Advances in Catalysis, Volume 41

ADVANCES IN CATALYSIS VOLUME 41 Advisory Board G . ERTL V. B. KAZANSKY BerlidDahlem, Germany J. M.THOMAS Evanston...

Author: D. D. Eley | Werner O. Haag | Bruce C. Gates

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ADVANCES IN CATALYSIS VOLUME 41

Advisory Board

G . ERTL

V. B. KAZANSKY

BerlidDahlem, Germany

J. M.THOMAS Evanston, Illinois

W. M. H. SACHTLER

Moscow, Russia

Evanston. Illinois

I? B. WEISZ State College, Pennsylvania

ADVANCES IN CATALYSIS VOLUME 41

Edited by D. D. ELEY

WERNER0. HAAG

The University Nottinghcrm. EnKland

Mohil Research and Development Corporation Princeron. New Jersey

ACADEMIC PRESS San Diego New York Boston London Sydney Tokyo Toronto

BRUCEGATES Uniwrsity of Cdforiiia Dcivis, CaliJiJrniu

Find Us on the Web! http://www.apnet.com

This book is printed on acid-free paper.

@

Copyright Q 1996 by ACADEMIC PRESS, INC All Rights Reserved. No part of this publication may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopy, recording, or any infonnation storage and retrieval system, without permission in writing from the publisher.

Academic Press, Inc. A Division of Harcourt Brace & Company 525 B Street, Suite 1900, San Diego, California 92101-4495 United Kingdom Edition published by Academic Press Limited 24-28 Oval Road, London NW I 7DX International Standard Serial Number: 0360-0564 International Standard Book Number: 0- 12-007841-4 PRINTED IN THE UNITED STATES OF AMERICA 96 97 9 8 9 9 00 01 Q W 9 8 7 6 5

4

3 2

I

CONTRIBLITORS ............................................................................................................................ PREFACE ..................... ....

ix xi

Vibrational Spectra of Hydrocarbons Adsorbed on Metals Part 1. Introductory Principles, Ethylene, and the Higher Acyclic Alkenes NORMANSHEPPARD AND CARLOS DE LA CRUZ

I. 11.

111.

I v. V.

VI. VII.

1 Introduction ..................................................................................... Experimental Considerations Relating to the Different Vibrational 3 Spectroscopic Techniques Available ............................................... Experimental Considerations Relating to Catalyst Preparation 7 or Sample-Handling Procedures ..................................................... Considerations Relating to the Interpretation of Vibrational 12 Spectra of Adsorbed Species .......................................................... Other Experimental Methods for Investigating the Structures of 26 Chemisorbed Hydrocarbons on Metals ........................................... Experimental Vibrational Spectroscopic Results on OxideSupported Metal Catalysts Classified by Alkene Adsorbate ........... 30 103 Conclusions ..................................................................................... 105 References .......................................................................................

Catalytic Chemistry of Heteropoly Compounds TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO I. 11. 111.

IV. V. V1.

Introduction ..................................................................................... Structure, Synthesis, Stability, and Characterization ...................... Acidic Properties ............................................................................. Acid-Catalyzed Reactions in the Liquid Phase ............................... Heterogeneous Acid-Catalyzed Reactions ...................................... Pseudoliquid Phase ......................................................................... Y

I13 1 18 139 150 I6 1 178

vi

CONTENTS

VII . VIII. IX . X. XI . XI1 .

XI11.

Redox Properties ............................................................................. Liquid-Phase Oxidation Reactions ................................................. Oxidation Catalyzed by Solid Heteropoly Compounds .................. Fine Chemicals Synthesis ............................................................... Hybrid Catalysts ............................................................................. Photocatalysis and Electrocatalysis ................................................ Conclusions ..................................................................................... References .......................................................................................

191 200 210 221 223 233 240 240

Microporous Crystalline litanium Silicates BRUNONOTARI

I. I1. 111.

I v. V. VI . VII . VIII .

Introduction ..................................................................................... Mixed Oxides of Ti and Si .............................................................. Titanium Silicates ........................................................................... Synthesis ......................................................................................... Catalytic Reactions ......................................................................... Catalytic Sites ................................................................................. Reaction Mechanism ....................................................................... Summary ......................................................................................... References .......................................................................................

253 257 267 288 293 317 318 326 327

Structural and Mechanistic Aspects of the Dehydration of Isomeric Butyl Alcohols over Porous Aluminosilicate Acid Catalysts KIRILL ILYCH ZAMARAEV AND JOHNMEURIC THOMAS I.

I1. 111.

IV. V. VI .

Introduction ..................................................................................... Characterization of Catalysts .......................................................... Kinetic Studies ................................................................................ Pathways of Butyl Alcohol Dehydration ........................................ The Nature of the +AI-O(R)-Si$ Reaction Intermediate ........... Conclusions ..................................................................................... References .......................................................................................

335 337 339 344 349 354 357

CONTENTS

vii

Thermal and Catalytic Etching Mechanisms of Metal Catalyst Reconstruction TA-CHINWE1 AND JONATHAN PHILLIPS

I.

I1. I11.

Introduction ..................................................................................... Thermal Etching ............................................................................. Catalytic Etching ............................................................................. References .......................................................................................

INDEX ............................................................................................................

359 362 383 415 423

This Page Intentionally Left Blank

Contributors Numbers in parentheses indiciite the pages on which the authors' contributions begin.

CARLOS DE LA CRUZ, Departmento de Quimica, Facultad de Ciencias, La Universidad del Zulia, Maracaibo, Venezuela ( 1 ) MAKOTOMISONO,Department of Applied Chemistry, The University of Tokyo, Tokyo 113, Japan ( 1 13) NORITAKA MIZUNO, Institute qf Industrial Science, The University of Tokyo, Roppongi. Tokyo 106, Japan ( 1 13) BRUNONOTARI, Notari Tecnologie, SnC, 20097 Sun Donato Milanese, Italy (253) TOSHIO OKUHARA, Division of Materials Science, Graduate School of Environmental Earth Science, Hokkaido University, Sapporo 060, Japan (1 13) JONATHAN PHILLIPS, Department of Chemical Engineering, The Pennsylvania State University, University Park, Pennsylvania 16802 (359) NORMANSHEPPARD, School of Chemical Sciences, University of East Anglia. Norwich NR4 7TJ, England ( 1 ) JOHN MEURIG THOMAS, Davy Faraday Research Laboratory, Royal Institution of Great Britain, London WIX 4BS, United Kingdom (335) TA-CHINWEI, Department nf Chemical Engineering, The Pennsylvania State University, University Park, Pennsylvania 16802 (359) KIRILLILYCH ZAMARAEV, Boreskov Institute of Catalysis, Novosibirsk 630090, Russia (335)

ix


With this volume of Advances in Carulysis, Herman Pines retires from the post of co-editor, which he has held since Volume 14. During this time, Herman’s popularity in the catalysis community, together with his perspicacity, has attracted many excellent contributions to Advances in Cutalysis. His own interests have been especially in the fields of organic chemicals and petroleum chemistry. It was in these areas that he achieved early and lasting fame for his work with Ipatieff, which began in 1932. Their catalytic process for producing “alkylate” for 100-octane aviation fuel played a major role in World War I1 (cf. Advances in Curulysis, Vol. I , pp. 27-28). We look forward to his continuing help as a member of the Advisory Board. At the same time we welcome our new co-editor, Bruce Gates, who will introduce the present volume. D. D. ELEY As this volume of Advances in Catalysis was going to press we received word that Herman Pines had passed away on April 10 at the age of 94. He leaves us with the legacy of his scientific work and the memory of an extraordinary man with a curious mind, boundless energy, independence of thought, human compassion, and a fine sense of humor. His early collaboration with Ipatieff began at Universal Oil Products and continued at Northwestern University where, from 1941 to 1952, he held a concurrent position as associate professor. In 1952 he became the Ipatieff Research Professor of Chemistry and director of the Ipatieff High Pressure and Catalytic Laboratory. The paraffin alkylation process mentioned above has become a major refinery process and is still practiced today on a very large scale, as is the process for the isomerization of paraffins which originated in Herman Pines’ laboratory. It was these discoveries and many others that led to his 145 patents and 265 publications, all of which were based on skillful experimental work. But at heart, Herman Pines was a theorist. As a professor he was an inspiring teacher and educator who set and demanded the highest standards. His research always aimed at new concepts and fundamental understandings. The initial work on acid-catalyzed reactions laid to rest the notion that paraffins are what their name implies, i.e., “parum affinis,” essentially inert. Imaginative studies on thermal, free-radical reactions-which led to a new manufacturing method in the perfume industry, among others-was followed by an in depth investigation of xi

xii

PREFACE

base-catalyzed hydrocarbon conversions. Most of what is known today about this subject is documented in his book, co-authored with Wayne Stalick, “Base-Catalyzed Reactions of Hydrocarbons and Related Compounds,” published by Academic Press in 1977. Few of us know that a key hydrocarbon intermediate in the manufacture of the widely used pain remedy lbuprofen is made in a basecatalyzed reaction first described by Herman Pines. The scope and depth of his work is reflected by the numerous awards he received, among them the Eugene H. Houdry Award in Applied Catalysis, the Chemical Pioneer Award, and the following American Chemical Society awards: The Fritzsche Award for his contributions to terpene chemistry, the Petroleum Chemistry Award, and the E. V. Murphree Award in Industrial and Engineering Chemistry. His colleagues, former students, and friends will miss him. W. 0. HAAG

The first chapter of this volume, by Sheppard and de la Cruz, addresses the application of vibrational spectroscopy for the characterization of adsorbed hydrocarbons. This chapter is a successor to the 1958 Advances in Catalysis chapter about infrared spectra of adsorbed species, authored by the pioneers Eischens and Pliskin. Vibrational spectroscopy continues to provide some of the most incisive techniques available for determination of adsorbate structures. The present chapter is concerned with introductory principles and spectra of adsorbed alkenes; a sequel is scheduled to appear in a subsequent volume of Advances in Catalysis. The heart of this volume consists of three chapters summarizing work on catalysts that are both industrially applied and structurally well-defined; the structural definition has allowed rapid progress in the development of relationships between structure and catalytic properties. Okuhara, Mizuno, and Misono report the catalytic properties of heteropoly compounds as exemplified by H,PW120,, and the anion [PW,,O,,,]’-. Some of these compounds are strongly acidic, and some have redox properties; the largescale applications involve acid-catalyzed reactions. The heteropoly compounds are metal oxide clusters, used as both soluble and solid catalysts. Their molecular character provides excellent opportunities for incisive structural characterization and for tailoring of the catalytic properties. Physical properties also affect catalytic performance. Catalysis sometimes occurs on the surface of the solid material, and sometimes it occurs in the swellable bulk. Notari summarizes the science and technology of catalysis by molecular sieves incorporating framework metal ions. The most important example in technology is Ti silicalite, a selective oxidation catalyst. Because the catalysts are crystalline materials, their structures are among the most well-known of any industrial catalysts, and catalytic sites such as Ti cations are identified. The work summarized in this chapter, almost all of it performed in just the preceding few years, gives an

PREFACE

...

Xlll

indication of how important molecular sieves may become as catalysts for reactions beyond those catalyzed by acids. Zamaraev and Thomas provide a concise summary of work done with a family of classic catalytic test reactions-dehydration of butyl alcohols-to probe the workings of acidic molecular sieve catalysts. This chapter echoes some of the themes stated by Pines and Manassen, who wrote about alcohol dehydration reactions catalyzed by solid acids in the 1966 volume of Advances in Catalysis. In the final chapter, Wei and Phillips tie together old and new results characterizing the processes of surface etching. They summarize evidence that chemical etching takes place by reactions of gas-phase free radicals. This subject pertains to catalyst redispersion and regeneration, and the chapter links the catalysis literature and literature less often consulted by catalytic scientists and engineers.

B.C.GATES


ADVANCES rh CATALYSIS. VOLUME 41

Vibrational Spectra of Hydrocarbons Adsorbed on Metals Part I. Introductory Principles, Ethylene, and the Higher Acyclic Alkenes NORMAN SHEPPARD School of Chemicul Sciences University (?f East Anglia Nonvich NR4 7TJ, England AND

CARLOS DE LA CRUZ Departmento de Quimicu Fucultad de Ciencias La Universidud del Zuliu Murucuiho. Venezuela

1.

Introduction

In 1897, Sabatier and Senderens ( I ) made a pioneering study of the use of a nickel as a catalyst for the hydrogenation of ethylene (ethene) to ethane. This investigation led to the award of the Nobel Prize to Sabatier in 1912. Since that time the importance of heterogeneous catalysts has continued to increase greatly, decade by decade, extending the boundaries of laboratory chemical researches and promoting new and more cost-effective processes within the chemical industry (2). The correct choice of a catalyst allows a desired reaction to proceed under milder conditions of temperature and pressure than would be Abbreviations: DRIFT(S)4iffuse-reflection infrared (Fourier-transform) (spectroscopy): EXAFS-xtended X-ray absorption fine structure; FTIR(S)-Fourier-transform infrared (spectroscopy); INS---inelastic neutron scattering; LEED-low-energy electron diffraction; MCT-mercury/ cadmium telluride photoconductive detector; MSSR-metal-surface selection rule; NEXAFSnear-edge X-ray absorption fine structure; NMR-nuclear magnetic resonance; PED-photoelectron infrared (spectroscopy): SERS-surface-enhanced diffraction; RAIR(S)-reflection-absorption Raman spectroscopy; STM-scanning tunneling microscopy; TPD-temperature-programmed desorption; UPS-ultraviolet photoelectron spectroscopy; VEEL(Stvibrationa1 electron energyloss (spectroscopy); XPS-X-ray photoelectron spectroscopy.

I Copyright U’ 1996 by Academic Press. Inc. All rights of reproduction in any form reserved.

2

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

the case otherwise. Often more important, catalysts make possible selectivity toward a particular desired reaction when a set of starting materials can give different sets of products. Until the 1950s the nature of the chemical surface species involved in the interactions between the surface of a catalyst and the reagents or products of a reaction-the reaction intermediates-could be discussed only as possibilities based on studies of the kinetics of the overall reaction, aided by results from isotopic variants among reactants and products ( 3 , 4 ) . Vibrational spectroscopy in the infrared region provided the first direct experimental evidence on the structures of adsorbed species themselves and has subsequently proved an excellent method for systematic investigations in both chemisorption and catalysis. Even today, vibrational spectroscopy in its various forms remains the most widely used physical method for identifying molecular surface species. It has the unique advantage that it can as readily be used to study adsorption on finely divided catalysts as on single-crystal surfaces used as simplified models. The single-crystal systems with surfaces of known atomic arrangements have the additional advantage that they can be studied by the numerous other spectroscopic and diffraction experimental techniques available in surface science. In 1954 R. P. Eischens, W. A. Pliskin, and S. A. Francis ( 5 ) of the Texaco Research Center in New York published the first infrared spectra of chemisorbed species, namely of carbon monoxide adsorbed on the silica-supported finely divided metal catalysts of Ni, Pd, Pt, and Cu. Also, in 1956, Pliskin and Eischens (6) were the first to obtain spectra of the hydrocarbons ethylene (ethene), acetylene (ethyne), and propene adsorbed on an oxide-supported metal catalyst, Ni/Si02. Eischens and his colleagues followed this up with further studies of chemisorbed n-alkenes and their surface-hydrogenation products on Ni/SiOz (7). Since these early days, many infrared studies of hydrocarbons adsorbed on oxide-supported metal catalysts have been carried out. Although much success has ultimately been achieved in identifying the chemisorbed species present, progress has tended to be slow for a number of reasons: 1. An observed spectrum is typically found to be derived from the presence

of several different adsorbed species. These arise from a plurality of different adsorption sites that can occur on particles of a given metal. Even a well crystallized particle could, for example, present sites associated with different types of facets, e.g., ( 1 1 l), (IOO), or ( I lo), and their various arrays of metal atoms (see Fig. 1 in Section IV.A.1). 2. Even when the spectrum of an adsorbed species proves to be that of only one of those anticipated, difficulties of recognition can arise because of the perturbing effects of the bonded metal atoms on the well known characteristic frequencies (wavenumbers) of hydrocarbon groups (8-10).

VIBRATIONAL SPECTRA OF HYDROCARBONS

3

3. Particular identification problems arise when unanticipated surface species are present. A case in point is the ethylidyne surface species, M3(CCH3) (M = metal atom), which has turned out to be of common occurrence from ethene adsorbed on metals near room temperature (11). About two decades after Eischens’ application of transmission infrared spectroscopy to the study of adsorption on high-area, finely divided metal catalysts, a new vibrational spectroscopic technique, termed electron energy loss spectroscopy (EELS), was developed. This technique proved to have the sensitivity required to obtain spectra from literally single monolayers of adsorbed species such as those occurring on the flat, low-area surfaces of single metal crystals. It was first applied to the study of adsorbed hydrocarbons in 1977 by H. Ibach and his colleagues, who studied ethene on Pt( 1 I 1 ) (12), and by J. C. Bertolini and colleagues, who investigated benzene on Ni( 1 1 1) and Ni(100) (13). The technique is variously described as high-resolution or vibrational electron-energyloss spectroscopy, abbreviated as HREELS or VEELS, respectively. We shall use the latter abbreviation because the EELS method is of low resolution relative to its competitor, infrared spectroscopy. The original adjective “high” relates to comparison with electronic EELS. Hitherto, in the form of reflection-absorption infrared spectroscopy (RAIRS), the infrared method had been capable of detecting single monolayers only in the exceptionally favorable (strong absorption) cases of carboxylate ions [Francis and Ellison (14)] or carbon monoxide [Chesters, Pritchard, and Sims ( 1 5 ) ] adsorbed on flat metal surfaces. The new challenge from VEELS provided the motivation for a search for improvements in RAIRS sensitivity, and this was very successfully achieved by M. A. Chesters and his colleagues through the introduction of Fourier-transform-based interferometric infrared spectroscopy (16). During the past two decades, the results obtained from hydrocarbons adsorbed on the simplified single-crystal metal surfaces by the use of VEELS or RAIRS have proved to be of crucial assistance in the interpretation of the more complex spectra obtained from finely divided metals. The VEELYRAIRS results have recently been reviewed by one of us (17); it is therefore timely to do likewise for the many results on finely divided metals now available in the literature. This is particularly so because there is a general need to revise many of the earlier interpretations owing to several improvements in our understanding of such spectra.

Ii. Experimental Considerations Relating to the Different Vibrational Spectroscopic Techniques Available A.

TRANSMISSION INFRAREDSPECTROSCOPY (18)

Infrared spectroscopy is capable of giving spectra of very good signahoise from adsorbed species on finely divided metal catalysts. Samples are usually

4


studied in the form of pressed disks which remain porous to gases. It is also capable of high resolution (1

High

Powdered metaVoxides Yes

Dipole change perpendicular to surface

Limited as yet

RAIRS

4000-800 (-200)

ca. 1

Low

Flat, preferably single-crystal, metal surfaces

Yes

Dipole change perpendicular to surface

Moderate

VEELS

40W200

High

Flat, preferably single-crystal, metal surfaces

No'

Dipole change perpendicular to surface (specular reflection)d. All modes'

Great

Raman spectroscopy

4W200

ca. 5

Low'

Finely divided' or flat surfaces

Yes

Polarizability change perpendicular to surface

Limited

SERS

4000-200

ea. 5

High

Mostly Cu, Ag, Au

Yes

Polarizability change perpendicular to surface

Moderate; great for Ag

IETS

4000-200

ea. 5

Moderate

Finely divided'

No

As for infrared and Raman

Limitedg

INS

4000-200

ea. 10

Moderate

Finely divided

Yes

Hydrogenic motions

Limitedh

IR transmission 4-a.

>30

"Limited by transmission of the oxide support; >I300 cm-' for silica, >I050 cm-' for alumina. 'Normally oxide-supported; see foomote a. 10 mbar. Dipolar mechanism. 'Impact mechanism. 'Enhanced sensitivity with some metals. 'Samples have to be at liquid-helium temperatures. Large samples needed, as well as access to an atomic pile as the neutron source.

' Ultra-high vacuum equipment needed; highest allowable gas-phase pressure

-

~'


9

It can be seen from Table I1 that, for the most part, metal loadings within the catalysts have been rather high (-10%). Such high metal loading was because of limited infrared sensitivity for the observation of adsorbed species. However, since the development of Fourier-transform spectrometers in the early 1970s, this has no longer been necessary. For financial reasons, most industrial catalysts formed from precious metals, such as Pt and Pd, have loadings of less than 1%; and, for the hture, it is to be hoped that more samples of this type will be investigated (60). High metal loadings tend to go with large metal particle sizes, which show well-developed facets under the electron microscope. Such preparations are likely to give spectra that are simplified by the effective operation of the metal-surface selection rule (see Section 1V.B). Rather more complex spectra are to be expected from the relatively few catalysts that have particle diameters less than 2 to 3 nm. Trenary (58) has emphasized that the catalyst reduction temperature is a particularly relevant variable in causing differently profiled spectra from the same adsorbate on different preparations of a given metal catalyst. In the initial work of Eischens et al. (5, 6) the metal/oxide catalyst was studied in the form of a powdered layer on a horizontal CaFz window. Little, Sheppard, and Yates (69) instead used a tube of porous glass as the oxide support, and this could conveniently be mounted vertically in the infrared beam. [Similarly, Terenin and Roev (70) used alumina gel as a metal support to study NO adsorption.] However, the small pores in the silica glass limited the amount of metal that could be incorporated, leading to weak spectra from adsorbed species. Furthermore, porous glass retains small amounts of reactive oxides of boron and other metals that can lead to spectral contamination. Sheppard and Ward (71), and also Dunken, Schmidt, and Hobert (39), therefore decided to form disks from (salt + oxide) by high pressure from a hydraulic press, starting from pure silica in the form of Cabosil or Aerosil. These disks, which were converted to metal-on-oxide through reduction with Hz in the evacuable infrared cell, also proved to be successfully porous to gaseous adsorbates. Many infrared spectra of adsorbed hydrocarbons have been obtained using such samples. The principal limitation of this method is the difficulty of obtaining uniform temperatures over the disk surface when temperatures significantly different from room temperature are required in the absence of a substantial gas phase. This is because the poor thermal conductivity of silica makes difficult the transport of heat to or from a source in contact with the edge of the thin disk. This difficulty can be avoided by spraying a slurry of the catalyst in a volatile organic solvent onto a heated CaFz plate so as to cause the solvent to evaporate rapidly. The powder usually adheres to the CaFz plate, providing good control for work at low or high temperatures. This method, adapted from Yang and Garland ( 7 2 ) , has been successfully used by Yates (50) and Trenary ( 5 7 ) for hydrocarbon work. An alternative approach to obtaining uniform catalyst temperatures over a

10

NORMAN SHEPPARD AND CARLOS DE LA CRUZ TABLE I I Churucterizution Dutu ,for O.ride-Supported Metul Catalysts Emplqved by Research Groups Studying the Adsorption of Hydrocurhons hy Infrured (IR)or Rumun (Ru) Spectroscopy

Research group

Metal loading Spectroscopy Catalysts Precursor

(wt%)

Eischens and Pliskin

IR

Ni/Si02

Ni(NO,)?

9

Little, Sheppard, and Yates

IR

Ni/SiO? Pd/SiOz CdSi02

Ni(N03)2 Pd(N0,)2 Cu(N03)~

2.3 5.8 9

Sheppard et a/."

IR

Ni/SiOz Pd/Si02 Pt/Si02 Ir/Si02 Rh/SiO?

Ni(N03)z PdCI2 H2PtClh H21rCl,, RhCI,

Pd/Si02 Pt/Si02

PdClz H2PtC16

2.5 2.5

Ni/SiO?

Ni(N03)2 Ni(N03)2

18.7 I0 5 5 15

Dunken, Schmidt, and Hobert

IR

Erkelens ul.

IR

Nils102

el

Cu/SiOz PdJSiO? Pt/Si02 Palazov ul.

IR

Metal Reduction particle temperature size (nm) ("C) References cu. 8

C4 Linear Alkenes

a. Single Crystal Results. Pent-1-ene seems to be the only molecule in this class that has been studied for adsorption on a metal single crystal. In their VEELS study on Pt( 1 1 l), Avery and Sheppard (210) qualitatively interpreted spectra obtained at 200 and 300 K as representing di-0 and pentylidyne adsorbed species, respectively. The relative simplicity of the 300 K spectrum suggested that, at the high coverage used, the pentylidyne surface species had adopted the planar zigzag conformation. The alternating relative intensities of the 6CH3s mode from ethylidyne (s), propylidyne (w), butylidyne (s), and pentylidyne (w) supported this supposition and, as will be discussed below, finds confirmation in the spectra on finely divided metal catalysts. At higher temperatures, decomposition pathways appeared to have common features with those from but- 1 -ene discussed above. b. Finely Divided Metals. Figure 22 compares the room-temperature alkylidyne spectra from ethene, propene, and but-1-ene on Pt/Si02 and from pent1-ene on Pd/Si02 (Fig. 2 of reference 257). This very clearly shows the alternation in intensities of the 6CH3 s mode noted on Pt( 1 1 1 ). Because of the much beffer resolution of these infrared spectra in comparison with VEELS, the analogous expected intensity alternation of the 6CH3 s and 6CH3 as modes is also evident in this figure. Figures 23A and B show the vCH regions of the near-room-temperature spectra of pent-1 -ene adsorbed on Pd/Si02 before and after hydrogenation ( 4 3 ) , and on Ni/Si02 after hydrogenation only (7). Another spectrum on Pt/Si02 with coadsorbed CO has been recorded by Palazov et al. ( I 71) but without a listing of band positions. The strongest absorption at 2968 cm- obtained after initial adsorption on Pd/Si02 is as expected for the planar zigzag pentylidyne species; but, unfortunately, we do not have a RAIR spectrum on Pd( 1 1 1) with which to determine which of the other absorption bands are associated with that species. The absorption at 3020 cm- I , however, clearly shows that some surface alkene groups are present, probably the n-adsorbed species. The spectrum after hydrogenation on Ni/Si02 exhibits a reasonable n-pentyl profile when the limited resolution available at that time is taken into account. On Pd/Si02, the initial addition of hydrogen led to a spectrum with a higher ratio of the 2945-cm- I compared with the 29 12-cm- I absorption, characteristic of CH3as and CH2 as modes, respectively, than would be expected for n-pentyl.

'

92


cm-’ I

1500

2800

3000 I

1300

I

t

c

.-0 ul .-ul

E

ul

c

0

L

I-

i

3000

2800

I0

cm-’ FIG.22. A comparison of the infrared spectra, all near room temperature for (A) ethene, (B) propene, (C) but-I-ene, all on WSiO,; and (D) for pent-I-ene on Pd/SiO2. The spectra show a clear alternation of intensities of the vCH&CH3s (0)and vCH3as/GCH3 us absorptions (0) as a function of the number of carbon atoms. The shaded spectra were obtained by RAIRS for the adsorption of the same hydrocarbons on Pt( I I I ) and attributed to alkylidyne surface species with planar zigzag carbon skeletons (257, 43. 170).

This is probably because of the physical adsorption of some n-pentane molecules with their two methyl groups. After evacuation and then readdition of hydrogen, a n-pentyl profile was again obtained and retained through repeated fH2 cycles. As discussed in the butene examples, the spectrum of the surface alkyl groups was much reduced (eliminated on the PdSi02 case) in evacuation.


0 N

::z - m N

3000

mnN

g"N

N

2800

3000

cm-,

93

'H2 2800

FIG. 23. Infrared spectra from pent-I-ene adsorbed near room temperature before and after hydrogenation on (A) Ni/Si02 (7) and (B) Pd/SiOz (43).

It was with pent-1-ene on Pd/Si02 that Avery demonstrated the zip-fastener relationship of the hydrogenatioddehydrogenation processes on surface n-alkyls (258). He also pointed out that hydrogenation led to a wavenumber lowering from 2967 to 2945 cm-I, and a broadening, of the absorption for the vCH3 as mode on hydrogenation of the initially adsorbed species (43). He attributed this effect, which seems to be at its most pronounced on Pd, to interaction of the CH3 group with the surface after hydrogenation. It is noted that, of the transition metals, it is Pd that most strongly absorbs hydrogen. It was the virtually identical infrared spectra from the hex- 1 -ene, hex-2-ene, and hex-3-ene isomers when adsorbed on Ni/Si02 (7) that first led Eischens and Pliskin to the conclusion that the finely divided metal catalysts involve isomerization of the monohexenes. Their mutual spectrum and that obtained after hydrogenation are shown in Fig. 24B. Similar spectra at higher resolution were obtained by Shopov, Andreev, and Palazov (162) and by Erkelens and Liefkens (262);these are illustrated in Figs. 24C and D, respectively. The spectrum by Shopov et al. shows rather more absorption above 3000 cm- attributable to alkene-type surface species; however, because probably all the spectra in Fig. 24 were initially recorded on substantially sloping backgrounds from the metal catalysts themselves, it is easy to over- or underestimate the strengths of weak absorptions in this region. Erkelens and Liefkens (Fig. 24D) obtained somewhat different spectra on H-covered and H-depleted surfaces. Weaker absorptions occur on the H-free surface at 2926 and 2851 cm-I, the v C H ~as and vCHzs absorptions of alkyl chains; but the vCH3as and vCH3s absorptions at 2963 and 2876cm-' are relatively unchanged. It seems that some additional nonterminal attachment of the C6 chain to the surface occurs on the H-free surface.

'

94

NORMAN SHEPPARD AND CARLOS DE LA CRUZ cm-1

3000

2800

3000

2800

cm-1 FIG. 24. Infrared spectra from hex-I-ene adsorbed near room temperature before and after hydrogenation on (A) Pt/A1203 ( / 7 / ) ; (B) Ni/Si02 (7); (C) Ni/Si02 (162); and (D) Ni/SiOz on a hydrogen-covered surface (upper spectrum) and on a hydrogen-depleted metal surface (lower spectrum) (262).

After hydrogenation and the pumping-off of any resulting n-hexane, a vCH band profile is retained, closely similar to that expected from a surface n-hexyl group. Erkelens and Liefiens (262) made the interesting observation that attempted “hydrogenation” with deuterium led initially to a large increase in intensity in the vCH region and a generation of only a weak broad absorption in the vCD region. The addition of hydrogen to the hex-I-ene initially absorbed on a D-covered Ni/Si02 surface also led to a large absorption increase in the vCH region to give a spectrum virtually identical with that in Fig. 24D obtained after hydrogenation. It was concluded that the formation of new CH bonds by hydrogenation was more efficient that that of CD bonds. It is possible that, in the experiment involving the addition of gas-phase D2, it is the surface hydrogen atoms originally formed from dissociate adsorption of the hexene that remain immediately adjacent so that, under pressure from more adsorbed D2 or H2, they


95

are most rapidly transferred back to the carbon skeleton. The hydrogeddeuterium exchange and addition processes are very slow, involving reaction for more than 20 hours. Similar spectra were obtained from hex- 1 -ene adsorbed on Pt/Si02 at 305 K coadsorbed with CO [Fig. 24A (I 7f)]or on Pt/AI203 at 473 K (48), although in neither case were precise band positions recorded. Spectra on Pt/A1203 at the higher temperature of 473 K by Baumarten and Weinstrausch (48) also showed absorption in the 175O-1400-cm-' region. After initial adsorption, followed by flushing with helium to remove physically adsorbed species, the spectrum showed (in addition to the usual three absorptions in the vCH region) a substantial absorption tail above 3000 cm-' and a broad absorption at ca. 1560 cm-I. These may be from n-adsorbed or dehydrogenated surface species. Additionally, this spectrum showed expected absorptions at ca. 1460 (ms) and 1380 cm-' from the 6CH&CH3 as and 6CH3 s modes, respectively. At room temperature on Ni/Si02 the corresponding absorptions were listed at 1466 and 1376 cm-', but not illustrated. Spectra from hept- 1-ene, hept-2-ene, and hept-3-ene have also been obtained by Shopov et al. (161, f 6 2 ) on Ni/SiOz. 4.

Linear Dienes

a. Buta-1.3-diene. It seems that the only study of buta-1,3-diene on a single-crystal surface has been on Pt( 1 11) investigated by VEELS (253). Two overlapping sets of absorption bands varied in relative intensity between 170 and 300 K. One set was well assigned to vinyl and the other to an alkane group. It was suggested that at 170 K the surface species involved di-a bonding across one double bond, with the other remaining as a vinyl group. At 300 K it was considered that a greater proportion of a di-(di-a)-bonded, i.e., 1,2,3,4-a-bonded (but still nondissociatively adsorbed) species was present. At 385 and 450 K, spectra were shown similar to the analogous spectra from the linear butenes. Among the finely divided catalysts, Ward studied buta- 1,3-diene adsorbed on Pt/SiO2 and Ni/SiOz (71); Avery on Pd/Si02 (43); Soma on Pd, Ni, and Co/A1203 (263); and Basu and Yates on Rh/A1203 (264). The results are somewhat different on the different metals, and more work needs to be done before definitive structural assignments can be made. However, some regularities are worth noting at this stage. The room-temperature spectra are most easily discussed. On Ni and Co (7f,263), it is clear that considerable polymerization has occurred to give lengthening polymethylene chains evincing the usual strong absorptions at ca. 2925, vCH2 as; 2854, vCH2 s; and 1450 cm- 6CH2 in the region down to the 1300-cm-' SiOz cutoff. In addition, there is a set of absorptions on Ni characteristic of non-n-bonded vinyl groups (3075, v=CH2; 30 10, v=CH; 1650, vC=C; 1407, 6=CH2; 1308,6=CH) with analogous bands on Co/Si02.

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NORMAN SHEPPARD A N D CARLOS DE LA CRUZ

Hydrogenation led to substantial additional absorption from -CH2 - groups and the elimination of the vinyl absorptions. Weak bands at ca. 2960 and 1370 cm- from methyl groups also occurred. On PdSiO2 (43, 263) at room temperature no discernible spectra were obtained after adsorption; but after hydrogenation, strong absorptions occurred from surface-attached n-alkyl groups of limited chain length. As was mentioned above in the context of the adsorbed linear butenes (Section VI.C.2.c), it seems that in the absence of hydrogen in the gas phase, the strong capability of Pd to absorb hydrogen leads to dehydrogenation to give linear polyenes who spectra are very weak for MSSR reasons. Buta-l,3-diene on Pt/Si02 at room temperature ( 7 1 ) shows a poorly defined spectrum in the vCH region with alkene- and alkane-type components. On hydrogenation, well-defined n-butyl spectra appear (as also from but- 1-ene), showing that on this metal no polymerization has occurred. Rh/A1203 also showed no polymerization, but in fact the reverse (264). At room temperature the spectra did not differ greatly from those of adsorbed but1-ene in the vCH region; but absorptions at ca. 2880 and 1340 cm-', which grew in intensity up to 350 K, showed the presence of a growing fraction of ethylidyne, i.e., C2 surface species. Between the lower temperatures of 200 and 250 K, absorptions from this catalyst, such as those listed above, denoted the presence of non-n-bonded vinyl groups. In the same temperature range, additional absorptions at ca. 1580 (ms), 1430 (ms), and perhaps 1380cm-' (m) occurred at very similar positions to analogous absorptions in the spectra on Pd/A1203 and Ni/A1203at 241 and 238 K, respectively. A band at 1320 cm- in the spectrum on Rh/A1203at 225 K, associated with the other three by Basu and Yates (264), did not appear in the spectra on Pd or Ni. Because of the 1580-cm- absorption, it is most probable that the above trio of bands shows the presence of vinyl groups that are n-bonded to the surface. In their second paper on buta-1,3-diene adsorbed on Rh/AI203, Basu and Yates ( 2 5 8 ~ )concluded that self-hydrogenationldehydrogenation reactions occurred on the surface above 250 K, presumably involving different adsorption sites. They also showed that added hydrogen led to adsorbed n-butyl groups but only above 230 K. We infer from the spectra that the first self-hydrogenation product was adsorbed but- 1-ene and that the methyl-rich spectrum first observed on adding H2 to the system at 230 K was from physically adsorbed n-butane. SER spectra of butadiene on cold-evaporated Ag at 60 K (257a) have been assigned to a mixture of R- and di-n-bonded surface species. Hence a number of possible surface structures for adsorbed buta- 1,3-diene have been proposed, including mono-n or di-a adsorption involving one vinyl group, di-n or di-(di-a) (i.e., 1,2,3,4-tetra-a) absorption involving both vinyl groups, and possibly a metallocyclopentane after dehydrogenation which could lead to the formation of polymethylene chains. To these should be added a

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CH2=CH-CH2-CM3 structure, i.e., an unsaturated analog of butylidyne, the possible source of the absorptions from non-n-bonded vinyl listed above for Ni/Si02 and Co/Si02. Further progress in structural interpretations probably requires spectra on simplified single-crystal systems. Such low- and room-temperature spectra using the higher resolution of RAIRS would be particularly valuable on three of the metals, Pt, Ni or Pd, and Rh, which give different results at room temperature on the finely divided metals. b. Hexa-l.5-diene. Spectra from this nonconjugated diene have been reported on Ni/SiOz by Shopov et al. (162) and by Erkelens and Liefiens (262). The spectra at room temperature were said to correspond closely to the same spectra obtained from each of the hexenes, with the presence of methyl absorptions proving evidence for isomerization. However, a well defined spectrum from the di-n-bonded form of hexa-1,Sdiene has been observed on Ag( 110) at 300 K ( 2 5 2 ~ ) .

D. BRANCHED-CHAIN ALKENES 1. 2-Methylpropene (Isobutene) This first of the branched alkenes is capable of giving symmetrical adsorbed species and hence spectra that are not too complex. Its spectra therefore merit individual attention. a. Single-Ciystal Results. 2-Methylpropene has been studied by VEEL on Pt( 1 1 1 ) between 170 and 420 K (210) and on Ni( 11 1) between 80 and 180 K (265). It has also been investigated by RAIRS on Pt( 1 1 1 ) at 90 K and room temperature (266) and on Ru(0001) between 90 and 300 K (254). The on-specular VEEL spectra at 170 K on Pt( 1 1 1) and at 80 K on Ni( 1 1 I ) have similar patterns of prominent bands [2920 (m), 1470 (m), 1090 (s), 800 (m), and 460 cm-I (vs) on Pt(ll1); 2910 (m), 1450 (m), 1055 (s), 760 (m), and 445 cm- (vs) on Ni( 11 l)]. The RAIR spectrum at 90 K on Pt( 1 1 1) has better resolution than in VEELS and has corresponding prominent features at 2910 (s), 1426 (mw), and 1062 cm-l (s) in the accessible region down to ca. 850 cm-l (266).That on Ru(0001) at 90 K has a closely similar pattern of features [2886 (s), 1046 (s)]. It seems probable that the spectra on all three metals at these low temperatures are derived from the same type of adsorbed species. By analogy with the cases of adsorbed ethene on Pt( 1 1 1) and Ni( 1 1 l), and by taking into account the relative simplicity of the present spectra, it was suggested by Avery and Sheppard (210) that this is the di-o species. Other authors have followed this assignment, but Hammer et al. (265) have pointed out, with the help of

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analogous spectra from (CD3)2C=CH2, that both alkene CH bonds appear to exhibit “soft” vCH modes which show up prominently in off-specular spectra. “Soft” modes are associated with “hydrogen-bond-like” interactions with metal atoms and the spectra from (CD3)2C=CH2 show no features from “free” vCH modes. Hammer et al. conclude that in some way the CCH2 of the molecule must be tilted downward toward the surface. It is difficult to envisage this occurring sufficiently from a di-a on the flat (1 11) surface, at least without surface reconstruction. Such tilting could more readily occur about the single metal atom to which a R- or dimethyl-substituted metallocyclopropane species could be coordinated. In the present context, the vCH frequencies of the terminal CH2 group seem to be remarkably low, at 2580/2660 cm- to be from a n complex. Furthermore, no absorption has been observed in the 1650-1450 cm region for attribution to a a-bonded C=C group. Steric effects between the surface metal atoms and the bulky methyl groups could cause some tilting of the C=CH2 group toward the surface, but seem unlikely to be sufficiently strong to cause such a strong lowering of the soft K H modes. Also in this connection it should be recalled that the spectrum of ethene itself on Ni( 1 11) shows such soft modes. Could there be a coexistence of metallocyclopropane and di-a species from ethene on this surface? Or could type I spectra-of which that from ethene on Ni( 111) is a typical example-be simply an extreme form of the type I’ spectra that we have associated with metallocyclopropane structures, and not from a di-a structure after all? We recall the earlier ambiguities in relating the spectra of the model di-a or metallocyclopropane osmium-based model compounds to typical type I spectra (Section IV.C, p. 26). At the higher temperature of 300 K on Pt( 1 1 1) a different, but still relatively uncomplicated, spectrum was obtained that Avery and Sheppard found to be consistent with the presence of the 2-methylpropylidyne (isobutylidyne) structure. This species has no conformational isomers associated with the carbon skeleton and the prominent absorption bands (at 2970, v C H as; ~ 1460,6CH3 as; and 1010 cm-I, CH3 rocking) are as expected according to the MSSR for the alkylidyne. The RAIR spectrum on Pt( 1 11) at room temperature was in agreement. Similar spectral features were observed on Ru(0001) and Ni(ll1) at the lower temperature of 180 K, particularly the dominant CH3 as absorption in the RAIR spectrum of Ru(000 1). However, the VEEL spectrum at this temperature on Ni(ll1) lacked the strong feature at 800cm-’ observed in the 300 K spectrum on Pt( 11 1). Hammer et al. doubted the retention of any vCH absorption in the VEEL spectrum from (CD3)2C=CH2 at 180 K and observed a weak band at 1595 cm- I , off-specular. They therefore preferred the interpretation of the spectrum on Ni( 11 1 ) in terms of a d i - a h [(CH3)2C=C] surface species. The RAIR spectrum on Ru(0001) at 300 K exhibited a single prominent absorption at 1366 cm- I , suggesting the occurrence of skeletal decomposition

’,

~

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99

to the ubiquitous ethylidyne. The spectrum from 2-methylpropene on Pt( 1 1 1 ) at 420 K was tentatively interpreted, in view of a TPD-estimated overall surface composition of C4H4.5, in terms of a CH(CHM2)3 species with 3-fold axial symmetry. b. Finely Divided Catalysts. Figure 25 shows collected spectra from 2-methylpropene adsorbed on SO2-supported catalysts of several metals [Pt (71, 269, Pd (43), Ni (71), Ir (36)]. Shahid and Sheppard (267) have most recently discussed the spectrum on Pt/SiO2 in detail. The room-temperature spectrum is shown in Fig. 25A. On heating to 373 K, this spectrum substantially simplified and gave dominant hydrocarbon absorptions at 2960 and 1450 cm- I , as expected under the MSSR for a 2-methylpropylidyne species. The 373 K spectrum on Pt/SiO2 is also closely similar to those obtained at room temperature on Pd/Si02 and Ni/Si02 (Figs. 25B and C). Additional absorptions on 2800

3000

N

3000

cm-'

1600

1400

"N"8

%

N

m

tH2 2800

1600

1400

cm-'

FIG.25. Infrared spectra from 2-methylpropene (isobutene) adsorbed near room temperature. and with hydrogenation spectra inset, on (A) Pt/Si02 (267); (B) Pd/SiOz (43);(C) Ni/Si02 (71); and (D) Ir/Si02, reprinted with permission from (36).copyright 1972 American Chemical Society.

