3 , respectively) were similar with samples calcined at 653 K. However, with increase in the calcination temperature to 693 K the conversion with K0 increased whereas that with K1 and K2 decreased and K3 remained relatively unchanged. With further increase in the calcination temperature to 713 K conversions with all samples decreased as did the selectivities although the conversion with K0 remained higher than that found with the remaining catalysts. After calcination at 763 K the K0 catalyst generated a relatively high conversion and an MAA selectivity of approximately 50%. These results, in particular those with the ammonium salt, appear consistent with those obtained with the cesium=ammonium salts, discussed earlier in this chapter. In situ powder XRD has been employed to follow the structural changes of H4 PVMo11 O40
32H2 O during heating and the conversion of IBA to MAA.…81† With feed mixtures of IBA=water equal to 1=9 and at 553 K a conversion of 50% and selectivities of 50% to MAA and 42% to acetone were measured together with small selectivities to propene. Although more detailed discussion of this work is included in the chapter on stability, it should be noted here that the authors ®nd a new cubic phase of a water-free vanadyl salt of the acid to be responsible for maximum conversions. Cesium salts of molybdovanadophosphoric acids (Csn H3‡xÿn PMo12ÿx Vx O40 ), with values of n from 0 to …3 ‡ x† and those of x from 0 to 2 have been studied for the ODH of IBA.…83† In the absence of cesium the heteropoly acid HPMo produced, at 623 K, the highest conversion (approximately 50%) but the selectivity to MAA was less than 40%. However, with HPMo11 VO40 the selectivity to MAA exceeded 50% while that to acetone was 5%. The conversions and MAA selectivities with the cesium salts of PMo11 VO40 were maximized at a value of n equal to 2.75 (Fig. 10.33). Potassium=ammonium salts of HPMo with added antimony were employed for the ODH of IBA.…84† Potassium=ammonium salts of HPMo were prepared by the addition of HNO3 to an aqueous solution of (NH4 †6 Mo7 O24
4H2 O, …NH4 †H2 PO4 , and KCl. The preparation of the antimony salts employed an

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Figure 10.33. Oxidative dehydrogenation of isobutyric acid at 623 K on Csn H4ÿn PMo11 VO40 for n equal to 0 to 4. s, methacrylic acid; u, acetone; n, propylene; m, COx.…83† Reprinted from Catalysis Today, 33, Lee et al., p. 183, copyright 1997, with permission from Elsevier Science.

aqueous solution of K(SbO)C2 H4 O6 , …NH4 †6 Mo7 O24
4H2 O, and …NH4 †H2 PO4 followed by the addition of HNO3. These authors employed FTIR, ion chromatography, diffuse re¯ectance UV±visible spectroscopy, and powder XRD for characterization purposes. The authors conclude that the antimony replaces the ammonium cation in the secondary structure of the K1 Sby (one atom of potassium and y atoms of antimony per Keggin anion) compounds. The addition of antimony has a markedly bene®cial effect on the thermal stability with decomposition beginning at 550±600  C for the K3 Sby series of catalysts. The introduction of antimony also leads to the reduction of molybdenum from Mo6‡ to Mo5‡ . Migration of antimony into the secondary structure during calcination results in the release of nitrogen with concomitant reduction of molybdenum. As a consequence of the changes in the redox properties of the molybdenum, the activity and selectivity to MAA in the IBA process are decreased. Cupric salts of H4 PMo11 VO40 have been tested for their activity and selectivity in the ODH of IBA.…85† The highest selectivities to MAA were obtained with x ˆ 0:5±0.7 …Cux H4ÿ2x PMo11 VO40 † (Fig. 10.34) although the highest conversion was found with x ˆ 0:3 at 593 K. Deactivation occurred after 200 days. Evidence is provided for the formation of vanadium species exterior to the Keggin anion some of which assume a PMo12 O40 3ÿ stoichiometry and structure, together with the agglomeration of vanadium and copper atoms into small oxide clusters, which can participate in oxygen and electron exchange with Keggin anions.

OXIDATION PROCESSES

283

Figure 10.34. Oxidative dehydrogenation of isobutyric acid on Cux H4ÿ2x PVMo11 O40 (593 K) (time on stream: 3 days).…85† Reprinted from Applied Catalysis, 178, Blouet-Crusson et al., p. 69, copyright 1999, with permission from Elsevier Science.

Further work on ammonium=cesium salts of HPMo for application in the ODH of IBA has been reported.…86† The salts were prepared by solid-state cationic exchange. The authors conclude that the activity of the cesium salts decreases as the quantity of the surface layer acid phase decreases, whereas with the NH4 ‡ =Cs‡ salts the activity increases with cesium content up to 3 but decreases between 3 and 3.1 while acetone becomes the predominant product.

10.8. n-PENTANE The oxidation of n-pentane on 12-molybdovanadophosphoric acids (H3‡n PVn Mo12ÿn O40
xH2 O, with n equal to 0 to 3) was studied at 543± 603 K.…87† The activity increased linearly with increasing n, except for the sample with n equal to 3 which had a lower activity than expected, attributed to partial decomposition of the catalyst. Only maleic anhydride was formed and its selectivity also increased linearly with n. The authors suggest that V±O±Mo bridged oxygen atoms are involved in the oxidation process.