100


Pt/Si02 at room temperature at 2970, 2925, 1045, and 1375 c m - ' were attributed to the presence of R- and di-a-bonded adsorbed species. The strong 6CH3 s feature at 1375 cm-' is particularly as expected for the di-a species, with its CH3C groups oriented at a high angle with respect to the surface. It may have been the conversion of di-a to the alkylidyne species, such as occurs for ethene on Pt/Si02 somewhat below room temperature, that accounts for the intensification of the 2960-cm- ' absorption from the alkylidyne on heating to 373 K. The profile of the spectrum at room temperature from Ir/SiO2 (Fig. 25D) resembles that from Pt/Si02 except for a more prominent vCH absorption at 2920 cm- ', possibly from the presence of a higher proportion of di-a species. As in the case for adsorbed but-1-ene on Pt/SiO2, the broad absorption at 1600cm-' in Fig. 25A is attributed to the presence of chemisorbed bridged hydrogen. It intensifies on heating up to above 473 K, and this would be consistent with additional surface hydrogen from thermal decomposition of the hydrocarbon species. At the same time, the alkyl absorptions become weaker and broader while a broad absorption from v=CH and/or v=CH2 grows to become the strongest spectral feature at 573 K. The inset spectra in Fig. 25 were observed at room temperature in the vCH region after hydrogenation ( + H2), followed by evacuation ( - H2) in order to remove physically adsorbed 2-methylpropane. Unlike the case of the linear butenes, considerable intensity is retained after the 4zH2 procedure and the species in question is clearly still a methyl-rich one, probably the alkyl group 2-methylpropyl. On Pt/Si02, hydrogenation was also studied at 373 K and above. In this case, the residual adsorbed species after 2-methylpropane removal at 373 K was weak and no longer methyl-rich. Readdition of hydrogen at 473 K led to a substantial spectrum of an unbranched n-alkyl nature, indicating the occurrence this time of skeletal surface isomerization. The same type of spectral removal ( - H2) and restoration ( + H2) thereafter occurred as for the adsorbed linear alkenes. Once again, some methane production, implying C -C hydrogenolysis, occurred at 573 K and above. 2.

Other Brunched-Chain Alkenes

The only higher branched alkene that appears to have been studied on a single-crystal plane is 2,3-dimethylbut-2-ene on Ni( 1 1 l), for which the VEEL spectrum has been measured at 80 K (255). At monolayer coverage, it gives a strong methyl-rich spectrum with absorptions at ca. 2870 (ms), 1470 (s), ca. 1400 (ms), 1060 (s), 690 (w), cu. 360 (s), and 320 cm-' (s). The first three prominent absorptions are those of the methyl groups and the strong 1060-cm absorption is probably from a coupled CH3 rocklvCC mode. It seems most probable that this is from the di-a species, although steric strain between vicinal methyl groups is more likely to give a surface species of C2 rather than C2,, symmetry.

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The adsorption of 2-methylbut-2-ene on Ni/SiO2 and PtISiO2 (71) and on Ir/Si02 (36), of 3-methylbut-1 -ene on Ni/SiO2, and of 4,4-dimethylpent-l -ene on Pd/SiOz (43) have been recorded, principally in the vCH region. For the most part, these give spectra dominated by absorptions from constituent CH3 groups, and it is difficult to draw specific structural conclusions from the data. Hydrogenation leads to slow desorption of the corresponding alkanes.

E. CYCLICALKENES The spectra of these will be discussed in Part I1 because of their intermediate status between cycloalkanes and aromatic hydrocarbons.

F. COMPARATIVE REACTIVITIES OF HYDROCARBON SPECIES ADSORBED ON DIFFERENT METALSURFACES Many of the same reactions of adsorbed hydrocarbon species occur on the surfaces of different metals but over differing temperature ranges. A lower temperature of completion of a given reaction implies a higher reactivity toward the product in question. In Table VIII are collected together the approximate completion temperatures of six different reactions on oxide-supported metals. A comparison is made between the analogous reactions involving C2 species derived from the adsorption of ethene and those involving C4 species from the linear butenes. In a few cases, analogous data are used from C3, Cs, or C6 species where the C4 data are incomplete. The experimental data are fragmentary and only approximate; nevertheless, some interesting trends can be discerned. It is satisfactory to note, for example, that the quoted temperatures for particular reactions involving C2 or C4 species are usually closely similar within the typical estimated uncertainties of f30 K. This implies that the reactivity on a particular metal depends principally on the functional group attaching the hydrocarbon species to the surface, as in M3CCH2R’ for the alkylidynes or MCH2R’ for the surface alkyls (R’ = alkyl). In the case of alkylidyne decomposition, it should be noted that different reactions are involved for the C2 and C4 species. It is apparent that Pt is rather generally the least reactive of the group VIII (IUPAC 8-1 0) metals, as its reaction-completion temperatures are substantially the highest. Only the temperature of alkylidyne formation is an exception to this generalization, where Pt is not notably different from the other metals. Palladium seems to be a particularly effective dehydrogenation metal in alkyl dehydrogenation on evacuation or, for the C4 species, alkylidyne formation from

102


di-a species. It has been suggested that this is related to the unique capacity of Pd to absorb, as well as to adsorb, hydrogen.

G . HIGHERALKENES ON METALS-AN OVERALL PERSPECTIVE Because relatively few VEEL or RAIR spectroscopic studies have been made for the higher alkenes (substituted ethenes) adsorbed on metal single crystals, the situation for the identification of the chemisorbed surface species is much less complete than is the case for ethene as the adsorbate. Furthermore, the number of conceivable surface structures is much greater than in the ethene case including, for >C, species, the effects of isomerization of the adsorbing species. Nevertheless, substantial progress has been made. The VEELS and RAIRS studies of propene on several metal surfaces, albeit all close-packed ones [more work on fcc (1 10) and (100) planes would be valuable], seem to have established sound criteria for identifying alkylidyne species on single-crystal or finely divided metals. A similar study of a series of alkenes on Pt( 1 I l), again by VEELS and RAIRS, also helped to generalize the conclusions from propene to the longer-chain alkenes. However, as hydrocarbon species on Pt seem to exhibit reactivity at higher temperatures than on most other metals (Section F above), similar studies on several single-crystal planes of other metals such as Ni and Pd would be very worthwhile. Nickel is likely to illustrate greater general reactivity, and Pd a greater propensity for dehydrogenation processes. At lower temperatures, spectra taken on these three metals, Pt, Pd, and Ni, might also permit the identification of the several possible substituted (CH2CH2) types of surface species, i.e., of the n, metallocyclopropane, or di-a types. Our earlier discussion of the case of 2-methylpropene adsorbed on Ni( 1 1 1) highlighted the continuing spectroscopic uncertainties in distinguishing between these species derived from the higher alkenes. For this purpose the capacity of both VEELS and RAIRS to give results in the fingerprint region below 1300 cm- is of importance. For finely divided metals the use of the more transparent A1203 support would be advantageous. For these further explorations the new higher-resolution techniques of VEELS (although unlikely to compete with RAIRS in this one respect) are of complementary importance for sensitivity reasons, and because off-specular studies can provide information about the modes with vibrational dipole changes more parallel than perpendicular to this surface. The off-specular studies are also more widely useful for the identification of vCH “soft” modes, with their implications about the nature and orientation of species with respect to this surface. A careful choice of the adsorbates to be studied could accelerate our understanding of this research field. Propene and 2-methylpropene have the advantages that they offer no spectroscopic complications from conformational isomerism within the hydrocarbon part of the adsorbed species, except perhaps

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in relation to nonplanarity of the CzMz dimetallocyclobutane skeleton in the case of the di-o species. In general, the spectra become more difficult to interpret as the number of carbon atoms increases, and even the linear butenes bring into play conformational mobility and/or isomerization reactions. Clearly, the most interpretable results are likely to be obtained from studies of the linear butenes, in normal and selectively deuterium-substituted forms. Although at sufficiently low temperatures all the individual linear alkenes preserve their identity as chemisorbed species, one of the clear-cut conclusions from work to date is that-even for the relatively unreactive Pt surface-at room temperature the finely divided metals cause virtually complete isomerization between but-1-ene and the cis- and trans-but-2-enes. The extent to which this can even occur on the close-packed Pt( 1 1 1) face at lower coverages at 300 K is a current question. Attempts to identify intermediates in the isomerization process, possibly of an allylic nature, would be a matter of priority on single crystals or finely divided metals. The conditions under which the hydrogenation of the higher alkenes leads to gas-phase alkanes or to surface-anchored alkyl species remain to be identified. High coverage may be a factor in inhibiting the replacement of the last carbonmetal bond by CH in order to give the gas-phase product, but this needs to be investigated further. Also on single-crystal surfaces, using RAIRS or VEELS, it should be possible to confirm or deny the alkane-to-polyalkene transformation that has been postulated as a means of accounting for the experimentally wellestablished, and highly-efficient, zip-fastener type of dehydrogenation of surface n-alkyl groups on the removal of gas-phase hydrogen.

VII.

Conclusions

In this article (Part I) we have comprehensively reviewed the structural implications of the vibrational spectroscopic results from the adsorption of ethene and the higher alkenes on different metal surfaces. Alkenes were chosen for first review because the spectra of their adsorbed species have been investigated in most detail. It was to be expected that principles elucidated during their analysis would be applicable elsewhere. The emphasis has been on an exploration of the structures of the temperature-dependent chemisorbed species on different metal surfaces. Particular attention has been directed to the spectra obtained on finely divided (oxide-supported) metal catalysts as these have not been the subject of review for a long time. An opportunity has, however, also been taken to update an earlier review of the single-crystal results from adsorbed hydrocarbons by one of us (N.S.) (I 7). Similar reviews of the fewer spectra from other families of adsorbed hydrocarbons, i.e., the alkynes, the alkanes (acyclic and cyclic), and aromatic hydrocarbons, will be presented in Part 11.

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Knowledge of the chemisorbed species present on a surface provides an essential database for the investigation of catalytic and other reaction mechanisms in which they are involved. The elucidation of reaction mechanisms in turn requires studies of chemical kinetics. An increasing number of such kinetic studies are now being made which involve vibrational spectroscopy and hydrocarbon adsorbates, and these will also be reviewed in Part 11. Overall perspectives of the results from ethene and the higher alkenes have been attempted in Sections VI.B.6 and V1.G. What has become clear, particularly in the context of hydrocarbon adsorption, is that the study of spectra on single-crystal surfaces is of great assistance in finding the correct interpretation of the more complex multispecies spectra obtained from finely divided metal catalysts. This has only become possible by the development of VEELS and RAIRS, the latter allied with the Fourier-transform methods that have also transformed the quality of the spectra from metal-particle catalysts obtained by transmission infrared spectroscopy. The use of RAIRS in turn has emphasized the general significance of the MSSR. The majority of identifications of the structures of adsorbed species have to date been made by vibrational spectroscopy. This is because, once spectral patterns have been assigned to particular species, the analysis is very rapid and can be carried out in the presence of more than one type of surface complex. The assignment of spectral patterns has in the past been made using the infrared spectra of ligands of known structures on metal-cluster model compounds. This method finds its ultimate origin in the use of diffraction methods, in these cases usually X-ray diffraction. It would be even more satisfactory if model spectra could be assigned from structures that are directly determined on metal surfaces. So far, relatively few such hydrocarbon structures have been determined in this way by LEED (83, 84). However recent advances in diffraction techniques, including tensor-LEED and photoelectron diffraction (PED), are likely to greatly improve the situation in the near future. The newer PED technique is particularly welcome as it provides local-site information and is not dependent on the occurrence of regular arrays of adsorbates. It is some forty years since Eischens, Pliskin, and Francis in 1954 made the break-through of obtaining spectra from chemisorbed monolayers on metal/ silica catalysts. The subsequent incorporation of three experimental developments (FTIR, VEELS, and RAIRS) and a theoretical understanding (MSSR) has led in the intervening years to tremendous advances in the field of spectroscopic research that they pioneered-so much so that a suggestion in 1954 that we might have reached our present capability by 1995 would almost certainly have been met with disbelief. But much remains to be done in the vibrational spectroscopy of chemisorbed hydrocarbons because Nature has the capacity to match scientific advances by revealing sophisticated new phenomena for investigation. In this review, we have been at pains not only to summarize what has


105

been achieved but also to point to some of the more promising pathways ahead. ACKNOWLEDGMENTS The authors are very grateful to the following persons who have given permission to us, for this review, to reproduce spectra from their publications or theses redrawn to a uniform format for comparison purposes: Dr. N. R. Avery, Prof. E. Baumgarten, Prof. G. Blyholder, Prof. R. P. Eischens, Prof. J. G. Ekerdt, (the late) Dr. J. Erkelens, Dr. G. Ghiotti, Prof. H. Knozinger, Dr. D. 1. James, Dr. G. S. McDougall, Prof. B. A. Morrow, Dr. J. D. Prentice, Dr. M. Primet, Mrs. A. Lesiunas, Prof. D. Shopov, Dr. Y. Soma, Prof. M. Trenary, Dr. J. W. Ward, and Prof. J. T. Yates, Jr. We are also indebted to the following copyright holders of the publications concerned for their permission to reprint the spectra in this review with references cited in figure captions: Academic Press (Figs. 7B, 9A, 13A. 13B, 14B, 15A, 15B, 178, 17C, 19A, 19C, 23A, 23B, 24B, 24D, 25D); Akademiai Kiado, Hungary (Figs. IC and 24A); the American Chemical Society (Figs. 6D, 7F, 7H. IOD, IOE, 12C, 12D, 12E. 13D, 14C, 14D, 14E, 15C, 15D, 19H, 25D); the American Institute of Physics (Fig. 17C); Baltzer Scientific Publishing Co. (Figs. 14F, 17E, 19D); Bulgarian Chemical Communications (Fig. 24C); Elsevier Science (Figs. 71, 75, 13E, 15E, 17A); Journal de chimie physique (Fig. IOF); the National Research Council of Canada (Fig. 25A); the Royal Society of London (Figs. 6A, 7A, 7E, IOB, 12A, 13C, 18A, 18C, 19B); and the Royal Society of Chemistry, London (Figs. 6B, 6C, 7G, 9A, 19F, 20, and 21). One of us (N.S.) thanks the U.K. Science and Engineering Research Council for a series of research grants that supported the work of his laboratories at the Universities of Cambridge and East Anglia in this research area. We are both very grateful to the Royal Society, London, and to the Consign0 Nacional de lnvestigaciones Cientificas y Tecnologicas of Venezuela for supporting two transatlantic exchange visits which assisted greatly in the writing of this review article. REFERENCES 1. Sabatier, P., and Senderens, J. B., Compf. Rend. 1358 (1897). 2. Bond, G. C., “Heterogeneous Catalysis; Principles and Applications,” Oxford Univ. Press, Oxford, 1974. 3. Kemball, C., Ah: Cafal. 11, 223, (1959). 4. Burwell, R. L. Jr., Acc. Chem. Res. 2, 289 (1969); Catal. Left 5, 237 (1990). 5. Eischens, R. P., Pliskin, W. A,, and Francis, S. A,, J. Chem. Phys. 22, 1786 (1954). 6. Pliskin, W. A,, and Eischens, R. P., J. Chem. Phys. 24, 482 (1956). 7. Eischens, R. P., and Pliskin, W. A,, Adv. Catal. 10, 1 (1958). 8. Bellamy, L. J., “Infrared Spectra of Complex Molecules,” Vol. I , 3rd Ed. Chapman & Hall, London, 1975. 9. Lin-Vien, D., Colthup, N. B., Fateley, W. G., and Grasselli, J. G., “Infrared and Raman Characteristic Frequencies of Organic Molecules,” Academic Press, New YorWLondon, 1991. 10. Nakamoto, K., “Infrared and Ramam Spectra of Inorganic and Coordination Compounds,” 4th Ed. Wiley-Interscience, New York, 1986. I / . Skinner, P., Howard, M. W., Oxton, 1. A,, Kettle, S. F. A,, Powell, D. B., and Sheppard, N., J. Chem. Soc. Faraday Trans. 2, 77, 1203 (1981). 12. Ibach, H., Hopster, H., and Sexton, B., Appl. Sur- Sci. I, I (1977). 13. Bertolini, J. C., Dalmai-lmelik, G., and Rousseau, J., Stir- Sci 67, 478 (1977). 14. Francis, S. A,, and Ellison, A. H., J. Opt. SOC.Amer. 49, 131 (1959). 15. Chesters, M. A,. Pritchard, J., and Sims, M. L., J. Chem. SOC.Chem. Commun. 1454 (1970). 16. Chesters, M. A,, J. Electron Spectrosc. Relat. Phenom. 58, 123 (1986).

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201. Cooper, E., and Raval, R, Surj.' Sci.. 331-333,94 (1995). 202. Karlsson, P. A,, Martensson, A-S., Andersson, S., and Nordlander, P., Sur$ Sci. 175, L759 ( 1986). 203. Zhu, X.-Y., and White, J. M., Surf: Sci., 214, 240 (1989). 204. Lapinski. M. P., and Ekerdt, J. G., J. Phys. Chem. 92, 1708 (1988). 205. Primet, M.. and Sheppard, N., unpublished work. 206. Sheppard, N., J. Electron Spectrosc. Relat. Phenom. 38, 175 (1986). 207. Primet, M., and Sheppard, N., J. Catal. 41,258 (1976). 208. Blyholder, G., Shihari, D., Wyatt, W. V., and Bartlett, R., J. Catal. 43, 122 (1976). 209. McDougall, G. A., personal communication. 210. Avery, N. R., and Sheppard, N., Proc. R. Sac. London A405, I (1986). 21 1. Krasser, W., and Renouprez, A. J., J. Raman Spectrosc. 1 I, 425 (1981 ). 212. Dubois, L. H., Castner, D. G., and Somojai, G. A,. J. Chem. Phys. 72, 5234 (1980). 213. Koel, B. E., Bent, B. E., and Somorjai, G. A,, Surf: Sci. 146,211 (1984). 214. Slavin, A. J., Bent, B. E., Kao, C-T., and Somorjai, G. A., Surf: Sci. 206, 124 (1988). 215. Slavin, A. J., Bent, B. E., Kao, C-T., and Somorjai, G. A,, Surf: Sci. 202, 388 (1988). 216. Eady, C. R., Fernandez, J. M., Johnson, B. F. G., Lewis, J., Raithby, P. R., and Sheldrick, G. M., J. Chem. Sac. Chem. Commun. 421 (1978). 217. Blyholder, G., and Wyatt, W. V., J. Phys. Chem. 78, 618 (1974). 218. Sakakini, B., Swift, A. J., Vickerman, J. C., Harendt, C., and Christmann. K., J. Chem. SOC. Faraday Trans. 1 83, 1975 (1987). 2 1 8 ~ .Sakakini, B., Havendt, C., and Vickerman, 1. C., Spectrochim. Acta 43A, 1613 (1987). 219. Mims, C. A,, Weisel, M. D., Hoffmann, F. M., Sinfelt, J. H., and White, J. M., J. Phys. Chem. 97, 12656 (1993). 219a. Ransley, I. A., Ilharco, L. M., Bateman, J. E., Sakakini, B. H., Vickerman, J. C., and Chesters, M. A,, Surf: Sci. 298, 187 (1993). 220. Henderson, M. A., Mitchell, G. E., and White, J. M., Surf: Sci. 203, 378 (1988). 221. Marinova, T. S., and Kostov, K. L., Surf: Sci. 181, 573 (1987). 222. Marinova, T. S., and Chakarov, D. V., Surf Sci. 192, 275 (1987). 223. Chakarov. D. V., and Marinova, T. S., S u r - Sci. 227,297 (1990). 224. Nyberg, C., Tengstal, C. G., Andersson, S., and Holmes, M. W., Phys. Chem. Lett. 87, 87 (1982). 225. McCash, E. M., Vacuum 40, 423 (1990). 226. Jenks, C. J.. Bent., D. E., Bemstein, N., and Zaera, F., Surf Sci. 277, L89 (1992). 226a. Kubota, J., Kondo, J. N., Domen, K., and Hirose, C., J. Phys. Chem. 98, 7653 (1994). 227. Ertiirk, U., and Otto, A., Surf: Sci.. 179, 163 (1987). 228. Slater, D. A,, Hollins, P., and Chesters, M. A., Surf: Sci. 306, 155 (1994); J. Electron Spectrosc. Relat. Phenom. 64165, 95 ( 1993). 229. Moskovits, M., and DiLella, D. P., Chem. Phys. Lett. 73, 500 (1980). 230. Wood, T. H., and Zwemer, D. A,, J. Vac. Sci. Technol. 18,649 (1981). 231. Ertiirk, U., and Otto, A,, Chem. Phys. Lett. 149, 284 (1988). 232. Brings, R., Mrozek, I., and Otto, A,, J. Roman Spectrosc. 22, 1 19 (1991). 233. Wolkow, R. A,, and Moskovits, M., J. Chem. Phys. 96, 3966 (1992). 233a. Patterson, M. L.. and Weaver, M. J., J. Phys. Chem. 89, 1331 (1985). 234. Seip, U., Tsai, M.-C., Kiippers, J., and Ertl, G., Surf: Sci. 147, 65 (1984). 235. Bernasek, S. L., Ann. Rev. Phys. Chem. 44, 265 (1993). 2 3 5 ~ .Demuth, J. E., and Ibach, H., SurJ Sci. 78, L238 (1978). 235h. Paul, D. K., Beebe, T. P., Jr., Uram, K. J., and Yates, J. T., Jr., J. Amer. Chem. Sac. 114, 1949 ( 1992).

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245. Gardner, P., Ph.D. thesis, University of East Anglia, 1988. 246. McCash, E. M., Ph.D. thesis, University of East Anglia, 1987. 247. Dent, A. L., and Kokes, R. J., J. Amer. Chem. Sac. 92, 6709 (1970). 248. Shahid, G., and Sheppard, N., Spectrochim. Acta 46A, 999 (1990). 249. Marshall, P. R., McDougall, G. S., and Haddon, R. A,, Top. Catal. I, 9 (1994). 250. Sheppard, N., Pure Appl. Chem. 4, 71 (1962). 251. Munro, S.. and Raval, R., Surf Rev. Lett. 1, 645 (1994). 252. Delbecq, F., and Sautet, P., Cutul. Lett. 28, 89 (1994). 252a. Carter, R. N., Anton, A. B., and Apai, G., J. Amer. Chem. Soc. 114,4410 (1992). 253 Avery, N. R., and Sheppard, N., Proc. R. Sac. A405,27 ( I 986). 254. Chesters, M. A., Horn, A. B., Ilharco, L. M., Ransley, I. A., Sakakini, B. H.. and Vickerman, J. C., J. Electron Spectrosc. Relat. Phenom. 54/55, 677 (1990). 254a. Chesters, M. A., Horn, A. B., Ilharco, L. M., Ransley, 1. A,, Sakakini, B. H., and Vickerman, J. C., Surf Sci. 2511252, 291 (1991). 255. Fricke, A,, Graupner, H., Hammer, L., and Miiller, K., Surf Sci. 272, 182 (1992). 256. Morrow, B. A,, and Sheppard, N., Proc. R. Sac. A311,415 (1969). 257. Shahid, G., and Sheppard, N., J. Chem. Sac. Faraday Trans. 90, 507 and 512 (1994). 257a. Itoh, K.,Tsukada, M., Koyama. T., and Kobayashi, Y.,J. Phys. Chem. 90, 5286 (1986). 258. Avery, N. R., J. Catal. 24,92 (1972). 258a. Basu, P., and Yates, J. T., Jr., J. Phys. Chem. 93, 61 10 (1989). 259. Shirakawa, H., and Ikeda, S., Polym. J. 2, 231 (1971). 260. Eischens, R. P., and Mertens, F. P., J. Electrochem. Sac. 115, C81 (1968). 261. Selwood, P. W., J. Amer. Chem. Soc. 79, 3346 (1957). 262. Erkelens, J., and Liekens, Th. J., J. Catal. 27, 165 (1972). 263. Soma, Y., Bull. Chem. Soc. Jpn. 50,2119 (1977). 264. Basu, P., and Yates, J. T., Jr., J. Phys. Chem. 93, 2028 (1989). 265. Hammer, L., Dotsch, B., Brandenstein, F., Fricke, A,, and Miiller, K., J. Electron Spectrosc Relat. Phenom. 54/55, 687 (1990). 266. Pudney, P., PhD. thesis, University of East Anglia, 1989. 267. Shahid, G., and Sheppard, N., Canad. J. Chem. 69, 1812 (1991).

ADVANCES IN CATALYSIS, VOLUME 41

Catalytic Chemistry of Heteropoly Compounds TOSHIO OKUHARA Graduate School of Environmenral Earth Science Hokkaido University Sapporo 060, Japan

NORITAKA MIZUNO Institute of Industrial Science The University of Tokyo Roppongi, Minaro-ku, Tokyo 106. Japan AND

MAKOTO MISONO Deparrment of Applied Chemistry Graduate School of Engineering The University of Tokyo Bunkyo-ku. Tokyo 113. Japan

1. A.

Introduction

HETEROPOLYCOMPOUNDS AS CATALYSTS

The catalytic properties of heteropoly compounds have drawn wide attention in the preceding two decades owing to the versatility of these compounds as catalysts, which has been demonstrated both by successhl large-scale applications and by promising laboratory results. Heteropolyanions are polymeric oxoanions formed by condensation of more than two different mononuclear oxoanions, as shown in Eq. (1): 12WO:-

+ HPOi- + 23H'

-+PW120:i

+ 12H20

(1)

Heteropolyanions formed from one kind of polyanion are called isopolyanions, as shown in Eq. (2): 7Mo0:-

+ 8Hf

+

Mo,O:i

+ 4H20

(2)

Acidic elements such as Mo, W, V, Nb and Ta, which are present as oxoanions I13 Copyright 0 1996 by Academic Press. Inc. All rights of reproduction in any form reserved.

114

TOSHlO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO

in aqueous solution, tend to polymerize by dehydration at low pH, forming polyanions and water (1-3). The term “heteropoly compound” is used in this review for the acid forms, e.g., H3PW12040.and their salts, e.g., Cs3PW1204~.Catalysts of which the main components are heteropoly or heteropoly-derived compounds are referred to here as “heteropoly catalysts,” and they are the subject of this review. Heteropolyanion-derived compounds are, for example, organic and metallo-organic complexes of polyanions (see Section 1.D for the terminology and nomenclature). Although there are many kinds of heteropolyanions (Section II), heteropolyanions having the Keggin structure are the most widely investigated as catalysts because of their stabilities and ease of synthesis. However, other heteropolyanions are also expected to be recognized as good catalysts. Heteropoly catalysts can be applied in various ways (4-1 0). They are used as acid as well as oxidation catalysts. They are used in various phases, as homogeneous liquids, in two-phase liquids (in phase-transfer catalysis), and in liquidsolid and in gas-solid combinations, etc. The liquid-solid and gas-solid combinations are represented by the classes of catalysis shown in Fig. 1 and described in the following sections. The advantages of heteropoly catalysts stem from the characteristics summarized in Table I. As excellent candidates for design at the atomic or molecular level, heteropoly catalysts have proven to be of value in fundamental studies as well as practical applications. But it is also true that much remains to be done. Efforts to establish methodologies for design of practical catalysts are still under way. The acid strength and acid site density can be controlled quite well both in solution and in the solid state, but the redox properties in the solid state are much less well understood because of the lack of sufficient thermal stability of mixedmetal (mixed-addenda) heteropolyanions. The acid strengths of some solid heteropolyacids have been suggested to reach the range of superacids, but they reactant 0

product -&

reactant product

0

4

\\ !/

f

product

(

.... .. .. .. .. . . . . . . . . . . . . .. .. .. .. .. .. .. .. .. .. .. .. .. . . . .. .. .. .. .. .. ....... .. .. .. .. .. . . . .. .. .. .. .. .. .. .. .. .

u

.............. .............

Solid (Surface catalysis)

reactant

Pseudoliquid (Bulk type I catalysis)

Solid (Bulk type I1 catalysis)

FIG. I . Three types of catalysis by heteropoly compounds.

CATALYTIC CHEMISTRY OF HETEROPOLY COMPOUNDS

Advaniages

115

TABLE I Heieropoly Caialysis

OJ

1. Catalvst design at aromic/molecular levels based on the following: 1-1. Acidic and redox properties

These two important properties for catalysis can be controlled by choosing appropriate constituent elements (type of polyanion, addenda atom, heteroatom. countercation, etc.). 1-2. Multifunctionality Acid-redox, acid-base, multi-electron transfer, photosensitivity, etc. 1-3. Tertiary structure, bulk-type behavior, etc., for solid state These are well controlled by countercations. 2. Moleculariiy-meial oxide clusier 2- I . Molecular design of catalysts 2-2. Cluster models of mixed oxide catalysts and of relationships between solid and solution catalysts 2-3. Description of catalytic processes at atomicimolecular levels Spectroscopic study and stoichiometry are realistic Model compounds of reaction intermediates.

3 . Unique reaciionfield 3- I. Bulk-type catalysis "Pseudoliquid" and bulk type I1 behavior provide unique three-dimensional reaction environments for catalysis. 3-2. Pseudoliquid behavior This makes spectroscopic and stoichiometric studies feasible and realistic. 3-3. Phase-transfer catalysis 3-4. Shape selectivity. 4. Unique basicity of polyanion 4-1. Selective coordination and stabilization of reaction intermediates in solution and in pseudoliquid phase, and possibly also on the surface 4-2. Ligands and supports for metals and organometallics.

are still weaker acids than sulfated zirconia. Unique complexing or basic properties of polyanions have not been clarified sufficiently, although it appears that they play important roles in industrial liquid-phase processes. The efforts to describe catalytic processes at the molecular level have also made significant progress in the preceding decade, but the number of well-elucidated reactions remains very small. Early attempts to use heteropoly compounds as catalysts are summarized in reviews published in 1952 (IZ) and 1978 (I). The first industrial process using a heteropoly catalyst was started up in 1972 for the hydration of propylene in the liquid phase. The essential role of the Keggin structure in a solid heteropoly catalyst was explicitly shown in 1975 in a patent concerning catalytic oxidation of methacrolein. Systematic research in heterogeneous catalysis with these materials started in the mid-1970s and led to the recognition of quantitative relationships between the acid or redox properties and catalytic performance

1 16

TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO

(4-9). Pseudoliquid-phase catalysis (bulk type I catalysis) was reported in 1979, and bulk type I1 behavior in 1983. In the 1980s, several new large-scale industrial processes started in Japan based on applications of heteropoly catalysts that had been described before (5, 6, 12): namely, oxidation of methacrolein (1982), hydration of isobutylene (1984), hydration of n-butene ( 1985), and polymerization of tetrahydrofuran (1987). In addition, there are a few small- to medium-scale processes (9, 10). Thus the level of research activity in heteropoly catalysis is very high and growing rapidly. One of the authors of this chapter has previously reviewed heterogeneous catalysis by heteropoly compounds (4-6). Catalysis in solution has also been described (7-10). In this chapter, we critically survey the literature and attempt to describe the essence of the catalytic chemistry of heteropoly compounds in solution and in the solid state. We have attempted to highlight the advantages of heteropoly catalysts as described in Table I.

B.

CLASSES OF CATALYSIS BY

HETEROPOLY COMPOUNDS

As will be described in more detail in later sections, in acid and oxidation catalysis by solid heteropoly compounds, that is, gas-solid and liquid-solid systems, there are three different classes of catalysis: (1) surface catalysis, (2) bulk type 1 (pseudoliquid catalysis), and (3) bulk type I1 catalysis, as shown in Fig. 1. The latter two have been specifically demonstrated for heteropoly catalysts, and they could be found for other solid catalysts as well. Surface-type catalysis is ordinary heterogeneous catalysis, whereby the reactions take place on the two-dimensional surface (on the outer surface and pore walls) of solid catalysts. The reaction rate is proportional to the catalyst surface area. Bulk type I catalysis was found in acid catalysis with the acid forms and some salts at relatively low temperatures. The reactant molecules are absorbed between the polyanions (not in a polyanion) in the ionic crystal by replacing water of crystallization or expanding the lattice, and reaction occurs there. The polyanion structure itself is usually intact. The solid behaves like a solution and the reaction medium is three-dimensional. This is called “pseudoliquid” catalysis (Sections 1.A and VI). The reaction rate is proportional to the volume of the catalyst in the ideal case; the rate of an acid-catalyzed reaction is proportional to the total number of acidic groups in the solid bulk. Bulk type 11 catalysis was discovered later for some oxidation reactions at high temperatures. Although the principal reaction may proceed on the surface, the whole solid bulk takes part in redox catalysis owing to the rapid migration into the bulk of redox carriers such as protons and electrons (Sections VII and IX). The rate is proportional to the volume of catalyst in the ideal case.


117

These three classes of catalysis are distinctly different from each other in the ideal cases. But the extent of the contribution of the inner bulk of the catalyst depends on the rate of the catalytic reaction relative to the rate of diffusion of reactant and product molecules in bulk type I catalysis and on the rate of reaction relative to the rate of diffusion of redox carriers for the bulk type I1 catalysis. c . CATALYST DESIGNBASEDON CRYSTALLINE MIXEDOXIDES

To develop efficient catalytic technology capable of solving contemporary problems related to energy and resource limitations, synthesis of materials, and environmental protection, novel concepts for catalyst design are needed. Catalyst design at the atomic level utilizing the techniques of advanced surface science is one of the possibilities; but this can be applied only for model catalysts, and the syntheses of industrial catalysts by this method are not yet realistic (6). Alternatively, we have attempted the molecular design of mixed-oxide catalysts by using crystalline mixed oxides whose bulk structures are known and whose potential for practical use is good. Heteropoly compounds, perovskites, and zeolites are the candidate catalysts. Since we believe that the relationships in Scheme 1 are useful for the design of catalysts (13), we place stress in this chapter on these relationships at atomic/ molecular levels of heteropoly compounds. In our opinion, sufficient care must be taken on the structure and stoichiometry in order to design catalysts taking advantage of the molecular nature of heteropoly compounds. D. TERMINOLOGY A N D NOMENCLATURE 1.

Generic Terms

Various generic names have been used for oxoacids and oxoanions. Because there are many of them, it is difficult to define the terms unambiguously and consistently. But the following statements may be helpful. For acid forms, polyacids = polyoxoacids, including heteropolyacids (e.g., H3PW12040) and isopolyacids (e.g., H2Mo6OI9); and for oxoanions, polyanions = polyoxoanions = polyoxometalates, including heteropolyanions (e.g., PWf20:O) and isopolyanions (e.g., Mo60:9). Performance of Catalyst

*

Chemiaal and Physical Properties

Composition

* Ekture

SCHEME I

Method of Synthesis

1 18


In addition, such generic terms as metal-oxygen cluster ion, metal-oxide molecule, etc., are used for polyanions (and polyacids). Since the traditional “heteropoly-” and “isopoly-” are unsatisfactory terms to express a variety of polyacids and polyanions, the terms “polyoxometalates” and “polyoxoacids” have been used recently (3). The terminology is still changing, thus reflecting the rapid expansion of the chemistry. Nonoxygen elements in the inner part of polyanions (usually P, Si, As, etc.) are called heteroatoms (in some cases, central atoms) and those in the peripheral part (usually Mo, W, V, Nb, etc.) are called addenda atoms or polyatoms (Section 1I.A). We use more or less conventional terminology here. 2. Nomenclature The rigorous and systematic nomenclature addressed by IUPAC ( 1 4 , in which all atoms and their topological connections are defined unambiguously, is too complicated here. Thus we use traditional names. But the semi-systematic nomenclature accepted by IUPAC (15) is mentioned briefly. For example, a heteropolyacid, H4SiMo12040,is called tetrahydrogen hexatriacontaoxo(tetraoxosilicato)dodecamolybdate(4-) [hydrogen nomenclature] or tetrahydrogen silicododecamolybdate [abbreviated semi-trivial name]. Or this is called 12-(or dodeca)molybdosilicicacid for the acid form and 12-(or dodeca)molybdosilicate for the anion [recommendations of IUPAC, 1971 ( I @ ] .

II. Structure, Synthesis, Stability, and Characterization A.

PRIMARY,

SECONDARY, AND TERTIARY STRUCTURES

Heteropolyanions and isopolyanions are polymeric oxoanions (polyoxometalates) (2, 3, 5, 6). The structure of a heteropolyanion or polyoxoanion molecule itself is called a “primary structure” (5, 6, 17). There are various kinds of polyoxoanion structure (Section 1I.A. 1). In solution, heteropolyanions are present in the unit of the primary structure, being coordinated with solvent molecules a n d or protonated. Most heteropolyanions tend to hydrolyze readily at high pH (Section 1I.C). Protonation and hydrolysis of the primary structure may be major structural concerns in solution catalysis. Heteropoly compounds in the solid state are ionic crystals (sometimes amorphous) consisting of large polyanions, cations, water of crystallization,and other molecules. This three-dimensional arrangement is called the “secondary structure.” For understanding catalysis by solid heteropoly compounds, it is important to distinguish between the primary structure and the secondary structure (5, 6, 17). Recently, it has been realized that, in addition

119


?d

P b

nn

."I& a@ .............._,

d a

..'& , .. ..............: ..' I.

I

cs+

PW,~O~,,~-

C

FIG. 2. Primary, secondary, and tertiary structures of heteropoly compounds. (a) Primary structure (Keggin structure, XM12040); (b) secondary structure (H3PW 12040. 6H20); (c) secondary structure (CS~PWI~O~O); (d) tertiary structure [Csz.sHosPWIz04O, cubic structure as in (c)].

to these structures, tertiary and higher-order structures influence the catalytic function (6). These structures are exemplified in Fig. 2 (3, 18). 1.

Primary Structure

a. Keggin Structure (1-3, 18, 19). Figures 3a and 3b illustrate the Keggin anions, which are the most popular heteropolyanions in catalysis. The ideal Kegging structure of the a type has G symmetry and consists of a central X 0 4 tetrahedron (X = heteroatom or central atom) surrounded by twelve M 0 6 octahedra (M = addenda atom). The twelve M06 octahedra comprise four groups of three edge-shared octahedra, the M3013 triplet (19), which have a common oxygen vertex connected to the central heteroatom. The oxygen atoms in this structure fall into four classes of symmetry-equivalent oxygens: X-0,-(M)3, M-Ob-M, connecting two M3013 units by comer sharing; M-0,-M, connecting two M3013 units by edge sharing; and Od-M, where M is the addenda atom and X the heteroatom. This Keggin structure is called an a isomer (18).

120


d

b

a

e

f

FIG. 3. Primary structures of heteropoly and isopolyanions. (a) Keggin structure, a-XM 120'& (the fourth M3013set and the X04 tetrahedron are not shown for clarity) (from Refs. 2 and 19); (b) Keggin structure, /?-XM120;; (the fourth M301) set and the XO4 tetrahedron are not shown for clarity) (from Refs. 24 and 25); (c) lacunary Keggin anion (the central XO4 tetrahedron is not shown) (from Ref. 2); (d) Dawson structure, X2M180gi (from Refs. 29 and 30); ( e ) Anderson structure, XM602/- (shaded tetrahedron indicates the heteroatom site) (from Refs. 18 and 33); (0 XMPO;Q (from Ref. 2); (g) isopolyanions, W,,O:; (from Ref. 2).

The known addenda and heteroatoms incorporated in heteropolyanions are summarized in Table I1 (20). The structures in this table of polyanions with Se(IV), Te(IV), Sb(III), Bi(III), Ti(IV), and Zr(IV) still need to be confirmed, since tetradendral coordination of these ions with oxide ions is seldom observed (2). In Table 111 (21, 22), the bond distances in various heteropoly corn ounds having the Keggin structure are listed. Bond lengths in PW120;for H3PW12040.6H20 are 1.71, 1.90, 1.91, and 2.44A for Od-W, 0,-W, Oh-W, and 0,-W bonds, respectively. The existence of isomers has been established for the Keggin anion. Figures 3a and 3b show the a-and /I-isomers. They can be separated by fractional crystallization (X = B, Si) or prepared separately (X = Si, Ge) (23). In /?-SiWlzO:,, one of the three edge-shared W3OI3 triplets of the a structure is rotated by 60°, thereby reducing the symmetry of the anion from to CJ, (24, 25). The other isomers involving the


121

TABLE I1 Known Addenda-(0) and Hetero-(0) Atoms Incorporated in Heteropolvacid.7

Pr Nb Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

Pu Am Cm Bk Cf Es Fm Md No Lr

60” rotation of two, three, and all four W3013 groups are called the y, 6 , and c structure (26), respectively. Compounds containing fluoride ions in metatungstate have been synthesized (27): (A) ( W I Z F Z O ~ R H Z(B) ) H ~(W12F3037HW4, , (C) ( W I ~ ~ ~ R F Z Hand ) H (D) S, ( W I ~ O ~ ~ F Hwhere ~ ) H the ~ , central atoms are protons. b. Lacunqv Keggin Anion (2, 28). In solution, several species are present in equilibrium, the composition depending on pH. Figure 4 shows an example of an aqueous solution containing MOO:- and HPOt- in a molar ratio of 12: 1 (2, 28). The reactions to form polyoxoanions other than Keggin anions are shown by Eqs. (3)-(5). 17H’ + I1MoO:- + HPOSP M o I , O : ~+ 9H20 (3)

-.

17H+ + 9M00:8H‘

+ HP0:-

+ 5Mo0:- + 2HPO:

+

+

PMO&I(OH~):P2MosO;;

+ 9H20

+ 5H20

(4) (5)

In the case of PW120:0, the lacunary P W I I O : ~is formed at pH = 2. The to give PW90i, occurs at pH > 8 (2). These lacunary degradation of PWI or defect derivatives of the Keggin structure are illustrated in Fig. 3. c. Dawson Structure. The Dawson structure, M2XIgOgZ, is shown in Fig. 3d (29, 30). Two PW90:, units, “lacunary Keggin anions,” fuse to form a Dawson structure. Three isomers exist, depending on the number of rotated

122


Bond Lengihs in

Compounds H3PW12040.6H2O H3PM01204.13H20 H3PM012040.30H20 H&M012040.13H20 &SiW12040. 16H20

Mob

TABLE 111 and Won Group in Heteropolyanion (A) (21, 22)

M-Od

M-0,.

M-oh

MhO,

XhO,

1.71 1.66 1.68 1.67 1.68

I .90 I .96 1.91 1.94 1.91

1.91 I .97 I .92 I .96 I .96

2.44 2.43 2.44 2.35 2.38

I .53 1.53 1.54 1.62 1.63

M3013groups. H6As2MoI8O62 and S2M0180:2 have also been synthesized (2, 31). A complex containing F ions, H Z F ~ N ~ W I ~(32), O : is ~ isostructural with the Dawson species, P2W180gY.

-

PH

FIG. 4. Distribution diagram for species present in fresh solutions containing MOO:- and HP0:- in a molar ratio of 12 : I at different pH values. (From Ref. 28.)


123

d. Anderson Structure. The Anderson structure, XM60;4, comprises seven edge-shared octahedra (Fig. 3e) (18, 33). 2. Secondary Structure of Solid Heteropoly Compounds In heteropoly acids (acid form) in the solid state, protons play an essential role in the structure of the crystal, by linking the neighboring heteropolyanions. Protons of crystalline H3PW12040. 6 H 2 0 are present in hydrated species, H50:, each of which links four neighboring heteropolyanions by hydrogen bonding to the terminal W-Od oxygen atoms, and the polyanions are packed in a bcc structure (Fig. 2b) (21). Various heteropolyacid hydrates that differ in the number of waters of crystallization have been reported (34-36); H3PW12040.nH20 (n = 14, 21, 24, and 29). The loss of water brings about changes in the anion packing; n = 29 [cubic (diamond-like)], n = 21 (orthorhombic), n = 6 [cubic (bcc)]. Cs3PW12040,in which the Cs ions are at the sites of H50: ions of hexahydrate (Fig. 2b), has a dense secondary structure and is O~~ anhydrous (Fig. 2c) (34). The lattice constants of C S J P W ~ ~and H3PW12040.6H20 are 12.14 and 11.86 A, respectively (21, 26, 37). Secondary structures containing organic molecules are known. H4SiW12040 * 9DMSO [DMSO = (CH3)2SO] contains nine molecules of DMSO in a unit cell, where there are weak hydrogen bonds between methyl groups and oxygen atoms of the heteropolyanion, p~lyanion--(CH~)~SO--H+ --OS(CH3)2--polyanion (38). Eight independent DMSO molecules join in four pairs of cations [H(O=S(CH3)2)2]+ by strong hydrogen bonds, and one DMSO molecule is weakly bonded. Another example is PW12040. [(C5H5N)2H]3,which is obtained by the reaction of anhydrous H3PW12040 with pyridine (39). As shown in Fig. 5 , six pyridine molecules lie almost in a plane, and the pyridine molecules

0 FIG. 5. Structure of [(C5HSN)2H],[PW12040]. (From Ref. 3Y.)

124


are paired, forming N--H--N hydrogen bonding. Similar structures are also formed by the contact of pyridine vapor with heteropolyacids (5). Other examples are PW12040 (H+-quinolin-8-01)3 4C2H50H* 2H20 (40) and PM012040[H+-TMU213(TMU = 1,1,3,3-tetramethyl urea) (41). 3.

Tertiary Structure of Solid Heteropoly Compounds Tertiary structure is the structure of solid heteropoly compounds as assembled

(5, 6). The size of the primary and secondary particles, pore structure, distribu-

tion of protons and cations, etc. are the elements of the tertiary structure. a. Group A and B Salts. Countercations greatly influence the tertiary structure of a heteropoly compound. The salts of small ions such as Na+ [classified into group A salts (42)] behave similarly to the acid form in several respects. The group A salts are highly soluble in water and other polar organic solvents. The surface areas of group A salts are usually low. Polar molecules are readily absorbed in interstitial positions (between polyanions) of the secondary structure (Section VI). On the other hand, the salts of large cations such as NH: and Cs+ (classified as group B salts) are insoluble in water and exhibit low absorptivity for polar molecules. Low solubility is due to the low energy of solvation of large cations. The surface areas of group B salts are usually high due to the smaller sizes of the primary particles, giving favorable properties for heterogeneous catalysis (43-46a). The thermal stability of most group B salts is relatively high, which is also important in heterogeneous catalysis. b. Surface Area and Pore Structure. The surface area and pore structure are closely related. Figure 6 shows the surface areas as a function of the extent 200

0 : cs

0

1

2

3

x in MxH3.xPW12040

as a function of the extent of Na or FIG.6. Surface areas of Na or Cs acidic salt of H3PW12040 Cs substitution. (From Refs. 46 and 47.)


125

of Na or Cs substitution of H3PW12040 (46, 47). As the Na content increases, the surface area decreases monotonically (46b).The change of the surface area with increasing Cs content is remarkably different. The surface area increases significantly when the Cs content, x in Cs,H3-,PW12040, changes from x = 2 (1 m2 g - ' ) to x = 3 (156 mz g-'), although it decreases slightly from x = 0 (6 m2 ! - I ) to x = 2. The surface area increases significantly to more than 130m g-I when the Cs content exceeds 2.5. Cs2.5Ho.5PWIz040 (and also C S ~ P W ~ consists ~ O ~ ~of) very fine particles (8-10 nm in diameter) (Fig. 2d). Pore structure is an important property of solid catalysts. Gregg and Tayyab ( 4 3 ) reported that (NH4)3PW1 2 0has 4 ~a microporous structure (pore diameter 20A) (44c) as well as micropores. They proposed that these pores exist in the crystal structure. Mizuno and Misono (37) examined the tertiary structure of C S ~ P W , ~byOestimating ~~ the surface area with three different methods: particle-size distribution measured by TEM (assuming spherical particles), the pore-size distribution measured by N2 adsorption (assuming cylindrical pores), and the BET equation. The three values are in good agreement with each other, showing that this material is composed of fine primary particles observed by TEM, with the pores being intercrystalline, not

0

0.5

1.o

PPO FIG.7. Nitrogen adsorption4esorption isotherms (77 K) for Cs3PW1204,,and (NH4)3PW12040, (From Ref. 44c.)

126

TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO TABLE IV Adsorption Data for C s , H j - , P W , ~ 0 4 0(x = 2.1. 2.2, and 2.5) (48)

Molecule c.s.'

Kinetic diameter

(A')

(A)

N2 (16.2) Benzene (30.5) Neopentane (37.2) 1,3,5-TMBd (41.1) I ,3,5-TIPBr (59.4)

3.6 5.9 6.2 1.5 8.5

Adsorption amount (Imol ti-7

137 (0.18) 21 (0.20) 99 (0.19) 0.6 (0.20) 9 x lo-' (0.20)

77 300 273 300 300

Cs2.1

Cs2.2

Cs2.5

Ratio"

487

861 124 179 I1 15

1648 232 390 237 236

0.52 0.53 0.46 0.05 0.06

10

5 -

-

"Cross section calculated from the molecular weight and density of liquid. hThe ratio of the partial pressure introduced (P) to the saturated vapor pressure (PO). 'Adsorption amount on Cs2.2 divided by that on Cs2.5. ,I 1,3,5-Trimethylbenzene. 'I ,3,5-Triisopropylbenzene.

intracrystalline. Considering the size and shape of the Keggin anion and the structure of C S ~ P (Fig. W ~2c),~there ~ ~are~no open pores in the crystal through which an N2 molecule (3.6 8, in diameter) can pass. W ~ ~ O ~ as ~ Cs2.2) is Recently, it was found that C S ~ . ~ H O . ~ P(abbreviated microporous and, according to adsorption experiments, has effective pores of about 7 A (48). Table IV is a comparison of the adsorption capacities of Cs2.2 (abbreviated as (32.5) for various molecules. Benzene and C S ~ , ~ H ~,2040 .~PW (kinetic diameter = 5.9 A) and neopentane (kinetic diameter = 6.2 A) are adsorbed on Cs2.2 and Cs2.5, and the relative adsorption capacity of Cs2.2 and Cs2.5 are similar to the corresponding ratio for N2 adsorption. On the other hand, both benzene and neopentane are little adsorbed on Cs2.1Ho.9PW12040 (Cs2. l), indicating that the effective pore size of Cs2.1 is less than 5.9 A. Of particular interest are the results observed with l13,5-trimethylbenzene (kinetic diameter = 7.5 8) and 1,3,5-triisopropylbenzene (kinetic diameter = 8.5 A). These two alkylbenzenes are adsorbed significantly on (32.5, but little on (32.2, so that the pore size of Cs2.2 is in the range of 6.2-7.5 A and that of Cs2.5 is larger than 8.5 8. This result demonstrates that the pore structure can be controlled by the substitution of H+ by Cs' (48). B. SYNTHESIS Heteropolyacids are prepared in solution by acidifying and heating in the appropriate pH range (I, 49-54). For example, 12-tungstophosphate is formed according to Eq. (1). Free acids are synthesized primarily by the following two methods: (1) by extraction with ether from acidified aqueous solutions and (2) by ion exchange from salts of heteropolyacids. Dawson-type heteropolyanions,


127

X2W180:2, are isolated as soluble ammonium or potassium salts or as free acids by extraction into ether (29, 30). Mixed addenda heteropolyanions with regiospecific substitution need careful preparation by use of lacunary heteropolyanions. If they are prepared from aqueous solutions of corresponding oxoanions, the products are usually mixtures of heteropolyanions having different compositions of addenda atoms. General procedures for the syntheses of various kinds of heteropolyacids are described in the literature (51-54).

C. STABILITY Particular attention should be paid to both the stability in solution and the thermal stability. The condesation-hydrolysis equilibria of heteropolyanions in aqueous media are shown in Fig. 8. Each heteropolyanion is stable only at pH values lower than the corresponding solid line (55). Some solid heteropolyacids are thermally stable and applicable in reactions with vapor-phase reactants conducted at high temperatures. The thermal stability is measured mainly by X-ray diffraction (XRD), thermal gravimetric analysis, and different thermal analysis (TG-DTA) experiments. According to Yamazoe et al. (56), the decomposition temperatures of H3PMo12040 and its salts depend on the kinds of cations: Ba2+, CoZf (673 K) < Cu2+, Ni2+ (683 K) < H', Cd2+ (693 K) < Ca2+, Mn2+ (700 K) < Mg2+ (710 K) < La3+, Ce3+ (730 K), where the

60 80 Cornposition/%

20 40

FIG.8. Stabilities of aqueous heteropolyacids (from Ref. 55): ( I ) PMoI20:{, (2) PW,,O:,, ( 3 ) GeMo120:o, (4) GeW120&-, (5) P2W180:2, ( 6 ) SiW120:i, (7) PMoIIO:9, (8) P2M0502;, (9) HzWi202Rr (10) P W I I O ; ~ .

128


decomposition temperature (in parentheses) is estimated by the exothermic peak in DTA. Herve et al. (57) investigated the thermal changes of structures by means of XRD and TG-DTA for Keggin-type heteropolyacids and proposed Scheme 2. Infrared spectroscopy of H4PMoI1VO40 showed the release of vanadium atoms to form and vanadium phosphate species (58). Exposure to water vapor induces the decomposition of the latter (indicated by the disappearance of a band at ca. 1037-1030 cm-I) (58). Results from TG and DTA show the presence of two types of water in heteropoly compounds, i.e., water of crystallization and “constitutional water molecules” (59). Loss of the former usually occurs at temperatures below 473 K. At temperatures exceeding 543 K for H3PMo12040or 623 K for H3PW12040,the constitutional water molecules (acidic protons bound to the oxygen of the polyanion) are lost. Data obtained by in situ XRD, 3 1P NMR, and thermoanalysis show that thermolysis of H3PMo12040 proceeds in two steps, as shown by Eq. (6) (60).

-

473 - 623 K

H3PMo12040-nH20

-

nH20

658 K

L

HxPM012038.5+x/2

-1.5H20

(x = 0.01)

H3PM012040

673 K

(PMo12038.5)~

Mo03(0.01P)

(6)

I>723K

+ (MbO2I2P2O7

The Moo3 phase appears at temperatures higher than 573 K. Thermal gravimetric analysis of H3PW12040and of C S ~ . ~ H O . ~ P showed W~~O~O that entire water molecules of crystallization are lost at temperatures as low as 573 K, and acidic groups are removed as water is formed from protons and lattice oxygens at temperatures exceeding 623 K. The numbers of protons lost, were 0.24, 0.31, and 0.32 after treatment at 623, x, in 673, and 773 K, respectively, whereas infrared spectra of Cs2.sHo.sPW1 2 0 4 0 remained unchanged at temperatures up to 773 K ( 6 1 ~ ) Similar . removal of protons of K~.5Ho.5PMo12040 begins by 500 K (61b).

D. CHARACTERIZATION OF HETEROPOLY COMPOUNDS 1 . Infrared Spectroscopy Infrared (IR) spectroscopy is a convenient and widely used method for the characterization of heteropolyanions. Keggin, Dawson, and lacunary heteropolyanions can be distinguished by their characteristic bands. a. Keggin Structure. In Table V (62-64), a partial list of the reported IR bands is given with their assignments. IR spectra of XWI2O;O, XMO,~O:~, and

129

CATALYTlC CHEMISTRY OF HETEROPOLY COMPOUNDS

curc

298 K

298 K

I

H4PMollV040- 13H20 triclinic T

H3PMo12040*13H20 triclinic T

I I H~PMo~~VO~O-~H~O

I I H3PM01204~7-8H20 333

- 353 K

333 - 353 K

unstable (cubic?)

unstable (cubic?)

I I

373

I I (0.5P205) 12 MOO3

723 K

undetected orthorhombic

-I353 K

I

H3PW120qty 6H20 cubic C 453 623 K

I

I

H3PMol 1vo40 tetragonal

t

333

-I

- I623 K

H3PMo12°40 tetragonal

H3PW12040 tetragonal

I

723 K

I

H3PW120q0.13H20 triclinic T

I I (0.5P205)

823 K

I

(0.5P205 + 0.5V205) t 12 Moo3 undetected orthorhombic

t 12 Wo3 undetected orthorhombic

SCHEME2

metatungstate are shown in Fig. 9 (62a).The vibrational mode of XO, is almost independent of the others for X = P, but it is mixed with other vibrational modes for X = Si, Ge, B, etc. (64).The X - 0 stretching bands (Si, 923 cm-'; Ge, 830 cm- I ; P, 1080 cm- I ) for XW,,O;, show higher frequencies than those TABLE V Infrared Absorption Bands of Heteropolyacids, em-' (62-64)

1080 982 893 812

926 980 878 779

I070 965 870 790

818 978 886 765

910 958 860 780

914 960 902 810

445 960 895 738

802 955 875 765

1091 962 914 780

v,,(X-0) v,,(W=O) v.,(W-0-W) v,,(W-O-W)

(X-0) v.,(Mo=O) v(Mo-0-Mo)

V",

130


FIG.9. Infrared spectra of XW12O;O and XMo120;U. (From Ref. 62a.) of XOl- anions, suggesting higher Ir-bond character of X - 0 bonds in the Keggin anion. Cation size influences the v(W-Od) frequencies for b-SiW120:O; the value of v(W-Od) decreases as the van der Waals radii of tetraalkylammonium cations increase (65). Due to the loss of hydrogen bonding between Od and water, the M-Od band (M = Mo, W) shifts to a higher frequency and the P-0 and M-Od peak intensities change (5). The W-Od ~ O ~into O doublets, suggesting a direct band for anhydrous C U ~ . ~ P W ~splits interaction between the polyanion and Cu2+ (66). In Table VI, IR bands of lacunary X M I 1 0 & are summarized (67).The P-0 stretching band for PWIIO:9 is split into 1085 and 1040 cm-’. It is believed


131

that the change of the symmetry from 5 (XMI2) to C,y(XMII) leads to the broadening of the band and sometimes causes bond splitting (67). The isomers, (Y- and P-PMo120:0 or SiW1203,, can also be distinguished by IR spectroscopy (68). The effects of solvent on the frequency have been reported (69). IR spectra of hydrated H3PW12040include a broad OH stretching band and two OH bending bands, at 1610 and 1720 cm- I . The latter two correspond to water and protonated water, respectively ( I ) . b. Dawson Structure (70). The IR spectrum of ( Y - P ~ W I ~resembles O~~ that of PW120:0. Three IR bands are observed in the PO4 stretching region, at 1090 (s), 1022 (w), and about 975 cm-I (sh) due to D3h symmetry of two PO4 groups. IR bands at 960, 912, and 780 cm- I are assigned to v(W-Oh), v(W-O,), and v(W-0b). respectively. 2. Raman Spectroscopy Vibrational frequencies in the Raman spectra of X M l l and XM12(X = Si, P; M = Mo, W) are summarized in Table VII (62, 65, 67, 71).The X - 0 vibration in Td symmetry of X04 is Raman-inactive. Among M - 0 bonds, M-Od is Raman-active. Raman spectra indicate that with an increase in pH, the structure of PMo120:i in aqueous solution changes, as shown by Eq. (7).

The states of hydrates of heteropolyacids are best distinguished by Raman spectroscopy, since Raman spectra give better resolved OH stretching bands than IR spectra (62, 72). H3PW12040.29H20has bands at 3570, 3525, 3490, 3450, 3205, and 3140 cm-', which correspond to 0-H distances of 2.94, 2.89, 2.86, 2.84, 2.72, and 2.65, respectively. The bands at 3205 and 3 140 cm- I are assigned to v(0H) of H(H2O);. The OH frequencies of H3PMo12040and H4SiW12040appear at 3570 and 3525 cm-', respectively. TABLE VII Raman Vibrational Frequencies of Keggin Anions (Aqueous Solutions), e m - ' (62. 65.67, 71)

pv,(M-OJ

shows polarized vibration.

132


3. NMR Spectroscopy a. Solid-state ' H NMR Spectroscopy. The broad-line NMR spectrum of H3PMo12040 29H20 shows that H30f and H20 become indistinguishable at temperatures greater than 298 K (73, 74). The NMR spectrum of anhydrous H3PMo12040 includes a narrow resonance (width PMol20;0 > PVl2O:0 parallels the decrease in the IR frequency of PO4, with both reflecting P - 0 n-bonding character. Five resonances observed for H5PV2M010040 correspond to the possible five isomers, in which the locations of two vanadium atoms are different (80). P M O ~ W ~ Oin: ~aqueous solution gives 13 peaks expected from the statistical distribution of P W 1 2 - x M ~ x 0 4(x0 = 0-12) (81). Figure 11 shows the solid-state 31 P NMR spectrum of Cs3PMollW040 (82) prepared from mixed aqueous solutions of H3PMo12040, H3PW12040, and CsN03. Five resonances agree well with the spectrum observed from the reaction mixture in solution and are assigned to a statistical mixture of (x = 0-4).M3PMo12040 (M = H, Na, K, Cs, and NH4) give similar chemical shifts. The 31PNMR chemical shift is greatly dependent on n in H3PW12040*nH20 (47, 77), the values being - 15.1 to - 15.6ppm for n = 6 and - 11.1 to - 10.5 ppm for n = 0. This difference is explained as follows: In the former, protonated water, H(H20):, is connected with the heteropolyanion by

134


0 FIG. 1 I .

5 Chemical ShiWppm

10

"P MAS NMR spectrum of C S ~ P M O ~ ~ W (From O ~ ~Ref. . 82.)

hydrogen-bonding at terminal oxygens, and in the latter, protons are directly attached to oxygen atoms of the polyanion. c. I83 W, "Mo, and *'Si NMR Spectroscopy. 183W Chemical shifts are sensitive to the heteroatom (83), being - 130.4 ppm for BWl20;0, -111.3ppm for H2WI20:O, -103.8ppm for SiW120:0, and -98.8ppm for PWl20:O. SiWIIO!J has five pairs of structurally identical W atoms and one unique W atom. Accordingly, SiW, gives six clearly separated resonances having the intensity ratio of 2 : 1.6 : 1 : 2 : 2 : 2, which is close to the expected ratio of 2 :2 : 1 : 2 :2 : 2. Similarly, P2Wl&ir which has twelve equivalent W atoms around its belt and six equivalent W atoms capping its ends (26), has two resonances; the larger one at - 170.14 ppm and the smaller doublet at - 124.87 ppm. The 183WNMR spectrum of 6-electron reduced cr-SiW120:0 shows three narrow resonances with the intensity ratio 1 : 1 : 2 (Fig. 12) (84). Two (3W, 6W)


135

have chemical shifts very close to that in the oxidized state and assigned to W(VI), whereas the third ( - 1500 ppm) is attributed to a W(IV) anion. This assignment is supported by the chemical shift of the W(IV) cation, w~o~(H~o):+. It is suggested that the protonated w6(v)w6(vI) polyanion undergoes intramolecular disproportionation to give W3(IV)W9(VI) species. The structures of positional isomers, a-1,2-XV2Wl~O:0, and a-l,2,3XV3W90kngf1)- (X = Si, P), are established from the 2-D connectivity pattern and 1-D spectra of 183W(85). The spectra of XV3W90:; (X = Si, P) are exclusively of the a-1,2,3-isomers. a- And p-SiWlzO:O can be also distinguished by 183WNMR spectroscopy; a-SiW120:0, gives one singlet (12 equivalent W atoms), and b-SiWI2O:[ three resonances in the ratio 1 : 2 : 1 (86). d. 170 NME Spectroscopy. 170 NMR spectroscopy gives information about the bonding nature of oxygen atoms. Figure 13 shows the assignments of the chemical shifts (87-89). There is a correlation between the downfield shift and the decreasing number of metal atoms to which the oxygen atom is bonded. The chemical shifts for the SiMoI20:0, SiW120:6, PMo120;;, and PW120:O are summarized in Table IX. 4. Electronic Spectra

Electronic absorption spectra give information about the electronic states of heteropolyanions (2). PWI20;; shows absorption at about 38,000 cm- I , due to

-1000 -I

FIG. 13. Assignments of " 0 NMR shifts due to various salts of mixed oxides. (From Refs. 87-8Y. )

136

TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO TABLE IX 0 NMR Spectral Data for Heteropolyanions (87-89)

17

Chemical shift (ppm) Anions

M-Od

SiMo120!; S~W 12~:” SiMoWIIO:; PMo120:R Pwt2Ok-

M-0h-M.

92 8 76 I 929(Mo-O~) 726(W-Od) 936 769

M-Oc-M

M-O,-(M),

580, 555 427,405 504,469

41 21 27

583, 550 431,405

78 -

ligand (oxygen) to metal (W) charge transfer (LMCT) (90). In most cases of XW’ Wl (one-electron reduction heteropoly blues), three absorption bands are observed at 8000-10,000 cm-l (band A), 13,000-16,000 cm-’ (band B), and ca. 20,000 cm-l (band C) in addition to that at 38,000 cm-I. The bands B and A are the results of intervalence charge-transfer transitions (IVCT) between metal atoms. Two types of transition are possible for reduced Keggin anions, that is, within an edge-shared group of M 0 6 octahedra (“intra” transition) and those between metals atoms linked by corner sharing. The band A has tentatively been assigned to an “intra” and the band B to an “extra” transition ( 2 ) . The UV band positions of various heteropolyacids are summarized in Table x (9&93). +

TABLE X UV Absorption Bands of Heteropolyanions Compounds

Absorption“ (kK)

Half-wave potentialh (V)

38.0, 50.0 38. I 37.8 39.0 38.3 34.0, 39.0 3 1.O, 47.0 32.3, 47.2 20.0, 25.0 20.2

- 0.023 -0.187 - 0.349 - 0.520 -0.510 + 0.02“ - 0.55“

See Refs. 9&93. Pope, M. T., and Varga, G. M., Inorg. Chem. 5, 1249 (1966); Prados, R. A., and Pope, M. T., Inorg. Chem. IS, 2547 (1976).


137

5 . Others a. Extended X-Ray Absorption Fine Structure (EXAFS) Spectroscopy. EXAFS spectra of PMoI20:; and PW120:; as amine salts have been measured (94, 95). PW120:; gives distinct peaks due to W-Od ( 0 . 1 6 ~ )and W-0-W (0.196 nm), whereas Mo-0 peaks of PMo120:i are not discernible. For the latter, a large Debye-Waller factor for Mo-0 bonds and also multiple scattering effects are presumed.

b. Scanning Tunneling Microscopy (STM) and Transmission Electron Microscopy (TEM). H3PWI2O4~ deposited on freshly cleaved and highly oriented pyrolytic graphite in air has been measured by STM (96). A fairly regular periodic pattern is observed, suggesting that the individual heteropoly species were directly imaged. Individual anions of the Dawson-type cyclopentadienyl titanium (CpTi) heteropoly compound, K7(C5H5)TiP2W17061 , were observed by TEM (97).The size of the anion is estimated to be 1 .&IS nm. This is consistent with a size determined from X-ray crystallographic data indicating an ellipsoid of about 1.0 X 1.5 nm. c. Electron Spin Resonance (ESR) Spectroscopy. ESR spectra give information about mixed-valence structures of reduced heteropoly compounds. ESR data for Keggin-type molybdates reduced by one electron (1e- reduction) in solution are shown in Table XI (2, 92, 98, 99). A Mo5+ signal with hyperfine at temstructure due to Mo5+ ( I = 5/2) is observed for PMo5+Mo7:O:; peratures less than 40 K (99). At higher temperatures line broadening is observed, and the hyperfine structure disappears. At room temperature, no Mo5+ signal is observed. These results indicate the rapid hopping of an electron among TABLE XI ESR Parameters of Some Reduced Polymolybdates (from ReJ 2) Anion 1.916 1.917 1.938 1.913 1.935 I .93 1 1.914 I .935

1.930 1.924 1.949 1.939 1.948 1.944 1.93 I 1.951

138

TOSHIOOKUHARA, NORITAKAMIZUNO, AND MAKOTO MISONO

I2 equivalent Mo atoms in a Keggin anion at higher temperature, and the mixed valence behavior is classified as class I1 (99, 100). The extent of electron delocalization increases as the number of molybdenum atoms in a Keggin anion increases. The order of extent of electron hopping is estimated to be PM0120:; > GeMo120g> Mo60:, (99). A similar electron hopping is estimated by the very weak signal intensity of MoS+ for solid H3+xPM~1204~ reduced by H2 at 423 K or room temperature (101, 102). Upon 2e- reduction, about 60% and 100% of the electrons are paired for SiMo120:0 and PWIZO:U, respectively, and the latter shows diamagnetism (2). Solid H3PMo12040 reduced by H2 at a lower temperature shows a very weak ESR signal intensity of Mo", probably because most of the Mo5+ ions are not detectable due to the rapid hopping of electrons. Heat treatment, which eliminates oxide ions from the heteropoly anion, leads to development of the Mo5+ signal, indicating the localization of electrons (101, 102). Early reports of ESR are likely due to these species. Several different spectra of reduced H3PMo12040 species are observed in highly reduced samples. The extent of reduction of H3PM0120~during the oxidation of methacrolein has also been investigated by application of ESR spectroscopy for detection of MoS+ (103). The states of V in the mixed-valence Keggin anion and Cu countercation were also investigated. The results show that more reducible countercations or addenda atoms such as Cu2+ and V5+ are reduced first, and an electron is localized on them (104-106). d. X-Ray Photoelectron Spectroscopy (XPS)). XPS gives information about electron density of solid heteropoly compounds (107-115). It was reported that Mo5+ was fairly uniformly present throughout the bulk of H3PMo12040formed by H2 reduction, whereas the surface was preferentially reduced by cyclohexane (101). The oxidation states of countercations Pd, Zn, and Cu were investigated by XPS in relation to redox properties and oxidation catalysis. Cu2+ and Pd2+, having higher electron affinities than Mo6+,in 12-molybdophosphatesare easily reduced by H2 and act as electron reservoirs (107-109). The following observations have been made for bulk heteropoly compounds. 1. The binding energy of 0 1 s electrons decreases with increasing negative charge of the heteropolyanion: PW12O:O > SiW120:; 2 BW120:; > H2WI20:; > PWIIO& > SiWl1O;9 > H2WI20::- (113). 2. For isostructural heteropolyanions, the 01s binding energy is higher for tungstates than for molybdates (I 13). 3. In PMo120:0 and SiMo120:U (2e- reduced state), two Mo atoms are in the oxidation state of 5 and ten Mo atoms are in the + 6 state (113, 114). A similar result was obtained for 2e--reduced PW120:; (113, 114).

+

139


4. A good correlation exists between the acid strength and the difference between 01s and W4f binding energies of HjFWl2040 and its salts (115).

111. A.

Acidic Properties ACIDITYIN

SOLUTION

Typical heteropolyacids having the Keggin structure, such as t -,PW12040and H4SiW 1 ~ 0 4 0are , strong acids; protons are dissociated completely from the structures in aqueous solution (8, 116). The dissociation constants, pK,, of heteropolyacids depend on the solvent. These are summarized in Table XU, together with pK, values for mineral acids ( I 17-1 19). Heteropolyacids are much stronger acids than H2S04, HBr, HCl, HN03, and HC104. For example, in acetic acid, the acid strength of H3PW12040is greater than that of H2S04 by about 2 pK, units. In acetic acid, which is less polar than water, heteropolyacids behave as relatively weak 1-1 electrolytes. The effect of solvent is evident for mixed solvents (120); as the concentration of water in aqueous acetic acid changes from 100 vol% to 4 vol%, the Hammett acidity function, Ho, for H3PW12040(0.1 M) decreases from 0.01 to - 1.78, TABLE XI1 Acid Constants of Heteropolvacids in Nonaqueous Media at 298 K ( I 1 6 1 19) Acetone Acid

PKI

PK2

Ethanol pK3

PKI

pK2

pK3

(3.0)

4.1

-

-

Acetic acid PKI

~~

(1.6) (1.8) -

(2.0) (2.0) (2.1) (2.1) (2.1) . -

(3.0) (3.2)

3.98 4.37

(1.6) -

-

-

-

-

-

3.61 3.62 3.69 3.73 3.90

(5.3) (5.3) (5.5) (5.6) (5.9)

(2.0) (1.8) (1.9) (1.9)

3.96 3.41 3.77 3.74

(6.3) (5.1) (5.9) (5.8) -

-

-

-

-

-

-

4.77 4.74 4.78 4.91 4.70 -

4.68 4.78 4.25

140


indicating an increase in acid strength. In the case of concentrated aqueous solution of H3PW12040,the acidity hnction is less than that of H2SO4 by 1-1.5 units (120). Titration curves indicate that the three protons of H3PW1204"are equivalent in aqueous solution. Three protons of H3PW 1 2 0 4 0 dissociate independently in acetic acid, as measured by 13C NMR spectroscopy (121). The greater acid strength of heteropolyacids than that of mineral acids is explained as follows (122). Since in heteropolyanions the negative charge of similar value is spread over much larger anions than those formed from mineral acids, the electrostatic interaction between proton and anion is much less for heteropolyacids than for mineral acids. An additional important factor is possibly the dynamic delocalizability of the charge or electron. The change in the electronic charge caused by deprotonation may be spread over the entire polyanion unit. As for the acid strengths of heteropolyacids, the following order has been reported for the compound in acetone (Table XI) (116): H3PW12040> H4SiW12040 = H3PMo12040 > H4PMoIIVO40 > H4SiMo12040.The acid strength decreases when W is replaced by Mo or V and when the central P atom is replaced by Si. The effect of the central atom has been demonstrated for acetonitrile solutions of Keggin-type heteropolytungstates. As shown in Fig. 14, the acidity increases in general with a decrease in the negative charge of the heteropolyanion, or an increase in the valence of the central atom (the valence of the central atom increases in the order Co < B < Si, Ge < P) (63). This order is reasonable, since the sizes of the polyanions are nearly the same; thus the interaction between the proton and the polyanion would decrease as the negative charge of polyanion decreases. 4

3

r" 2

1 2 3 4 5 6 7 Negative charge of polyanion

FIG.14. Values of Hammett acidity function (Ho) of H,XWt204o as a function of the negative charge of the polyanion. (From Ref. 63.)


141

Pope et al. (123) measured the formation constant of the 1 : 1 complex of 1,l -dihydroxyl-2,2,2-trichloroethane(chloral hydrate) and polyanions in nitrobenzene by using NMR spectrometry [Eq. @)]: ,0--H,

C13CCH,/o-H + A-

C13CCH:

0-H

“O--

‘:A-

(8)

H’

The formation constants (in parentheses) of the complexes are as follows: PW120:o (1.30) < PMo120:0 (3.1 1) < SiMo1204, (24.7). These values represent the capability of the heteropolyanion to form hydrogen bonds. Thus the acid strength is in the order H3PW12040 > H3PMo12040 > H4SiMo12040.Izumi et al. (124) obtained the following order in acid strength by the same technique: H3PW12040 > H3PM012040 > H4SiW12040= H4GeW12040> H4SiMoI2O4,,> H4GeM012040. These orders are in genera! agreement with those obtained (118, 119) with indicator tests (Table XII). In addition to the acidity, the softness of the heteropolyanion is an important characteristic relevant in catalysis (124). The softness has been estimated by the equilibrium constant in aqueous solution of the following reaction [Eq. (9)] at 298 K. AgnX + nNa1

* nAgl + Na,X

(X

=

polyanion)

(9)

The order of softness was found to be the following: SiWl2O4, > GeWl20:; > PWl20:O > PMo120iU > SiMo120d, > SO:-. The softness greatly influences the catalytic behavior in concentrated aqueous solution or in organic solution, as described below.

B. ACIDITYI N

THE SOLID STATE

1. Acid Forms The strength and the number of acid centers as well as related properties of heteropolyacids can be controlled by the structure and composition of heteropolyanions, the extent of hydration, the type of support, the thermal pretreatment, etc. Solid heteropolyacids such as H3PW12040and H3PM012040 are pure Br~insted acids and are stronger than conventional solid acids such as Si02-A1203 (45, 125). According to an indicator test, H3PW12040 has a Hammett acidity function less than - 8.2 (126), and it has even been suggested to be a superacid (127, 128). A superacid is an acid with a strength greater than that of 100% H2SO4, i.e., a value of If0 < - 12 (129). Thermal desorption of basic molecules also reveals the acidic properties. Figure 15 compares the acid strengths of heteropolyacids and Si02-A1203 (125). Pyridine adsorbed on Si02-A1203 is

142

TOSHIO OKUHARA, NORITAKA MIZLTNO, A N D MAKOTO MISONO

.ss 4

%' 1

0

300

400 500 TemperatureK

600

FIG.15. Thermal desorption of pyridine from heteropolyacids (from Ref. 125). H3PW12 refers to WWi2040.

completely desorbed at 573 K. On the other hand, sorbed pyridine (see Section VI) in H3PW12040mostly remains at 573 K, indicating that H3PW12040is a very strong acid. The acid strength can also be demonstrated by temperature-programmed desorption (TPD) of NH3 (Fig. 16) (127). H3PW12040 gives a relatively sharp peak at about 800K, together with peaks for N2 and H20 at the same temperatures (not shown). N2 is formed by the reaction of NH3 with oxygen atoms in the heteropolyanions. NH3 adsorbed on Si02-A1203 and on HZSM-5 is mostly desorbed at temperatures less than 800 K (127, 130). Thus according to TPD results for NH3, the order of the acid strength is sulfated zirconia (SO;-/ZtQ) > H3PW12040 > HZSM-5 > SiO2-AI2O3. The temperature of NH3 desorption from (or decomposition of) ammonium salts of heteropolyacids is in the following order: (NH&PW1204o > (NH4)4SiW12040> (NH&PM012040 > (NH414SiMoI 2 0 4 0 (131). The strengths of heteropolytungstic acids have been determined more quantitatively by calorimetry of NH3 absorption (132, 133). Figure 17 shows the differential heats of NH3 absorption after treatment at 423 K. The initial heats of absorption of NH3 are as follows: 196 kJ mol-I for H3PW12040; 185 kJ mol-' for H4SiW12040; 164 kJ rno1-I for H6P2W210,1(H20)3; and 156 kl rno1-I for H6P~W18062. These values show that the order of acid strength is H3PW12040 > H4SiW12040> H6P2W21071(H20)3 > H6P2W18062, in afleement with the statements above. The data also indicate that the Keggin-type heteropolyacids are much stronger acids than Dawson-type heteropolyacids (132).

143

CATALYTIC CHEMISTRY OF HETEROPOLY COMPOUNDS 21

373

'

'

573

'

'

I

a

'

773 973 TernperaturdK

1173

~ ~H3PW120m. O~U. FIG.16. TPD profiles of NH3 from various solid acids: (a) C S ~ . S H ~ . ~ P W(b) (c) SO:-/Zr02, (d) SiO2-AI2O3. (e) H-ZSM-5. Solid line: NH3 (m/e = 17); dotted line: N2 (mle = 28). (From Ref. 127.)

Heats of NH3 absorption further confirm that the heteropolyacids are stronger acids than zeolites or simple metal oxides: 150 kJ mol- for HZSM-5 and 140 kJ mol-' for y-A1203(133). Table XI11 shows the strengths measured by Hammett indicators with pK, values ranging from -5.6 to - 14.5. As described above, dried H3PW12040 possesses superacidity (127). The order of the acid strengths agrees with that

'

r

200

E al c .- 100 -5 L

C

f

E

n 0 0

200

400

600

800

loo0

1200

pmol of NH3 9-1

FIG. 17. Differential heat of NH3 absorption as a hnction o f the ammonia uptake for heteropolyacids. (From Ref. 132.)

144

TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO TABLE XI11 Acid Strengths Measured by Hammett Indicators (127) pK. of indicatorh

Catalyst" C~2.sHo.sPWi20a(573 K) HiPWi2040 (573 K) HZSM-5 (808 K) SO: 12102 (643 K) SiO2-AI2O3(723 K) ~

-5.6

-8.2

-11.35

-12.70

-13.16

+

+ + + + +

+ + + + +

+ + + + +

+ + +

+ + + 4-

-

-13.75

-14.52

-

-

-

-

-

-

+-

+

-

-

a The figures in parentheses are the pretreatment temperatures. Acidic color of the indicator was observed ( + ), or not ( - ). Indicators: - 5.6. benzalacetophenone; - 8.2, anthraquinone; - I I .35. p-nitrotoluene; - 12.70, p-nitrochlorobenzene; - 13.16, m-nitrochlorobenzene; - 13.75. 2,4-dinitrotoluene; - 14.52. 2.4-dinitrofluorobenzene.

estimated from NH3 TPD (127). In Fig. 18, various superacids in both solution and solid are summarized with reference to the Ho function (128, 129). H3PW12040is stronger than Nafion (Ho= - 12), but weaker than SO:-/ZrOz and AIC13-CuS04.

FIG. 18. Acid strengths of liquid and solid superacids. (From Ref. 128.)


145

Viswanathan et al. (115) reported small differences in binding energy between 01s and W4f: H3PW12040 (495.5 eV) > All.5PW12040(495.4 eV) > Na3PW12040(495.1 eV). This is the same order as the order of acid strength: > Na3PW12040. H ~ P W I Z O>~ A11.5PWl2040 O The locations of protons in solid heteropolyacids have been studied by NMR and IR spectroscopies. The protons of H3PW12040 with a high water content exist in the form of dioxonium ions H50: or H+(HzO), (21). Data obtained by 17 0 NMR spectroscopy confirm that the bridging oxygen atoms are protonated in CH3CN solution (88). In the solid state of dehydrated H3PW12040, there are two possible sites for protonation: the terminal W=O oxygen atoms or the bridging W-0-W oxygen atoms. Kozhevnikov et al. (134) assumed on the basis of an 170 NMR study that the terminal W = O oxygen atoms are the predominant protonation sites. The resonance of the terminal W=O oxygen atom shifts upfield 60 ppm, relative to that of the solution spectrum, whereas no such large shift was observed for the bridging W-0-W oxygens. On the basis of IR band broadening of the W-0,-W band for anhydrous H3PW12040 and D3PW12040,Lee et al. (78) concluded that protons are located on the bridging oxygens of PWI2O:,j-. In contrast, there was no difference observed for the W-Od band. As described above, the most basic oxygen is the bridging oxygen according to 170 NMR spectra of the solution. It was reported in a crystallographic study that reduced PMo12040 is protonated at bridging oxygen ( 135). Quantum-chemical calculations at the EHMO level (136) and bond-lengthhond-strength correlations (13 7) indicated that the bridging oxygen atoms carry the highest electron density (are the most basic).