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10.0. BUTADIENE The oxidation of butadiene was investigated with salts of HPMo supported on pumice at reaction temperatures from 315 to 365  C.…88† The yield of furan reached a maximum at a conversion of approximately 70%, regardless of the reaction conditions and the composition of the catalysts. The cesium, sodium, and ammonium salts showed the highest catalytic activity and produced the highest yields of furan. The rate of conversion of butadiene was found to be proportional to the oxygen content and to increase with the addition of steam. The author considers that the oxidation is dependent on the activation of gaseous oxygen on basic sites. The yield of furan and of maleic anhydride are, for intermediate and higher conversions, approximately reciprocal, indicating the conversion of furan to form maleic anhydride. 10.10. 1-BUTENE As with many homogeneously catalyzed industrial processes, there is considerable impetus for the development of heterogeneous processes in the homogeneous Wacker process for the production of acetaldehyde from ethylene.…89;90† The heteropoly acids H3‡n PVn Mo12ÿn O40 with n ˆ 0, 2, 3, 8 and their Cu2‡ , Ni2‡ , and Cs‡ salts supported on silica have been tested as the redox components in the oxidation of 1-butene to butanone. The supported samples were impregnated with palladium sulfate. Previous work has shown that the oxidation of ethylene to acetaldehyde with heteropoly acid and PdSO4 proceeded at a rate signi®cantly faster than the conventional PdCl2 =CuCl2 combination.…91;92† The liquid-phase oxidation of 1-butene has also been shown to bene®t from the use of heteropoly oxometalates as redox components.…93† The reoxidation of the reduced heteropoly acids by O2 was found to be rate determining in the catalytic cycle…89ÿ93† : Pd2‡ ‡ C4 H8 ‡ H2 O ! Pd…0† ‡ C4 H8 O ‡ 2H‡ Pd…0† ‡ 2H‡ ‡ HPA ! Pd2‡ ‡ H2 HPA H2 HPA ‡ 1=2 O2 ! HPA ‡ H2 O where HPA and H2 HPA refer to heteropoly acids and their reduced forms. The steady-state activity of the catalyst is found to be dependent on the rate of reoxidation of the reduced heteropoly acid. …91;92† This rate can be increased by increasing the number of vanadium atoms per Keggin unit or by exchanging the protons of the heteropoly acids by copper or nickel. The selectivity to butanone was found to be dependent on the acidity of the protons of the heteropoly acid, becoming as high as 98% with copper or palladium as the cations.

OXIDATION PROCESSES

285

The Wacker oxidation of cyclohexane to cyclohexanone has also been studied with the palladium(II)±copper(II)±heteropoly acid combination.…94† The oxidation of 1-butene on the sodium salts of H3 PMo3 W9 O40 and H3 PMo12 O40 supported (15 wt %) on silica has been employed to provide further insight into structure sensitivity.…95† 10.11. CYCLOHEXANE Cyclohexane is a well-known source for the production of nylon-66 through its oxidation to form cyclohexanone and cyclohexanol which are converted to adipic acid. The oxidation process is usually catalyzed homogeneously by the introduction of cobalt or manganese compounds in a batch system into which oxygen is introduced. Although most of the research involving the use of heteropoly oxometalates as catalysts for the oxidation of cyclohexane has been done in a homogeneous medium, and therefore is of tangential relevance to a volume concentrating on heteropoly oxometalates as heterogeneous catalysts, nevertheless it seems worthwhile to include a brief discussion of this topic. Tetrabutylammonium salts of 12-tungstophosphoric acid into which one of TiIV , VV, CrIII, MnII , FeIII , NiII ; CuII , RuIV , or CoII has been substituted for one of the tungsten atoms have been examined as catalysts for the oxidation of cyclohexene with hydrogen peroxide.…96† The decomposition of H2 O2 , in the absence of cyclohexane, was found to occur rapidly with CoII and CuII , moderately with MnII , FeIII , and CrIII and to little or no extent with VV, NiII , and TiIV . Tetra-n-hexylammonium salts of SiM(H2 O†W11 O39 5ÿ where M is Ru or Rh were found to have higher catalytic activities than the corresponding Fe- and Cosubstituted analogues for the homogeneous oxidation of cyclohexane to cyclohexanol and cyclohexanone at 333 K with tert-butyl hydroperoxide as the oxidant.…97† The homogeneous oxidation of cyclohexane and adamantane at 20  C in CCl4 or CH3 Cl with tert-butylhydroperoxide as oxidant and with derivatives of HPW containing various transition metal ions as catalysts was shown to be relatively unaffected by the nature of the added metal ion (CoII , FeIII , and CrIII ).…98† HPW=SiO2 was employed as a heterogeneous catalyst in the photooxidation of cyclohexane at 20  C with oxygen.…99† Little or no carbon dioxide was formed. Tetra-n-butylammonium salts of HPW and of the latter with one tungsten atom substituted by a ®rst-row transition metal were employed as homogeneous catalysts for the oxidation of cyclohexane at 80  C in acetonitrile with H2 O2 as oxidant.…100† The anion with Fe substituted was found to have the highest catalytic activity.

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10.12. AMMOXIDATION Although not a purely oxidation process, ammoxidation is an important process with considerable catalytic rami®cations.…101† However, relatively little work on the application of heteropoly compounds has appeared to this date.…102† The catalytic properties of H3‡n PMo12ÿn Vn O40 (n ˆ 0, 1, 2) and the sodium salts of the acid with n ˆ 4 and 6 have been investigated for the ammoxidation of methylpyrazine. The selectivity to cyanopyrazine (at 90% conversion) on the acids increases with n to values of approximately 82%. At the highest conversions (98±99%, 400  C) the selectivities with the sodium salts are slightly smaller than those obtained with the n ˆ 2 acid. The highest yield (77.3%) of cyanopyrazine was obtained with the n ˆ 2 catalyst at 380  C. With times on stream up to 36 h, a reaction temperature of 390  C, and the HPMo catalyst, neither the activity nor the selectivity changed signi®cantly. This stability is attributed to the formation of ammonium salts.

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66. K. KuÈrzinger, G. Emig, and H. Hofmann, in: Proc. 8th Int. Congr. Catal., Vol. V, p. 500, Verlag Chemie, Frankfurt am Main (1984). 67. M. Ai, J. Catal. 98, 401 (1986). 68. V. Ernst, Y. Barbaux, and P. Courtine, Catal. Today 1, 167 (1987). 69. T. Haeberle and G. Emig, Chem. Eng. Technol. 11, 392 (1988). 70. G. B. McGarvey and J. B. Moffat, J. Catal. 128, 69 (1991). 71. G. B. McGarvey and J. B. Moffat, J. Catal. 130, 483 (1991). 72. D. Lapham, G. B. McGarvey, and J. B. Moffat, Stud. Surf. Sci. Catal. 73, 261 (1992). 73. G. B. McGarvey and J. B. Moffat, J. Catal. 132, 100 (1991). 74. J. G. High®eld and J. B. Moffat, J. Catal. 88, 177 (1984). 75. J. G. High®eld and J. B. Moffat, J. Catal. 89, 185 (1984). 76. J. B. McMonagle and J. B. Moffat, J. Colloid Interface Sci. 101, 479 (1984). 77. G. B. McGarvey and J. B. Moffat, Catal. Lett. 10, 41 (1991). 78. C. Marchal, A. Davidson, R. Thouvenot, and G. HerveÂ, J. Chem. Soc. Faraday Trans. 89, 3301 (1993). 79. A. Aboukis, D. Ghoussoub, E. Blouet-Crusson, M. Rigole, and M. Guelton, Appl. Catal. A 11, 109 (1994). 80. S. Albonetti, F. Cavani, F. Tri®ro, M. Gazzano, M. Koutyrev, F. C. Aissi, A. Aboukais, and M. Guelton, J. Catal. 146, 491 (1994). 81. T. Ilkenhaus, B. Herzog, T. Braun, and R. SchloÈgl, J. Catal. 153, 275 (1995). 82. L. Weismantel, J. StoÈckel, and G. Emig, Appl. Catal. A Gen. 137, 129 (1996). 83. K. Y. Lee, S. Oishi, H. Igarashi, and M. Misono, Catal. Today 33, 183 (1997). 84. F. Cavani, A. Tanguy, F. Tri®ro, and M. Koutyrev, J. Catal. 174, 231 (1998). 85. E. Blouet-Crusson, M. Rigole, M. Fournier, A. Aboukais, F. Daubrege, G. Hecquet, and M. Guelton, Appl. Catal. A Gen. 178, 69 (1999). 86. C. Marchal-Roch, N. Laronze, R. Villanneau, N. Guillou, A. TeÂzeÂ, and G. HerveÂ, J. Catal. 190, 173 (2000). 87. G. Centi, J. L. Nieto, C. Iapalucci, K. BruÈckman, and E. M. Serwicka, Appl. Catal. 46, 197 (1989). 88. M. Ai, J. Catal. 67, 110 (1981). 89. A. W. Stobbe-Kreemers, R. B. Dielis, M. Makkee, and J. J. F. Scholten, J. Catal. 154, 175 (1995). 90. A. W. Stobbe-Kreemers, G. van den Lans, M. Makkee, and J. J. F. Scholten, J. Catal. 154, 187 (1995). 91. K. I. Matveev, Kinet. Katal. 18, 716 (1977). 92. K. I. Matveev, E. G. Khichina, N. B. Shitova, and L. I. Kuznetsova, Kinet. Katal. 18, 320 (1977). 93. S. F. Davison, Ph.D. thesis, University of Shef®eld (1982). 94. Y. Kim, H. Kim, J. Lee, K. Sim, Y. Han, and H. Paik, Appl. Catal. A 155, 15 (1997). 95. B. N. Racine, C. M. Sorensen, and R. S. Weber, J. Catal. 142, 735 (1993). 96. N. I. Kuznetsova, L. G. Detusheva, L. I. Kuznetsova, M. A. Fedotov, and V. A. Likholobov, Kinet. Katal. 33, 516 (1992). 97. Y. Matsumoto, M. Asami, M. Hashimoto, and M. Misono, J. Mol. Catal. A 114, 161 (1996). 98. M. R. Cramarossa, L. Forti, M. A. Fedotov, L. G. Detusheva, V. A. Likholobov, L. I. Kuznetsova, G. L. Semin, F. Cavani, and F. Tri®ro, J. Mol. Catal. A 127, 85 (1997). 99. A. Molinari, R. Amadelli, L. Andreotti, and A. Maldotti, J. Chem. Soc. Dalton Trans. 1203 (1999). 100. M. M. Q. Simoes, C. M. M. Conceicao, J. A. F. Gamelas, P. M. D. N. Domingues, A. M. V. Caveleiro, J. A. S. Cavaleiro, A. J. V. Ferrer-Correia, and R. A. W. Johnstone, J. Mol. Catal. A 144, 461 (1999). 101. R. K. Grasselli, Catal. Today 49, 141 (1999). 102. V. M. Bondareva, T. V. Andrashkevich, L. G. Detusheva, and G. S. Litvak, Catal. Lett. 42, 113 (1996).