2 . Salts Acidic properties of heteropoly compounds in the solid state are sensitive to countercations, constituent elements of polyanions, and tertiary structure. Partial hydrolysis and inhomogeneity of composition brought about during preparation are also important in governing the acidic properties. There are several possible types of origins of acidity (5): 1. Dissociation of coordinated water: for example, Ni(H20)b+ +Ni(H20),,,- ,(OH)+ + H i

2. Lewis acidity of metal ions. 3. Protons formed by the reduction of metal ions: for example, Ag'

+ tH2+Ago + Hf

4. Protons present in the acidic salts: for example, Cs,H3 -.rPW12040. 5. Partial hydrolysis during the preparation process: for example, P W 1 2 0 g-PW,,O:C

+ W0:- + 6H'

146


As for type ( I ) , Niiyama et al. (42) proposed that protons are generated by dissociation of water and that the equilibrium of the dissociation is a function of the electronegativity of the metal cations. Formation of Brensted acid sites in the aluminum salt of H3PW12040as a result of exposure to water vapor at 573 K was confirmed by IR spectra of sorbed pyridine (138). The presence of Lewis acidity [type (2)] or Brensted acidity [type ( l ) ] is revealed by the IR spectrum of pyridine sorbed on All,5PW12040 (138, 139). IR spectra of sorbed of NH3 show that C U ~ . ~ P W Ihas ~ OLewis ~ O acidity as well as Brensted acidity (140). Ghosh and Moffat (141) measured the acidities of several salts of H3PW12040 with Hammett indicators. The acid strength increases with an increase in the calculated charge on the peripheral oxygen atom of the polyanion: Zr > A1 > Zn > Mg > Ca > Na (141). Proton formation from H2 [type (3)] has been demonstrated by IR and NMR spectroscopies for Ag3PW12040(142). Ag3PW 1 2 0 4 0 treated with hydrogen at 573 K and then pyridine gives IR bands of pyridinium ion. Such bands are not observed for Ag3PW12040which had simply been evacuated at 573 K prior to sorption of pyridine. The former sample (in the absence of pyridine) was catalytically much more active than the latter. As shown in Fig. 19, the introduction of H2 to Ag3PW12040at 488 K leads to three resonances at 9.3, 6.4, and 4.1 ppm in the 'H NMR spectrum (143-145). The resonance at 4.1 ppm is due to H20 molecules formed by the hydrogen treatment. That at 6.4 ppm increased as the H2 pressure increased and disappeared as a result of the addition of deuterated pyridine (C5D5N) at 303 K. The peak at 9.3 ppm was nearly unchanged during the experiments. These results indicate that the acid strength of the proton-containing groups indicated by the resonance at 6.4 ppm is higher than that of the groups indicated by the resonance at 9.3 ppm. This difference in acid strengths is attributed to the differences in the interaction between proton and polyanion. Group B salts, incorporating acid groups, have high surface areas and high thermal stabilities. In TPD of NH3, gives a desorption peak in the temperature range 633-923 K, and the peak temperature is close to that for H3PWIZ040,indicating that the acid strength of Cs2.5Ho.5PW12040 is similar to that of H3PW12040(127); C S ~ . ~ H ~ has . ~a slightly P W ~broader ~ ~ ~distribution ~ of acid strengths than H3PW12040 (Fig. 16). In group A salts, the acid amount exceeds the number of protons expected from the formula for Na salts of H3PW12040. Partial hydrolysis has been proposed [type ( 5 ) ] (46a). Solid-state 31PNMR spectroscopy was used to demonstrate that all protons of C S ~ . S H O . ~ P Ware ~ ~distributed O~O randomly through the bulk of the material. Figure 20 is a 31P NMR spectrum for C S ~ . ~ H ~ . ~ P(47). W ~ ~Special O ~ O care was necessary in the measurement to protect the sample from moisture. Four resonances ( - 14.9, - 13.5, - 12.1, and - 10.9 ppm) appeared when


147

. 6.4

I

I

30

20

10

0

-10

Shift I ppm FIG. 19. ' H NMR spectra (room temperature) in reduced Ag3PW12040.(a) Ag3PW12040reduced with 40 kPa of H2 at 488 K and recorded in the presence o f H2; (b-g) recorded after evacuating H l at 77 (b), 303 (c), 333 (d), 488 (e), 523 (0 and 623 K (g). Reprinted with permission from Ref. 143. Copyright 1993 American Chemical Society.

Cs~.5Ho.sPW12040 was evacuated at 573 K. Cs3PW12O40, which has no proton, gives a resonance at - 15.3 ppm, which is close to the first resonance ( - 14.9 ppm). Anhydrous H3PW12040,which has three protons directly attached to the polyanion, has a resonance at - 10.9 ppm. Thus the chemical shift is determined by the number of the protons attached to the polyanion, and the four

148


FIG. 20. "P NMR spectrum (bottom) and one of the primary particles (top) of Csz.sHo.sPWlzOao. The dotted lines in the spectrum show the relative peak intensities expected for the statistical distribution of protons. (From Ref. 47.)

peaks correspond to the Keggin anions having 0, 1, 2, and 3 protons, respectively. The relative intensities of these peaks for C S ~ . ~ H ~ . ~are P in W good ~ ~ O ~ ~ agreement with those expected from a random distribution of protons (broken line, Fig. 20). The number of protons on the surface (shown in Fig. 21) was estimated by multiplying the formal concentration of protons on the surface by the quantity of polyanions present on the surface calculated from the surface area. As the Cs content increases, the number of surface protons at first decreases, but greatly increases when the Cs content exceeds 2 . The increase in the number of surface protons is mainly due to the sharp increase in the surface area (47). When x

x in C~,H3.~PW12040

FIG.21. The amount of surface protons (surface acidity) as a function of Cs content for C S , H I - , P W , ~ O ~(From ~ . Refs. 47, 48.)


149

increases from 2.5 to 3.0, since the formal concentrations of protons become low or zero, the number of surface protons decreases greatly. As described in Section V, the catalytic activity is in close correspondence to this surface acidity.

3 . Supported Heteropolyacids Supporting heteropolyacids on solids with high surface areas is a useful method for improving catalytic performance. It is necessary to pay attention to the changes in the acid strength, the structures of the aggregates, and the possibility of decomposition. In general, strong interactions of heteropolyacids with supports are observed at low loadings and the intrinsic properties of heteropolyacids prevail at high loadings. Basic solids such as A1203 and MgO tend to decompose heteropolyacids ( 146-1 49). Microcalorimetry of NH3 absorption (adsorption) reveals that when H3PW12040is supported on S O 2 , the acid strength decreases, as shown in Fig. 22 (150). The acid strength of H3PW12040diminishes in the sequence of supports Si02 > A1203 > activated charcoal. The acid strengths of heteropolyacids supported on Si02 (151) measured by NH3 TPD is in the following order: H3PW12040(865 K) > H4SiW12040(805 K) > H3PMo12040(736 K) > H4SiM~12040 (696 K), where the figures in parentheses are the temperatures of desorption. The interaction between H3PW12040 and the surface OH groups of Si02 has been detected by ' H and "P MAS NMR spectroscopies (152). The OH resonance of Si02 at 1.8 ppm (relative to tetramethylsilane) becomes smaller as the amount of H3PW12040on Si02 increases. At 20-50 wt% H3PW12040,a new type of proton appears at 5 ppm. At loadings >50 wt%, a resonance at 9.3 ppm

170

-

150

T

130

3 110 I

g

90

70

0

0.5

1.o

mmOl Of NH3 g 1

FIG.22. Differential heats ofabsorption of NH, on ( 1 ) H3PW12040and (2) 20 wt% H,PWl,O,$ SiO2. (From Ref. 150.)

150

TOSHIO OKUHARA, NORITAKA MIZUNO, A N D MAKOTO MlSONO

due to crystalline H3PW12040appears. At 20 wt%, 31P NMR also gives a broad resonance at - 15.8 ppm, which is different from the - 12.4 ppm resonance of crystalline H3PW12040.The interaction between H3PWI2O40and OH groups on SiOz is assumed to be as shown in Eq. (10). H3PW12040+ mOH-Si

+

H3-,,,PW12040-Si

+ mH2O

(10)

Lefebvre (153) reported for 1 3 4 7 % H3PW12040 supported on Si02 that two 31P resonances are due to the bulk H3PW12040 (15.1 ppm) and to SOH:H2PWI2O4&,( - 14.5 ppm). 31PSpin-lattice relaxation of H3PMo12040/Si02gives 155). For unsupported an estimated coverage of H ~ P M O on ~ ~ Si02 O ~ (154, ~ H3PMo12040.13H20, TI is 20 s; but TI drastically decreases as a result of dispersion on the support: 2.5 s (loading amount as Mo, 35 wt%) to 430 ms (loading amount as Mo, 2 wt%). Raman spectroscopy shows that the structure of P M o 1 2 0 ~is~preserved at high coverages on Si02 (156). When the loading decreases, the v(Mo=O) frequency decreases, which is probably due to weakening of the anion-anion interaction. On heating to temperatures >573 K, 23 wt% H3PMo12040/Si02 produced an unidentified species andor Moo3 (157). Upon exposure to water vapor, the unknown species was reconverted to the Keggin anion. When Moo3 was supported on Si02 and treated with H20, H4SiMoI2O40formed on the S i 0 2 surface (I58). Soled et al. (159) prepared a unique Cs acidic salt of H3PW1 2 0 4 0 , which is present in a ring 10 pm thick located about half-way into an S i 0 2 extrudate -1.5 mm in diameter.

-

IV. Acid-Catalyzed Reactions in the Liquid Phase

Heteropolyacids are more active catalysts for various reactions in solution than conventional inorganic and organic acids, and they are used as industrial catalysts for several liquid-phase reactions (6, 12). Important characteristics accounting for the high catalytic activities are the acid strength, softness of the heteropolyanion, catalyst concentration, and nature of the solvent (6, 7, 9, 116, 160, 161).

Figure 23 shows a correlation between the catalytic activity and the Hammett acidity function (Ho)of Keggin-type heteropolyacids in CH3CN. The catalytic activity increases with the acid strength (63). Softness of the polyanion is also important in bringing about unique activity. It has been suggested that reaction intermediates like oxonium ions or carbenium ions are stabilized on the surfaces of heteropolyanions (see below), probably by virtue of the softness (162). lzumi et al. (124) observed that H4SiW12040 was more active for the reaction of dibutyl ether with acetic anhydride than H3PW12040,although H4SiW12040 is a weaker acid than H3PW12040. This difference was explained by the stabilization of the cationic intermediate


7 -

151

1.0

0

E

4

I

c

zt

B m 3

0.5

0.5

1.5

2.5

-H0

FIG.23. Rate constants (per mole of protons) for decomposition of isobutyl propionate as a function of Hu values of HnXW12040 (acetonitrile solutions). (From Ref. 63.)

through the formation of the intermediate-polyanion complex due to the softness of the polyanion, which is in the order SiW120:0 > PWl2O;O ( 9 ) .The effect of the softness becomes significant for reactions in aqueous solutions, in which the influence of the difference in the acid strength is slight, since most heteropolyacids are completely dissociated. The concentration of the heteropolyacid is sometimes critical. Excellent performance was demonstrated for the hydration of butenes when heteropolyacids were used in a highly concentrated solution (163). As shown in Fig. 24, the rate increases remarkably with an increase in the concentration of

0 Concentrationof H3PM0~~0~drnol.drn-3 FIG. 24. Dependences of the initial rate of TBA formation and ratio of DIP/TBA on H~PMO~IO concentration ~O in the hydration of butenes. TBA, tert-butyl alcohol; DIB, diisobutylene. 357 K, isobutylene/l-butene = l / l (molar). (From Ref. 163.)

152

TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO TABLE XIV Acid-Catalyzed Reactions in the Liquid Phase Reaction

-

c C=C

/

\

+ H2O C

+ HzO

C=C-C-C

Catalyst

Remarks"

Ref.

c I I

C-C-C

-

H3PW12040

T = 313-353 K

162. 166

HiPW12040

T = 473 K, 200 atm

12

OH C-C-C-C

I

OH

+ 2HCHO

Ph-CH=CHz

K2KIC-CH2

\ /

-

P h p O OJ

+ KIOH

H~PW12040 T = 328 K H3PW12040: H2SO4 = 300: I (ratio of activity)

162

H,PWlz04o

T = 318 K H,PW1204o : HlS04 = 30: I

124

H3PW12040

T = 368 K 124 H,PW1204o : BF3 ' E t 2 0 = 50:l

H,PWl2040. nH20 (n = 0 6)

T = 333 K MW = 3000

H,PW

T = 298 K, Y

0

/

\

R,O

(2

+ R2Rl-C-CH2

RzRI-C-CHz

+ AcOH

-

OH

-

I

\

HO

OR,

AcO-(CH~)~-OAC

HOf(CH2)4-OkH (PTMG)

-

Cyclotrimerization of propionaldehyde

+ CHIOH 0

II

-

+o/

HbP~W1~062T = 315 K. S

- cR -A - -

C,H,OCH,

85%

176

z

100%

17'7'

II

HlPW12040

T

363-393 K

161

H3PW12040

T = 373 K H3PW 12040: HlS04 = 60:l

63

COR II

0

+ CH,COOH

A O C O C H ,

Ph -CH,OH

=

0

W O + 2ROH C II

0

12040

I64

f CnHa-CH,t

+ CH,COCl

C&0CH1(COCH,)

H4SiMo12040T = 298 K

I80

H4SiM012040 Reflux

180

continued

I53

CATALYTIC CHEMISTRY OF HETEROPOLY COMPOUNDS TABLE XIV-Continued Reaction

Catalyst

0 II

~~

"

Remarks"

H3PWI2O4,) T = 333 K. Y

=

72% (p,p)

Ref.

"

~

7' = temperature, Y = yield. S = selectivity " Asahi Chem Co , LTD , JP 1990-45439

H3PMo1204n.In a certain range of concentration of H3PW120jo,polymerization of THF proceeds efficiently, whereby the molecular weight of the product is also well controlled (164). In some cases, the effective acid strengths of heteropolyacids level off due to the leveling effect of reactant molecules (63). For example, the difference in activity between heteropolyacids and HzS04 is not significant for ester exchange and esterification because alcohols are good proton acceptors ( 6 3 ) . In organic reactants or solvents, high activities of heteropolyacids are often realized (Y). Typical reactions catalyzed by heteropolyacids are listed in Table XIV. A. 1.

REACTIONSIN AQUEOUS SOLUTION

Hydration and Dehydration

Tokuyama Soda commercialized a catalytic process for propylene hydration catalyzed by aqueous solutions of heteropolyacids such as H3PW I 2 0 4 0 ( 165). In Table XV, the activities of various acids are compared at a constant proton concentration. H3PW12040is two or three times more active than H2S04 or H3P04. The reason for the high activity is assumed to be the stabilization of intermediate propyl cations by coordination. Izumi et af. (162, 166) observed high activities of heteropolyacids for the hydration of isobutylene in dilute solution, as follows. The activation energy is 4 kcal mol-I lower for H3PW12040than for HN03. The reaction order with

154

TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO TABLE XV Effect of Anion on Rate of Propylene Hydration" (165)

Anion PMolzO:i pw120:;

so: Po:

-

Concentration of anion (IO~-'moldm-') 1 .oo I .oo I .50 1.50

Relative specific rate 3.4 I .7 I .o 1.o

"423 K, I4 atm, [H,O+]; 3.0 X lo-' mol dm-'

respect to heteropolyacid is about 1.5, whereas that for HN03 is 1. The proposed reaction mechanism is shown in Scheme 3. Path A corresponds to specific acid catalysis. Path B involves an intermediate n-complex formed from protonated isobutylene and a heteropolyanion. The overall rate is expressed by the sum of the rates of paths A and B [Eq. (1 1a)], where kl and kll are the rate constants, PR is the partial pressure of isobutylene, and [heteropolyanion] is the concentration of heteropolyanion, respectively. r = kl(Ps[H~O+])+ k~l(P~[H~O'])[heteropolyanion]

( 1 la)

The existence of a significant interaction between the heteropolyanion and the carbenium ion is supported by the presence of alkyl-heteropoly complexes. b o t h and Harlow (167) synthesized 0-alkylated compounds such as [(CZH5)30]3PWIZ037, where a CzHs cation is bonded to the polyanion. Farneth et al. (168) and Lee et al. (169) reported the formation of methyl groups and

Path A

SCHEME 3


155

ethyl groups directly attached to the heteropolyanion (methoxy and ethoxy), respectively. A commercial process for the separation of isobutylene from a mixture of isobutylene and n-butenes through direct hydration of isobutylene to give tertbutyl alcohol has been established by use of a concentrated solution of heteropolyacids (6, 163, 170). The reaction order in the heteropolyanion varies from 1 to 2 as the concentration of heteropolyanion increases from 0.05 to 10 mol dm-3. This increase corresponds to a change from the first to the second term in Eq. (Ila). At concentration of the heteropolyanion greater than 0.5 mol dm-3, path B in Scheme 3 becomes dominant. In addition, the solubility of isobutylene increases linearly with the concentration of H3PWI2O4,,whereas the solubility is little affected by the concentration of HN03 (169). At a 1.5 mol dm-3 concentration of heteropolyanion, the solubility of isobutylene is about 2.3 times higher than for mineral acids. Hence the high catalytic activities of heteropolyacids (about 10 times higher than that of HN03) are explained by the combination of the three effects of (a) the high solubility, (b) the coordination of isobutylene to the heteropolyanion, and (c) the high acid strength (162, 166). Isobutylene in a mixture of isobutylene and 1 -butene is very selectively hydrated in the concentrated solution; diisobutylene formation is negligible, and sec-butyl alcohol concentrations are < 100 ppm. An apparently different explanation for the catalysis has been advanced by Kozhevnikov et al. (171, f 72). The rate constant, k (min-I) for H3PW12040and that for H4SiW12040at room temperature show linear relationships with the Hammett acidity function (H,) as shown by Eq. ( I Ib). log k

= - 1.04H0 -

3.46

(1 lb)

The data of H2SO4, HCI, H N 0 3 , and HC104 also fit this equation. On this basis, they suggested that the hydration of isobutylene in the presence of heteropolyacids and inorganic acids proceeds via a common mechanism, in which the rate-limiting step is the conversion of the n-complex into a carbenium ion (Scheme 3). The complexation effect as described above is possibly included in the value of Ho according to this explanation. Heteropolyacids are much more active than H2SO4 and HC104 for hydration of phenylacetylene [Eq. (12)] ( I 73). Also in this case, the rate of reaction in the presence of heteropolyacids shows an approximately second-order dependence on the catalyst concentration. This observation suggests that this reaction proceeds by a mechanism similar to that of Scheme 3: PhCECH

+ HzO

+

PhCOCHi

(12)

The higher activities of heteropolyacids are also found for the dehydration of lP-butanediol to give tetrahydrofuran (THF) (I 74). The activity order is H4SiW12040> H3PW1204,> H3PMoI2O4,. Rate data are summarized by

156

TOSHIO OKUHARA, NORITAKA MIZUNO, A N D MAKOTO MISONO

Eq. (13), and the reaction probably proceeds by a mechanism similar to that shown in Scheme 3. I' =

( k , + k2[SiW120:i])[ 1,4-butanediol][Ht]

(13)

2 . Prins Reaction

Heteropolyacids catalyze the Prins reaction of alkenes [Eq. (14)] more efficiently than H2SO4 and p-toluenesulfonic acid (PTS). For example, H3PW12040 is 10-50 times more active than H2SO4 or PTS (162). In this reaction, oxocarbocations may be stabilized through complexation with heteropolyanions: PhCHzCH2

+ 2HCHO

-

,CHz-CH, PhCH, O+H2

\

(14)

/O

B. REACTIONSIN ORGANIC SOLUTION 1.

Ether Cleavage Reactions

H3PW12040shows a higher activity for the alcoholysis of epoxides [Eq. (15)] than H2SO4, PTS, or HC104 (9, 124, 175). While rapid deactivation is observed with H2SO4, which is probably due to the formation of an alkyl sulfate, H3PW12040maintains its high catalytic activity.

Tetrahydrofuran is cleaved with acetic acid to give 1,Cdiacetoxy acetate selectively (124) (Table XIV). The activity order is H3PW12040> H4SiW12040> H3PMo12040%- BF3 * EtzO > PTS. The catalytic activities for the reaction of THF with acetic anhydride are in the order H4SiW12040= H3PW12040> H4GeW1204~> H4SiMo12040 > H3PMo12040 S H2S04 (124). The cleavage of dibutyl ether with acetic acid is also catalyzed by heteropolyacids (124). H4SiW12040 is the most active among the heteropolyacids, probably due to its higher softness. 2 . Polymerization of Tetrahydrofuran

Aoshima et al. (164) found that THF can be directly polymerized to give polyoxotetramethyleneglycol (PTMG) is a highly concentration solution of heteropolyacids [Eq. (16)].


157

H@/H3PW12040 (Mol ratio)

FIG.25. Influence of the molar ratio of H20/H,PW12040 on THF polymerization (333 K). ( A ) Solid-liquid phase, (B) two-liquid phase, (C) homogeneous liquid phase. (From Ref. 1 6 4 . )

PTMG

The water content greatly influences the activity and the molecular weight of PTMG, as shown in Fig. 25. At the ratio of H20/PW120:; = 10, the reaction mixture consists of two liquid phases; the upper phase is mainly THF and the lower phase is the complex of H3PW12040 and THF (in the catalyst phase). The THF polymer is formed in the catalyst phase and is transferred to the THF phase. This “phase-transfer polymerization” is illustrated in Fig. 26. IR spectra of the C-0-C stretching region of the catalyst phase are shown in Fig. 27. The peaks a, b, and c were assigned by Aoshima et al. (264) to uncoordinated THF, hydrogen-bonded THF, and protonated THF, respectively. The polymerization activity is associated with the amount of protonated THF.

3 . Condensation Reactions Table XVl is a summary of typical results observed for cyclotrimerization of propionaldehyde to give 2,4,6-triethyl- 1,3,5-trioxane ( I 76). In catalysis by H ~ P M O ~ the ~ Oreaction ~ ~ , mixture separates into two phases during the course of the batch reaction. The products are present in the upper layer and the catalyst in the lower layer, so that the catalyst solution can be used repeatedly without a catalyst isolation step. Selectivities exceeding 97% and turnovers exceeding 300 moles of product per mole of catalyst have been obtained. Polymerization of 1,3-trioxane, a cyclic trimer of formaldehyde, is catalyzed by H3PMoI2O40 (277). The polymerization is very fast at concentrations of mol dm-3. To obtain comparable rates using BF3 H3PMo12040as low as catalyst, a BF3 concentration of mol m-3 is required.

158


FIG.26. Reaction model of phase-transfer polymerization. (From Ref.

164.)

Condensation of acetone to give mesitylene is catalyzed by H3PW12040at room temperature ( I 78).

4. MTBE Synthesis Table XVlI is a comparison of the catalytic activities for liquid-phase MTBE synthesis from isobutylene and methanol (I 79). The catalyst structure and composition have a strong effect on the activity. The highest activity per proton was obtained with a Dawson-type heteropolyacid, H6P2W18062,although the acid strength of H6P2WI8O62is lower than that of the Keggin-type H3PW1204,) (Section 111). Water added to the mixture has little effect on the reaction rate at water concentrations less that 2 wt%, but at 5 wt% the rate is less by a factor of 2.5. At the same time the selectivity is less due to the formation of tert-butyl alcohol.


159

v(WE0) A

1100

1070

1040

1010

980

950

Wave number/cm-’

FIG.27. Infrared spectra of C-0-C stretching vibrations of THF in catalyst phase and their deconvolution results. ( I ) H4SiW12O4dH2OiTHF= 1/7/20 (molar); (2) H3PW1204dH20/THF= 1/7/20 (molar). (From Ref. 164.)

TABLE XVI Liquid-Phase Cyclotrimerization of‘ PropionaldehyJL.“ ( I 76)

Catalyst HiPMo I 2 0 4 ” HiPMo I 2 0 4 / Hd’W12040 H4SiW ,?04” Si02-A1203 AICIj ZnC12 ZnCIg P-TsOH P-TSOH’ H3Po4’

Amount (mmol)

Time (h)

Conversion (%)

Selectivity” (mol%)

0.4 0.4 0.3 0.3

2 2 2 2 2 2 2 12 2 2 4

87. I 84.7 86.3 66.2 2.8 91.5 87.6 69.2 78.5 47.3 58.2

97.2 97.5 97.2 97.3 3. I 88.6 91.5 92.5 4n.9 97.6 97.3

~

7.5 7.3 -

5.3 -

8.7

TON‘ 350 340 490 370 -

19 20 15 13 I5 II

“Reacted at room temperature, I g of catalyst to 10 g of propionaldehyde. Selectrivity to 2,4,6-triethyl-I ,3,5-trioxane. “Turnover number (molar ratio of aldehyde reacted to trioxane to catalyst). For twelfth run. Second run. ’Fifth run. ‘Contained I5 wt% of water.

160


Activiry

TABLE XVlI Methyl tert-Bury1 Ether' ( I 79)

01Heteropolyacid in Svnthesis of

2.0 x 10' 2.0 x 10' 1.5 x 10' 0.86 x 10' 1.9 x 10' 0.72 x 103 0.6 x 10'

2.0 3.3 2.3 1.8 3.3 2.0 2.3 5.5 3.1 0.76 9.0

4.4 x lo3 1.6 x 10' 0.55 x 10' 0.45 x 10'

Reaction conditions: 315 K, 12 ml CHIOH, 1 g cat., P

= 100 kPa.

5 . Esterijication and Ester Decomposition The addition of carboxylic acids to olefins proceeds in solution in the presence of 10-4-10-2 mol dm-3 of HjPWl2040 at 2 9 3 4 1 3 K with a selectivity of 100% (7). H2SO4 is a less active catalyst than H3PW12040by a factor of 30-90. Esterification of p-nitrobenzoic acid with ethanol has been carried out by using H2SO4 catalyst in an industrial process (160). In the presence of H3PWl2040, this reaction takes place with a yield >99% (160). At the end of the reaction, the reaction solution separates into two phases. The upper layer contains ethyl-p-nitrobenzoate, toluene, and ethanol, and the lower layer consists of a solution of the heteropolyacid in ethanol. Consequently, the catalyst can be readily separated and reused. The activities of heteropolyacids for the decomposition of isobutyl propionate were found to be 60-1 00 times higher than those H2S04 and p-toluenesulfonic acid (63). The activity increases with increasing acid strength of heteropolyacids. 6. Alkylation and Dealkylation Nomiya et al. (180) demonstrated that alkylation and acylation [Eqs. (1 7) and (18)] proceed in the presence of H4SiM012040: nPhCH2OH ArH

+ CH3COCI

+

(-C~H~-CHZ-)"

(17)

-+

ArCOCH3

(18)


161

In an acetic acid solution of H3PW 1 2 0 4 0 , dealkylation of 2,6-di-tert-butylphenol to give 2-tert-butylphenol takes place at 357 K. The catalytic activity of H3PW12040is more than 100 times greater than that of H2S04(160). 7. Others In an acetone solution, decomposition of isopropylbenzene hydroperoxide to give phenol and acetone is catalyzed by H3PMo12040(181). The heteropolyacid is 3 times more active than H2SO4. Synthesis of I , 1 -diacetate from aldehyde and acetic anhydride has been attempted by using H4SiW12040and the zeolite HZSM-5 (182). Both catalysts gave more than 98% yield of the diacetate from benzaldehyde. However, the reaction rate was far higher for H4SiW12040than for HZSM-5.

V.

Heterogeneous Acid-Catalyzed Reactions A. GENERAL CHARACTERISTICS

Among heteropolyacids, polytungstic acids are the most widely used catalysts owing to their high acid strengths, thermal stabilities, and low reducibilities. As summarized in Section 111, the acidic properties (acid strength, acid amount, type of acid, etc.) are controlled by (i) the structure and composition of the heteropolyanion, (ii) the formation of salts (or the countercations), (iii) the tertiary structure, and (iv) supporting onto carriers such as Si02 and active carbon. Factors (iit(iv) are specific to the solid state. The acid strength can be controlled mainly by (i), and the number of acid sites is greatly influenced by (ii) and (iii). In this section, these influences will be described. Besides the acidic properties, the absorption properties of solid heteropolyacids for polar molecules are often critical in determining the catalytic function in “pseudoliquid phase” behavior. This is a new concept in heterogeneous catalysis by inorganic materials and is described separately in Section V1. With this behavior, reactions catalyzed by solid heteropoly compounds can be classified into three types: surface type, bulk type I, and bulk type I1 (Sections VII and IX). Softness of the heteropolyanion is important for high catalytic activity, although the concept has not yet been sufficiently clarified. The influence of the heteroatom on the acid strength is shown, for example, in Fig. 28. Here the rates of alkylation of 1,3,5-trimethylbenzene and the decomposition of cyclohexyl acetate catalyzed by the acid forms of several 12-tungstates are plotted against the negative charge of polyanion for solid heteropolyacids (63, 183). The catalytic activities correlate well with the acid strength in solution (Fig. 14). This correlation indicates that the acid strength of

162


2

3 4 5 6 7 Nagative charge of polyanion

FIG. 28. Catalytic activities (per surface proton) as a function o f the negative charge of polyanion, XW 120;". (0)Alkylation of I ,3.5-trimethylbenzene with cyclohexene; ( 0 )decomposition of cyclohexyl acetate. (From Ref. 183.)

the acid form in the solid state reflects that in solution and decreases with increasing negative charge of the polyanion (Section 111). A correlation between the acid amount of the surface and the catalytic activity for the Cs salts of H3PW12040is shown in Fig. 29 (128). The number of surface

a 0 R

7

0 0.5

-

J

-

. ms

a"

0

20

40 60 200 240 Surface Acidity/pmol g-1

280

FIG.29. Surface acidity and catalytic activities for alkylation of 1.3.5-trimethylbenzene with cyclohexene (closed symbols) and decomposition of cyclohexyl acetate (open symbols). (From Ref. 128.)


163

acid sites was estimated by multiplying the formal concentration of protons on the surface by the specific surface area [e.g., (0.5) X (number of polyanion on the surface) X (number of protons per polyanion)] (6, 47). The catalytic activities correlate linearly with the number of surface acid sites. These results are reasonable because the acid strengths of acidic Cs salts are similar to that of the acid form (127).The high catalytic activity of C S ~ . ~ H O . ~ P W(Cs2.5) I ~ O ~that ~) was reported previously (46a, 48, 127) is thus inferred to be due to the high surface acidity (high surface area and presence of protons). As described in Section 11, the pore size of the acidic Cs salts can be controlled by the Cs content. C S ~ . ~ H O . ~ P(Cs2.2) W ~ ~ has O ~pores ~ having an effective size of 6.2-7.5 A, and the pore size of Cs2.5 is larger than 8.5 A. Figure 30 shows the catalytic activities of Cs2.1 (the pore size is less than 5.9 A) and (32.2 for each reaction relative to the activity of Cs2.5 (48).Cs2.5 catalyzed all the reactions with considerable activities (the reaction rates are shown in parentheses in Fig. 30). On the other hand, although Cs2.2 was found to be as active as Cs2.5 for dehydration of 2-hexanol (molecular size, 5.0 A) and decomposition of isopropyl acetate (molecular size, 5.0 A), it was much less active for the decomposition of cyclohexyl acetate (molecular size, 6.0 A) and alkylation of 1,3,5-trimethyIbenzene (molecular size, 7.5 A).Therefore, Cs2.2 is

1-

Relative Activity I Reaction

Dehydration :

.. .. .. .. .... . . . . . .. ..... .. . . . .. .. .. .. ....(r:4 . . . . :. .. .. ..... .)id): ... Y

Decomposition :

* ' . .""'L

.. .. .. .. . . . . . . . . . . . . . . .. .. ... .446): . . . . . . :. . :.:.: .:(w ...

COCH,

I

Akylation:

~ O2.2, ~ ~and 2.5) for various kinds of FIG. 30. Relative activities of C S ~ H ~ - ~ P (xW=~2.1, reactions in liquid-solid reaction systems. Catalytic activity was estimated from the initial rate of the reaction. The activity of Cs2.5 for each reaction is taken to be unity. The figures in parentheses are reaction rates in units o f mmol g- h - I. (From Ref. 48.)

'

164


active only for smaller molecules; the reactions are influenced by reactant or transition-state shape selectivity. As for Cs2.1, it is active for the dehydration of 2-hexanol but inactive for other reactions, notwithstanding its high surface area (55 m2 g-I). To our knowledge, this is the first example of shape-selective catalysis by a solid superacid (47, 48). A remarkable effect of the countercation is demonstrated in Fig. 3 1, where the rates of several reactions are plotted against the extent of neutralization by Na or Cs, that is, x, in M,H3-.PW12040 (M = Na or Cs) (46, 128). In the case of Na salts, the rates decrease more or less monotonically as the Na content increases. On the other hand, peculiar changes in activity are observed for the Cs salts; activity maxima occur at x = 0 and x = 2.5. The activity of Cs2.5 relative to that of H3PW12040changes depending on the reaction. As will be described in more detail in Section VI, the activity pattern is principally governed by the extent of the contribution of the bulk type I catalysis. As was described in Section 111, the acid strength usually decreases when the catalyst is supported on SiOz. Figure 32 shows the influence of the loading of a

1.o

0.8

0.6

.b >

‘ i ;

0

0.4

m

$ m

c

0.2

m u

O

.z m -

1.0

0,

d

0.8 0.6 0.4

0.2

0 0

1

x in M,Hs-,PW

3

2 12040

FIG. 31. Catalytic activities of acidic Na or Cs salts of H3PWI2O4,)as a function of Na or Cs content. (a) M = Na: (0) dehydration of 2-propanol, (A)decomposition of formic acid, (0) conversion of methanol, (W) conversion of dimethyl ether. (b) M = Cs: (0) dehydration of 2-propano1, (m) conversion of dimethyl ether, (A) alkylation of 1,3,5-trimethylbenzene with cyclohexene. (From Refs. 46 and 128.)


roo^ 2.3 ,6.9, I

165

Mo (wt%) 12.1 ,

2?3

26;2

Coverage, fraction of monolayer

30;7

1 .o

FIG.32. Dependence of selectivities in methanol oxidation on surface coverage and weight YOof Mo for H4SiMo12040. (a) (CH3)20,(b) HCHO, (c) (CH30)2CH2.(d) HC02CH3. (From Ref. IX4.)

H4SiM~12040 on the support for catalytic oxidation of methanol (184). At coverages larger than 0.25 monolayer, the selectivities remain constant; dimethy1 ether is the main product (about SO%), showing the acidic character of the catalyst. Below 0.25 monolayer, the acidic character is lowered, and dimethyl ether rapidly disappears. Typical examples of reactions catalyzed by heteropoly compounds in the solid state are summarized in Table XVIII. B. DEHYDRATION A N D HYDRATION In catalytic dehydration of alcohols, pseudoliquid phase behavior (bulk type I reaction) of heteropolyacids has been demonstrated (Section VI). The high catalytic activity is associated with this behavior and the strong acidity. Unique pressure dependences of the catalytic activity and selectivity are found for H3PW12040 due to the pseudoliquid phase (Fig. 40). H3PWIZ040was found to be much more active for the dehydration of 2-propanol than Si02-Al203 (by a factor of about 30 per weight and about 2000 per surface area) at 398 K (185).The activities of heteropolyacids decrease in the sequence PW12 (H3PWI2040)> SiW12 > PWlOV2 > PMo12 > PMoIOV2 = SiMol2, which is close to the acid-strength series in solution. Niiyama et al. ( 4 2 ) found that the catalytic activities of the water-soluble salts correlate with the electronegativity of the metal cations. Dissociation of coordinated water [type (1) in Section 1111 has been proposed to account for the acidity.

166

TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO TABLE XVllI Heterogeneous Acid-Catalyzed Reactions

>OH

-

=( + H 2 0

Cata Iyst

Reaction

-

A1203 = 30 I (ratio of activity/g) H4SiWI2O40/ Amberlyst I5

+OH

-

Cl& Hydrocarbons

+ CzHsOH

-CHACOOC~H~

O O C O C H ,

-,&+A

Ref.

+ H,O

CH30H(CH30CH3)

CHiCOOH

Remarks"

+ HzO

- 0+

-

+,

CHICOOH

etc.

T = 343 K, S = 98.6% (1 1 % conv.)

I86

H3PW12040 T = 573 K I 89 C~z.sHo.sPW1204o T = 563 K, S = 74% 195 (Cz-C4 olefins) H3PW1~040/Si02T = 423 K S = 91% (90% conv.)

I51

C~2.sHo.sPWi2040T = 373 K csz 5 : so: -/zro2 = 43 : I (activity/g)

I98

H3PW12O4dSi02 T = 308 K, sensitive to pretreatment

207

H3PW12040

209

T = 298 K

I YY

continued


167

TABLE XV I I I-Contin ued Reaction

~

C

CHiNH: NH,

O

C

Catalyst

Ref.

d

+I

H4SiWI2O4,JSiO2 T = 263 K S = 94.1% (30% conv.) T = 323 K HcP~WIHO~

-HCN

+ CHIOH

Remarks"

T = 113 K

226

( N H ~ ) I P W I Z O ~ T~ = 150 K

227

CSIsH,~ P W I Z O T~ =~ 413 K s = 97%

h

H1PWi20411

(CH3).NH,,,

21x 220

(94% conv.) ~

"

T = Temperature. Y

=

~~~

yield, S = selectivity. *Sumitorno Chem. Co., LTD., JP 1991-300150.

The Cs salts of H3PW12040show a unique activity pattern for dehydration of 2-propanol (Fig. 32) (46a). Although the change in the catalytic activity resembles that of the surface acidity (Fig. 21), the activity of H3PW12040 for this reaction is relatively high. The high activity is explained by the pseudoliquid phase behavior of H3PW12040and its affinity for the polar molecule, 2-propanol. H ~ P W , Z OH4SiW12040, ~~, and the supported heteropolyacids catalyze the hydration of isobutylene (186). The supports are listed as follows in the order of activity: Amberlyst 15 [porous sulfonated poly(styrene-divinylbenzene)] = activated carbon > SiOz > TiOz. The maximum activity of H4SiWI2O4dArnberlyst 15 was obtained at about 30 wt% H4SiW12040. It was found that polypyrrole and polyacetylene-supported H3PMo12040are much more active than the unsupported parent acid (187). The activity per unit mass of heteropolyacid of the supported catalyst is about 10 times higher for the production of ethylene and ether than that of the unsupported catalyst. The dehydration of ethanol catalyzed by a membrane comprising H3PW12040 and polysulfone was reported (188). The polysulfone is more permeable for

168


ethylene than for diethyl ether. Thus the selectivity for ethylene in the membrane reactor was higher than that in a fixed-bed reactor containing the same catalyst. c . CONVERSION

OF

METHANOL INTO HYDROCARBONS

Ono et al. (189) reported that heteropolyacids such as H3PW12040 and catalyze the conversion of methanol into hydrocarbons, although H4SiW12040 the activities are less than that of HZSM-5. In contrast to HZSM-5, the main products observed with heteropolyacids are aliphatic C hydrocarbons, the selectivities for aromatic hydrocarbons being small (Table XIX). Countercations influence the rate and selectivity of this reaction. The activity order, as for cations, was found to be Ag > Cu, H > Fe > A1 > Pd > La > Zn (1YO). The distributions of product hydrocarbons were found to be similar to those observed for H ~ P W , Z (Table O ~ ~ XIX), suggesting similar reaction mechanisms. Ag and Cu salts of H,PW12040 are much more active than the acid form catalyst. Protons generated by the reaction of Ag' with H2 are presumed to give the more active catalyst (191). Hayashi and Moffat (192) reported that the A1 salt and the NH4 salt were effective catalysts for the conversion of methanol to hydrocarbons. They claimed that the N& salts show high catalytic activity and selectivity for the formation of saturated hydrocarbons rather than olefins. The salts of organic TABLE XIX Product Distribution in Conversion of Methanol into Hydrocarbons (189) Catalyst HTP"

CuTP

AgTP

HTS"

CUTS

AgTS

Product distribution (%)* MeOH I .3 MeOMe 38.6 Hydrocarbons 60.1

I .2 37.3 61.5

0 2.0 98.0

3.9 57.5 38.6

4.7 35.5 50.8

0 20. I 79.9

Hydrocarbons distribution (%)h CH4 3.7 C2H4 11.3 C2H6 0.9 C3H6 8.3 C3Hs 16.1 c4 39.3 cs 12.5 c 6 7.9

5.2 9.5 0.8 8.5 13.4 39.2 13.7 9.1

9.0 9.0 5.2 3.8 34.0 26.1 7.2 5.7

I .6 10.3 0.5 8.3 9.8 41.8 15.8

7.3

3.2 10.3 1.2 8.3 2 I .9 36.4 13.3 5.4

11.9

11.2

0.5 8.7 14.5 35.1 15.1 7.6

TP and TS indicate dodecatungstophosphateand dodecatungstosilicate, respectively. 'Calculated on a carbon-number basis.


169

bases are effective for the formation of olefins (193, 194). In the case of an acidic Cs salt, Cs2.sHo.sPW12040, the selectivity for C 2 4 4 olefins increased continuously as the Cs content increased, e.g., being 43% for H3PW1204o and 64% for S2.5Ho.5PW12040(195). Pseudoliquid phase behavior is important in this reaction. The kinetics observed with H3PW12040is quite different from that observed with HZSM-5. No induction period in the rate of hydrocarbon formation was observed in the former, in contrast to the latter, suggesting that the methanol conversion catalyzed by H3PW12040 is not autocatalytic (196). The results were explained by pseudoliquid phase behavior of H3PWI2040 (Section VI). The selectivity to lower olefins is improved by control of the pseudoliquid phase behavior. The olefidparaffin ratio in the product hydrocarbon depends inversely on the ability of the heteropoly compounds to absorb reactant dimethyl ether. The ratio is greatly increased as the contribution to catalysis of the bulk phase of the catalyst (pseudoliquid phase) decreases (Section VI) (195, 197). It was confirmed that the acid strength is not significant in influencing the selectivity of this reaction.

D. ESTERIFICATION A N D ESTER DECOMPOSITION Okuhara et af. (198, 199) found that C S ~ . ~ H ~ . S P(Cs2.5) W I ~ Ois~much ~ more active for the decomposition of cyclohexyl acetate in liquid-solid mixture than Nafion-H (sulfonated fluorocarbon resin), HY zeolite, HZSM-5, SiO2-AI2O3, and SO:-/Zr02. Figure 33 demonstrates the superiority of Cs2.5 for this reaction as well as the alkylation of 1,3,5-trirnethylbenzene.The activity of Cs2.5 is more than 200 times as high as that of Si02-A1203. The difference is much greater than that observed for gas-solid mixtures (46a, 15)s).Cs2.5 works as an insoluble catalyst in esterification of acetic acid with ethanol (200). The Activity I mmol I I g-l h-l

Catalyst

2TL7T-i

Cs2.5H0.5PW12040 H3PW12040 S042-/Zr02 Nafion-H Si02-Al203 HZSMQ H2S04

FIG.33. Catalytic activities of various acid catalysts for liquid-phase reactions. TMB: Alkylation of 1,3,5-trimethylbenzene with cyclohexene; CA: Decomposition of cyclohexyl acetate. (From Ref. 47.)

170

TOSHIO OKUHARA, NORITAKA MIZUNO. AND MAKOTO MISONO

order of catalytic activities for the esterification at 333 K (acetic acid/ethanol/ catalyst = 100/100/1 by weight) is as follows: Amberlyst 15 (1.6 X > Cs2.5 (0.52 X > HZSM-5 (0.03 X where the -I figures in parentheses are rate constants in units of min- g . In the hydrolysis of ethyl acetate (ethyl acetate/water/catalyst = 1080/28/0.1 by weight), the activities are in the order Amberlyst 15 > Cs2.5 > HZSM-5. Izumi et al. (201) found that activated carbon firmly entraps H3PWI2O40and the acid is resistant to extraction with hot water or hot methanol. The entrapped H3PW12040 showed a catalytic activity for the esterification of acetic acid with ethanol comparable to that of Nafion-H. With H3PW12040supported on carbon, a selectivity to ethyl acetate of 99.5% at 95% conversion was obtained. Heteropolyacids supported on Si02 or carbon are more active than Si02-A1203 (202), but A1203 is not suitable as a support, because A1203 readily decomposes heteropolyacids due to the high basicity of the surface. Cs2.5 in water is hardly separable by filtration because of its very fine particle size. Cs2.5/Si02 prepared by the hydrolysis of ethyl orthosilicate in the presence of colloidal Cs2.5 is efficient for the hydrolysis of ethyl acetate, although it is less active than Amberlyst 15 (200). The catalytic activity of H3PW 12040 for esterification of phthalic anhydride with 2-ethylhexanol is higher than those of conventional soluble acids such as p-toluenesulfonic acid (202).

E. ALKYLATION A N D DEALKYLATION Conventional soluble catalysts such as H2SO4 and AIC13 have been used in industry for alkylation reactions, but these are characterized by operational problems of corrosion, catalyst removal, waste formation, etc. Insoluble solid catalysts are desirable for these reactions. H3PW, 2 0 4 0 catalyzes the monoalkylation of p-xylene with isobutylene (203) [Eq. (19)]. The product tert-butyl-p-xylene (BPX) is an important precursor for liquid crystalline polyesters and pol yamides having low melting points and good solubilities. The introduction of the tert-butyl group in the position ortho to the methyl group of p-xylene takes place very slowly, with H2SO4 or AlC13 as the catalysts (whereas alkylation at the meta or para position of 0- or m-xylene occurs rapidly):

(BPX) The results of Table XX demonstrate that H3PW12040is an excellent catalyst for this reaction. The selectivities are in the order H3PW12040(75%) >

171

CATALYTIC CHEMISTRY OF HETEROPOLY COMPOUNDS TABLE XX Selectivity for Alkylation of p-Xylene with lsoburylene by Various Acid Cutulysts (203) Selectivity" (%) Conversion

Mass balance (%)

Catalyst

BPX~

Oligomers'

OPX'

PXM'

(TO)

Hd'W12040 H4SiW 1 2 0 4 0 HsBW 1 2 0 4 0 CSZ5HosPW12040

74.6 26.6 6.1 50.5 50.9 4.2 1.3

11.5 70.2 93.9 48.0 26.0 95.8 98.7 92.8 I00

6.3 0.4 0 0.8 3.4 0 0 0 0

7.6 2.8 0 0.7 19.7 0 0 0 0

90 91 97 89 86 92 75 86 10

so: /zro2 ~

Amberlyst-I 5 SiO2-AI2O3

1.2

CFjCOOH

0

96 I08 I I9 I10 86 109 101

68 81

" Mol%. 'f-Butyl-p-xylene. 'Sum of dimer, trimer, and tetramer o f isobutylene. "Octyl-p-xylene. 'Di-p-xylylmethane. The reaction was performed at 303 K for 30 min. p-Xylene, 0.28 mol; isobutylene, 0.37 mmol m i - - ' ;catalyst, 0.45 g.

H4SiWI2O40(26.6%) > HSBW1204,, (6.1%) (203). The results indicate that the acid strength is essential for the selectivity; that is, strong-acid sites are effective for the alkylation to BPX, but weak-acid sites preferentially catalyze the oligomerization of isobutylene. Although SO:-/Zr02 is a stronger acid than H ~ P W I ~ itO is~ less ~ , selective. SO:-/Zr02 also has a large number of weakacid sites which are probably active for the oligomerization. Alkylation of p-cresol by isobutylene is an important reaction in the synthesis of phenolic antioxidants. The activity of H3PW12040 for this reaction is greater by four orders of magnitude than that of H2SO4 (160). Alkylation of p-(tevtbuty1)phenol (TBP) with cyclohexene, 1-hexene, styrene, or benzyl chloride proceeds in the presence of H3PW12040 at 3373423 K (204). Alkylation with styrene gives 2,6-dialkyl TPB with a selectivity of 90% at 100% conversion. When the alkylation of TBP is completed, an excess of o-xylene is introduced into the reaction system, and 2,6-dialkylphenol is obtained through the trans alkylation without the need for separation of 2,6-dialkyl-4-tert-butylphenol(160) [Eq. (20)].

$+--

6

+

%But

(20)

R R Dealkylation of 2,6-di-tert-butylphenol takes place at 4 0 3 4 2 3 K in the presence of solid H3PWI2O4,,. H3PW12040is two orders of magnitude more active than aluminum sulfate for this reaction (205).

172


60

5 40 .Q

Y?

I

5 20 0 0

30

60

90

120

Timdmln FIG. 34. Time course of alkylation of 1,3,5-trimethylbenzene with cyclohexene at 343 K.

(0) C S ~ ~ H " ~ P W , ~ O ~ (0) ~ ( Csecond S ~ . ~ run ) , for Cs2.5, ( 0 )third run for Cs2.5. (A) SOi-/ZrOz, (V)Nafion-H, (M) HY zeolite. (From Ref. I Y Y . )

Figure 34 shows the time course of alkylation of 1,3,5-trimethylbenzene with cyclohexene (199). In the presence of Cs2.5, the reaction rate reached a steady value after approximately 1.5 h. When Cs2.5 was filtered and reused (in the second and third runs), the rates were nearly equal to the steady-state value of the first run, indicating that there was little catalyst deactivation and that the reaction did not take place in solution. The primary reason for the high activity of Cs2.5 is both its strength of acid sites and acid strength (Section 111). The specific activities of heteropoly catalysts (rates per acid group) are much greater than the activities of Si02-A1203, SO:-/Zr02, or zeolites. This result cannot be explained simply on the basis of the acidic properties. There are additional effects such acid-base bihnctional acceleration by the cooperation of proton (acid) and polyanion (base) groups (127). A proposed reaction model is illustrated in Scheme 4. Supported heteropolyacids are also used for alkylation reactions. Alkylation of benzene with I-dodecene was examined with H4SiWI2O40/SiO2as the catalyst (206). Si02 is a better support for the catalyst in this reaction than AI2O3 or Si02-Al203 ; H3PWI2O40/SiO2 catalyzes the reaction at room temperature (207). The pretreatment temperature of the catalyst has a significant effect on the activity. As shown in Fig. 35, the maximum conversion of 1-octene was obtained when the catalyst was treated at 423 K. Pretreatment at a lower temperature such as 373 K gives a reduced acid strength, probably because of the remaining water of crystallization. H3PW1204dMCM-41 exhibits a higher activity than H2SO4 in liquid phase alkylation of TBP with isobutylene or


-

173

Me Me-

Me

Me

SCHEME 4

styrene (208). Supported heteropolyacids are also active catalysts for alkylation of benzene with ethylene at 473 K in the presence of vapor-phase reactants (151). Recently, Soled et al. (159) reported that C2,5Ho.sPW1204~ supported on Si02 extrudate is effective for this reaction. The alkylation of isobutane with n-butenes to give C8 alkylate [Eq. (21)] is a widely used and increasingly important process in petroleum refining. The

40

20 0

274

374

474

574

674

Calcination ternperature/K FIG. 35. Alkylation ofbenzene with I-octene with H3PW12040(303 K). ( 0 )Supported on silica A (15 wt%); (0)supported on silica B (15 wt%), (A) unsupported. (From Ref. 207.)

I74

TOSHIO OKUHARA, NORITAKA MIZUNO. AND MAKOTO MISONO

commercial catalysts are HF or H2SO4; HF has the disadvantages of being highly toxic and corrosive, and H2SO4 processes produce large amounts of spent acid. Both liquid acids constitute environmental hazards because of potential spills. -

9

(21)

Therefore, new solid catalysts to replace these liquid acids are desirable. W ~ ~ O ~this O alkylation Okuhara et al. (209) reported that C S ~ . ~ H O . ~ P catalyzes reaction at room temperature and that it is much more active than SO:-/Zr02. Yield and selectivity data are summarized in Table XXI. The catalytic activities are in the following order: Cs2.5 > H3PW12040> SO:-/Zr02. The selectivity to C8 alkylates in the total products was 73.3% for (32.5. A patent from ldemitsu Co. (210) also describes the high activity of (32.5 for this reaction. Another patent from Mobil gives data for H3PW1204dMCM-41 (211). The alkylation of toluene with methanol is catalyzed by the NH4 salt of H3PW12040 (212). TABLE XXI Yields and Selectivities of Producis in Alylation of isobuiane with I-Buiene Catalyzed by Solid Acids ar Room Temperature (209)

Total yield (w%)“ Selectivity (wt%)* 224-TMP [RON: 223-TMP [RON: 234-TMP [RON: 233-TMP [RON: 23-DMH (C8 alkylates) C547‘‘ Dimersd C9-C 12‘

1001

1101

1031 1061

19.4

25.1

23.0

0.3 24.1 23.6 14.5 10.8 (73.3) 1.5 13.9 11.3

0.6 18.4 15.2 13.9 8. I (56.2) 0.8 8.5 34.0

1.6 28.0 13.9 10.9 7.2 (61.6) 0.9 9.2 26.6

“Yield (wt%) is defined by 100 X [the weight of products divided by the weight of I-butene charged]. ”224-TMP = 2.2,4-trimethylpentane, 23-DMH = 2,3-dimethylhexane, etc. Figures in parentheses are research octane number (RON). ‘Hydrocarbons containing 5-7 carbon atoms. dOctenes. Hydrocarbons constaining 9-12 carbon atoms. Catalyst, 1 .O g; I-butene, 0.94 g; isobutane, 9.4 g. All data were collected at 7 h.

175


F. ACYLATION Industrial acylation of alkylaromatics is generally carried out with acid chlorides as reactants or stoichiometric amounts of AlC13 (catalyst). Solid acid catalysts would be desirable. Izumi et al. (207) found that Si02-supported H3PW12040 and H4SiW12040 are good insoluble catalysts for acylation, with anhydrides or chlorides as acylating agents [Eq. (22)].

It was confirmed that the reaction did not proceed in the liquid phase. A rapid decline in catalytic activity was observed. The deactivation is probably caused by strong adsorption of the product benzophenone on the catalyst; the acylation was retarded when the reaction was started in the presence of benzophenone. The reaction also proceeded with H3PMoI2O40/SiO2as the catalyst, but the H3PM012040 decomposed during the reaction. It was presumed that the catalytically active species was not the heteropolyacid on the support, but was probably Mo chloride instead (213). C~2.5Ho.sPW12040 (Cs2.5) catalyzes the acylation of aromatic compounds with benzyl chloride, benzoyl chloride, benzoic anhydride, benzoic acid, or acetic acid (214). As shown in Table XXIl, H3PW12040 is usually less active than TABLE XXll Friedel-Crajis Acylation Catalyzed by C SsH,, ~ .TPW,,O;" (214)

Substrates

Product yeld' (%)

Acylating agent

Aromatic compound

CszsHn.sPWizO40

HJ'Wi2040

(PhC0)20 (PhC0)20 (PhC0)20 PhCOzH PhCOzH Ac~O AcOH n-C7HISCOCI

p-xylene anisole chlorobenzene p-xylene anisole anisole anisole mesitylene

57

3 69'

85 0

I I" 3 89

0 8" 4'

sd

16

I 5'

80

44'

Yield is based on acylating agent. Acylating agenuaromatic compoundcatalyst = 5/100/0.1 mmol; reflux 2 h. ' Catalyst was dissolved. "Acylating agenuaromatic compound catalyst = 5/100/0.05 mmol. The water liberated was continuously removed by means of DeanStark equipment. 'Acylating agentlaromatic compoundkatalyst = 5/100/0.10 mmol. 'Catalyst was partly dissolved.

176


Cs2.5 for the acylation. Anisole and p-xylene are acylated with benzoic anhydride and acetic anhydride in the presence of Cs2.5 without the dissolution of this catalyst. Carboxylic acids are much less reactive as acylating agents than the corresponding anhydrides because of the liberation of water. But when the water is removed, the acylation proceeds smoothly (214). Although the reaction of benzene with acetic acid is attractive in prospect, there is no report of heteropoly compounds as catalysts for this reaction. G . SKELETAL ISOMERIZATION

OF

ALKANES

The skeletal isomerization of straight-chain paraffins is important for the enhancement of the octane numbers of light petroleum fractions. The isomerization of n-butane to isobutane has attracted much attention because isobutane is a feedstock for alkylation with olefins and MTBE synthesis. It is widely believed that the low-temperature transformation of n-alkanes can be catalyzed only by superacidic sites, and this reaction has often been used to test for the presence of these sites. Nowinska et al. (215) reported the isomerization of n-hexane catalyzed by ~ by H3PW12040/Si02.H3PWI2O4dSiO2has an appreciable (NH&PW 1 2 0 4 and activity at 423 K, but the activity of H3PW12040 is low, and Si02-A1203 is inactive at this temperature. Cs2.5 catalyzes the isomerization of n-butane at 573 K, and the rates of isobutane formation and the selectivity were much higher than those of SOi-/ZrO2, as shown in Table X X I I I (216, 217). The initial activity of SO;-/ZrO2 is very high, but the conversion decreases considerably during the initial stage of reaction. In contrast, the deactivation is relatively small for Cs2.5. Figure 36 shows the effects of reaction temperature for catalysis by Cs2.5. Deactivation was observed at 473-573 K, but not at 423 K. At temperatures less than 473 K, S0:-/Zr02 is more active than (32.5. The TABLE XXlII Activiw and Selectivi@,for Skeletal lsomerization of n-Butane" (21 7)

Selectivity' (mol%) CataI y st C~z.sHo.sPWi2040 H3PWizOw

so: - lZr02

H-ZSM-5 H-Y"

C,

loMX Rateh

CI

2.0 0.4 0.4 2.9 0.03

1.1 1.1

3.1 0.7 15.8

C2 + C; 2.0 2.4 9.1 2.3 33.3

Cs

+ C;

8.5 11.7 23.0 74.5 18.1

Isobutane 83.1 80.9 60.8 14.1 11.1

C; 0.8 0 0 0.4 18.1

C,(+) 4.5 3.9 4.0 8.0 2.9

"573 K, butane 5%. 'Formation of isobutane, mol g - l s - ' . ' C , , CH4; C2 + C;, C2H4 + CzH6; + CT, C3H6+ C,Hg; C4=,CdHg; Cs( +), hydrocarbons containing more that 5 carbons. d673 K.


0 0

--- ---A

1

A

#A

A

2

3

A

A

4

A

177

A

5

Tim& FIG. 36. Time course of n-butane isomerization catalyzed by Csz sHo 5PW12040 at various 523 K, (W) 473 K, (0) 423 K, (A) 373 K. (From Ref. 217.) temperatures: ( 0 )573 K, (0)

skeletal isomerization of alkanes catalyzed by metal-promoted heteropoly compounds is described in Section XI.

H. MTBE SYNTHESIS Methyl tert-butyl ether (MTBE) is a good, widely used octane improver of gasoline. lgarashi et al. (218) reported that H3PW12040, H3PMoI2O40, etc. and their SiO2-supported analogs have catalytic activities superior to those of mixed oxides, fluorinated oxides, and mounted minteral acids for the MTBE synthesis from isobutylene and methanol in gas-solid reactors. Among the heteropoly catalysts, 20 wt% H4SiMoI2040/Si02was most effective; the selectivity was 95% at 30% conversion of isobutylene at 363 K. Ono and Baba (219) used Ag3PW12040 supported on carbon as a catalyst for MTBE synthesis. The treatment of the Ag salt with H2 greatly enhanced the activity, whereas no effect was observed for the acid form and Al salt. Effects of acid strength and structure of heteropolyanion on catalytic activity have been examined (220). The activity order is H&W18062 %- H3PW12040> H4SiW12040= H4GeW12040> H5BW12040> H6CoWI2O40, whereby the selectivity for MTBE exceeded 95%. The results are much better than those observed with SOi-/Zr02, Si02-A1203, and HZSM-5 (220). It is probable that the pseudoliquid phase behavior of H6P2W18062 is responsible for its high performance. By supporting H6P2W18062or H3PW12040on S O * , the yield of MTBE increased greatly and became comparable to that of the resin Amberlyst 15, which is representative of today's industrial catalysts for MTBE synthesis.

178


The reaction of tert-butyl alcohol and methanol to form MTBE is also catalyzed by heteropoly compounds (221-223). A relationship was found between the amount of pyridine sorbed in or on heteropoly compounds and tertbutyl alcohol conversion (221). The dependence of the rate on methanol partial pressure resembles that for the absorption of methanol in the bulk, suggesting pseudoliquid phase behavior (223).

I. OTHERREACTIONS A wide variety of acid-catalyzed reactions besides those described above have been investigated with heteropoly compounds as catalysts. A1203-supported H3PW1 2 0 4 0 (probably decomposed) catalyzed propylene-ethylene codimerization at 573 K to form pentenes with a selectivity of 56% (butenes 17%, hexenes 27%) (224). Propylene oligomerization proceeded on various kinds of salts of H3PWI2O40 (225). The activities of the salts decrease in the order A1 % Co > Ni, NH4 > H, Cu > Fe, Ce > K. The A1 salt gave trimers with 90% conversion at 503 K. The selectivities to trimer are about 40% for Al, Ce, Co, and Cu, while that of the acid form is 25%. Dehydrogenation of monomethylamine to give hydrogen cyanide is catalyzed by H3PW12040 at 773 K (226). A1203and Si02-A1203 have no activity for the reaction under the same conditions. (NH4)3PW12040 is active for synthesis of methylamines from ammonia and methanol (227). The formation of trimethylamine is suppressed with this catalyst, which is explained by the strong adsorption of trimethylamine. Cracking of paraffins (228, 229), isomerization of olefins (230, 231), transformation of alkylbenzene (232), etc. have also been reported.

VI.

Pseudollquld Phase

In ordinary heterogeneous catalysis of gas-solid and liquid-solid reactions, the reactions take place on the two-dimensional surfaces of solid catalysts (both on the outer surface and on the surfaces of pore walls). In contrast, the reactions of polar molecules in the presence of heteropoly catalysts often proceed not only on the surface but also in the bulk phase. We call this “pseudoliquid phase” behavior. The pseudoliquid phase is a unique reaction medium consisting of the three-dimensional solid bulk, as was first proposed in 1979 ( 1 7, 233, 234). Because of the flexible and hydrophilic nature of the secondary structures of the acid forms and group A salts (Section 11), polar molecules like alcohols and amines are readily absorbed into the solid bulk by substituting for water molecules and/or by expanding the distance between polyanions. The number of absorbed molecules is 10-102 times greater than the amount of monolayer


179

TABLE XXIV Types of Acid Catalpis by Salts of Heteropolyacids

Salts"

Polar molecules (e.g., dehydration of ethanol)

Nonpolar molecules (e.g., isomerization of butenes)

A salts (e.g., Na) B salts (e.g., Cs)

Pseudoliquid (bulk type I) Suface type

Surface type Surface type

Besides the above types, there is bulk type I1 (Section IX).

adsorption estimated from N2 adsorption. Heteropoly compounds absorbing significant amounts of polar compounds behave in a sense like concentrated solutions. As classified in Table XXIV, the pseudoliquid behavior is found for acid catalysis of reactions of polar molecules by group A salts at relatively low temperatures. Another bulk-type catalysis (type 11) is described in Section IX. A.

ABSORPTION OF POLARMOLECULES

When pyridine vapor at 298 K is brought in contact with solid H3PW12040or H3PMo12040after dehydration at 403 K, several molecules of pyridine per polyanion (in the entire bulk) are sorbed (5, 125). After evacuation at 298 K, the number of pyridine molecules per polyanion becomes about 6 (twice the number of protons). Single-crystal X-ray and IR analyses revealed that H3PW12040.6C5H5N consists of [(C5H5NH)3* PW120401(39). Considering the surface area of H3PW12040( 5 m'g-I) and the cross section of a pyridine molecule (3 1 A '), the uptake of 6 pyridine molecules per polyanion corresponds to about 80 times that corresponding to a surface monolayer (235). Some heteropoly compounds swell by absorbing a great number of polar molecules. Deliquescence is observed upon excess absorption. Large uptakes of alcohols into H3PW12040have also been reported (125, 235-238). In Table XXV, the rates and amounts of absorption of various molecules are summarized (235).Polar molecules such as alcohol, ether, and amine are readily absorbed. In contrast, nonpolar molecules like hydrocarbons are sorbed only on the surface. The initial rates of absorption of molecules are plotted against the molecular size in Fig. 37. The initial rates of alcohol sorption greatly decrease as the molecular size increases from 20 A2 (methanol) to 35 A2 (1-butanol). The rates are higher for amines than for alcohols, regardless of the molecular size. This difference is due to the greater basicity of amines. Thus, it may be stated that the rate is primarily determined by the basicity (or polarity) and secondarily by the molecular size (235). Diffusion coefficients of molecules in the lattice of H3PW12040are ca. lo3 times less than those of molecules in the micropores of zeolites (235).Niiyama et al. reported that the effective diffusion coefficient is in the order of

180

TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO TABLE XXV Rate and Amount ofdbsorption of Molecule into H j P W 1 2 0 4 ~(235) Absorption amounts‘

Molecule

(PK,)

Size”

p”

P‘

Rate*

d

1’

( - 2.0) ( - 3.0) ( - 3.0) ( - 3.2) ( - 3.8) ( - 3.8)

20 25 30 31 35 30 37

1.71 1.73 1.73 1.67

30 (20) 30 (43) 20 (80) 30 (57 4 (50) 760 (15) 350 (59) 30 (52) 30 (76) 30 (9) 30 (5) 30 (28) 15 (62) 65 (43) 60 (0.13) 0.8 (0.1)

5.2 5.5 3.8 3.1 2.0

6.0 14.1 14.0 12.1 5.4 4.0 6.3 4.9 8.3 12.0 9.5 15.2 8.5 0.50 0.04 2.8

3.1 (26) 6.1 (64) 6.2 (78) 6.1 (76) 2.8 (41) 2.6 (33) 5.8 (89) 4.6 (86) 4.8 (66) 7.0 (89) 6.9 (86) 8.6 (125) 6.0 (78) 0.10 (1.4) 0.03 (0.2)

( - 3.6)

( - 4.3) ( - 2.9)

(10.7) (10.6) (10.8)

(5.2)

45

1.81

1.30 1.17 I .20 I .40

33 32 33 35 31 34 16

1.32 2.32 0 0

13

-

1.34 -

1.6

6.2 1.1 2.2 3.1 4. I 6.0 4.0 -

-

-

“Cross section (A’). hDipole moment, Debye. “Introduced at 301 K, Torr. The figures in parentheses are the relative pressure (%). *The initial absorption rate; number of molecules (anion. 10 min)- I. ‘Number of molecule. anion- I. ’Saturated amount. ‘Irreversible amount. The figures in parentheses are the numbers of surface layers (see text). ‘Feed gas: 0.1% NO + 5% 0 2 + 5% H20 + 89.9% He at 423 K (Yang, R. T., and Chen, N., Ind. Eng. Chem. Res. 33, 825 ( 1994)).

10-”-10-13 m s-I, much less than those in the gas phase but close to those in the liquid phase (239). The amounts of absorbed molecules in H3PW12040tend to be integral multiples of the number of protons (3, 6,9, etc.), suggesting that these molecules form stable secondary structures throughout the bulk. Transitions between the different absorption states take place as a result of pressure changes (235).These transitions are closely related to the catalytic behavior for the dehydration of alcohol, as described below.

B. EVIDENCE FOR

THE

PSEUDOLIQUID PHASE

There is circumstantial evidence in our early studies indicating the pseudoliquid behavior: (1) rapid absorption of a large quantity of polar compounds as described above; (2) expansion of the solid volume of the materials upon absorption; and (3) quite high catalytic activity of H3PWI2O4,,, H3PMo12040, etc., despite their low surface areas (234). Firm evidence that catalytic reactions


.,

10

20

30

40

181

50

Molecular Cross Section /A2 Initial rates of absorption of various compounds by H3PW12040rt ted to cross section FIG. of molecules. ( I ) Ethylene, (2) dichloroethane, (3) benzene, (4) toluene, ( 5 ) methanol, (6) ethanol, (7) I-propanol, (8) 2-propano1, (9) 1.4-dioxan. (10) I-butanol, ( I I ) I-propanamine, (12) 2-propanamine, (13) I-butanamine, (14) pyridine. (From Ref. 235.)

take place in the pseudoliquid phase has been obtained by a transient response analysis using isotopically labeled reactants (236, 240). Figure 38 illustrates a typical result observed for the dehydration of 2-propanol catalyzed by H3PW12040,The feed gas was instantaneously changed from 2-propanol-do to 2-propanol-dx after a steady state of the reaction had been attained. 2-Propanoldo at the outlet was slowly replaced by 2-propanol-dx (solid lines, and open and solid circles in Fig. 38), whereas the change was rapid in the absence of the do-4 100

4

80

. g

E

60 ..4-

v)

5 0

40

20 0

0

5

10 15 Time I min

20

30

FIG.38. Transient response in the gas-phase composition resulting from a change of feed from 2-propanol-do to -4in the dehydration catalyzed by H3PW1~0m at 353 K (2-propanol, 3.4%; flow rate, 100 cm3 min-'; 50 mg of H3PW12040).(From Ref. 236.)

182


catalyst (broken lines). The amount of 2-propanol-do held by the catalyst under the reaction conditions corresponds to the shaded area, and the amount of 2-propanol-d8 newly absorbed equals the dotted area. The agreement between these two areas led to the estimate of about 7 molecules per polyanion. This value corresponds to about 100 surface layers, which shows that 2-propanol was mostly in the bulk during the reaction. Furthermore, the rates of absorption and desorption were estimated from these data to be 50 times the reaction rate. It is noteworthy that the concentration of 2-propanol in the bulk (6 X mol ~ m - under ~ ) the reaction conditions is comparable to that of liquid mol cm-’). These results justify the term “pseudoliquid phase.” It was confirmed by the same methods that dehydration of ethanol also proceeds in the pseudoliquid phase of H3PW12040(240). Saito and Niiyama (241) investigated the transient behavior of ethanol dehydration catalyzed by Ba1.5PW12040. When the ethanol feed was stopped after a steady state had been attained, ethylene continued to form for a prolonged period, whereas ether, formation decreased rapidly. This transient behavior, as well as the kinetics under stationary conditions, was well simulated with a model based on the assumption that the ethylene and ether are formed by unimolecular and bimolecular reactions in the bulk, respectively.

c.

UNUSUAL

KINETICSIN

PSEUDOLIQUID PHASE

There is evidence that at least two different pseudoliquid phases may be present, even during catalytic reactions, and these may change reversibly with changes in the reactant partial pressures, as shown for dehydration of 2-propanol in Fig. 39 (242). A small change in reactant partial pressure led to an abrupt

B 6 -

4 2 0-

0.2

0.5

1

2

5

10

102 x Pressure of a - ~ r ~ ~ a t r n

FIG.39. Pressure dependence of the catalytic reaction rate and the amount of absorbed propanol in the dehydration of 2-propanol catalyzed by H3PW12040at 353 K. (From Ref. 242.)


183

change from a high- to a low-activity state. The transition partial pressure tends to be higher as the reaction temperature increases. The amount of 2-propanol absorbed, which was determined by the transient response method, changed along with the reaction rate. At 353 K, the numbers of molecules absorbed per anion were 3 and 8 for the high- and low-activity states, respectively. At 373 K, these two states vaned reversibly and rapidly upon the change in the pressure of 2-propanol; but at a lower temperature, the transition was slower upon the change of the partial pressure. The unusual pressure dependences of the rate and selectivity associated with the pseudoliquid that were observed for ethanol dehydration catalyzed by H3PW12040 are shown in Fig. 40 (169, 243). The rates of ether and ethylene

-1

1

0

2

WPlkPa)

FIG.40. Rates of formation of diethyl ether and ethylene from ethanol catalyzed by H3PW12040 as well as the amount of absorbed ethanol under the working conditions as a function of the partial pressure of ethanol at 403 K. (From Refs. I I Y , 243.)

184


formation increased at first with increasing ethanol partial pressure, but decreased markedly at higher ethanol partial pressures. The maximum rate was observed at a higher pressure for the ether formation than for the ethylene formation. The amounts of ethanol absorbed corresponded to 4-80 times the monolayer value (20-400 times the number of the surface polyanions), demonstrating that ethanol was absorbed in the bulk. These results are contrasted to the pressure dependences observed for ordinary solid acids; on A1203and SiO2-AI203 the formation of ethylene is usually zero-order in ethanol and the formation of ether is zero- to first-order in ethanol (244). Furthermore, it is emphasized that the activity of H3PW12040is lo2 times greater than that of Si02-AI203. Since ethylene is formed from one molecule of ethanol and ether from two molecules, it is understandable that ethylene is preferentially formed when the ratio of ethanol to protons in the pseudoliquid phase is low and ether is favored as this ratio increases. Equations (23)-(25) represent a possible mechanism that explains the essential trend in Fig. 40. C ~ H ~ O+HH + P C ~ H ~ O H-'c?H~ ; +H ~ O C~HSOH+ C~HSOH; P (C~HSOH); + (C2H&O

(n - 2)C2HSOH + (C2HSOH)2HCP C2HSOH),Ht

(23)

+ H20

(24)

(not reactive)

(25)

Simulation of the pressure dependence-assuming that the reactions of the first steps of Eqs. (23) and (24), and of Eq. (25) are in equilibrium-reproduced essential trends of the rates and the amounts of absorption.

ANALYSISOF PSEUDOLIQUID PHASE D. SPECTROSCOPIC Pseudoliquid phase behavior facilitates the spectroscopic investigation of the catalysts as the phenomena occur nearly uniformly in the bulk. The IR spectrum of absorbed diethyl ether and its changes during the thermal desorption are shown in Fig. 41 (78). A distinct peak at 1527 cm-' is characteristic of protonated dimer species { v(0-H) mode of [(C2H5)20..*H...0(C2H5)2]} (245). Moffat et al. also detected this peak by IR-photoacoustic spectroscopy (PAS) (238). It is reasonable that this band was not observed for diethyl ether adsorbed on A1203 or SiOz due to the absence of strong protonic acids on the surfaces of these solids, or for D3PWI2O4". 6(C2H&O due to the formation of [(C2H5)20...D-..0(C2H5)2] instead of [(C2H5)20***H..*O(C2H5)2]. A protonated monomer, [(C2H5)20H]3PW12040 (ethedproton = I), remained after evacuation at 328 K, and the dimer (indicated by the band at 1527 cm- I ) disappeared. Further evacuation at 423 K gave four peaks (2954, 2924, 2897, and 2870 cm- ') due to an ethoxide bonded to the polyanion.


185

7

I

I

1700 1500 Wave nurnberlcm'

I

1300

FIG. 41. Changes in the IR spectra of diethyl ether absorbed in H3PWIZ040during stepwise heating in vacuum. (From Ref. 78.)

NMR spectroscopy also provides useful information. As shown in Fig. 42, H3PW12040.6C2H50Hgives three 'H NMR resonances at 9.5,4.2, and 1.6 ppm, which are assigned to OH, CH2, and CH3, respectively (169, 246). This is probably the first observation of a well-resolved solid-state 'H NMR spectrum for protonated organic compounds in this catalyst system. The relative intensities (OH/CH2/CH3 = 1.45 : 1.83 : 3.0), as well as the stoichiometry of ethanol to proton (2 : I), are consistent with the protonated dimer, (C2H50H)2H+.The high resolution is explained by the high mobility of the protonated ethanol and the homogeneity of the bulk phase. The chemical shift of the hydroxyl proton of H3PW12040* 6C2H50H (9.4 ppm) is close to those reported for protonated ethanol in superacids: 8.3 in HF-BF3,9.3 in FS03F-SbF5-SO2, and 9.9 ppm in HS03F (247) (as compared with a value of 1.0ppm for a dilute ethanol

186


1 6b*9d 31P

'H

l3C

CH2

-15

PPm

CH3

CH3

ppm

ppm

Reprinted I ~ ~ ~ ~ with . FIG. 42. "P, I3C, and 'H MAS NMR spectra of [ ( C ~ H S O H ) ~ H ] ~ P W permission from Ref. 169. Copyright 1992 American Chemical Society.

solution). Hence, the pseudoliquid phase of H3PW12040may be regarded as a superacidic medium. The chemical shifts of I3C (61.9 and 17.2ppm for CH2 and CH3, respectively) are different from those of pure ethanol (57.0 and 17.6ppm, respectively). Changes in the 13C NMR spectra upon heat treatment are shown in Fig. 43. Peaks at 65.0 and 16.8 ppm detected at 333 K indicate * 3C2H50H(169). Further heating gave a new set of peaks at 82.1 H3PW12040 (CH2) and about 14.3 ppm (CH3). The peak at 82.1 ppm is assigned to an ethoxide. The shift from 65.0 to 82.1 ppm is of similar magnitude to the transformation observed for transformation of adsorbed methanol to methoxide on K3-,H,PM01204~ (from 51 to 75 ppm) reported by Farneth et al. (168). These shifts are significant but smaller than that observed for alkyl cation formation. For example, sec-propyl and tert-butyl cations in a superacid solution show a downfield shift of about 260 ppm. Hence, the species at 82 ppm is more like ethoxide than ethyl cation (247, 248). Main reaction paths of the thermal desorption of ethanol are proposed in Scheme 5. The species observed directly by NMR spectrscopy are surrounded by broken lines. At temperatures less than 323 K, the dehydration did not proceed, and only reversible desorption took place. The protonated ethanol dimer is transformed into protonated ether at temperatures exceeding 323 K. Diethyl ether is formed only in the gas phase by replacement with ethanol. Protonated ethanol monomer probably gives ethylene via the ethoxide at temperatures exceeding 333 K (169). E. SURFACE AND BULK TYPEI REACTIONS As shown in Fig. 44a, the catalytic activity of NaxHj-,PW12040 for dehydration of 2-propanol, conversion of methanol, and decomposition of formic acid decreased monotonically with the Na content in the salts. The activities for these


I87

12.0 i

b

-

100

60

20

-20

P P FIG.43. Transformation of protonated ethanol dimer in HjPW120~by heat treatment. Solidstate "C CP/MAS NMR spectra were obtained by using high purity ''C ethanol: (a) Dimer, (b) 333 K, (c) 343 K, (d) 363 K, (e) 373 K, (f) 423 K . Reprinted with permission from Ref. 169. Copyright 1992 American Chemical Society.

reactions correlate well with the bulk acidities measured by the thermal desorption of pyridine (46b).All the reactants are polar molecules, so that the reactions proceed in the bulk. On the other hand, the activity pattern for butene isomerization is quite different (Fig. 44b). Butene is nonpolar and not absorbed, and it reacts only on the catalyst surface. The irregular variations probably reflect the surface acidity which changed depending on the Na content and pretreatment. Niiyama et al. (223) found that the reaction rate characterizing MTBE synthesis from methanol and tert-butyl alcohol catalyzed by H3PW12040 increases in proportion to the amount of methanol absorbed in the bulk of Hd'W12040.

188


The influence of bulk-type behavior also exists in supported heteropoly catalysts. Changes in the activity as a function of the loading of heteropolyacids depend on the reaction type (151). At low loadings, the rates of both the bulk and the surface reactions increase as the loading increases because the

Number of pyridine molecules anion-1

FIG. 44. Relationships between catalytic activity and bulk acidity. (a): (0) Dehydration o f 2-propanol, (A) decomposition o f formic acid, (0)conversion o f methanol. (b): (M) lsomerization of cis-2-butene after treatment at 423 K, ( 0 )isomerization of cis-2-butene after treatment at 573 K. (From Ref. 46h.)


189

:tivity/ mmol g-lh-l

FIG. 45. Catalytic activities of solid acids for reaction of liquids: Alkylation of 1,3.5-trimethylbenzene with cyclohexene (373 K);alkylation of phenol with I -dodecene: rearrangement of benzopinacol. (From Ref. 249.)

dispersion of the heteropolyacid becomes high. At high loadings, the dispersion becomes low, so that there is little increase in the rate of the surface reaction at increasingly high loadings. Pseudoliquid phase behavior is also evidenced by the liquid-solid reactions of polar compounds (249). An example is shown in Fig. 45. Heteropoly compounds are much more active than HzSO4 and other solid acids. The rate of alkylation of 1,3,5-trimethylbenzene increases in proportion to the surface acidity. Thus Cs2.5H0.5PW12040is much more active than H3PW12040. On the other hand, H3PW12040 is more active for the alkylation of phenol and the rearrangement of benzopinacol than C S ~ . ~ H ~1.2~0 4P0W . Since the latter two reactants are polar, the pseudoliquid phase is formed for H3PW12040 in the latter reactions. The activity of H3PW12040relative to that of Cs2.5Ho.5PW12040 decreases for these reactions in the order alkylation of trimethylbenzene > alkylation of phenol > pinacol rearrangement. The results indicate that the pseudoliquid phase is most important for the first reaction in this series and least important for the last. When the rearrangement of pinacol in I ,2-dichloroethane was examined in detail at 323 K, the amount of pinacol held by the catalysts was about 20 times the amount corresponding to a surface monolayer for H3PW12040rwhereas this ratio was less than 1 for the other catalysts. Niiyama et al. (250) attempted to separate isobutylene from a mixture of isobutylene and I-butene by using a H3PW12040-porousglass hybrid membrane, in which H3PWI2040 was loaded into the pores of the glass. A mixture of butenes and water vapor was brought in contact with one side of the membrane. Since the hydration of isobutylene takes place preferentially, tert-butyl alcohol formed on one side of the membrane is absorbed in the H3PW12040 and diffused through the bulk to the opposite side. After the tert-butyl alcohol reached the surface of H3PW12040 at the opposite side, it decomposed to give isobutylene and water.

190

TOSHIO OKUHARA. NORITAKA MIZUNO, AND MAKOTO MISONO

F. CONTROL OF ABSORPTION PROPERTIES AND CATALYTIC REACTIONS Absorption is influenced significantly by the cations in the salts (235). Typical examples for absorption of ethanol are shown in Fig. 46. It is clear that the absorption of H3PW12040is greatly suppressed when H + is replaced by Cs', while the absorptivity vanes little when H + is replaced by Na'. In contrast, there is little difference between the Na and Cs salts when they are characterized by the acid amounts measured by pyridine absorption. The pressure dependences as well as the amounts of ethanol held by CsxH3-.PW12040 (CsX) during the catalytic dehydration of ethanol were measured experimentally (240). The pressure dependences for Cs2 and Csl resemble those of the low- to intermediate-pressure region observed for ethanol in H3PW12040,suggesting that the pseudoliquid phase was also present in the acidic Cs salts. Slightly higher pressures are required for the latter to give the same trend because of the lower absorption capability of the Cs salts. The selectivity is affected strongly by the presence of the pseudoliquid phase. An example is given in Fig. 47. The oiefin-to-paraffin ratio in the product of the conversion of dimethyl ether increased markedly as the absorptivity of the heteropoly compounds decreased (195, 197). The ratio observed for catalysis by Na,H3-,PW12040 was similar to those observed for catalysis by H ~ P W I ~ O ~ ~ . The acidity governs the reaction rate but plays a minor role in determining the selectivity. The change in the olefidparaffin ratio can be explained by the formation of the pseudoliquid phase as follows. When olefins are produced in the bulk phase, they have a relatively high probability of undergoing hydrogena

x in MxH3-xPW12040(M = Na,Cs)

x in MxH3-xPW12040(M = Na,Cs)

FIG.46. Total amount (a) and absorption capacities (b) of Na and Cs salts of H3PW12040: (0) Na, (0) Cs. The acid amounts were measured by pyridine absorption (adsorption). (From Ref. 235.)


191

- 4

-3 - 2

10

20

30

40

Uptake of DMVnumbers of surface layers

FIG.47. Olefin-to-paraffin ratio in the product hydrocarbons from dirnethyl ether conversion as a function of absorptivity of the heteropoly compounds, expressed in term of the absorption of dirnethyl ether in surface layer units: (1) PW12 (H3PW12040), (2) NaH2PWI2, (3) Na2HPWI2, (4) C&PWiz, (5) C S ~ S H O S P W I( Z 6 ). CszssHo I s P W I ~(7) . ( N H ~ H Z P W I(8) ~. (NW~HPWIZ. (9) (NH4)2sHo5PW12, (10) 1,4-diazine, ( I I ) 1,3-diazine, (12) 1,4-bis(aminornethyl)benzene, (13) triazine. (From Ref. 197.)

transfer reaction to form paraffins and coke before they desorb. On the other hand, when the reaction occurs near or on the surface, the olefins formed may desorb rapidly into the gas phase without undergoing significant hydrogentransfer reactions. Similar selectivity patterns associated with the pseudoliquid phase were observed in the dehydration of ethanol to give ethylene and diethyl ether.

VII. A.

Redox Properties

REDOXCHEMISTRY IN SOLUTION

A general feature of heteropolymolybdates and heteropolytungstates is their high reducibility. Electrochemical investigations of Keggin-type heteropolyanions in aqueous or nonaqueous solutions have revealed sequences of reversible one- or two-electron ( l e - or 2eC) reduction steps [Fig. 48 (3)] which yield deeply colored mixed-valence species (“heteropoly blues”). Electronic spectra of the reduced heteropolyanions show intensified d-d bands in this visible region and intervalence charge-transfer (IVCT) bands in the near-1R region. Depending on the solvent, the acidity of the solution, and the charge of the polyanion, the reductions involve either single-electron or multi-electron steps, often accompanied by protonation. In protic solvents, the Keggin anions exhibit

192


EN FIG. 48. Cyclic voltammogram of a-PMol2OG in 0.1 M HCI (50% H20/EtOH). (From Ref. 3.)

electrochemical reductions by two electrons. The reduction potentials depend on pH as a result of protonation [e.g., Eq. (26)] (251): PMo120&

-+

H2PMo120:;

-t

H4PMo120&

(26)

Acidification of nonaqueous solutions of an unprotonated 1e --reduced species causes disproportionation to a 2eC-reduced anion and an oxidized anion (251). At higher pH values, polarograms show a series of le- reductions (252). Reduction of Keggin anions beyond 2e- reduction leads to modest changes in electronic and molecular structure, although the reductions remain reversible. For example, in the case of molybdates, a 4e--reduced P-Keggin structure is stabilized, with the bridging oxygen atoms being partially protonated [Eq. (27)] (3, 253, 254): a-PMo120+ ~ ~4e-

+ 3H’

-+/Y-H3PMol20;;

(27)

Electronic and NMR spectra of these complexes, together with X-ray structural determinations, indicate that the valence is completely averaged over at least six Mo atoms (251, 253, 254). ESR and I7O NMR spectra of le--reduced SiW120:0 demonstrate that the unpaired electron is weakly trapped on a W atom at low temperatures but undergoes rapid hopping (intramolecular electron transfer) at room temperature (Section 11). Anions generated by 2e- (and 4e-) reduction are ESR-silent, but 17 0 and Ig3WNMR spectra show that the additional electrons are fully delocalized (on the NMR timescale) at room temperature and generate “ring currents” analogous to those produced by the n-electrons of benzene. In contrast, in the case of le--reduced PMoWIIO:i, the electron is localized on a more reducible Mo atom at room temperature (251). The reduction potentials of heteropolyanions containing Mo and V are high, and they are easily reduced. Oxidizing ability decreases in general in the order V- > Mo- > W-containing heteropolyanions (Fig. 49) (8, 91). As for heteroatoms, the reduction potential (or oxidizing ability) decreases linearly with a


193

Gi

x

. '0

Y

-3

-4

Anion

-5

-6

-7

charge

FIG. 49. Dependence of reduction potentials on anion charge: ( 0 ) XMWllO&; (0) XMMolIO[;R;SCE = saturated calomel electrode. (From Ref. 91.)

decrease in their valence or an increase in the negative charge of the heteropolyanion (Fig. 49) (91). For polyanions with mixed-addenda atoms, the reduction potentials have been reported to be PMoloV20:G > PMollVO:o > PMo120:,; and PMo6W60& > PMo120i0 (2). XWI IMn039(X = Si, Ge) forms an oxygen adduct in nonpolar solvents. This was claimed to be the first example of an inorganic oxygen camer (255). Recently, it was reported that XWI ICrO39 reacted with PhIO, H 2 0 2 , or NaOCl to form XW11(CrO)O39 [Eq. (28)]: PWIICr3+(H?O)O: PMo120i0 > PMoIOV20:0 > PWI~O:;, in parallel with the reduction potentials in solution (except for PMoloV2O:; ) (276). 2. For a given polyanion, the effects of metals are divided into two groups: (a) Transition metals play roles in the redox processes; e.g., they activate reducing agents and molecular oxygen and possibly provide reservoirs of electrons (108, 272, 273, 276, 278). (b) Alkali and alkaline-earth metals are not reduced. Oxidizing abilities measured by the rate of reduction decrease upon the formation of alkali salts (Fig. 5 3 ) (258, 272, 273, 277). The reason for the decrease of the oxidizing ability with alkali content is not hlly understood, although suggestions have been made concerning the electronegativity of the cation and the role of protons in the reduction process. The mixed-addenda atoms affect the redox properties; mixed-addenda heteropoly compounds are used as industrial oxidation catalysts. For example, the rate of reduction by H2 is slower and less reversible for solid P M o 1 2 - , V , 0 ~ ~ " ) than for solid PMoI20:;, although the former are stronger oxidants than the latter in solution (279, 280). The effects of substituting V for Mo on the catalytic activity are controversial (279, 281-284). Differences in redox processes between solutions and solids, the thermal or chemical stability of the heteropoly compounds, and the effects of countercations in solids have been suggested to account for the discrepancies. (M = Mo, W) is It has been demonstrated that V 5 + in H3+xPMLZ-xVx040 eliminated from the polyanion framework upon thermal treatment or during catalytic oxidation, and the V 0 2 + salt of H3PMI2O40is formed (284). It has been reported (I03) that H3PMo12040 is re-formed from thermally decomposed H3PMoI2O4,,under the conditions of methacrolein oxidation.

VIII.

Liquid-Phase Oxidation Reactions A.

OXIDATION WITH

DIOXYGEN

Typical examples of liquid-phase oxidation with molecular oxygen catalyzed by heteropoly compounds are listed in Table XXVI. Introduction of VSf or


20 1

other transition-metal ions such as C 0 2 + or Mn2+ as addenda atoms usually enhances the catalytic activity, reflecting changes in the reduction potential. Recently, paraffin oxidation in the liquid phase, which does not proceed with the V' and mono-Co2+-substituted heteropolyanions, has become possible, catalyzed by heteropolyanions having tri-transition-metal (Fe2Ni or Fe3) sites. The use of heteropoly compounds in combination with transition metals as in Wacker-type reactions is described in Section XI. PV2MoloO:i has usually been used in the acid form. H5PV2MolOO40catalyzes aerobic oxidative cleavage of cycloalkanes, 1-phenylalkanes, and ketones. For example, the oxidation of 2,4-dimethyl cyclopentanone and 2-methylcyclohexanone gives 5-0x0-3-methylhexanoic acid and 6-oxoheptanoic acid, respectively, in yields higher than 90% (285, 286). Bromination of arenes with HBr (289, oxidative dehydrogenation of cyclohexadiene (288, 289) and a-terpinene (290), oxidation of 2,4-dimethylphenol (291) and sulfides (292) are other examples. The mechanism of aerobic oxidative dehydrogenation of a-terpinene to give p-cymene catalyzed by PV2MoloOii has been investigated (290). On the basis of kinetics along with the use of UV-visible, ESR, 3'P-NMR, and IR spectroscopies, a reaction scheme was proposed, as shown in Fig. 55. In this scheme, a stable a-tetrapinene'-H6PV s+ V4 + M010O40 complex (b) is formed via an electron transfer complex (a), which is a reduced heteropolyanion attached to an oxidized cation radical of a-terpinene. A doubly protonated reduced heteropolyanion and p-cymene are generated via the intermediate. On the basis of the kinetics (zero-order reaction in a-terpinene, second-order reaction in PV2MoIoO:~,and first-order reaction in 02)and the IR and NMR data, the ratelimiting step is proposed to be the reoxidation of the catalyst. In the reoxidation heteropolyanions are reoxidized in a 4estep, two reduced PV:!'MoloO:i redox reaction via a p-peroxo intermediate (c). Co2+-Substitution at the addenda atoms gives catalysts for the epoxidation of olefins in the presence of aldehyde (293). PWII-Co is the most active among the mono-transition-metal-substituted polyanions; the order of activity is PWII-Co 9 -Mn 2 -Fe 2 -Cu > -Ni. Here, PWII(M"+)O\~-")-(M = Co2+, Cu2+, Fe3+,Ni2+, Mn2+) is denoted by PWII-M. The same order was observed for the oxidation of isobutyraldehyde, suggesting that the oxidation of aldehyde to give peracid is an important step in the reaction. It has been reported that substitution of V5+ for Mo6+ in PMo120:i gives a good catalyst for epoxidation and the Baeyer-Villiger reaction (294). Styrene and 1-decene are selectively epoxidized, as shown in Table XXVII (293). The rates observed for PWIl-Co are greater than those observed for Ni(dmp), and Fe(dmph, and the selectivities are comparable or higher for the former (295). It is also remarkable that PWII-Co polyanion exhibits a steric effect comparable to that of a moderately hindered TTMPP ligand in the +-

202

TOSHIO OKUHARA. NORITAKA MIZUNO. AND MAKOTO MISONO

TABLE XXVI Liquid-Phase Oxidation Reactions with Molecular Oxygen Catalyzed by Heteropoly Compounds

-

Reaction

~~

PhCOCHzPh + 0

R2S

Catalyst

Temp. (K)

Ref.

+ PhCHO

HSPMOIOVZO~O

333

285

6

H~PMoYV~O~O

363

287

HSPMOIOVZO~O

293

289

HSPMOIOVZO~O

343

288. 289

HSPMOIOVZO~I

298

290

HsPMoioVz040

333

291

PW 1coo:; PMo,VhOI,

303

293,294

- A +A

H7PWyFezNi037

423

297

-

(TBA)4HJ’WyFezNi037 TBA, ( ~ - C ~ H Y ) ~ N

355

299

PdC12 + Na,.H3 + .-,PMoIz (X = 2-3)

393

See Section IX

2

PhCOOH

-

RzSO or RzSOZ

6

+ HBr + 1 / 2 0 ,

-

+ H20

Br

$

($+02-

&+oz-& 0

)=( + 0, + aldehyde 0

n+ 0 2 O + O 2

HzC=CHz

+0 2

-% OH

6 . 6 CH3CHO

-x V , O ~

continued


203

TABLE XXVl-Continued Reaction

Catalyst

Temp. (K)

Ref.

N ~ , P M O ~ V ~ O ~ ~293

0

SO2

+ 112 0 2 + H2O

-

HzS04

Hd'M07Vs040

298

"J. E. Lyons et al., Stud. SurJ Sci. Catal. 67, 99 (1991). h R . Neumnann er a/.. Dioxygen Activation and Homogeneous Catalysis Oxidation L. 1. Simanded, Ed. Elsevier, Amsterdam, 12 1 ( 1991). ' B. S. Jumakaeva et a/., J. Mol. Catal. 35, 303 ( 1986).

epoxidation of (R)-(+ )-limonene, 4-vinyl- 1-cyclohexene, and I -methyl- 1,4cyclohexadiene (296). Activation of paraffins by heteropolyanions has been attempted by using tritransition-metal-substituted heteropolyanions. Propane, ethane, and methane are oxidized to the corresponding alcohols and ketones in the presence of tritransition-metal-ion (Fe2Ni)-substituted PW I 20:;, although the composition and structure of the catalyst are still to be examined (297). Recently, Mizuno et al. (298) characterized its structure, and confirmed that it catalyzed the oxygenation of adamantane, ethylbenzene, and cyclohexane with molecular oxygen alone (299). For example, adamantane was catalytically oxidized mainly to 1-adamantanol with small amounts of 2-adamantanol and 2-adamantanone. The total number of turnovers was 25, calculated on the basis of the bulk polyanion, or 3750, calculated per surface polyanion. This number of turnovers is the highest for the dioxygen oxidation of adamantane on p-0x0 di- or tri-iron and Ru complexes with or without any reductants (300, 301). Fe=O, Fe-OH, or Fe-OOH species are assumed to play an important role in the reaction. B. OXIDATION WITH OXIDANTS OTHERTHANO2

Hydrogen peroxide and alkyl hydroperoxides are important oxidants in organic synthesis, but they usually need to be activated by catalysts such as tungsten, molybdenum, and titanium oxides. Heteropoly compounds are good catalysts for oxygenation of olefins or paraffins and oxidative cleavage of vic-diols.

204


ICatalyst Reduction)

batalyst Reoxidation

h

FIG. 55. Mechanistic scheme for oxidative dehydrogenation of a-terpinene catalyzed by H S P M O ~ O V (From ~ O ~ ~Ref. . 290.)

1. Hydrogen Peroxide

Typical examples are collected in Table XXVIII. Mixed addenda heteropolyanions show unique catalytic properties. For example, in the oxidation of cyclopentene to glutaraldehyde with hydrogen peroxide, the catalytic activity of

205

CATALYTIC CHEMISTRY OF HETEROPOLY COMPOUNDS TABLE XXVII Expoxidation of Olefins Catalyzed by PW,,-Co at 303 K" (from ReJ 293)

Alkene

Conversion" PA)

Selectivity"(%)

Cyclohexened Styrene' 1 -Decene/

78

90

100

91

73

80

" Olefin, 250 pmol; isobutyraldehyde, 1000 pmol; solvent, 1.2-dichloroethane; 1 h. Based on the starting olefin. Based on the converted olefin. "PWII-Co,1.2 pmol; 1.3 h. "PWII-Co,4.5 pmol; I h. 'PWII-Co, 1.2 pmol; 4 h.

"

H3PW6M06040 was observed to be the highest among H3PW12-xM~r040 species, as shown in Fig. 56 (302-304). A theoretical explanation of this activity pattern or synergistic effect has been attempted by accounting for excess electronic energy of nonbonding levels (305). However, there is a possibility that the active peroxo catalysts are formed by the degradation of the Keggin structure (304). Phase-transfer catalysis has been developed by the combination of Keggintype heteropolyanions and quaternary countercations such as tetrahexylammonium or cetylpyridinium ion. The oxidations of alcohols (306), ally1 alcohols ( 3 0 3 , olefins (308), alkynes (309),j3-unsaturated acids (310), vic-diols (311), phenol (312),and amines (313) are the examples. Ishii et al. (306, 307, 310, 311, 313) and Venture110 et al. (308) have developed various oxidation reactions using organoammonium salts of PM ,O:, (M = Mo, W) for phase-transfer catalysts. A reaction mechanism of the epoxidation of olefin is proposed in Scheme 7. In epoxidation of long-chain olefins, epoxides are produced in the organic phase by peroxo species which are only slightly soluble in the organic phase. Since the salts dissolve only slightly in the organic phase, they do not catalyze undesirable ring opening of epoxides. At a conversion of olefin of 82-98%, the selectivity to epoxide is more than 98%. A simple molybdenum compound, H2MoO4, shows no catalytic activity under these reaction conditions. It is probable that neither metal oxides nor peroxomolybdates are soluble in the organic phase (288). Recently, it was claimed that the active catalysts are not the starting Keggin-type heteropolyanions, but rather PO4[WO(O2)2]: and/or [W203(02),( H20)2l2 -, which are peroxo compounds formed in aqueous solution (308, 314-316). P04[WO(02)& was two orders of magnitude more active than [W203(02)4(H20)2]2- for olefin epoxidation. It is suggested that the catalytic cycle is mainly PO4[WO(O,)2]:[P,W,Os(O2),],- [r is 4 or 3 (316)], although the composition of the species formed by the reaction depends on the ratio [H2O2]/[H3PWI2O40](316, 31 7). ~

-

206


TABLE XXVlll Liquid-Phase Oxidation Reactions with Hydrogen Peroxide Catalyzed by Heteropoly Compounds Reaction

Catalyst

Temp. (K)

298

Ref.

307a

OH

--

F OH

0

continued

207

CATALYTIC CHEMISTRY OF HETEROPOLY COMPOUNDS TABLE XXVlII-Coniinued Catalyst

React ion

0-0. 0"-

Temp. (K)

Ref.

303

312

r.t.

313

27 I

319

reflux

(1

D C O C H 3 353

h

0 -

-d 0

" D. Attanasio el al., J. Mol. Caial. 51, LI ( I 989). S. Sakaguchi er al., J. Org. Chem. 59, 568 I (1 994).

75

,

I

0

3 6 9 12 x in H Q P M O ~ ~ - ~ W ~ O U

FIG.56. Effects of addenda atoms on the catalytic oxidation of cyclopentene by H202 at 303 K. ( 5 ) H3PMol2040 + H3PW12040.Reaction time, 3 h. (From Ref. 304.)

208

TOSHIO OKUHARA, NORITAKA MIZUNO, AND MAKOTO MISONO H3pw12040

Aqueous phase

..................................................

P04[W0(02)21~ _._____. (w203(02)4(H20>2)2-

SCHEME 7

Cetylpyridinium salts of H3PM12040 (M = W, Mo) also catalyze the oxidation of secondary alcohols or alkynes to give the corresponding ketones, but they are not active for primary alcohols (308b, 309b). Schwegler (308c)reported that lacunary heteropoly- 1 1-tungstates are better catalysts than tetrahexylammonium 12-tungstophosphate for the oxidation of cyclohexene in biphasic systems. As described above, catalytic oxidation with hydrogen peroxide is usually limited to polyoxometalates containing only metals in the d o state. Recently, however, a tetrairon-substituted heteropolyanion, Fe4(PW9034):o-, and dimanganese-substituted heteropolyanion, [ W Z I M ~ ~ ( Z ~ W ~ O (sandwich~~)~]'~type compound with a WZnMn2 ring between two B-ZnW9034 units), have been reported to catalyze selectively the epoxidation of olefins (328). Table XXIX summarizes the oxidation of various reactants with hydrogen peroxide and [WZnMn2(ZnW9034)2]12-. At 275 K, the selectivities were more than 99% in all cases. High yields of more substituted olefins indicate that the reactivity of

209

CATALYTIC CHEMISTRY OF HETEROPOLY COMPOUNDS TABLE XXIX Oxidation of Various Reactants with 30% H202 Catalyzed by pZnMn2(Zn W90,4)J’2- at 275 K“ (f?om ReJ 318) ~~

Substrate Cyclohexene

I-Octene 2-Methyl-I -heptene trans-2-Octene Cyclohexanol

Products

Turnovers

Cyclohexene oxide Cyclohexen-2-01 Cyclohexen-2-one I-Octene oxide 2-Methyl-I-heptene oxide trans-2-Octene oxide tran+2-Octen-4-one Cyclohexanone

1450

0 5

25 190 340 0 510

“Catalyst, 0.2 p o l ; reactant, I mmol; H202. 2 mmol; 5 pmol, methyltricaprylammonium chloride.

the reactant is a function of the nucleophilicity of the carbon-carbon double bond.

2. tert-BuWl Hydroperoxide and Others Mixed addenda or transition-metal-ion-substituted heteropolyanions containing Co, Mn, and Ru are catalysts for oxidation reactions with tert-butyl hydroperoxide and other oxidants. Typical examples are listed in Table XXX. TABLE XXX Liquid-Phase 0-ridarion Reurtions with rert-Burylhydroperoxide o r Other Oxidants Catalvzed bv Heteropoly Compounds Reaction

0

+

fert-BuOOH or PhIO

0 +PhIO

-

00”

0

0

Catalyst

+0

0 PWI~COO:~

PWlIMO:9 PW17MO:;(M = Co, Mn)

D. Mansuy et al., J. Amer. Chem. SOC.113, 7222 (1991).

Temp. ( K ) Ref.

298

320, 32 I

297

321“

2 10

TOSHIO OKUHARA. NORITAKA MIZUNO. A N D MAKOTO MISONO

Strongly acidic H3PM12040(M = Mo, W) species catalyze oxidation of thioether into sulfoxide and sulfone with 98-99% and 1-3% selectivities, respectively. The V5+ substitution increases the selectivity to sulfone up to >99% (319). Cobalt- or manganese-substituted PWl20:O and SiWl1039Ru(OH2)S catalyze the oxidation of paraffins such as cyclohexane and adamantane (320, 321) as well as the epoxidation of cyclohexene with tert-butyl hydroperoxide, iodosylbenzene potassium persulfate, and sodium periodate (321, 322). The reactivity depends on the transition metals. In the case of epoxidation of cyclohexene with iodosylbenzene, the order of catalytic activity of PWII(M)O:; is M = Co > Mn > Cu > Fe, Cr. A Ni-containing sandwich complex, Ni3(a-PW9039)2,is a better catalyst than PWll(M)O:; (M = Co, Mn) for the formation of N-alkylacetamide from adamantane and isobutane with tert-butylhydroperoxide as the oxidant (323). As for the mechanism of oxygenation of paraffins with oxygen donors (DO) such as iodosylbenzene and potassium persulfate, Eqs. (33) and (34) have been proposed, whereby oxgenation of metal centers by oxygen donor is followed by 0x0 transfer from the transition metal to CH bonds of paraffins: ~

IX.

LM+DO

+

RO+D

LM=O t RH

+

ROH

+ LM

(L = heteropolyanion)

(33) (34)

Oxidation Catalyzed by Solid Heteropoly Compounds

Solid heteropoly compounds are suitable oxidation catalysts for various reactions such as dehydrogenation of 0- and N-containing compounds (aldehydes, carboxylic acids, ketones, nitriles, and alcohols) as well as oxidation of aldehydes. Heteropoly catalysts are inferior to Mo-Bi oxide-based catalysts for the allylic oxidation of olefins, but they are much better than these for oxidation of methacrolein (5). Mo-V mixed-oxide catalysts used commercially for the oxidation of acrolein are not good catalysts for methacrolein oxidation. The presence of an a-methyl group in methacrolein makes the oxidation difficult (12). The oxidation of lower paraffins such as propane, butanes, and pentanes has been attempted (324).Typical oxidation reactions are listed in Table XXXI and described in more detail in the following sections. Keggin-type heteropoly compounds having Mo and V as addenda atoms are usually used for such oxidations. The catalysts reported in patents often contain several elements other than Mo, V, and P. An excess amount of P is added to stabilize the structure, and the presence of additional transition elements like Cu improves redox reversibility. Supported heteropoly catalysts are also important for industrial applications and have been characterized (69, 325, 326). To understand oxidation catalysis by solid heteropoly compounds, the contrast between surface and bulk type I1 catalysis, and acid-redox bifunctionality

21 1

CATALYTIC CHEMISTRY OF HETEROPOLY COMPOUNDS TABLE XXXl Heterogeneous Oxidation Reactions Catalyzed by Heteropoly Compounds

Reaction

Catalyst

Temp. (K) 533

= + 02

-

CH3COOH

+ O2 N

-

C

I

12

573

333

573

34P

Pd + H4SiW1204dSi02

423

h

PbFeBiPMo120,

673

C=C-CHO

- no -

563

J

+ 0 2

0

0

CH4 + NzO HCHO, CH3OH CHxCHO CzH6 + N 2 0 ( 0 2 ) -C2H4, ---+

- - no A -A

Ref.

+ O2

CH2=CHCOOl

c

H3PMo12040/Si02 H3PMo1204dSi02

843 540

I

Hd’M01z040 ( +As)

613

352, 353

633

324, 350

623

354-356

583

106

473-563

345

443

110

+ 0 2

0

0

C

I

+ o2

+o,-

CH3OH

+0 2

-OH

+ 0,

H3PM012040

C=C-COOH

0a 0

-

0

HCHO, (CH3)2O, etc.

CH3CH0,(C2H5)20

H P M o 12040 H~PMoIzO~O + polysulfone)

(

OM. Akimoto et al., J. Cural. 89, 196 (1984). ’T. Suzuki ef al., US Patent, 5405996 (1995). ‘T. Ohara, Shokubai (Catalysf) 19, 157 (1977). ’M. Ai, J. Catal. 85, 324 (1984). ‘S. S. Hong ef a/., Appl. Catal. A 109, I17 (1994). ’S. Kasztelam et al., J. Cats/. 116, 82 (1989).

of heteropoly catalysts must be properly taken into account, along with relationships between the oxidizing properties of catalysts and their catalytic activities (5, 6, 258, 266, 327, 328). A. 1.

CONCEPT OF SURFACE AND BULK TYPEI1 CATALYSIS AND REDOX (MARS-VANKLEVELEN)MECHANISM

Bulk Type II Catalysis

As described above, heterogeneous catalytic reactions on heteropoly compounds are classified into three different types, surface, bulk type I (pseudoliquid phase), and bulk type I1 (Fig. 1). The surface reactions are typical of

212


ordinary heterogeneous catalysis, and therefore the surface acidity and oxidizing ability are important. Bulk type I reactions proceed within the catalyst. In the bulk type I1 reduction, the rapid migration of redox carriers enables the whole catalyst to be reduced (260). The differences between the noncatalytic surface and bulk type I1 reduction are reflected in catalytic oxidations proceeding via redox mechanisms, as described below (258, 260-262, 329). If a catalytic oxidation proceeds by the cyclic reduction and reoxidation of the catalyst (a redox mechanism), the rate of catalytic oxidation and the rates of reduction and reoxidation of the catalyst must coincide if they are measured in the stationary state of catalytic oxidation. Measurements were made for the oxidation of H2 catalyzed by H ~ P M Ohaving , ~ ~ different ~ ~ specific surface areas (Fig. 57). While the degree of reduction in the stationary state differs among the catalysts, the rates of reduction of the catalysts by H2, the rates of reoxidation, and the rates of catalytic oxidation agree quite well. The agreement indicates that the catalytic oxidation of H2 proceeds by a redox mechanism. A similar redox mechanism was confirmed for H3PW12040 and alkali salts of H3PMol2040. In addition, the rate of reduction depends only slightly on the surface area (Fig. 5 2 ) , although the rate of reoxidation is proportional to the surface area. If these rates are compared for a particular degree of reduction at the steady state of the catalytic oxidation, the differences in rates from one catalyst to another

d

0

.CI

Y

I

;; f

cw

-a

41

0

,

,

0.1 0.2 Degree of reduction 1 e- anion1

FIG. 57. Rates of catalytic oxidation of H2. reduction by H2. and reoxidation by O2 for H3PMo12040catalysts having different specific surface areas: Dependence on the degree of reduction of each catalyst at the stationary state. The numbers next to the symbols indicate the surface areas (after reaction) in m' g I . ( 0 )Catalytic oxidation; (-) reduction; (0)reoxidation. Reaction temperature, 573 K. (From Ref. 262.)


213

are small. Consistent with this pattern, the rate of the catalytic reaction is only weakly dependent on the surface area, as shown in Fig. 58. On the other hand, the rate of catalytic oxidation of CO is proportional to the specific surface area, as shown in Fig. 58. This dependence indicates ordinary heterogeneous catalysis. The linear dependence in Fig. 58 can also be explained on the basis of the redox mechanism, as both the rate of CO conversion and the rate of O2 conversion are proportional to surface area (Fig. 53). A weak support effect, as shown in Fig. 59, is another indication of bulk type 11 behavior (327). With an increase in the loading of the heteropoly compound on the support, the rate of bulk type I1 catalysis increases to high loading levels, whereas the rate of surface catalysis shows saturation at relatively low loadings because of the decrease in the dispersion of the heteropoly compound on the support (327). Good correlations are observed between the oxidizing abilities of catalysts and the catalytic activities for oxidation, provided that the bulk and surface catalysis are properly accounted for. Examples are shown in Figs. 60a and b. A linear correlation is observed between the rates of catalytic oxidation of

Surface area / m 2 gl

FIG. 58. Rates of catalytic oxidation of H2 and o f CO in the presence of catalysts having different surface areas. (a) H3PMo12040 at 573 K, (b) Na2HPMo12040 at 623 K. ( 0 )H 2 4 2 ; (A)C G 0 2 (From Ref. 262.)

214


100

0

30 Amount of H,PMo,,O, 10

20

40

50

loaded / wt 8

FIG.59. Catalytic oxidative dehydrogenation of cyclohexene (0, surface catalysis) and oxidation of acetaldehyde (0.bulk-type 11); the catalyst was H~PMOIZON supported on Si02. Masses catalyst: 0.2 g for cyclohexene and 0.1 g for acetaldehyde. (From Ref. 327.)

acetaldehyde (surface reaction) and the rate of reduction of catalysts by CO (indicating surface oxidizing ability). A similarly good relationship for oxidative dehydrogenation of cyclohexene (bulk type I1 reaction) and the rate of reduction of catalysts by H2 (bulk oxidizing ability) has also been found. However, the

~1

tb

0

1

Na2-2,3,4

2

FIG. Correlations between catalytic activity and oxidizing ability for (a) oxidation a acetaldehyde (surface reaction) and (b) oxidative dehydrogenation- of cyclohexene (bulk-type I1 reaction). (From Ref. 327.) r(aldehyde) and r(hexene) show the rates of catalytic oxidation of acetaldehyde and oxidative dehydrogenation of cyclohexene, respectively. (From Ref. 337.) r(C0) is the rate of reduction of catalysts by CO; r(H2) is the rate of reduction of catalysts by H2. M, denotes M,Hj-xPMo1204~.Na2-I, 2, 3, and 4 are Na2HPMoI2O4,,of different lots, of which the surface areas are 2.8, 2.2, 1.7, and 1.2 m2 g-', respectively.


215

correlation is poor if the rate of the surface reaction is plotted against the bulk oxidizing ability. The following are typical reactions that have already been found to be described by either of the two types of catalysis by H3PMoI2O4"and its alkali salts: Surface type: Oxidation of CO, acetaldehyde, and methacrolein. Bulk type II Oxidation of H2, oxidative dehydrogenation of cyclohexene, isobutyric acid.

Since the classification is essentially based on rates of catalytic reactions relative to rates of difhsion of redox carriers, there are oxidation reactions that are intermediate between the two limiting cases. We note that neither the molecular size nor the polarity of reactant molecules is the principal characteristic determining the type of catalysis. Although oxide ions migrate rapidly in the bulk, bulk type 11 catalysis is not observed for oxidation catalyzed by Bi-Mo oxides. In this case the rate-limiting step is a surface reaction. It is noteworthy that in the two industrial processes to produce methacrylic acid, both involving catalysis by H ~ P M O and ~ ~ its O alkali ~ ~ salts, one involves bulk type I1 catalysis and the other, surface type catalysis, as described in the following section.

I

CH,=C-CHO

B.

SurfacewF

OXIDATION OF ALDEHYDES

Methacrylic acid has been used for the synthesis of poly(methy1 methacrylate). It has been synthesized industrially via a reaction of acetone with hydrogen cyanide (12, 17, 330, 331). However, the process produces ammonium bisulfate and uses the toxic hydrogen cyanide. Recently, an alternative, a twostep oxidation of isobutylene, has been developed. The first step is the oxidation of isobutylene to methacrolein, and the second is the oxidation of methacrolein to methacrylic acid: CH3 I CH~=C-CHJ

CH3 I + CHz=C-CHO

+

CH3 I CH2=C-COOH

(36)

1. Mechanism and Roles of Acidity and Oxidizing Ability In the second step of Eq. (36). methacrolein is oxidized by heteropoly catalysts, of which the active component is essentially H3PMo12040 or its

216

TOSHIO OKUHARA, NORITAKA MIZUNO, A N D MAKOTO MISONO

equivalent. Therefore, the oxidation mechanism was investigated with H3PMo12040and its salts. It was concluded that the reaction proceeds by a redox mechanism according to the reaction scheme showin in Eq. (37): RCHO F & RCH(OMo)(OM)

RCOOMo

-

RCOOH (M = H or Mo)

(37)

Equation (37) was derived on basis of the following experimental facts: (i) A fair correlation between the rate of catalytic oxidation and the oxidizing ability of the catalyst measured by the reduction with CO was observed (Fig. 61) (17, 332). This result shows that the rate-limiting step is part of the second reaction, that is, the oxidative dehydrogenation of the intermediate. The first reaction requires acidic sites, as nonacidic catalysts were inactive. But the rate-determining step is inferred not to be part of the first reaction because there was no parallel between the acidity and the rate (332). (ii) Catalyst oxygen is involved in the reaction, since the reaction continued and the selectivity remained essentially the same for a prolonged period after the supply of oxygen was stopped, with the catalytic reaction proceeding in the stationary state (17).

Na3 Cs3 C s l

N a l Cs2.86 Na2 \ / H

. C

0

0

1

2

3

lo6 x Rate of reduction by CO / mol rnin-lg-l FIG.6 I. Correlation between the conversion of methacrolein and the rate of reduction of catalyst by CO. M, denotes MxH3--xPMo12040.(From Ref. 332.)


217

oxid

CO,

acid SCHEME 8

(iii) Rapid and direct exchanges of isotopic oxygen between either two of the species, methacrolein, water, and M,H3- rPMolOOOO(M = Na, Cs; x = 0-3.15), were confirmed by elaborate pulse-mass spectrometry experiments ( I 7, 332). In addition, the rate of oxygen exchange between methacrolein and the polyanion increased with the number of Brsnsted acid sites. This result indicates that the first reaction in Eq. (37) is catalyzed by Brsnsted acids. (iv) The presence of water vapor had a strong influence on the rate and selectivity. This effect was reversible and instantaneous (103). The roles of acidity and oxidizing ability were investigated in detail for the oxidation of acetaldehyde catalyzed by various salts of H3PMoI2O4(); the data were interpreted on the basis of the reaction scheme deduced (Scheme 8 ) (333). The rate of each reaction in the scheme was estimated from the rates of oxidations of acetaldehyde and acetic acid, and compared with the acidities and oxidizing abilities of the catalyst surfaces. The comparisons indicate that the oxidizing ability influences mainly the reactions of CH3CHO -+ CH3COOH and CHJCHO CO,, and the acidity accelerates CH3COOH CH30H + CO and CHJOH + CHjCOOH CH3COOCH3. Ai (334a) investigated the effect of countercation and additives for the oxidation of methacrolein. It was found that the catalytic activity of C S ~ H P M O ~is ~enhanced O ~ ~ by the addition of oxoanions such as BO:- and PO:-. The effect was explained on the basis of acidic and basic properties of the heteropoly compounds. Ai el al. (334h) also investigated the oxidation of crotonaldehyde to furan catalyzed by H3PMo12040 and its salts. The rate increased markedly with an increase in the steam concentration but was almost independent of the partial pressures of oxygen and crotonaldehyde. The reaction catalyzed by the Cs salt was faster than that catalyzed by the parent acid. This result was explained as follows. The addition of a basic species such as Cs ion reduces the acidity of the catalyst and the affinity for furan. which is basic. The weaker interaction facilitates the furan desorption, which is assumed to be ratedetermining. -+

+

+

2. Catalysts In industry, heteropoly catalysts of H3-,Cs,PMo12- ,,V,,04(,(2 < x < 3 ; 0 < v < 2) are used to oxidize methacrolein into methacrylic acid with 60-70%

2 18

TOSHIO OKUHARA. NORITAKA MIZUNO, A N D MAKOTO MISONO

yields (12, 335). Formation of Cs salts markedly increases the surface area and thermal stability of the catalysts, but the stoichiometric Cs salt is not catalytically active, probably because of the absence of acidity. Thus acidic salts that are nearly stoichiometric are preferred. The acidic salts are often mixtures or solid solutions of the acid form and salts. It is assumed in some cases that the acidic Cs or K salts are the acid form epitaxially formed as thin films on the . 337). surface of Cs or K salts ( 4 6 ~336, It has been claimed (335) that preparation of an acid form catalyst by the thermal decomposition of pyridinium salts results in a cubic crystal structure and increases the surface area and pore volume. For example, the surface area of H4PMoIIV040increases from 1.O to 5.3 m2g- by the creation of macropores having radii of 103-104 8. As a result of macropore formation, higher yields are obtained (Fig. 62). The formation of acetic acid, CO and CO2 at high conversion is suppressed by treatment of the catalyst with pyridine. The application of this method to acidic Cs salts further improves the activity and selectivity.

'

c.

DEHYDROGENATION OF ISOBUTYRIC ACID

This reaction is another possible route for the production of methacrylic acid, since isobutyric acid can be obtained by an 0x0 process from propene and CO. Heteropoly compounds and iron phosphates are so far the most efficient catalysts for the reaction. The favorable role of the presence of an a-methyl group is remarkable for oxidative dehydrogenation, as the heteropoly compounds are not good catalysts for the dehydrogenation of propionic acid (338, 339).

t 0 . 0 ' ' 0

' 20

'

' ' 40

"

60

I

'

80

100

Conversion of methacrolein / % and treated (0) FIG.62. Oxidation of methacrolein catalyzed by H4PMoIIV040.untreated (0) with pyridine. Catalyst, 10 cm3; reaction temperature, 553 K; SV, 1000 h- I. Reactant composition: methacrolein 2%. oxygen 6%0, water 20%. (From Ref. 335.)

219


The effects of countercations as well as heteroatoms have been investigated by Akimoto et al. (340). The catalytic activity increased in the order H < Li < Na < Rb < Cs at 573 K, and the reverse order was found at 523 K. The authors concluded that the reaction proceeds by a redox cycle involving heteropoly catalysts and that the reduction of the catalyst is rate-determining at 573 K, whereas the reoxidation is rate-determining at 523 K. It was supposed that oxygen atoms bonded to Mo become more reactive as the electronegativity of the countercation decreases. The authors hrther suggested that a metal ion having a high oxidation potential, such as Pd or Ag, acts as an electron reservoir and accelerates the reaction. It was reported that Cu2+ is also an effective cation (341). Recently, Lee et al. (342) have compared the catalytic performance with that of C S , H ~ - . ~ P M O ~ ~ ~ , .( xV, ,y, = O ~W) ~ and found that Cs2.75H1.2sPMoIIVO40 was effective for the oxidative dehydrogenation of isobutyric acid. For example, the selectivity to methacrylic acid was 78% at 97% conversion and 623 K. Herve et al. (284) demonstrated that the elimination of VS+ from H4PMoI IVO40 took place during the reaction. Notwithstanding the extensive investigations, a consistent explanation is lacking; further investigations are necessary. The heteropoly catalyst is deactivated during a prolonged reaction period by the loss of Mo to form volatile Mo-containing species via the interaction of isobutyric acid andor methacrylic acid with the catalyst (343).The deactivation was suppressed by presaturating the feed by flow over a bed of Moo3 (343). In the dehydrogenation of isobutyric acid, the by-products in addition to CO and C 0 2 are propylene and acetone. Two reaction mechanisms were proposed (340, 341) and the latter is shown in Scheme 9 (340). The formation of methacrylic acid and acetone involves a common intermediate: The El elimination of a proton from I yields the methacrylic acid while a nucleophilic SN1 attack of oxide ion produces COZ and acetone (344). On the other hand,

CH3 \CH-COOH CH3

'

CH3

I

CH2=C -COOH

[

-

CH3 \ cH3, CH COOH]

H'

C3H6'

co

ads

2-heptanol > 2-hexanolB 2-nonanol. A similar sequence was found for the 1-alcohols: 1-octanol > 1-heptanol > 1-hexanol > 1-nonanol. However, the reactivities of primary alcohols are much lower than the reactivities of secondary alcohols. While an increase in reactivity of 2-alcohols with increasing chain length can be expected on the basis of chemical reactivity, the decrease beyond CBmust have another origin, which may be reactant shape selectivity in the TS-1 catalyst. The 2-alcohol generally react faster than the 3-alcohol (Van der Pol et al., 1993b). Glycols undergo oxidation with H202 and titanium silicates, but it is also possible that some of the reactions observed proceed as noncatalytic reactions once the primary oxidation products are formed. Ethylene glycol is oxidized to glycolic acid: CH2-CH2 I I OH OH

+ H202

// -+

CH2-C I OH

0

+ H20

\

(14)

OH

With a 1 : 1 molar ratio of ethylene glycol and H202, the selectivity to glycolic acid was 86% at 44% conversion of the glycol and 99% conversion of the H202. Small amounts of glyoxylic acid (3%) and formic acid were also detected. Propylene glycol is oxidized to hydroxyacetone with high selectivity: CHj-CH-CH2

I

I

+ H202

OH OH

-+

CHj-C-CH2

II

I

+ 2H20

(15)

0 OH

With a molar ratio of propylene glycol to H202 of 2.5, the selectivity to hydroxyacetone at 32% conversion of the glycol was 94%, and the selectivity based on H202 was 85%. Small amounts of acetic acid and formic acid were detected. The initial oxidation proceeds with high selectivity for the secondary alcohol group. Further oxidation affords oxidative cleavage products rather than pyruvic acid, as is observed when the oxidation of hydroxyacetone is carried out with O2 and noble metal catalysts. More complex is the oxidation of 2,3-butandiol (I): the first oxidation product, acetoin (11), can be obtained at 333 K with 96% selectivity at 62% conversion if the solvent is water. When methanol is used as the solvent, the reaction rate is decreased, but the product distribution is the same. When acetone is used as the solvent, a larger amount of butane-2,3-dione (111) is obtained, as summarized in Table IV. C-C-C-C

I

I

OH OH I

C-C-C-C II I 0 OH 11

C-C-C-C

II II 0 0 111

CH3COOH IV

303

MICROPOROUS CRYSTALLINE TITANIUM SILICATES TABLE IV Oxidation of Butane-2.3-diol with H202 Catalyzed by TS-I‘ Selectivity (%)

Diol/HzOz ratio

I I I 4 4 4 4 4

Time

Conversionh

Solvent

(h)

(%)

I1

111

Hz0 Hz0 H2O Hz0 CHiOH CHiOH acetone acetone

3 5 6 4 2.5 4 2 4

40 60 62 16

98 96 96 96 94 91 47 48

2 4 4 4 6 9 53 52

11

16 10 15

~

~