11 ENVIRONMENTALLY RELATED PROCESSES

11.1. CONVERSION OF NITROGEN OXIDES 11.1.1. The Heteropoly Acids The reduction of nitrogen oxides (NOx ) in stationary plants is frequently accomplished with NH3 as the reductant and catalysts containing titanium and vanadium. Considerable work has been done to develop catalysts that would affect the decomposition of NOx without the use of ammonia, which is itself environmentally disadvantageous. The decomposition of NOx has been investigated on heteropoly oxometalates and the information obtained has both fundamental and practical implications.…1ÿ9† Although the conversion of NOx is primarily accomplished through reduction processes, the acidic properties of the heteropoly oxometalates appear to play an important role in the conversion. To assess the importance of the nature of the central atom and the peripheral metal atoms, the sorption and conversion of NO and NO2 on HPW, HPMo, and HSiW have been studied.…1† Substantial quantities of NO2 are taken up by the tungsten-containing acids but relatively small amounts by HPMo (Fig. 11.1). While a portion of the NO2 is converted to O2, NO, and HNO3, the remainder is trapped on the solid. Only insigni®cant quantities of N2 were formed under any of the experimental conditions. The maximum quantities of NO2 sorbed at 150  C on the tungsten-containing solids show a one-to-one correspondence with the number of protons (Table 11.1). In view of the markedly higher uptake of NO2 on the tungsten-containing solids relative to HPMo and the correspondence between the NO2 sorbed and the protons present in the Bronsted acidic solids, the sorption of NO2 appears to be 289

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Figure 11.1. Losses of NO2 from the gas phase in a sequence of pulses of NO2 on HPW, HPMo, and HSiW. Pretreatment and reactor temperature of 150  C. Mass of solid in reactor, 0.075 g; pulse size, 17.0 mmol NO2.…2†

related to the acidic strengths of the materials. This appears to be in accord with the results of earlier extended HuÈckel calculations which have shown that the protons in the tungsten-containing heteropoly acids are more mobile and hence more acidic than those in the acidic solids formed from molybdenum.…10† The adsorption±desorption isotherm of NO2 on HPW at 25  C shows, in agreement with the aforementioned observations, that approximately 3 molecules of NO2 per anion are irreversibly sorbed on the solid after evacuation again corresponding to 1 molecule of NO2 per proton (Fig. 11.2).…9† Evidently NO2 is TABLE 11.1 Sorbed NO2 a…2† Temperature Catalyst

150  Cb

300  Cb

HPW HPMo HSiW

2.9 0.1 2.5

0.8 0.0 0.3

a b

Maximum molecules NO2 per anion. Pretreatment (helium, 1 h) and sorption temperatures.

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291

Figure 11.2. Adsorption±desorption isotherm of NO2 on HPW at 25  C. The sample of HPW (0.19 g or 66 mmol) was pretreated in helium at 150  C. The number of molecules of NO2 is shown per anion of the acid.…9† Reprinted from Catalysis Today, 40, Belanger and Moffat, p. 297, copyright 1998, with permission from Elsevier Science.

penetrating into the bulk of the solid between the protons and anions in order to interact with the interior protons as well as those on the surface. As noted elsewhere in this volume, similar phenomena have been observed from photoacoustic FTIR spectroscopy with other compounds, such as ammonia, pyridine, and methanol.…11ÿ14† In contrast to the observations with NO2, only a small quantity of NO is sorbed (Fig. 11.3).…2;5† However, with the increase in presorbed NO2 the quantity of NO sorbed increases, passes through a maximum, and decreases. At the maximum the quantity of NO taken up is approximately equal to that of presorbed NO2. The sorption of NO on HPW containing presorbed NO2 increases with temperature, passes through a maximum at approximately 150  C, and decreases with further increase in temperature to become insignificant at 360  C.