~~~~~

~

“Sheldon and Dakka (1992). ’At a molar ratio of 4 : I , the maximum conversion is 25%.

Acetic acid (IV) is formed when acetoin or butane-2,3-dione is oxidized in acetone, even at low temperatures. The use of acetone significantly changes the relative rates of the different oxidation reactions. The possibility that acetone takes part in promoting the oxidation of acetoin deserves further investigation. b. Oiejins and Diolefins. The oxidation of the lower olefins with H 2 0 2 catalyzed by pure crystalline titanium silicates selectively produces the corresponding epoxides. With some olefins, the reaction can proceed further to give glycols and eventually C=C cleavage products. Olefin oxidation can be carried out with a polar solvent so that a single reactant phase including H 2 0 can be obtained. Methanol or methanol-water mixtures are convenient because, in the presence of olefins, the rate of methanol oxidation is negligible. Ethanol can also be used, but its own oxidation should be evaluated. Ethylene is oxidized to give ethylene oxide, and propylene is oxidized to give propylene oxide (PO): CHI

\

CH =CH2

+ H202

-

CH,

\

CH -CH2 \

/

+ H20

(16)

0

In a subsequent reaction with the solvent, PO produces small amounts of propylene glycol or glycol monoethers. The rate of formation of PO from propylene, H 2 0 2 , and TS-I in methanol at 313 K is shown in Fig. 22. The selectivity to PO is 85%, but it can be improved by silanization of the catalyst, which reduces the hydrolytic activity, or by addition of small amounts of sodium acetate to the reaction mixture. With these modifications, at 97%

304

BRUNO NOTARI

Oo

zo

40

m

M)

100

izo

140

ie

Time I rnin FIG. 22. H 2 0 2 epoxidation of propylene catalyzed by TS-I in methanol-water at 313 K: (1) propylene oxide; (2) H202; (3) propylene glycol and glycol ethers. (From Clerici ef a/., 1991b.)

H202conversion, the selectivity based on PO reaches 97%, and glycol formation is reduced to 3% without impairing the overall catalytic activity. This result is interesting in view of the reported negative effect of sodium acetate on the catalytic activity for n-hexane oxidation (Clerici, 1991a). Clearly, the concentration of the sodium acetate is critical; small amounts reduce the activity of SiOH groups that catalyze hydrolysis of the epoxide initially formed, while larger amounts interfere with the catalytic activity of the titanium centers. When ethanol is used as the solvent, it is oxidized to acetaldehyde at a rate close to that of propylene oxidation, showing that both reagents compete for the catalytic centers. The oxidation of propylene can also be carried out in other solvents, such as methyl acetate, acetonitrile, or tert-butyl alcohol, but the rates are lower than the rate in a methanol-water mixture. The reaction in tert-butyl alcohol is hrther complicated by the fact that, in the presence of TS- 1, the alcohol and H202 react to give TBHP (Maspero et al., 1989). The oxidation of the different butenes gives the corresponding epoxides with high selectivity; the by-products glycol and glycol ethers can be reduced to very low concentrations, as in the case of propylene oxidation. An interesting feature is the retention of stereochemical configuration: cis-2-butene gives exclusively the cis-epoxide, and trans-2-butene gives exclusively the trans-epoxide (Clerici et al., 1989b, 1989~). The sequence of reactivity for the butenes is cis-2-butene (1 6) > 1-butene (6) > isobutylene (4.7) > trans-2-butene (1 .O). A higher reaction rate for the cis isomer was also observed in the epoxidation of 2-hexenes, as was retention of stereochemical configuration (Tatsumi et al., 1990a). In the epoxidation reactions with H202 or organic hydroperoxides catalyzed by group IV-VI metals, the active oxidizing species act as electophiles. The reactivity is determined by the electron density at the double bond, which

305

MICROPOROUS CRYSTALLINE TITANIUM SILICATES

increases with alkyl substitution because of the electron-donating properties of alkyl groups. The sequence of reactivity is isobutylene > cis-2-butene > trans-2-butene > 1-butene. It is reasonable to consider that in titanium silicate-catalyzed reactions the oxidizing species also acts as an electrophile. The different order of reactivity of the C4 olefins in the presence of titanium silicates relative to that observed with soluble catalysts must therefore arise from the fact that alkyl substitution at the double bond is responsible not only for inductive effects, but also for increases in the size and the steric requirements of the molecules. Since the rates of difhsion of the different butenes cannot be the cause of the different reaction rates, a restricted transition-state selectivity must be operating. The presence of fluorides, either HF or NH4F, decreases the rate of 1-butene epoxidation. Fluoride ions coordinate strongly at the titanium centers (the enthalpy of formation of TiF4 is - 1550 kJ/mol) and in this way inhibit the interaction with the reagent molecules. The oxidations of c5-C~ olefins, allyl chloride, allyl alcohol, and allyl methacrylate with H202 give the corresponding epoxides. Data characterizing oxidations camed out in methanol solvent at an olefin concentration of 0.90 M, in the presence of TS-I (6.2 g/L), are given in Table V. As discussed for the epoxidation of propylene, as well as in the case of allyl alcohol epoxidation, the selectivity can be modified by treatments of the catalyst. Addition of Na2C03 drastically reduces the catalytic activity, while with addition of sodium azide a satisfactory level of catalytic activity is maintained with 88% selectivity to the epoxide glycidol. When aluminum-containing TS- 1 is used, no glycidol is formed; only products arising from secondary solvolysis of glycidol are obtained, the solvolysis being attributed to the acidic character of aluminum-containing TS-1 (Hutching et al., 1995). TABLE V Oxidation of CrC8 Olefns and Ally1 Compounds in Methanol"

Olefin I-Pentene I -Hexene Cyclohexane I-Octene Ally1 chloride Allyl alcohol Allyl methacrylate a Tatsumi

Concentration of H202 Temperature Time, I (M) (K) (min) 0.18 0.18 0.18 0.17 0.18 0.14

298 298 298 318 318 318 338

60 70 90 45 30 35 -

H202 Selectivity conversion based on H202

(%)

(%)

(min)

94 88 9 81 98 81 -

91 90

5 8

h

-

91 92 12 77

5 7

et al. ( 199 I ) and Clerici et al. ( 1993b). Not determined.

16

-

3 06

BRUNO NOTARI

Evidence of variables that influence the relative rates of reaction of olefins and alcohols was obtained from experiments with compounds that have both olefinic and alcoholic functions and by the competitive oxidation of mixtures of olefins and alcohols. The data of Table VI show that when the double bond has no substituents, as in ally1 alcohol, but-3-en- 1-01, or 2-methylbut-3-en- 1-01, only the epoxide is formed; but when the double bond has substituents, the epoxidation rate is decreased and ketone and aldehyde products from the oxidation of the OH group are formed. This effect is more pronounced with a greater degree of substitution. Since the double bond and the OH group are part of the same molecule, the difference must arise from the different abilities of the reactants to coordinate and react at the titanium center; restricted transition-state shape selectivity is a possibility. The terminal double bond, sterically less hindered, interacts strongly with titanium, preventing coordination of the competing OH TABLE VI Oxidation of Unsaturated Alocohols in the Presence of TS-I" Product yield [mol (mol Ti)-'] Reactant

Ketonelaldehyde

Epoxide 19 16

30

-OH

"Tatsumi et al., 1993.

31

95

37

4

7

27

43

65

44

141

98

94

18

10

75

17


307

group. Because of steric hindrance, this interaction is weaker in substituted olefins, allowing the OH group to undergo oxidation. The steric interactions between the reagents and the titanium centers embedded in the zeolitic structure are demonstrated by the relative rates of reaction of two closely related molecules: C

\

c

/

C

\

c

/

O

but-2-en- 1-01

H

C

\

c

/

C

\

c

/

OH

I C 3-methylbut-2-en- 1-01

Both undergo oxidation with epoxide and aldehyde formation, but the former gives an epoxidatiodalcohol oxidation ratio of 3, whereas the latter gives an epoxidatiodalcohol oxidation ratio of only 0.2. That this difference is due to the steric requirements inside the pores of the catalyst is demonstrated by the fact that the relative rates of epoxidation and alcohol oxidation of the same molecules on a large-pore Ti02/SiOz catalyst are almost the same (Tatsumi et al., 1993).

The epoxidation of cis- and trans- 1-hydroxy-3,7-dimethylocta-2,6-diene with H202 and TS-1 is chemoselective at the 2-position and stereoselective: no epoxidation takes place at the 6-position, and reactant molecules retain their structure in the products. The interesting results have been described as hydroxy-assisted epoxidation. The role of the OH group in the reaction is confirmed by the fact that in the epoxidation of cyclopent-2-en- 1-01, the product with the 0 cis to the OH group is favored over the one in trans by a factor of 9 : 1. The same steric preference was found in the epoxidation of cyclo-hex2-en-1-01 (Kumar et al., 1995). The competitive oxidation of olefins has also been investigated in the presence of alkanes. As discussed below, alkanes are oxidized by H202 in the

308

BRUNO NOTARI

presence of titanium silicates. When the oxidation of 1-octene was carried out in the presence of n-hexane under conditions that would lead to the oxidation of each if it were used separately, it was observed that 1-octene is preferentially oxidized by a factor greater than 40. (Huybrechts et al., 1992). Cyclohexene is oxidized very slowly in the presence of TS-1; little if any epoxide could be obtained under conditions of rapid oxidation of 1- and 2-alkenes to the corresponding epoxides. This low reactivity has been ascribed to the molecular dimensions of cyclohexene, which cannot enter the channel system of TS-1. Evidence for this suggestion was obtained by elution chromatography; when TS- 1 was loaded in a chromatographic column and a mixture of cyclohexene and 2-hexene injected, the retention time for cyclohexene was much less than that of linear 2-hexenes, despite the higher boiling point of cyclohexene (Tatsumi et al., 1990a). Cyclohexene can be oxidized by H 2 0 2 if Ti-beta is used as the catalyst (Corma et al., 1994). A comparison of the performances obtained with the two catalysts, TS-1 and Ti-beta, for the oxidation of cyclohexene and other olefins is given in Table VII. The results confirm the previously reported low reactivity of cyclohexene when the catalyst is TS-1 and indicate that Ti-beta is active for the oxidation of cyclohexene and other bulky olefins. However, for cyclohexene and the linear olefins, the major reaction products formed in the presence of Ti-beta are glycols and glycol ethers, whereas in the presence of TS-I, epoxides are predominantly formed. Also in this case, the epoxides initially formed in the presence of Ti-beta undergo secondary reactions catalyzed by the acidic centers associated with the aluminum in the material, as previously seen for ally1 alcohol and for the epoxidation of 1 -butene on aluminum-containing TS-1 (Bellussi et al., 199la). A different product composition was observed for cyclododecene, TABLE VII Selective Oxidation of Olefns Catalyzed by TS-I and Ti-Beta” TS- I H202

Olefin 1-Hexene Cyclohexene I -Dodecene Cyclododecene

TON”

conversion

(s-I)

(%)

50

98 83 26

I 110

5

Ti-Beta Selectivity

conversion (%)

EP

GJ

GE‘

(s-I)

96

-

100

23 34

4 -

12 14 87 20

77 66

-

Selectivity

H202

TON

80 80 80 41

E

G

GE

12

8

80

-

-

100

-

100

-

80

20

-

“Coma er al. (1994b). ’TON = turnover number. ‘E = epoxide. d G = glycol. ‘GE = glycol ethers.


309

which in the presence of Ti-beta gives a large amount of epoxide. This result indicates that the reaction of cyclododecene does not proceed inside the pore structure, since in that case the same secondary reactions with glycol formation that are observed for the other alkenes would occur. Thus, for each catalyst the cyclododecene reaction likely takes place on the outer surfaces of the crystals, giving, with some difference in rates, a similar product distribution for each. The fact that the reactivity of cyclododecene is higher than that of cyclohexene in the presence of TS-1 is surprising, because its dimensions are less favorable than those of cyclohexene; reactions on the outer catalyst surface are very likely responsible for these results. In some cases, oxidation of double bonds does not stop at the epoxide, but proceeds further to oxidative cleavage of the double bond. It was reported that the reaction of a-methyl styrene with H202 in the presence of TS-1 or TS-2 produces a-methyl styrene epoxide (15%), a-methyl styrene diol (10 4 0 % ) and acetophenone (40-60%) (Reddy, J. S. et al., 1992). However, results similar to those obtained with titanium silicates were obtained for many other catalysts, such as HZSM-5, H-mordenite, HY, A1203, HGa-silicalite-2, and fumed SO?. These materials have much different properties and differ significantly from titanium silicates; thus, the results cast some doubt on the role of the catalyst in this reaction. Furthermore, the oxidation of styrene is reported to proceed with C=C cleavage and formation of benzaldehyde, in contrast to previous reports of the formation of phenylacetaldehyde with 85% selectivity (Neri et al., 1986). The reinvestigation of this reaction has shown that styrene epoxide is formed as the initial product, but it rapidly undergoes isomerization to phenylacetaldehyde. In the presence of solvent methanol, however, addition of the solvent to the epoxide produces 32.8% of 2-methoxy-2-phenylethan01,decreasing the phenylacetaldehyde yield. In all cases the reaction is accompanied by further oxidation to benzaldehyde (Kumar et al., 1995). The oxidation of butadiene, diallyl carbonate, or diallyl ether gives the products of monoepoxidation with selectivities of 85% and higher when the ratio diene/H202 is 2.5 (Table VIII). The diepoxides are formed in larger amounts when the diene/H202 ratio is 1. TS-1 does not catalyze oxidation reactions with TBHP or other hydroperoxides (Romano et al., 1990; Tatsumi et al., 1991), but other titanium silicates have been found to be active. Interesting results have been obtained with Ti-beta. In the oxidation of I-octene with TBHP catalyzed by Ti-beta, together with the epoxide substantial amounts of glycol, glycol ethers, octanal, and octan-2-one are formed, and even heptanal deriving from the oxidative cleavage of the double bond. However, when the acidity of Ti-beta is neutralized by exchange with alkali or alkaline-earth metal ions, the selectivity to epoxide is increased to 90-100%. Similar high epoxide selectivities have been reported for reactions of olefins and TBHP in the presence of Ti-beta when acetonitrile was

310

BRUNO NOTARI TABLE VIIl Oxidation of Polyunsaturated Compounds in ihe Presence of TS-1" H202 conversion

Yield based on H 2 0 2

Selectivity (%)

("/)

(%)

Monoepoxide

Diepoxide

Alkeneb

Solvent

T (K)

Butadiene Diallyl Carbonate Diallyl ether

t-BuOH CHiOH

293 338

98 95

85 50

85 93

15

CHiOH

338

96

60

90

4

a

5

Romano el a/. (1990b). * DieneM202 ratio = 2.5.

used as the solvent (Corma et af., 1995; Sat0 et al., 1995). Very likely, the weakly basic acetonitrile preferentially coordinates to the acid sites of the catalyst, thus preventing acid-catalyzed ring opening or rearrangements of the epoxide initially formed. Ti-MCM-41 is also an active catalyst with TBHP and olefins; norbomene gives the corresponding epoxides with 90% selectivity in the presence of Ti-MCM-41. However, when H 2 0 2 is the oxidant, conflicting results have been reported for the oxidation of 1-hexene (see Section M E ) . c. Phenols. Phenol is oxidized to give catechol and hydroquinone in almost equal yield and with high selectivity . The reaction is schematically represented as follows:

After the initial discovery of the catalytic activity of TS-1 for this reaction (Esposito et al., 1983,1985), other titanium containing zeolites such as TS-2 have been found to be effective catalysts; indeed, this reaction has become a standard for measurement of the activity and selectivity of different titaniumcontaining zeolites. Suitable solvents are H20-methanol mixtures or H20-acetone mixtures. The best results, 94% selectivity based on phenol and 84% based on H202at 30% conversion of the phenol, can be obtained only when all the reaction conditions are optimized. The temperature must be controlled within few degrees of 363 K, the addition of H202 must be gradual, and the catalyst must be a pure phase (Romano et al., 1990). The results are strongly influenced by the dimensions of the catalyst crystallites, the range 0.2-0.3 pm giving the best performance, as discussed above. In a number of papers it has been reported


311

that benzoquinones were formed as major products in the oxidation of phenol, especially at low conversions of H202, but that they were not present in the final reaction mixture when all the H202 had been consumed (Thangaraj et al., 1990, 1991a, 1991b; Reddy, J. S . et al., 1990, 1991d, 1992a). However, it has recently been shown that an error in the analytical method was responsible for these results (van der Pol et al., 1993a). The oxidation of phenol with H202 catalyzed by TS-1 obtained by secondary synthesis by reacting Tic& with preformed ZSM-5 has led to conflicting results. In one case, the catalytic activity was found to be identical to that of TS-1 (Kraushaar et al., 1988, 1989); in another, low activity and formation of tarry products were reported and ascribed to the residual acidity of the starting material and/or the presence of extra-framework Ti02 (Huybrechts et al., 1991b). The oxidation of phenol has also been investigated with Ti02 deposited on silicalite as the catalyst (Ferrini et al., 1990, Kooyman et al., 1992). The performances of TS-1 and TS-2 catalysts for this reaction were shown to be identical when the titanium contents and crystallite dimensions were equal (Tuel et al., 1993a). Because the oxidation of phenol is sensitive to the purity of the titanium silicate catalyst, it has been used as a test reaction to evaluate the purity of the catalytic materials. A standard material called EURO TS-1 has recently been prepared and evaluated in several laboratories (Martens et al., 1993). The high selectivity by which dihydroxybenzenes can be obtained from phenol and H202 when titanium silicates are used as catalysts has led to the development of a new process and the construction of a plant capable of producing 10,000 tons of catechol and hydroquinone per year. This plant, built by Enichem in Ravenna, Italy, includes a catalyst production and regeneration unit. Having started operations in 1986, this facility is working with excellent technical and economical results. The advantages of the new process relative to previous technologies include low H202 decomposition and high phenol conversions, with the formation of < 10% high-boiling by-products. These characteristics minimize the power requirements and make the process economics very competitive (Notari, 1988). Phenol derivatives have also been oxidized; anisole undergoes substitution both in the ortho (30%) and the para position (70%) (Romano et al., 1990). d. Aromatic Hydrocarbons, Alkanes, and Cycloalkanes. Aromatic hydrocarbons are oxidized by H202with titanium silicate catalysts (Esposito el al., 1983, 1985; Thangaraj et al., 1990). Saturated hydrocarbons can also be oxidized (Clerici et al., 1989, 1991;Huybrechts et al., 1990; Tatsumi et al., 1990b). The reactions take place readily at 323 K, but many investigations have been carried out at 353-373 K. The oxidation of these hydrocarbons is quite unusual, considering their chemical inertness. More surprisingly, methanol can be used as

3 12

BRUNO NOTARI

a solvent without undergoing oxidation itself. This uncommon result is probably the consequence of the hydrophobic nature of titanium silicates and the adsorption of the nonpolar hydrocarbons in preference to the more polar methanol. Benzene is oxidized to phenol with high selectivity at low conversions, but at higher conversions reaction of the phenol gives dihydroxybenzenes, as stated above. When silicates containing titanium and aluminum are exchanged with alkali metals and used under carefully selected conditions (40 wt% catalyst based on benzene, 293 K, 24 h reaction time), benzene can be oxidized to phenol at almost total conversion with a selectivity of 95% (Nemeth et af., 1 993). Little information is available characterizing the titanium-containing phases present in these catalysts, but from the preparation procedure it can be inferred that they likely contain highly dispersed Ti02 in the preformed zeolitic structure (Section II1.G). Toluene is oxidized to cresols (ortho : meta : para ratio = 5 : 1 : 4) and not to the benzyl derivatives, despite the low dissociation energy of the benzylic C- H bond. p-Xylene undergoes oxidation to 2,5-dimethylphenol. Ethyl benzene undergoes oxidation both in the aromatic ring and in the side chain, producing ethyl phenols (40%), acetophenone (56%) and 1 -phenyl ethanol (4%). 4-Methylethylbenzene undergoes oxidation in the aromatic ring, giving 2-methyl-5-ethyl phenol (3 IYo), and in the side chain, giving l-ethanol4-methylbenzene (29%) and 1-ethanone-4-methylbenzene (40%) (Romano et af., 1990b; Khouw et al., 1994). Cumene does not undergo oxidation at a measurable rate. I-Butylbenzene undergoes oxidation mainly in the side chain, with traces of aromatic ring oxidation, producing 1-phenyl-1-butanol, 1-phenyl-3-butanol, and the corresponding ketones (Clerici, 1991). Alkanes are oxidized to alcohols and ketones. Linear alkanes are oxidized to secondary alcohols and ketones, with good selectivity based on hydrocarbons and H 2 0 2 (Table IX). For C5 and higher alkanes, the oxidation at the 2-position is favored over that at the 3- and 4-positions. The efficiency of H 2 0 2 utilization is influenced by the purity of the titanium silicates; pure phases give efficiencies of 80% or more, whereas with materials containing TiO2, almost 30% of the H202 is lost by decomposition to O2 (Huybrechts et al., 1992). The reaction proceeds through the oxidation of the hydrocarbon to the alcohols, which are subsequently oxidized to ketones. This sequence is indicated by the fact that with increasing conversion. the selectivity for ketones increases at the expense of that for alcohols. The absence of isomerization products was confirmed by the formation of only 2-hexanone when 2-hexanol was the reactant; the equivalent statement pertains to the 3-compounds (Parton et al., 199I). Conflicting results have been reported for the products obtained in the oxidation of branched hydrocarbons. For 2-methylpentane, reaction at C2 with

313

MICROPOROUS CRYSTALLINE TITANIUM SILICATES TABLE IX Oxidation of n-Alkanes in 95% Methanol" Selectivity based on H202 Hydrocarbon

(%I

Propane n-Butane n-Pentane n-Hexane n-Heptane n-Octane n-Decane

35 69 82 86 75 63 56

Product distribution (mot%) 213 ratio'

4.5 2.6 1.9 2.6 1.1

2-01 66.2 55.0 34.3 32.1 33.7 30.1 11.5

3-01

16.1 25.9 29.2 20.5 20.5

4-01

%one

3-one

4-one

6.2 12.5 36.2

33.8 45.0 47.4 39.8 28.1 32.8 16.5

2.1 2.0 2.8 3.0 4.5

trace 1.o 10.8

"Clerici (1991a) and Huybrechts er a/. (1991a). *Represents the moles of oxygenated products obtained per 100 moles H202 reacted. 'Ratio between 2- and 3-compounds.

the formation of the tertiary alcohol as the major or exclusive reaction product has been reported (Parton et al., 1991; Clerici, 1991a); reaction at C4 with formation of 2-methylpentan-4-one has also been reported (Tatsumi et ai., 1990). In 3-methylpentane oxidation the preference for the tertiary carbon decreases in favor of oxidation at C2, as in the linear alkanes. As shown by other examples, the C-H bond strength is only one of the factors determining the course of the reaction; steric factors are also important. As was mentioned in Section V.C.3 .b, when competitive oxidation of I -octene and n-hexane is carried out, the alkene is preferentially oxidized. Correspondingly, alkenes react at lower temperatures than alkanes. It is therefore surprising that under noncompetitive reaction conditions, the rate of oxidation of n-hexane is higher than that of 1-octene (Huybrechts et al., 1992). One possible explanation for this observation is that the reaction conditions were different (Clerici et al., 1993b). At 373 K titanium peroxo compounds decompose, thereby giving rise to radical chain reactions that are negligible at lower temperatures. Thus there could be a different mechanism for low- and high-temperature oxidations made more complex by secondary uncatalyzed oxidation of initial products (Spinact et al., 1995). e. NHj and Nitrogen Compounds. Titanium silicates catalyze the oxidation of ammonia by H202; and if the reaction is camed out in the presence of a ketone, the corresponding oxime is formed. When the ketone is cyclohexanone, cyclohexanone oxime is formed (Roffia et al., 1987):

314

BRUNO NOTARI TABLE X Catalysis of Cyclohexanone Ammoximation’

Catalyst None Si02 (amorphous) Silicalite Ti02/SiO2 Ti02/Si02 TS- I

Cyclohexanone Reactant Catalyst Ti H2O~/cyclohexanoneConversion Oxime selectivity Oxime yield content (%) molar ratio (”/) (%) (YOon H202)