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Figure 11.3. Nitric oxide sorbed on HPW containing quantities of previously sorbed nitrogen dioxide. Temperature: 150  C.…5† Reprinted from the Journal of Molecular Catalysis, 114, Belanger and Moffat, p. 319, copyright 1998, with permission from Elsevier Science.

TPD pro®les of HPW previously exposed to NO2 show two major peaks, one at approximately 300  C resulting from the desorption of molecularly bound water (Fig. 11.4).…9† Another peak at 500±550  C previously attributed to the associative desorption of water resulting from the extraction of terminal oxygen atoms of the anions by protons…15† overlaps with one formed from the desorption of NO2 and=or HNO3. The lower temperature water peak diminishes in intensity as a result of the reaction of NO2 with water to form nitric acid which desorbs from the solid. The processes involved in the sorption of NO2 are apparently of the form H‡ ‡ NO2 ! HNO2 ‡ 2NO2 ‡ H2 O ! HNO3 ‡ HNO2 3NO2 ‡ H2 O ! 2HNO3 ‡ NO After exposure of HPW to NO2 the structure of the anion remains unchanged as evidenced by the set of ®ve or six bands in the 1200±800 cmÿ1 region characteristic of the Keggin structure (Fig. 11.5).…5† As noted elsewhere in this volume, bands at approximately 1080 and 980 cmÿ1 are attributed to the triply degenerate asymmetric stretch of PO4 and the stretching vibration of the tungsten-terminal oxygen bond.…11† A band at 2264 cmÿ1 characteristic of the nitronium ion appears after exposure of HPW to NO2.…16† Bands at 1708 and 3200 cmÿ1 due to the hydronium ion (H3 O‡ ) and water, respectively, decrease in intensity after exposure of the acid to NO2 (Fig. 11.5).…5†

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293

Figure 11.4. Temperature-programmed desorption of HPW (A) pretreated in helium at 150  C for 1 h, (B) then exposed at 150  C to one pulse of NO2, (C) two pulses of NO2, (E) seven pulses of NO2. Helium ¯ow rate 45.0 ml minÿ1 , ramp rate 60  C minÿ1 , mass of solid in reactor 0.165 g, and pulse size 17.0 mmol NO2. Legend: 1, H2 O; 2, NO2.…2†

Although as noted earlier NO is not signi®cantly sorbed by HPW at 150  C, after presorption of NO2 at 150  C, the solid acid sorbs NO reaching a maximum at one NO molecule taken up per molecule of NO2 presorbed. IR spectra show that a new band appears at 1304 cmÿ1 attributed to N2 O3 (Fig. 11.5C)…16† resulting from NO ‡ NO2 ! N2 O3 1

H MAS NMR spectra of HPW show that the chemical shift of the hydrogen peak has decreased signi®cantly on exposure to HPW while the intensity has increased (Fig. 11.6).

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Figure 11.5. IR spectra of HPW before (A), after exposure to NO2 (B), and after exposure to NO2 followed by NO (C). Peaks due to the presence of nujol are marked with an asterisk.…5† Reprinted from the Journal of Molecular Catalysis, 114, Belanger and Moffat, p. 319, copyright 1998, with permission from Elsevier Science.

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295

Figure 11.6. 1 H MAS NMR of HPW, before (A) and after (B) exposure to NO2.…5† Reprinted from the Journal of Molecular Catalysis, 114, Belanger and Moffat, p. 319, copyright 1998, with permission from Elsevier Science.

The band at 2264 cmÿ1 tentatively attributed to the nitronium ion (NO2 ‡ )…16† may be rationalized through the disproportionation of NO2 to form NO2 ‡ and NO2 ÿ…16;17† : 2NO2 ! NO2 ‡ ‡ NOÿ 2 The nitrite ion may form HNO2 from its interaction with H‡ while NO2 ‡ substitutes for the proton to provide a charge balance. Alternatively the band at 2264 cmÿ1 may result from the association of the protons with NO2 to form HNO2 ‡ : H‡ ‡ NO2 ! HNO2 ‡ The shift to 2264 cmÿ1 from the reported asymmetric stretch for NO2 of 2375 cmÿ1 …16† may result from perturbation of the electron density of NO2 by the proton. The n3 (antisymmetric) band in NO2 ‡ of inorganic salts has been found in the range 2360±2392 cmÿ1 .…18† The one-to-one association of the proton

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Figure 11.7. 15 N{1 H} MAS NMR of HPW previously exposed to 15 NO2 at 150  C.…9† Reprinted from Catalysis Today, 40, Belanger and Moffat, p. 297, copyright 1998, with permission from Elsevier Science.

and NO2 provides further evidence for HNO2 ‡ and the formation of a nitronium salt. Con®rmation of the association of one NO2 with each proton in a monomeric, nonlinear, and charged form is found from proton-decoupling NMR spectra with HPW exposed to 15 NO2 which shows one peak at 4.14 ppm (Fig. 11.7).…9† Heteroatom-proton coupling spectra also show a single peak, indicating that the proton is associated with the oxygen atom of NO2 (Fig. 11.8).…9† 11.1.2. Ammonium 12-Tungstophosphate Ammonium 12-tungstophosphate (abbreviated as NH4 PW) may be prepared by precipitation from an aqueous solution of the parent acid to which an aqueous solution of an ammonium salt has been added. A high-surface-area (150± 200 m2 gÿ1 ) microporous solid is formed.

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297

Figure 11.8. 15 N MAS NMR of HPW previously exposed to 15 NO2 at 150  C.…9† Reprinted from Catalysis Today, 40, Belanger and Moffat, p. 297, copyright 1998, with permission from Elsevier Science.

Exposure of NH4 PW to NO2 results in the formation of N2, NO, and N2 O, in proportions which are temperature dependent (Fig. 11.9).…4ÿ9† As the temperature increases, the selectivity to N2 increases while that to NO decreases. At a temperature between 150 and 200  C, NO2 is converted predominantly to N2 . Although NH4 PW produces N2 in contrast to the observations with its parent acid, HPW, the possibility that NH3 may be released into the atmosphere from NH4 PW at suitable NO2 reduction temperatures would diminish its practical advantages. However, TPD experiments performed on NH4 PW show that temperatures as high as 400  C are required for the release of ammonia and only small quantities are evolved up to 500  C (Fig. 11.10A,D). The peak of desorbed water at 175  C increases substantially after saturation of NH4 PW with NO and NO2 at 30  C (Fig. 11.10B,C), particularly after NO2 is added, showing that water forms during the reduction process. After pretreatment of NH4 PW and exposure to NO, both at 150  C, the water peak at 175  C in the TPD pro®le has

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Figure 11.9. Products from NO2 injected onto NH4 PW.…7† Reprinted from Applied Catalysis, B13, Belanger and Moffat, p. 167, copyright 1997, with permission from Elsevier Science.

increased and several overlapping peaks are evident at temperatures from 300 to 500  C (Fig. 11.10E). In contrast, after pretreatment of NH4 PW and saturation with NO, both at 150  C, a large peak with a maximum at approximately 475  C is observed, resulting from the desorption of NO2 (Fig. 11.10F). After saturation with NO2 the latter peak is again present but the intensity of the water peak at 200  C is increased. The NO2 that desorbs after saturation with NO probably results from the extraction of anionic oxygen atoms by the NO. No evidence was found for the retention of N2 at any temperature. It is clear from the TPD pro®les that NH3 is not signi®cantly desorbed at temperatures of 400  C or less. Thus, at the optimum reduction temperatures of 150±200  C, NO2 and NO are reduced by NH3 held on the solid as NH4 ‡ and little or no NH3 is released into the gas phase as a contaminant. The IR spectra of NH4 PW, in addition to the characteristic IR bands for the Keggin structure in the 800±1200 cmÿ1 region, contain bands at 3200 and 1421 cmÿ1 attributed to the ammonium ion (Fig. 11.11A). After exposure to NO2 a small band appears at 2263 cmÿ1 and intensi®es with the sample pretreated and exposed to NO2 at 300  C. On saturation of HPW pretreated at 150  C with NO2 at 150  C the 1 H chemical shift increases with little or no diminution in the intensity of the protium peaks (Fig. 11.12A, 1 and 2), showing that protons are not lost. In contrast, the 1 H MAS NMR spectra for NH4 PW contain four peaks, the largest at 4.85 ppm, but after saturation with NO2 at 300  C only the peak at 7.89 ppm remains, indicating that the reduction of NO2 on NH4 PW restores the parent acid, HPW.

ENVIRONMENTALLY RELATED PROCESSES

299

Figure 11.10. TPD of NH4 PW (A) pretreated in helium at 30  C for 1 h; then (B) saturated with NO at 30  C for 1 h; (C) same pretreatment as in (A) but saturated with NO2 at 30  C for 1 h; (D) NH4 PW pretreated at 150  C for 1 h; then (E) saturated with NO at 150  C for 1 h; (F) NH4 PW pretreated at 300  C for 1 h, then saturated with NO at 300  C for 1 h; (G) same pretreatment as in (D) but saturated with NO2 at 150  C for 1 h. Mass of solid in reactor 0.075 g (38.8 mmol). Pulse size 17 mmol NO2. Helium ¯ow rate 45 ml minÿ1 . Temperature rate 60  C minÿ1 . Legend: 1, H2 O; 2, NH3 ; 3, NO2 ; 4, N2 O.…4† Reprinted with permission from Belanger and Moffat.…4† Copyright 1996 American Chemical Society.

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Figure 11.11. IR spectra of (A) NH4 PW at 25  C, then (B) pretreated and exposed to 20 pulses of NO2 at 150  C. Same salt as in (A) but (C) pretreated and exposed to 14 pulses of NO2 at 300  C. * ˆ nujol bands.…9† Reprinted from Catalysis Today, 40, Belanger and Moffat, p. 297, copyright 1998, with permission from Elsevier Science.

Although NO is not reduced at 150  C on NH4 PW, after pretreatment of the solid with either NO2 or HNO3, the reduction of NO occurs. NO2 has been shown to enter the bulk structure of both HPW and NH4 PW. In the latter case its microporous structure will augment this process. The protons in both HPW and NH4 PW play active roles, in the ®rst case to interact directly with the NO2 and in the second to bind NH3 within the solid so that the reduction of NO2 with NH4 PW occurs on and in the solid with little or no participation of a gas-phase reaction. Other workers have studied the conversion of nitrogen oxides on heteropoly oxometalates.…19ÿ22† It is claimed that NO is absorbed on HPW at 150  C and, on heating at 450  C, 68.3% of the absorbed NO is converted to N2 .…19;20† However, the authors note that O2 and H2 O are needed for NO absorption. Because NO and O2 would be expected to produce NO2, it appears that, as with the previously noted work, NO apparently does not sorb. However, the conversion of NO2 to N2 on HPW seems doubtful.…19ÿ22†

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301

Figure 11.12. 1 H MAS NMR of (A): (1) HPW pretreated at 150  C; (2) HPW pretreated and saturated with NO2 at 150  C; (B): (1) NH4 PW pretreated at 300  C; (2) NH4 PW pretreated and saturated with NO2 at 300  C…9† ; and (C): (1) NH4 PW (ÿ 15%) pretreated at 300  C, (2) NH4 PW (‡ 15%) pretreated at 300  C. Reprinted from Catalysis Today, 40, Belanger and Moffat, p. 297, copyright 1998, with permission from Elsevier Science.

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11.2. HYDRODESULFURIZATION (HDS) The removal of sulfur from petroleum-based fuels is of increasing importance as a result of recently imposed government regulations regarding sulfur content. Although cobalt (nickel)±molybdenum catalysts are well known for HDS, the role of heteropoly oxometalates is attracting interest.…23ÿ27† One of the earliest reports compared catalysts prepared by the impregnation of alumina with an aqueous solution of the ammonium salt of nickel heteropolymolybdate, …NH4 †4 NiMo6 O24
xH2 O, with other Ni=Mo samples.…23† The authors concluded that the existence of a bond between nickel and molybdenum was responsible for an increase in the catalytic activity in thiophene (as a test material) HDS in comparison with a catalyst not containing nickel. A higher activity per gram of molybdenum was found with samples prepared by impregnation with the nickel oxometalate or with catalysts prepared under conditions expected to be suitable for the formation of a heteropoly compound. Because the authors have not indicated the source of the nickel heteropolyoxometalate, its composition and structure before use and after pretreatment at 723 K for 3 h and use in the thiophene conversion process at 673 K are unknown. Maitra and co-workers have pointed out the dif®culties arising in attempting to prepare HDS catalysts by impregnating porous supports such as alumina.…24† They note that because anionic base elements require acidic conditions for adsorption on alumina whereas impregnation of cobalt or nickel salts is best done in basic solution, two separate impregnation steps, which may not be conducive to uniformity of deposition, may be necessary. These authors employed the anions …H6 XY6 O24 †xÿ with XˆCo or Ni and YˆMo or W. …NH4 †4 …NiII W6 O24 H6 †
5H2 O, and …NH4 †4 …NH4 †4 …NiII Mo6 O24 H6 †, II …Ni Mo3 W3 O24 H6 † were prepared by a method reported in the literature…25† as was that for …NH4 †3 …CoIII Mo6 O24 H6 †
7H2 O…26† and …NH4 †8 …CoII2 W11 O40 H2 †
13H2 O.…27;28† The parent free acids were prepared by cation exchange.…29† The authors conclude that heteropoly anions containing promoter and base metals in a single complex are appropriate for use in the impregnation of porous supports and result in catalysts that are effective for HDS and hydrogenation but not hydrodenitrogenation. Titania-supported HPMo and its cobalt and nickel salts have been tested for the HDS of thiophene.…30ÿ32† Although the source of HPMo was not indicated, the salts were prepared by use of the method of Tsigdinos.…33† IR and X-ray photoelectron spectroscopy were employed to characterize the samples before and after reduction. The highest initial conversion of thiophene was obtained with HPMo supported on titania although the deactivation of this catalyst occurred more rapidly than that of those containing Co and Ni. HPMo was employed as a dispersed catalyst for the HDS at 340  C of a Cold Lake crude straight-run diesel fraction with boiling range from 161 to 343  C.…34†

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303

The catalyst was converted into a mixture of oxides and sul®des during the reaction. Two cobalt salts of HPMo supported on g-alumina have been tested for the HDS of thiophene,…35† nonreduced Co3=2 PMo12 O40 , and reduced Co7=2 PMo12 O40 . The nonreduced compound was decomposed on impregnation on the alumina but produces a well-dispersed surface polymolybdate phase. The authors contend that the use of the reduced heteropoly salts for the preparation of the HDS oxidic precursors enhances the interaction between the promoter atom and the oxomolybdate in the impregnating solution as well as avoiding the formation of the surface polymolybdate phase obtained with the nonreduced compound. 11.3. POLYMERIZATION Heteropoly acids have been explored for use as initiators in polymerization processes. As early as 1935 HSiW was tested as a catalyst for the polymerization of propylene to low-molecular-weight oligomers and found to be considerably more active for this process than phosphoric acid on kieselguhr.…36† Bauxite and low-surface-area aluminas were the preferred supports whereas silica gel resulted in poor activity. A separate study of the same process found that the most active catalysts were obtained with silica±alumina, natural clays, and alumina, in decreasing order.…37† Benzyl alcohol and its 4-chloro-, 4-methyl-, and 2,4,6trimethyl derivatives formed polyaromatics with HPMo and HSiMo as catalysts.…38† No polymerization was obtained with triphenylmethanol and 1and 2-phenylethanols. More recent reports have focused on the heteropoly acids as initiators of cationic polymerization.…39† These authors note that superacids when employed in polymerization often suffer from the reactivity of the anions which is less likely to occur with heteropoly acids. Further, other acidic catalysts may be dif®cult to obtain in pure form and consequently must be prepared in situ, a disadvantage not found with heteropoly acids. The latter are nonvolatile and soluble in many organic liquids, of obvious importance in polymerization processes. HPMo was shown to be an effective initiator in the cationic polymerization of 1,3-dioxolane, 1,3-dioxepane, and 1,3,5-trioxane with each of the three protons initiating one chain. However, oxidation±reduction processes of the Keggin anions were observed. Aqueous formaldehyde is trimerized in the liquid phase to produce trioxane, an intermediate to polyoxymethanes.…40† In the vapor phase, with 1-vanado-11molybdophosphoric acid supported on silicon carbide at 102.5±110  C and 0.7± 1.15 bars the reaction was found to increase stepwise with increase in the partial pressure of formaldehyde which the authors attribute to the diffusion of formaldehyde into the bulk of the catalyst. This appears to be analogous to the

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observations with photoacoustic FTIR spectroscopy showing that, for example, gaseous pyridine and various alcohols penetrate into the bulk structures of the heteropoly acids at rates considerably less than those of smaller molecules, such as ammonia.…11ÿ14;41† HPW has been investigated as an initiator in the polymerization of tetrahydrofuran with ethylene oxide as the promoter.…42† Considerably lower concentrations of HPW than those reported in earlier work…39† were successfully employed where the molecular weight of the products was controlled by the concentration of water or glycol.

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INDEX

Acid-catalyzed processes alkylation processes benzene, 214 isobutene and 2-butene, 212 isobutene and methanol, 209 conversion of butane, 194 of butene, 196 of butanol, 208 of C5±C8 alkanes, 216 of C6±C8 alkenes, 217 of ethanol, 189 of isobutane, 201 of isobutene, 206 of methanol, 175 of propene, 192 of propanol, 193 Friedel±Crafts reactions, 215 ring contraction of cyclohexane, 218 ring expansion of methylcyclopentane, 218 Acidity adsorption±desorption, 150 calorimetry, 144 electrical conductivity, 143 probe reactions, 160 solid±liquid titrations, 147 spectroscopy, 150 temperature-programmed desorption, 156 theoretical, 147 Ammoxidation, 286 Characterization electrochemical methods, 24

elemental analysis, 25 EDTA titration, 26 HPLC, 25 ion chromatography, 25 other techniques electronic, 23 electron paramagnetic resonance, 23 infrared, 13 nuclear magnetic resonance, 18 photoacoustic FTIR, 14 scanning tunneling, 24 spectroscopic methods, 13 tunneling, 24 x-ray photoelectron, 23 History, 1 Hydrodesulfurization, 302 Microporosity salts of monovalent cations alkali metals, 98 ammonium, 98 ion exchange, 121 group 11 and 13 elements, 110 sorption and diffusion, 119 stoichiometry, 105, 111 surface areas and pore structures, 98, 116 Nitrogen oxides conversion of, 289 Oxidation of butadiene, 284 307

308 of of of of of of of of

butane, 260 butene, 284 cyclohexane, 285 ethane, 252 isobutane, 263 isobutyric acid, 268 methacrolein, 267 methane with N2O and O2, 227 effect of gas phase additives, 248 of pentane, 283 of propane, 259

Polymerization, 303 Shape selectivity, 139 Stability thermal, 41 pH molybdate/phosphate system, 62 regeneration, 67 tungstate/phosphate system, 62 Structure anions, 29 cations, 34

INDEX crystallographic, 34 ®lms Langmuir, 37 Supported catalysts supports alumina, 83 carbon, 71 clays, 88 formation of heteropoly acids from oxides supported on silica, 91 heteropoly salts, 88 MCM-41, 87 MgF2, 84 polymers, 90 SiC, 86 silica, 74 titanium dioxide, 73 zeolite Y, 87 ZrO2, 86 Synthesis of 12-heteropoly acids mixed addenda, 8 molybdenum-containing anions, 7 other heteropoly acids, 7 tungsten-containing anions, 6