0 0 I .5 9.8 I .5

53.7 55.7 59.4 49.3 66.8 99.9

I .07 I .03 1.09 I .04 I .06 1.05

0.6 1.3 0.5 9.3 85.9 98.2

0.3 0.7 0.3 4.4 54.0 93.2

~~~

a

Roffia et al. (1990). Reaction at 353 K; reaction time, 1 h.

The performances of different titanium-containing catalysts are summarized in Table X. Many titanium-containing materials, even Ti02/Si02, are active for this reaction, but the oxime yields based on HzOz are always much lower than with TS-1, which reaches 93.2% at total conversion of cyclohexanone. Two sequences of reactions have been proposed to interpret the experimental facts: 1. The preliminary condensation of the ketone with ammonia, followed by the catalyzed oxidation of the imine to the oxime: \

C=O

+ NH3

,C=NH + H202

\ +

\

\ +

C=NH

+ H2O

C=NOH

+ H2O

(19) (20)

2. The catalytic oxidation of ammonia to NH20H, followed by the noncatalyzed condensation of NHlOH with the ketone: NH3 + H202 \

/C=O

+ NHzOH

--t

NHzOH \

+

+ H20

,,C=NOH

+ H20

(21) (22)

The first sequence was proposed on the basis of spectroscopic evidence and by-product formation (Thangaraj et al., 1991; Reddy, J. S. et al., 1991; Tvaruzkova et al., 1991, 1992). The fact that TS-1 is an efficient catalyst for the selective oxidation of ammonia to NH20H in the absence of ketones strongly favors the second pathway (Mantegazza et a)., 1991). Further evidence was

315

MICROPOROUS CRYSTALLINE TITANIUM SlLICATES

obtained from the observation that ketones such as cyclododecanone and tertbutyl cyclohexanone, which are too large to be adsorbed in the pore structure of TS- 1 , give high yields of the corresponding oximes, consistent with the hypothesis that the catalytic oxidation of ammonia to NH2OH takes place inside the pores, whereas the noncatalyzed condensation of N H 2 0 H and the ketone takes place in solution (Zecchina et al., 1992). When a sample of TS- 1 interacts with a stoichiometric amount of H 2 0 2 , a strong band appears at 25,000 cm-' in the UV-visible spectrum; upon addition of a stoichiometric amount of ammonia, the band shifts to 27,500 cm-' and slowly declines. This sequence has been interpreted as evidence of formation of a mixed complex in which both NH3 and the peroxo group are coordinated to the Ti'". This mixed complex could be the precursor of N H 2 0 H (Geobaldo et al., 1992). Primary amines are oxidized to the corresponding oximes. The sequence of reactions closely parallels the sequence observed with other mono oxygen donors, i.e., oxidation to alkyl hydroxylamines (V) followed by oxidation to alkylnitroso compounds (VI) which, via a prototropic shift, rearrange to the oximes (VII): \

,CH-NH~

)CH-NHOH V

)CH-NO VI

-

\

(23)

,C=NOH VII

The major difference with respect to other oxygen donors is the high selectivity to the oxime, for many (but not all) of the amines, and particularly the limited formation of other oxidation products such as nitro compounds, imines, and alkylnitroso dimers, which easily form in the presence of other oxygen donors by reaction of alkylnitroso compound (VI). In the absence of a catalyst, lower conversions and selectivity are observed (Table XI). TABLE XI Oxidation of Primav Amines Cata1,vzed by TS- I"

Amine

Solvent

CH3NH2 CH3NH2 n-C3H7NH2 i-C3H7NH2 i-C3H7NH2 i-C3H7NH2 CdiiNHz C6HiiNHz CGHsCHzNH2

CH3OH CH3OH CHIOH CHxOH I-BUOH~ I-BuOH CHIOH I-Bu-OH CH30H

Conversion

Oxime selectivity

Catalyst

(%)

(%)

(YO)

TS-I none TS-I TS-I TS-I TS-2 TS-I TS-I TS-I

40 3 32 38 29 31 3 3 20

88 0 73 I7 14 84 33 32 82

90 0 86 88 85 90 8 8 55

"Reddy, J. S. ef al. (1993b). *f-BuOH = tert-butyl alcohol.

HzOz efficiency

316

BRUNO NOTARl

The constrained environment in which electrophilic oxidation occurs has a strong influence on the exclusion of bulky product formation, but it is also important that the oxidation potential of titanium peroxo compounds is adequate, as demonstrated by the fact that the same reactions carried out in the presence of V-silicalite, which has the same pore structure, give rise to imine and nitro compounds as major reaction products (Reddy, J. S. et al., 1994). The oxidation of aniline with H202 gives rise to a number of products, some of which are accounted for by the reactivity in solution of the primary products. A possible reaction scheme is the sequential oxidation of aniline (VIII) to phenyl hydroxylamine (IX) to nitrosobenzene (X). The condensation of unreacted VIII with X results in the formation of azobenzene (XI), while reaction of IX with X produces azoxybenzene (XIII). Only when excess H202 is used does nitrobenzene (XII) form. Many titanium silicates are active for this reaction, Ti-beta and Ti-HMS being the most active. The results demonstrate that in catalysis by TS-I and by medium-pore zeolites, the reaction is limited by diffusion. TS-48, which has been found to be inactive in other oxidation reactions, is an active catalyst for the oxidation of aniline (Gontier et al., 1994; Sonawane et al., 1994). Tertiary amines are oxidized to the corresponding nitrogen oxides. Tosyl hydrazones of ketones and aldehydes and imines are oxidized to the corresponding carbonyl compounds. Reactions have been carried out with small molecules and also with molecules that would not diffuse into the pore structure of the titanium silicates. As in the case of C=C bond cleavage, it is possible that these reactions take place on the outer surface of the catalyst crystals. Toluene p-sulfonyl hydrazones (tosyl hydrazones) undergo C =N oxidative bond cleavage with H202 in the presence of TS-1, to give high yields of the corresponding carbonyl compounds; very likely, the reaction proceeds through

N

m SCHEME 2


317

the intermediate formation of oxaziridines followed by hrther oxidation to carbonyl compounds (Kumar et al., 1993). f. Su[fides. Sulfides undergo oxidation with silicates to form sulfoxides and sulfones: R \ /

s

-

R \

H202

R'

/

s=o

-

H202

R

R'

O \

H202

catalyzed by titanium

/

s4

(24)

\

R'

0

Typical results for different sulfides in acetone at reflux temperature are given in Table XII. Under similar conditions, diphenyl sulfide, C~HS-S-C~HS, was found to be unreactive, indicating that the reactions listed in Table XI1 take place inside the pore structure, which is not accessible to bulky molecules such as diphenyl sulfide. The selectivity to sulfoxides (Table XII) is the result of a competition between sulfides and sulfoxides for the catalytic site: dimethyl sulfide competes effectively, and the selectivity to sulfoxide is 97% with only 3% sulfone produced. The other reactant molecules give larger amounts of sulfones. However the reaction of dimethyl sulfide was camed out at 298 K, whereas the other sulfides reacted at the reflux temperature of acetone; the temperature difference may explain part of the differences shown in Table XII.

VI.

Catalytic Sites

The catalytic activity of titanium silicates must be ascribed Ti" sites, because pure crystalline silicas are totally inactive. As was discussed in Section 111, Ti" is present in the crystalline structure at random. Very likely, the random distribution that is obtained in the precursor reagents is maintained in the solid. Being dilute, each Ti" is expected to be surrounded by OSi" groups and isolated from other Ti" ions by long 0-Si-0-Si-0 sequences. It has been TABLE XI1 Oxidation Oj'SuiUlfideswilh H202 Catalyzed by

TS-2"

Selectivity (%) Conversion Reactant

(%)

Sulfoxide

'Reddy, R. S. et al. (1992). Reaction at 298 K

Sulfone

318

BRUNO NOTARI

proposed that Ti", in this state of dispersion and tetrahedral coordination, has properties different from those of other materials having Ti" sites with octahedral coordination and are not isolated from each other. It has been proposed that isolated Ti" sites have a low activity for H202decomposition (Notari, 1988). However, isolation alone is not sufficient to explain all the observed properties. The hydrophobic environment prevailing inside the catalyst pores and the multiple Ti-0- Si bonds that allow interaction with reactants but prevent complete hydrolysis of the Ti" sites in titanium silicates must also play a role in stabilizing the dispersed Ti" sites.

Vii.

Reaction Mechanism

Investigation of mechanisms of reactions catalyzed by titanium silicates has been limited to oxidation reactions with H 2 0 2 as the oxidant, as described below. As was previously discussed, elements different from titanium and silicon in the catalyst materials change their properties. Catalytic activity of doubly substituted materials such as Ti-beta, H[Al,Ti]-MFI and -MEL, and H[Fe,Ti]-MFI and -MEL is considered separately because the acidic properties associated with the added element affect the composition of the reaction products. Mechanistic information is difficult to obtain when the catalytically active titanium centers are present in a dilute matrix of silica. Only few techniques can be applied, and the available information does not allow discrimination between possible mechanisms. Consequently, it is necessary in this discussion to rely on analogies with the known chemistry of titanium compounds. Once the crystalline structure and the distribution of titanium therein are established, one of the first issues that must be considered concerns the interactions of the isolated titanium sites with water, H202, and other reactants. Adsorption of a number of compounds on the titanium sites leads to an increase in the coordination of titanium which, as previously noted, is reversible. By reaction with water, Ti-0-Si bonds are hydrolyzed, forming TiOH and SiOH groups. The hydrolysis process cannot be extensive, because Ti'" species separated from the matrix would undergo rapid sintering, as is observed for Ti'" dispersed on amorphous silicas. Therefore, in titanium silicates the Ti" must maintain a number of bonds to the crystalline lattice (Deo et al., 1993). To simplify the graphic presentation, hydrolyzed groups are indicated as TiOH and Ti(OH)2. This latter group can be considered equivalent to a titanyl group, Ti=O. Spectroscopic properties of the Ti(OH)2and Ti=O may be different, but their chemical properties are substantially identical; therefore, their notation will be used interchangeably. The scheme shown in Eq. (25) indicates the hydrolysis of the Ti-0-Si bond and further interactions with water.


3 19

OF CRYSTALLINE TITANIUM SILICATES A. ACIDITY

In Section 11, the acidity of TiOz-SiOz was discussed and it was shown that no significant acidity is present in the Si02-rich region. This is very likely due to the fact that, in the presence of excess Si02, Ti" assumes a tetrahedral coordination with the consequence that no charge imbalance exists and no Brcansted acidity is created. The same is true for crystalline titanium silicates that contain small amounts of Ti" and, as demonstrated in Section 111, assume tetrahedral coordination. The chemical behavior is as expected; many acidsensitive compounds like the epoxides can be obtained in high yield without undergoing major hydrolysis or solvolysis. The fact that the limited solvolysis observed for propylene and ally1 alcohol can be reduced by silanization or by the addition of controlled amounts of bases is consistent with the hypothesis that the solvolysis is due to silanol groups and disappears when they are transformed or neutralized. Further evidence comes from the observation that if the same epoxidation is carried out with titanium silicates containing traces of trivalent elements, hydrolysis takes place (Section V). Furthermore, TS-1 is inactive for acid-catalyzed hydrocarbon reactions; but when A13 is present, TS- 1 becomes active for xylene isomerization (Reddy, J. S. er al., 1994). As for Ti02-Si02. the adsorption of water, ammonia, pyridine, and pyrrole indicates that Ti" in titanium silicates behaves as a Lewis acid; and the fact that these adsorbates are removed by simple evacuation leads to the conclusion that Ti" is a fairly weak Lewis acid (Bittar et al., 1992; Liu et al., 1994). The change in Ti" coordination that is observed as a result of adsorption has no effect as far as the development +

320

BRUNO NOTAM

of protonic acidity is concerned, since the molecules that are coordinated are neutral. To produce the charge imbalance in the solid that is at the origin of acidity, charged 0 2 -groups are necessary. Acidity in crystalline titanium silicates has been observed only when a titanium-containing zeolite interacts with H 2 0 2 , but this is due to the formation of peroxo compounds, as discussed below. B. TITANIUM PEROXO COMPLEX The addition of H202 to titanium silicates brings about the formation of a titanium peroxo complex, which, as indicated in Section I, has been proposed as the active species for the transfer of oxygen from the oxidant to the reactant. From the chemistry of Ti" compounds in the liquid phase, it is known that H202 acts as a bidentate ligand and displaces other ligands to form the very stable sidebonded (or q2-02)peroxo species XIV. Weakly bonded neutral molecules of the solvent S are also coordinated at the Ti" center. The alkyl hydroperoxides form peroxo complexes displacing only one ligand to give the end-on hydroperoxo complex XV. In some cases, the alkyl hydroperoxo complex may adopt the sidebonded structure XVI, displacing a weakly bonded neutral molecule of solvent:

The stability of the peroxo complexes formed with H202, measureL -y the association constant, is much higher than the stability of hydroperoxo complexes. The strong repulsion between formally unshared electrons in planar H202 can be reduced by transition-metal ions such as Ti", as they accept electron density from the filled antibonding orbitals of H202 interacting with the empty metal d orbitals of appropriate symmetry. It is for this reason that even hydroperoxo complexes may prefer the side-on configuration that can provide the added stability (Conte et al., 1992). The formation of a peroxo complex between H202 and a titanium silicates has been demonstrated in several ways, the most convincing being the appearance of an absorption band in the UV-visible spectra at 26,000 cm- when H202 is added to a titanium silicate. A band at the same frequency is present in the UV-visible spectra of the peroxo complex [TiFs(02)l3-, and the absorption has been attributed to a charge-transfer process 0:- -+ Ti4+ (Geobaldo et al., 1992). The stability of these complexes is limited to a temperature of 333 K they decompose rapidly at 373 K (Huybrechts et al., 1991). The thermal stability of the peroxo complex formed on TS-1 is markedly increased in the presence of

'


32 1

bases, the decomposition temperature being shifted to 523-673 K. Stable species are formed which have been characterized by both physical chemical methods and catalytic activity tests. Under carefully controlled conditions the ratio "active O"/Ti approaches unity, indicating that every Ti" site in the solid has been transformed to a peroxo complex. The increased stability of the titanium peroxo complex in the presence of bases could be the reason for the observed deactivation of these catalysts caused by alkalies (Clerici et al., 1993). In the [TiFS(O2)]'- ion, the peroxo group is bonded to Ti" side-on, and therefore this could also be the structure of the complex formed on titanium silicates. However, the possibility of a hydroperoxo species bonded end-on cannot be ruled out, because the side-on structure requires a deeper degree of hydrolysis to give the Ti(OHh group, whereas the hydroperoxo can form on a TiOH group, which is more easily obtainable in a material resistant to hydrolysis. The two forms can be represented as follows:

The titanium peroxo complex under neutral conditions oxidizes alkenes, giving the epoxides with no evidence of acid-catalyzed reactions. But when the oxidation reaction is complete and the epoxides are exposed to an excess of H 2 0 2 , a hydrolytic reaction of the epoxides is observed. The rate of this reaction is similar to the rate obtained when the same epoxides are exposed to 0.1 M formic acid. By contrast, silicalite-1 is inactive and unaffected by the presence of H202. The addition of NaOH reduces the hydrolytic activity of the system TS- 1 H 2 0 2 , indicating that an acidic species is responsible for the hydrolysis. The formation of protonic acidity by H202 has been ascribed to the interaction of the titanium peroxo complex with a donor hydroxyl moiety of a molecule such as H 2 0 coordinated on Ti", resulting in the formation of a cyclic structure. The stabilization provided by the cyclic structure would make the dissociation and the protonic acidity possible. When alcohols, and particularly methanol, are used as solvents, the coordinated OH group could be that of the alcohol (Bellussi et al., 1992):

+

The fact that the acidic properties and the hydrolysis reaction are absent during epoxidation must mean that other donor molecules like alkenes are more strongly coordinated at Ti" and prevent the formation of the complex responsible for generating protonic acidity. The formation of an acidic species could therefore also be explained as the result of the transformation of a

322

BRUNO NOTARI

hydroperoxo group into a side-bonded peroxo group, which becomes possible when more strongly coordinating molecules are absent:

This process would be stabilized by alkalies, and the overall process could be represented as follows [Eq. (28)]:

In the presence of excess alkalies these reactions could be inhibited by the formation of Si-0-Na groups (Khouw et al., 1995). C. MECHANISTIC PROPOSALS Many characteristics related to the particular structure of the material contribute to the final outcome of H202 oxidation reactions catalyzed by titanium silicates: The resistance of the titanium center to extensive hydrolysis. The thermal decomposition of the titanium peroxo compounds, with the consequence that different mechanisms can be operative at low and high temperatures. The selective adsorption of reactants, owing to the hydrophobic nature of titanium silicates. Reactant shape selectivity effects related to the dimensions of reactant molecules and catalyst pores, including restricted transition-state shapeselectivity effects as well as chemical and stereochemical selectivity. These effects were illustrated above. The proposals advanced to give a representation of the mechanism by which the 0 is transferred from the titanium peroxo complex to the reactant molecules can be classified as concerted mechanisms and radical mechanisms. Concerted mechanisms have been proposed on the basis of work carried out with soluble MeV', Wv’ and Ti” peroxo compounds. The experimental evidence is consistent with the hypothesis that these compounds act as oxidants in stoichiometric epoxidations and that the reactions involve electrophilic attack of the peroxo compound on the organic molecule or, what is equivalent, a nucleophilic attack of the organic molecule on the peroxidic oxygen, in a “butterfly” transition state. The reaction product is formed and, after desorption, the peroxo compound is regenerated by reaction of Ti” with H202;this accounts for the catalytic nature of the reaction (Amato et al., 1986). The same type of mechanism


323

has been proposed for the surface titanium peroxo complex formed on titanium silicates (Notari, 1988). In the case of alkene oxidation, the mechanism of oxygen transfer from the titanium peroxo complex to the alkene would be as follows:

A modification of the mechanism that involves the hydroperoxo titanium complex and one solvent molecule has been proposed that involves the formation of a stable cyclic titanium peroxo complex (Clerici et al., 1993). In this case, the two peroxo oxygens are not equivalent, and thus two intermediates would be possible:

+

\

/

,c=c,

The metallacycle mechanism can also be considered a concerted mechanism. It is analogous to the one proposed for metal peroxo complexes and is based on the assumed formation of a cyclic intermediate that includes the peroxo group, the reactant molecule, and the metal ion (Mimoun, 1982, 1987; Huybrechts et al., 1992). For alkene epoxidation, the sequence of events would be represented as follows [Eq. (31)]:

The radical mechanism has been proposed to explain the oxidation of saturated hydrocarbons. In the previous mechanisms, the electron density of the double bond or the aromatic ring is considered essential for the attack on the peroxidic oxygen. This condition is absent in saturated hydrocarbons, and considering their inertness, their oxidation probably requires a homolytic mechanism, proceeding through radical intermediates. By analogy with vanadium

324

BRUNO NOTARI

peroxo compounds, it has been proposed that the titanium peroxo complex gives rise to a radical by breaking of one Ti-0 bond. The radical thus formed abstracts a hydrogen atom from the saturated hydrocarbon, producing a radical that recombines with the OH group bonded to the titanium and forms an alcohol, regenerating the catalytic center (Mimoun, 1987; Huybrechts et al., 1990): /

-

'$ o. /

'0'

The radical mechanism has also been proposed as a general mechanism for oxidation of alkenes and aromatics, but several objections have been raised because of the absence of products typically associated with radical reactions. In classical radical reactions, alkenes should react also at the allylic position and give rise to allyl-substituted products, not exclusively epoxides; methyl-substituted aromatics should react at the benzylic position. The products expected from such reactions are absent. Another argument was made against the radical mechanism based on the stereoselectivity of epoxidation. Radical intermediates are free to rotate around the C-C bond, with the consequence that both cis- and trans-epoxides are formed from a single alkene isomer, contrary to the evidence obtained with titanium silicates (Clerici et al., 1993). Recent evidence seems to indicate, however, that radical reactions inside a zeolite do not necessarily follow the same pathways as classical radical reactions. The oxidation of the hydrocarbon ethylcyclopropane with H2O2catalyzed by TS-1, which should proceed through radical intermediates, gives only the alcohol (27%) and the ketone (73%) resulting from hydrogen abstraction at the carbon linked to the cyclopropane ring, and not the products that would be expected from the rearrangement of the radical intermediate:

Furthermore, isopropylcyclopropane gives the tertiary alcohol (50%), and unidentified products (50%), but not products with the rearranged structure. These results indicate that radicals formed on titanium and inside a zeolite can


325

be very short-lived, which prevents the rearrangements expected from classical radicals; alternatively, their movements may be restricted so that no rearrangement can occur (Khouw et d., 1994). From the evidence described above concerning the uncommon behavior of radical reactions in hydrocarbon oxidation at a titanium center inside a zeolite, the possibility that peroxo radicals are also involved in oxidation reactions of other compounds should be considered. In contrast to the mechanism discussed above for the hydrocarbons, in this case both the peroxo and the reactant molecule would be coordinated at the Ti" center, and the reaction would therefore take place between two coordinated species. The initial coordination of reactants has indeed been proposed to explain the selective oxidation of alkenes in the presence of saturated hydrocarbons. It was argued that, owing to the hydrophobic nature of titanium silicates, the concentration of both hydrocarbons inside the catalyst pores is relatively high and hence the alkenes must coordinate to Ti'". Consequently, the titanium peroxo complex will be formed almost exclusively on Ti'" centers that already have an alkene in their coordination sphere, and will therefore oxidize this alkene rather than an alkane which may be present in the catalyst (Huybrechts et al., 1992). Objections to this proposal are based on the fact that the intrinsically higher reactivity of alkenes with respect to saturated hydrocarbons is sufficient to account for the selectivity observed (Clerici et al., 1992). But coordination around the titanium center of an alcohol molecule, particularly methanol, is nevertheless proposed to explain the formation of acidic species, as was previously discussed. In summary, coordination around TitV could play a more important role than it does in solution chemistry as a consequence of the hydrophobicity of the environment where the reactions take place. For the oxidation of alcohols, two equally valid kinds of intermediates have been considered for the abstraction of the hydrogen atom that leads to the hydroperoxo and radical formation of aldehydes or ketones-namely, intermediates: H

H

In both cases, the preliminary coordination of the reactants is considered to explain the differences in rates of oxidation of 2- and 3-alcohols. The strainof-coordination bond angles of the alcohol in the transition state are affected by the length and size of the remaining groups and hence by the position of the OH group on the chain. It seems more difficult to account for this effect with an external nucleophilic attack from a noncoordinated molecule (Maspero et al., 1994).

326

BRUNO NOTAN

A radical mechanism involving species coordinated at the same Ti" center is schematically represented as follows for alkene epoxidation:

According to this hypothesis, the results are modified from what would be expected from classical radical reactions. The interest in this hypothesis is that, with the sole exception of saturated hydrocarbons, it could apply to all the compounds that can be coordinated at the Ti" center, such as alkenes, aromatics, alcohols, and sulfides. According to this hypothesis, the weak Lewis acidity of Ti" would help to bring the reactant into its coordination sphere. The initial coordination of the reactant would explain the oxidation of methylsubstituted aromatics in the aromatic ring and not in the side chain, even with a radical-type mechanism. A crucial point in this mechanism is the formation of the peroxo radical, which requires the reduction of the Ti4+ to Ti3+-a process that is easier with highly dispersed Ti" and should therefore be favored in titanium silicates in which the dispersion of Ti" is the highest ever observed. Radical reactions have been proposed to explain the mechanisms of many oxidations catalyzed by metals (Mimoun, 1987), but until now they were considered incompatible with the experimental evidence available for titanium silicate-catalyzed reactions. The recent results indicate that the rationalization of the observed facts with radical mechanisms is as plausible as that with other mechanisms. Indeed, for the oxidation of saturated hydrocarbons, the radical mechanisms are judged to account for the observations better than other mechanisms. Further investigations are needed to clarify these issues.

VIII.

Summary

The discovery of the new titanium silicates and of their catalytic properties in

H202 oxidation reactions has had a major impact in catalytic science and its industrial applications. One 10,000todyear plant for the production of catechol and hydroquinone has been operating since 1986 with excellent results. Moreover, successfir1 tests conducted on a 12,000-todyear pilot plant for cyclohexanone ammoximation (Notari, 1993b) could be followed soon by an industrial-size plant that would greatly simplify the synthesis of caprolactam. Both these examples are clear indications of the potentials of the new oxidation chemistry made possible by the new materials.


327

Because of stringent environmental constraints, it is difficult for many innovative chemistries to find industrial application. The new chemistry described here is applied successfully because it offers both ecological and economic advantages. Availability and cost of raw materials (especially H202), integration with other productions, value of by-products, price trend expected for products in the future, investments required, and risks involved in the new technology, all play a vital role in the evaluation. There seems to be no doubt that oxidations with H 2 0 2 or hydroperoxides could be advantageous in many other industrial processes of both basic and fine chemicals. The impact of this subject on the science of solid-state chemistry and catalysis has also been significant. Structures are now synthesized that were considered impossible just a few years ago. Oxidation reactions with zeolites were almost unknown before the discovery of titanium silicates, and the demonstration of the properties of these catalysts has stimulated the search for new materials with analogous catalytic properties. The development of this subject has been more successful than almost anyone would have predicted. With wide-pore titaniumcontaining crystalline silicas, with the possibility of eliminating the undesirable reactions due to the acidic properties of aluminum-containing titanium silicates, with the new Ti02-Si02 aerogel materials, with new materials incorporating titanium in microporous crystalline structures such as ALPO, and with materials incorporating other elements such as vanadium in a crystalline silica structure, it has been possible to carry out oxidations of even large molecules, using H 2 0 2 andor hydroperoxides, thus overcoming the most severe limitations of the smaller-pored TS-1 and TS-2 and expanding the range of oxidation reactions. Further developments are to be expected in this subject, driven by the combination of scientific and technological opportunities. REFERENCES Amato, G., Arcoria, A., Ballistreri, F. P.. Tomaselli, G.A., Bortolini, O., Conte, V., Di Furia, F., Modena. G., and Valle, G., J. Mol. Cafal.37, 165 (1986). Anpo. M., Nakaya, H., Kodarna, S., and Kubokawa, Y., J. Phys. Chem. 90, 1633 (1986). Armor, J. N., U.S. Pat. 4,163,756 (1979). Armor, J. N., J. Amer. Chem. SOC. 102, 1453 (1980). Armor, J. N., and Zambri, P. M., J. Catul. 73, 57 (1982) and references therein. Barrer, R. M., in Olson, D. and Bisio, A., Proc.6th Int. Conf. Zeolites, Reno, 1983, Butterworths Ltd., U.K. 870 (1984). Behrens, P., Assmann, S., Felsche, J., Vetter, S., Schultz-Ekloff, G., and Jaeger, N. I., Proc. 6th Int. Conf. X-ray Absorption Fine Structure, York, U.K., p. 552. Elsevier, Amsterdam, 1990. Behrens, P., Felshe, J., Vetter, S., Schultz-Ekloff, G., Jaeger, N. I., and Niemann, W., J. Chem. Soc.. Chem. Commun. 678 (1991a). Behrens, P., Felshe, J., and Niemann, Cafal. Today 8,479 (1991b). Bellussi, G., Clerici, M. G., Buonomo, F., Romano, U., Esposito, A,, and Notari, B., U.S. Pat. 4,701,428 (1987).

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ADVANCES IN CATALYSIS, VOLUME 41

Structural and Mechanistic Aspects of the Dehydration of Isomeric Butyl Alcohols over Porous Aluminosilicate Acid Catalysts KIRILL ILYCH ZAMARAEV Boreskov Instiiuie of Caialysis Russian Academy of Sciences Novosibirsk 630090. Russia

AND

JOHN MEURIG THOMAS Davy Faraday Research Laboraioiy Royal Insiiiuiion of Greai Britain 2I Albemarle Sireei, London WIX 4BS. United Kingdom

1.

Introduction

Seldom in the study of heterogeneous catalysis does it prove possible to ( 1 ) specify precisely the concentration and nature of the active sites, (2) test whether these sites are of comparable strength and are distributed in a spatially and chemically well-defined manner, and (3) explore the structural and mechanistic features of the system using a wide range of complementary techniques, many of them in situ. Even rarer are situations in which both the access to the active sites and the shape of the reactants may be systematically and subtly varied, so that one is able to compare the performance of the active site in a crystalline environment with an essentially identical one embedded in an amorphous solid. Over the past five years, we and our colleagues have undertaken an extensive study of the acid-catalyzed dehydration of the four isomeric butyl alcohols. In so doing, we compared the performance of crystalline, molecular-sieve acid catalysts (HZSM-5)in a range of crystal sizes (so as to vary diffusion path and active-site concentration) with that of amorphous aluminosilicate (AAS) gels in which the pore size is significantly larger. Our results, which permit the 335 Copynght 0 1996 by Academic Press, Inc. All nghu of reproduction in any form reserved.

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identification of reaction pathways and key intermediates, reveal the interesting catalytic consequences of pore confinement. They also point the way to a more rational design of new microcrystalline solid acid catalysts of the kind increasingly required in the petrochemical industry (1). Traditionally,the same overall mechanisms of acid catalysis invoking carbenium ions have been assumed to prevail both in heterogeneous (2) and in liquid homogeneous (3) systems. But these mechanisms do not adequately take into account the fact that adsorbed, rather than free, carbenium ions are formed in the pores of solid catalysts. Consequently, a quantum-chemical model that demonstrates how the interaction of carbenium ions with the sites of their adsorption can influence the reaction mechanism has been formulated by Kazansky ( 4 ) , taking double-bond-shift reactions in olefins as a particular example. According to this view, adsorbed carbenium ions are best regarded as transition states rather than reaction intermediates, a notion that had also been proposed earlier by Zhidomirov and one of us (5). Although theoretical and computational advances now afford powerful insights into the mechanisms of heterogeneous catalysis, especially on acidic, zeolitic solids (6a-d), experimental studies (7, 8) still hold sway. This we hope to demonstrate here by reference to the wide range of techniques-spectroscopic, kinetic, and analytical-that we have brought to bear in our studies of the catalytic dehydration of butyl alcohols. Dehydration of butyl alcohols over HZSM-5 and AAS acid catalysts provides unique opportunities to elucidate experimentally how confinement of the reagents, intermediates, and products inside the pores of the catalyst influences the reaction pathways. Indeed, in both HZSM-5 and A A S the dehydration reaction proceeds over the same active site, viz.

\

/ \ /

/

\

-Al \

H I 0

Si-

/

(henceforth written as /Al-O(H)-Si\) with carbenium ions functioning as reaction intermediates (2, 9). These catalysts have profoundly different pore structures: The diameter of the channels in HZSM-5 is very close (ca. 5.5 A) to that of the four distinct reactants, to that of the supposed key intermediate/ transition states (i.e., carbenium ions), and to that of the reaction products (butenes and ethers). The pore diameter of our AAS, on the other hand, is much larger and estimated to be 50 t 5 A. Moreover, the critical molecular diameters of the reactants range from a value that is smaller than 5.5 A for n-butyl alcohol to one that is significantly larger in the case of ferf-butyl alcohol.

DEHYDRATION OF ISOMERIC BUTYL ALCOHOLS

337

TABLE I Characterization of Aluminosilicate Samples Sample

Crystallite size (pm)

SVAI

N x IO-*’ Sites/g”

42 35

HZSM-5 1

3and5 4

NaHZSM-5 2 AAS

CH-*CH,-OH

The spectrum of Fig. 7A refers to the adsorbed isobutyl alcohol, with the line at 74.9ppm being characteristic of the carbon atom that is bound to the oxygen

I vv

>

PPM

""

FIG. 7. Variation of the I3C CP/MAS NMR spectrum of CH3 CH--*CH*-OH adsorbed in CH3 HZSM-5 upon successive heating of the sample for a certain time at various temperatures: (A) AAer adsorption at 298 K; (B) 70 min at 343 K; (C) 40 min at 398 K; (D) 60 min at 413 K (inserted is the contour plot of the 2D J-resolved "C MAS NMR spectrum in the vicinity of the signal at 30.5 ppm); (E) 60 rnin at 448 K, (F) 60 rnin at 448 K.


35 1

atom. Upon heating the sample for 30 min at 373 K, where we know from our previous kinetic studies that the alcohol molecules are dehydrated, we observe the line at 73.0 ppm (Fig. 7B) from the isobutyl silyl ether (IBSE) intermediate that is formed in reaction + I1 of Scheme 1: CH3\

/CH3 CH

I

*CH2

I

\

-A1 /

0

/ \ /

Si-

\

Further heating results in the selective transfer of the 13C label from the CH2 group of IBSE into the CH groups of two nonequivalent isobutyl fragments (lines at 30.5 and 31.5 ppm) and to a lesser extent into the CH3 group (the line at 19.2 ppm) of IBSE. The lines at 30.5 and 3 1.5 ppm were attributed to the CH group according to their 2D J-resolved "C MAS NMR spectra ( 8 i ) (insert, Fig. 7D). Molecular rearrangements that explain ( 8 i ) how the I3C label can be transferred to the CH3 and CH groups of IBSE are shown in Scheme 3. Unambiguously, they must involve the isobutyl carbenium ion (IBCI) as the reaction intermediate or transition state. This ion is shown in the upper left comer of Scheme 3. Note the selectivity in the transformation of the su posed IBCI inside HZSM-5. As suggested by the preferential transfer of the IpC label to the CH rather than to the CH3 group, adsorbed IBCI prefers to transform via pathway I1 rather than pathway I of Scheme 3 ; i.e., pore confinement in HZSM-5 kinetically favors the formation of linear, rather than branched, C4 carbenium ions. This conclusion is hrther supported (8e) by our solid-state *H NMR studies, according to which the transformation of isobutyl cation ion into tert-butyl cation is not favored inside HZSM-5, even though such rearrangement is favored both thermodynamically and kinetically in solutions. The skeleton of a carbenium ion bears a positive electric charge, and the surface atomic layer of the zeolitic channels is composed of negatively charged oxygen atoms. Ostensibly, a flat linear carbenium ion adheres more strongly to the channel walls in HZSM-5 than do the non-flat skeletons of the branched carbenium ions. This can be the driving force for the energetic preference for the formation of a linear ion. OR

Thus, for the dehydration of isobutyl alcohol, the 7 7 f / species has the structure of the isobutyl silyl ether, this ether being the true reaction intermediate, since 100% of the isobutyl alcohol reactant is dehydrated (8e) into it via reaction + I1 of Scheme 1. In the absence of the additional flow of isobutyl alcohol molecules, it is stable up to 398 K, but decomposes spontaneously upon

352

KIRlL ILYCH ZAMARAEV AND JOHN MEURlG THOMAS

-

Pathway I

\

0-

/ \

0

/ \ /

/% P i A

*\'

\

/ \

/ \ /

AOP\ 'O A

I

cH2

0- /0\ / \

\

/

Si\ A A Pathway I1

CH C H - & ~ - C H ~ \

3-+ 0I \

0

/ \ /

B'\ A A

tl

B SCHEME 3

fiu-ther heating to give butene oligomers. As shown in Scheme 3, the IBSE structure rearranges reversibly into the IBCI structure through which butene product is expected to form. For tert-butyl alcohol, we have observed (Sf)with I3C CPMAS and 2H NMR spectroscopy, the formation of tert-butyl silyl ether (TBSE)at temperatures as low as 296 K. In these experiments, 13C and 2H-labeledtert-butyl alcohols were used:


CH3, CH,-*CH-OH, /

CH,

*CH3 \ CHI-C-OH, /

CH3

and

353

CD, \ CD3-C-OH /

CD3

However, we found that TBSE is a rather stable species, decomposing only upon heating above 373 K. Thus, at 296 K it behaves as a side intermediate species, through which only a small fraction of tert-butyl alcohol molecules dehydrate. The main reaction stream bypasses the TBSE structure, proceeding presumably through the tert-butyl cation ion as the key intermediate (Scheme 4). It is interesting that-regardless of where the I3C label is initially placed (i.e., into the CH-OH or a CH3 g r o u p t a f t e r keeping tert-butyl alcohol within the HZSM-5 sample at 296 K for 15 h, a complete scrambling of the 13C label is observed among various positions in the TBSE intermediate and butene aligomers formed 8'( in HZSM-5 channels from butene under the conditions of our NMR and FTIR studies. Moreover, if D20and (CH3)3COH are simultaneously adsorbed at 296 K in HZSM-5, the deuterium is transferred from the D 2 0 molecule to methyl groups of the alcohol. Similarly, in the course of (CD3)&OH dehydration in HZSM-5, deuterated water is formed as a reaction product. Thus, scrambling of deuterium atom between water molecules and methyl groups of tert-butyl alcohol molecule is facile. Undoubtedly, such scrambling phenomena are possible only with the participation of carbenium ions as reaction intermediates or transition states. Molecular rearrangements that rationalize the scrambling of deuterium atoms are shown in Scheme 5. OR

T!T

Thus, reaction intermediate for dehydration of butyl alcohols can exist in two forms, i.e., butyl silyl ether (BSE) and adsorbed butyl carbenium ion (BCI). Our NMR and kinetic data imply the existence of reversible transformations between BSE, BCI, and adsorbed butene (Bua& that ar shown in Scheme \ / 6. BUads is butene that is hydrogen-bonded to /AI-O(H)-Si\. It has been

SCHEME 4

354

KIRIL ILYCH ZAMARAEV A N D JOHN MEURIG THOMAS

L

CHs-h-CH2D

+

+ 0‘ + DHO

I

I

CH3-C-CH2D

+ OD

I

I

OH

SCHEME 5

observed earlier on HNaY zeolite with IR spectroscopy by Pauhhtis et al. (5a) and on HZSM-5 and HY zeolites with I3C NMR spectroscopy by Lazo et al. (5b) at lower temperatures. BCI is a more labile chemical structure centrally involved in the main reaction stream. BSE is a more stable structure, but is sometimes still active enough to participate in the main reaction stream (the case of isobutyl alcohol). But sometimes it is too stable to be in that stream, as is the case for tert-butyl alcohol. Then it plays the role of the side intermediate or, under certain experimental conditions, may even become the deadend of the reaction similar to butene oligomerization. Bu,ds may be a precursor of butene production in the reaction stream.

VI.

Concluslons

The reaction mechanism for the catalytic dehydration of all four butyl \ / alcohols over the same Brensted acid yAl-O(H)-Siy sites in HZSM-5 zeolites with different crystallite sizes and in amorphous aluminosilicate (AAS) catalyst may be rationalized by one all-embracing scheme. Reaction pathways are found to be identical for all the butyl alcohols and the catalysts studied, with respect both to the main reaction stream and to the side reactions. However, depending on the particular structure of both the reactant (i.e., n-, sec-, iso-, or tert-butyl alcohol) and catalyst (i.e., amorphous AAS or crystalline HZSM-5, of specific crystallite size), the observed reaction rates and selectivity

B“

r

BC I SCHEME 6

‘“ads

355


toward various products can be dramatically different, even under similar reaction conditions. All these differences are, however, well understood in terms of the proposed and tested reaction mechanism. Three important general conclusions may be drawn from our studies. The first OC4b

refers to the nature of the key reaction intermediate nb , which clearly exhibits carbenium-ion properties, such as scrambling of carbon and hydrogen atoms over its skeleton. In this crucial respect, dehydration of butyl alcohols over solid HZSM-5 and AAS catalysts mechanistically resembles homogeneous acid-catalyzed reactions, whereby carbenium ions in their classical forms serve m4H9

as key reaction intermediates. But the nb intermediate does not coincide exactly with a classical carbenium ion for the following reasons: OC4H9

1. T/fT can exist in three different states: in an ion pair involving a butyl \ / carbenium ion and the yA1-0- - S i y residue of the solid catalyst; as a butyl \ / silyl ether; and as butene hydrogen-bonded to the ~ A l - O ( H ) - S i ~ group. Interconversions between these states (Scheme 6) as well as intramolecular \ / rearrangements in the BCI//Al-O--Si~ ion pair of the type shown in Schemes 3-5, proceed with finite rates that vary with the temperature. (We note in passing that catalytic chemists may fine-tune those rates in reaction engineering by, e.g., making some or all of them lower than the rates of formation andor decomposition of . In this way, reaction selectivity with respect to a desired product can be optimized.) OC4H9

2. Interaction of a carbenium-ion state of nb with the wall of the catalyst pore can kinetically favor reactions that are not favored for carbenium ions in solutions. For example, pore confinement in HZSM-5 favors the formation of linear C4H: (Section V), although branched C4H: are favored in solution both kinetically and thermodynamically. Our conclusions support the ideas of Derouane et al. (16) who, by analogy with enzymes, proposed that confinement effects must be important for catalysis with zeolites. We would expect peculiarities (1) and (2) to be also inherent in carbeniumion-type intermediates of other reactions (such as the cracking, alkylation, and isomerization of hydrocarbons) occurring in the pores of solid acid catalysts. OR

The knowledge of the properties of the key reaction intermediate nb alerts one to what may be done to improve the performance of acid catalysts in converting various organic feedstocks into desired products. One may, for

356

KIRIL ILYCH ZAMARAEV AND JOHN MEURIG THOMAS \

example, try to design catalysts where the /Al-O(H)-Siy

/

active site is

located in pockets the size and geometry of which match best those of the particular carbenium ion that leads to the formation of the desired product. In particular, we may look forward to designing a catalyst that would selectively convert a mixture of butene isomers into only one isomer (e.g., isobutylene) ( I 7) via the hydration of the butene feedstock into alcohols followed by their subsequent selective dehydration into the desired olefin only. The second conclusion that we draw is that, in the catalyst pores, reaction mixtures that are conventionally treated as gases may, in fact, exist in a liquidlike form. We believe that for catalysts with fine pores, this indeed must be a general phenomenon if only because of the ubiquity of capillary condensation. This being so, it follows that reactions in catalyst pores should demonstrate certain effects that can be neglected in gases but not in liquids. Among these are the cage effect, the influence of the dielectric properties of the liquid on reaction rates, solubility effects, and so on. Solubility effects may explain the formation of immiscible liquid phases by mixtures of polar and nonpolar substances, as well as segregation in catalyst pores of gaseous mixture of an organic compound with H2 or 0 2 , into a mixture of the gaseous H2 or O2 with the organic liquid where solubility of H2 or O2 is rather poor. Recently, such segregation of reagents in catalyst pores has been reported ( I d ) in hydrogenation and hydrodearomatization processes. It seems that such solubility effects may exert quite dramatic effects on the overall reaction kinetics. Deliberate, premeditated design of the porous structure can help to utilize the positive effect (e.g., easier separation of reaction products) as well as to suppress the negative effect (e.g., decrease of the reaction rate) of such segregation. It is interesting that, for reaction mixtures consisting of molecules with dimensions close to the cross-sections of the catalyst pores (as in the case of butyl alcohol dehydration in HZSM-5), the reacting mixture may be envisaged as a liquid with dimensions less than three. This, in turn, introduces additional factors with respect to the unanalyzed peculiarities of mass-transfer kinetics in the catalyst pores. The third conclusion that we draw is that a reactant may play a twofold role in catalysis, acting not only as a reactant but also as an agent that protects the active sites from becoming involved in undesirable size reactions. Once the importance of such an effect is recognized and understood, one may formulate other ways of protecting active sites from undesirable reactions, e.g., by intentionally adding to the reaction mixture some special substance (other than the reactant) that serves to protect active sites but is not consumed during the course of reaction. Such an approach resembles the strategy of choosing an appropriate solvent composition when carrying out homogeneous catalytic reactions in solution.


357

ACKNOWLEDOMENTS The authors are grateful to the Russian Academy of Sciences, the Royal Society, and the Science and Engineering Research Council for their financial support of the research summarized in this account. This review was prepared during the tenure of a Kapitza Fellowship by one of us (KIZ) at the Royal Institution. REFERENCES 1. Sanefield, C. N., “Heterogeneous Catalysis in Practice.” McGraw-Hill, New York, 1980. 2. Rabo, J. A., Ed., in Zeolite Chemistry and Catalysis: ACS monograph. Amer. Chem. SOC.,

Washington, D. C., 1976. 3. Olah, G. A., and Schlosberg, R. H., Eds. “Carbonium Ions.” Vols. I-IV. Wiley, New York, 1968. 4. Kazansky, V. B., Ace. Chem. Res. 24,379 (1991). 5. Zamaraev, K. I., and Zhidomirov, G. M., in Proc. 5th Intern. Symp. on Relations between Homogeneous and Heterogeneous Catalysis, Novosibirsk, 1986. (Yu. 1. Yermakov and V. A. Likholobov, Eds.), Vol. 1, p. 23. VNU Science Press, Utrecht, 1986. 60. Zhidomirov. G. M., and Kazansky, V. B., Adv. Caral. 34, 131 (1986). 66. Sauer, J., Chem. Rev. 89, 189 (1989). 6c. Jackson, R. A,, and Catlow, C. R. A,, Mol. Simulation I, 207 (1988). 6d. van Santen, R. A,, de Bruyn, D. R., den Ouden, C. J. J., and Smit, B., in “Introduction to Zeolite Theory and Modelling” (H. van Bekkum, E. M. Flanigen, and J. C. Jansen, Eds.), Introduction to Zeolite Science and Practice, Studies in Surf. Sci. Catal. Vol. 58, p. 317. Elsevier, Amsterdam, 199 I . 7a. Liengme, B. V., and Hall, W. H., Trans. Faraday SOC.62, 3229 (1966). 76. Lombardo, E. A,, Sill, G.A,, and Hall, W. K., J. Cafal. 22, 54 (1974). 7c. Kemball, C., Leach, H. F., Sundrich, B., and Taylor, K. C., J. Catal. 27,416 (1972). 7d. Kemball, C., Leach, H. F., and Moller, B. W., J. Chem. Suc. Faraday Trans. I624 (1973). 8a. Makarova, M. A,, Williams, C., Romannikov, V. N., Zamaraev, K. I., and Thomas, J. M., J. Chem. SOC.Faraday Trans. 86, 581 (1990). 86. Williams, C., Makarova, M. A,, Malysheva, L. V., Paukshtis, E. A., Zamaraev, K. I., and Thomas, J. M., J. Chem. SOC.Faraday Trans. 86, 3473 (1990). 812.Makarova, M. A,, Williams, C., Thomas, J. M., and Zamaraev, K. I., Cafal. Leff. 4, 261 ( 1990). 8d. Williams, C., Makarova, M. A., Malysheva, L. V., Paukshtis, E. A., Talsi, E. P., Thomas, J. M., and Zamaraev, K. I., J. Catal. 127, 377 (1991). 8e. Stepanov, A. G.,Romannikov, V. N.. and Zamaraev, K. I., Catal. Left. 13,407 (1992). ’8 Stepanov, A. G.,Zamaraev, K. I., and Thomas, J. M., Catal. Left. 13, 407 (1992). 8g. Stepanov, A. G.,Maryasov, A. G.,Romannikov, V. N., and Zamaraev, K. I., Proc. 10th In. Congr. Catal., Budapest, 1992 (L. Guszi, F. Solymosi, and P. Tetenyi, Eds.), Part A, p. 621. Akademiai KiadoElsevier, Budapest, 1993. 8h. Stepanov, A. G.,Maryasov, A. G.,Romannikov, V. N., and Zamaraev, K. I., Magn. Res. Chem. 32, 16 (1994). 8i. Stepanov, A. G.,and Zamaraev, K. I., Caral. Left. 19, 153 (1993). 8j. Makarova, M. A., Paukshtis, E. A., Thomas, J. M., Williams, C., and Zamaraev, K. I., J. Cafal. 149, 36 (1994). 8k. Makarova, M. A,, Williams, C., Zamaraev, K. I., and Thomas, J. M., J. Chem. Soc. Faraduy Trans. 90, 2147 (1994). 81. Makarova, M. A., Paukshtis, E. A,, Thomas, J. M., Williams, C., and Zamaraev, K. I., Caral. Today 9 61 (1991).

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8m. Thomas, J. M., and Zamaraev, K. I., Top. Curd 1, 1 (1994). 9u. Yue, P. L., and Olaofe, O., Chein. Eng. Res. Dev. 62, 81, 167 (1984). 9b. Aronson, M. T., Gorte, R. J., and Fameth, W. E., J. Cutul. 98, 434 (1986); ibid. 105, 455 (1987). 9c. Kalvachev, Yu.,Bezouhanova, C., and Lechert, H. Zeolites 11, 7 3 (1991). IOU. Olson, D. H., Kokotailo, G. T., Lawton, S. L., and Meier, W. M., J. fhys. Chem. 85, 2238 (1981). lob. Lermer, H., Draeger, M., Stefen, J., and Unger, K. K., Zeolites 5, 131 (1985). IOc. van Koningsveld, H., Jansen, J. C., and van Bekkum, H., Zeolites 10, 235 (1990). 11. Paukshtis, E. A., and Yurchenko, E. N., Usp. Khim. 52, 426 (1983 (in Russian). 12u. Paukshtis, E. A., Soltanov, R. I., and Yurchenko, E. N., React. Kinet. Cutul. Lett. 19, I19 (1982). 126. Cardona-Martinez, N., and Dumesic. J. A., J. Curul. 125, 427 (1990). 13u. Knozinger, H., and Kohne, R., J. Cutul. 5, 264 (1966). 13b. Knozinger, H., Angew. Chem., Int. 7 , 791 (1968). 13c. Jacobs, P. A., Tielen, M., and Uytterhoeven, J. B., J. Cutal. 50, 98 (1977). 13d. Moravek, V., and Kraus, M., Collect. Czech. Chem. Commun. 51, 763 (1986). 14. Zhorov, Yu. M., “Isomerisation of Olefins.” Khimia, Moscow, 1977 (in Russian). 150. Paukshtis, E. A.. Malysheva, L. V., Kotsarenko, N. S.,and Karakchiev, L. G., Kinet. Kutul. 21, 455 ( I 980) (in Russian). 1.56. Lazo, N. D., Richardson, B. R., Schettler, P. D., White, J. L., Munson, E. J., and Haw, J. F., J. fhys. chem. 95,9420 (1991). 16. Derouane, E. G . , Andre, J-M., and Lucas, A. A., J. Cutul. 110, 58 (1988). 17. Natarajan, S., Wright, P. A., and Thomas, J. M., J. Chem. Soc. Commun. 1861 (1993). 18. Ostrovskii, N. M., Bukhavtsova, N. M., and Duplyakin, V. K., React. Kine!. Cutul. Lett. 53,253 ( I 994).

ADVANCES Ih' CATALYSIS. VOLUME 41

Thermal and Catalytic Etching Mechanisms of Metal Catalyst Reconstruction TA-CHIN WE1 AND JONATHAN PHILLIPS The Pennsylvania State Universiv Department of Chemical Engineering 133 Fenske Laboratory Universiry Park, Pennsylvania 16802-4400

1.

Introduction

Metals can undergo two types of etching processes in single gases and gas mixtures: thermal etching and catalytic etching. Since the differences between these two types of etching are not clearly delineated in most studies, some confusion has arisen. Thus the first goal of the present review is to provide clear definitions of the two types of etching so that the processes can be readily distinguished. The mechanism of etching, particularly that of catalytic etching, is still in dispute. Therefore; the second goal is the presentation of a systematic and critical discussion of the various proposed etching mechanisms. The third goal is to demonstrate that the mechanism of catalytic etching is virtually identical to that of other types of etching. In particular, it will be demonstrated that plasma etching, etching in low-earth orbit, and catalytic etching all occur via related mechanisms. The first three goals focus on scientific aspects of etching. In contrast, the fourth goal is to demonstrate that etching is technologically significant. There are numerous areas of technology in which etching is critical. For example, the impact of catalytic etching on the properties of metal gauzes, used as catalysts in a number of large scale industrial processes, has been a recognized problem for more than 70 years. Catalytic etching results in the loss of valuable metal, active catalyst surface, and mechanical strength. Potential positive applications of catalytic etching exist as well. For example, some previously discovered catalytic etching phenomena may lead to novel plasma processes which could be of value in integrated circuit processing. Indeed, at present one difficulty with using copper as an interconnect material in integrated circuits is the fact that there are no suitable technologies for dry etching copper. Yet, copper is known to etch catalytically in hydrogen-oxygen mixtures. This fact suggests some new technological strategies for dry etching copper. 359 Copyright 0 1996 by Academic Press. Inc. All rights of reproduction in any form reserved.

360

TA-CHIN WE1 AND JONATHAN PHILLIPS

Thermal etching also has technological significance. For example, current crude oil refining technology practice involves the treatment of reforming catalysts, used in continuous beds, with chlorine. This treatment redisperses the metal via a type of thermal etching, as described below. Catalyst particle “cracking,” another form of thermal etching, is being adopted to assist in the redispersion of precious metals used in automobile exhaust abatement catalysts. This technology may permit less expensive metals to replace rhodium in this application. A.

PHENOMENOLOGY AND DEFINITION

Thermal etching is defined to be any metal surface reconstruction that can occur in the absence of chemical reactions involving gas-phase species. By this definition, reconstruction that occurs in the presence of reacting mixtures, but which mechanistically could occur in the absence of any reaction between gas phase-species, is still clearly thermal etching. In general, the phenomenology of thermal etching is the reconstruction of the surface to produce facets and/or pits with little or no mass change. However, when molecular gas-phase species interact with metal to form volatile species significant loss of mass can occur. Catalytic etching is defined to be any surface reconstruction that can take place only in an environment in which there are reactions between gas-phase species. Metal reconstruction is the result of interactions between species formed due to these reactions and the metal surface. Catalytic etching is normally characterized by severe corrosion at temperatures far below the melting temperature of the metal. Because it often involves significant metal volatilization, catalytic etching can be a major industrial problem. Generally, catalytic etching results in significant loss of mass, although in special cases it leads to the redeposition of the metal as particles without significant net metal loss. In all cases, catalytic etching leads to the production of highly irregular surface structures, which are not minimum-energy configurations. Some related processes are difficult to classify. One example is the formation of blisters and pits in metals which sometimes accompanies treatment of metals, either simultaneously or sequentially, with hydrogen and oxygen (embrittlement). Limitations of the classification scheme with respect to some reconstruction processes are discussed below.

B. MECHANISMS The above definitions are somewhat different fiom those given in earlier publications. In earlier studies, thermal etching has been defined to be etching that takes place in “nonreactive” or “mildly reactive” gases or under vacuum. Catalytic etching has been defined to be any etching process that takes place in

THERMAL AND CATALYTIC ETCHING

36 1

“reactive” gas environments. The revised definitions given here are believed to better reflect the significant differences in the mechanisms of the two types of etching revealed in recent studies. In particular, a significant shift in understanding catalytic etching has occurred in the last ten years. A number of mechanisms leading to reconstruction result from the interactions of single gases with metals, and all are considered thermal etching. First, reconstruction results from the evaporation of metal as the temperature approaches the melting point. Second, surfaces can reorganize into lower energy configurations via localized surface diffusion without metal volatilization or any loss of mass. Third, reconstruction sometimes results from evaporation of compounds that form due to interactions between gas-phase molecules and metal surface atoms. This includes the evaporation of metal oxide species in an oxygen atmosphere and the formation and evaporation of metal carbonyls in a gas environment containing significant amounts of carbon monoxide. Such processes can take place at temperatures far below the metal melting temperature. Fourth, supported (catalyst) particles sometimes “crack” as a result of strains induced by metal/gas interactions. There is a history of debate regarding which of the above processes is primarily responsible for thermal etching. It is argued in the present review that all of these processes take place, and that the dominant mechanism is a function of the particular system. In contrast to thermal etching, for which many mechanisms exist, recent investigations suggest that one particular mechanism explains the vast majority of reconstructions that can occur only in the presence of reacting gas mixtures. That is, catalytic etching generally results from the interaction of free radicals, formed via gas-phase reactions, with the metal surface. Homogeneously formed free radicals interact with the metal to form volatile, metastable metal-containing intermediates. This is consistent with the finding that etching often occurs near known explosion limits. In some cases, the volatile intermediates interact in the gas phase to produce metal particles. These particles are then redeposited on the metal surface, leading to the production of particle-covered surfaces. In other cases, the volatile intermediates “wash out” of the reactor, leading to rapid weight loss. The above model of catalytic etching is not universally accepted. Several older mechanistic models exist. For example, it was suggested that localized temperature gradients, induced by surface reactions, might lead to uneven rates of diffusion or volatilization, and hence catalytic etching. Evidence from investigations outside the field of catalysis tends to support the more recent model. Specifically, a great deal of work outside of catalysis shows that free radicals are responsible for many etching processes. For example, several studies designed to demonstrate the existence of free radicals show that methylene radicals will cause volatilization of a number of metals. Also, modern research into the mechanism of etching in plasmas (dry etching for integrated

362


circuit processing) indicates that free radicals produced in the plasma are often the sole agent responsible for metal etching. Thus, for many metal-etching applications the plasma simply serves as a convenient means to generate free radicals. Finally, a number of studies of material damage to spacecraft in low earth orbit indicate that free radicals, particularly oxygen atoms, are responsible for much of the damage/etching of polymer and metal components. The theory that homogeneously formed free radicals are responsible for catalytic etching is consistent with the conclusions reached in these other fields, namely, that free radicals can interact with metals to form volatile intermediates.

11.

Thermal Etching

Early studies of the etching of a number of metals, particularly silver, copper, platinum, and tungsten, have been thoroughly reviewed by Shuttleworth in 1948 (I), Moore in 1963 (2), and Flytzani-Stephanopoulos and Schmidt in 1979 (3). Little of the early work is repeated here; rather, the goal of the following review is to focus on the significant contributions and clarify fundamental concepts. A.

CLASSICAL THEORIES

Two theories of thermal etching emerged from early studies: namely, that thermal etching results either from (i) enhanced evaporation (kinetic control) of metal or metal complexes (e.g. metal oxides) from specific surface planes or from (ii) surface migration of species to reorganize the surface into a minimum energy state (thermodynamic control). In some cases it was suggested that both processes worked simultaneously, although the relative importance of each is not generally resolved. B. EARLY EXPERIMENTAL INVESTIGATIONS

Experimental investigations conducted before about 1970 were characterized by the use of a variety of optical microscopy techniques and thus were limited to relatively large samples, such as foils, wires, and gauzes. A number of general conclusions can be drawn about the phenomenology of the process of thermal etching from a careful consideration of the early work. First, the temperature at which facet formation is found to occur roughly relates to the volatility and/or the melting temperature of the metal. Second, the nature of the facet pattern is a function of the composition of the gas phase. Indeed, in many gases, facets do not form at all. Third, under low-pressure and vacuum conditions the samples tend to form pitted surfaces more readily or to be characterized by


3 63

“hill-and-valley structures.” Fourth, when evaporation is “inhibited,” surfaces are relatively smooth. 1. Unsupported Metals under Vacuum or in a Single-Component Gas The metal first and most frequently studied in conjunction with etching is platinum. One of the earliest and most complete investigations of platinum weight loss was conducted by Hulett and Berger (4), although qualitative observations were made as early as 1879 (5). Hulett and Berger observed measurable weight loss from platinum foils and wires following electric heating in flowing air at temperatures as low as 800°C. No weight loss was measured in experiments conducted at lower temperatures or when the experiments were “conducted in vacuo” (no details given), in agreement with earlier work (6). The authors also collected platinum crystals, 80 to 200 pm in diameter, on relatively cool surfaces about 1.5 cm from wires of platinum heated electrically in air. They conjectured that stable, volatile compounds of platinum form in the presence of oxygen at high temperatures but are unstable and decompose at lower temperatures. No information regarding the structure or appearance of the surface of the metal is available in these early reports. In 1902 Rosenhain (7) reported that grain boundary grooves form on platinum following high-temperature treatment. Later, Rosenhain and Humphrey (8) noted the same phenomenon for steel heated in air at elevated temperatures. The first report of “striations” came from Rosenhain and Ewen in 1912 (9) for silver heated in air (Fig. 1). Leroux and Raub (10) performed a detailed study of the etching of silver and silver alloys in oxygen and hydrogen at elevated temperatures. They were the first to propose that thermal etching is kinetically controlled. Specifically, they suggested that faceting results from differences in the rate of evaporation from different planes and that evaporation rates are influenced by adsorbed species. investigations of other systems supported the model of facet formation resulting from enhanced evaporation of metal oxides from selected surface planes. For example, Elam (11) suggested a similar explanation for the formation of facets on copper heated at elevated temperatures (950°C) in air at low pressures. Studies conducted in the 1950s of platinum surfaces treated in oxygen also indicate that the rates of evaporation and facet formation are enhanced for platinum heated in oxygen at temperatures exceeding 1200°C (12-14). Gwathmey and Benton (Z5-17), studied the behavior of copper spheres in different environments and found that the nature and velocity of the transformations is a function of the composition of the gas phase. The finding is consistent with etching resulting from the evaporation of volatile species. According to a second class of model, thermal etching is driven by a need to reduce total surface free energy. According to this theory, faceting will take place even in the absence of any net weight loss. The first to suggest a model of

364

TA-CHIN WE1 A N D JONATHAN PHILLIPS

FIG. I . Optical micrographs (300X) from an early study of thermal etching showing the striation structure on silver heated in air (9).

this variety was Johnson (18), who noted that the method of heating was important in determining the final structure of the surface. Direct current produced faceted surfaces on tungsten, tantalum, nickel, molybdenum, platinum, and iron surfaces. Thermal gradients were also shown to produce faceted surfaces. On the basis of these results, Johnson concluded that facets form via the surface migration of ionic species. Others who studied the effect of directcurrent heating on faceting supported the suggestion of a thermodynamic driving force for faceting (19, 20). On the basis of other etching experiments with silver, Chalmers et al. (21) suggested an explanation for thermal etching involving both of the earlier models. They concluded that thermal etching results from two processes. Specifically, grain boundary grooving was postulated to be thermodynamically driven, with the surface rearranging itself in the grain boundary region via a


365

process of atomic diffusion, tending toward the minimum energy configuration. On areas separate from the grain boundaries, they suggested (in agreement with the earlier workers) that facet formation results from a kinetic process in which silver oxide-like species evaporate at a higher rate from particular locations on the surface. The two models were tested by Hondros and Moore (22), who concluded that etching is in fact a kinetically controlled process. They heated a silver foil in air in a closed silver container to reduce net evaporation to zero. No weight loss and no evidence of the formation of facets or significant grain boundary grooving was detected. In contrast, in flowing air in a glass reactor, etched surfaces formed very quickly. They argued that if thermodynamics were driving the etching process, then faceted surfaces should have formed both with and without net evaporation. Moreover, they found that surfaces which were originally faceted by heat treatment in flowing oxygen became smooth after heat treatment in the closed silver box. On the basis of this finding, they raised a question regarding the true equilibrium surface stucture. Hondros and Moore (22) also found that heating under vacuum or in nitrogen did not lead to the formation of facets. The latter produced smooth surfaces, and the former led to pit formation. From these data and the studies of earlier workers, the authors concluded that the mechanism of silver etching is the preferential evaporation of silver oxide from certain surface planes. Rhead and Mykura (23) tried with limited success to repeat the results of Moore. They found that limited faceting took place when net evaporation was suppressed; however, the extent of the faceting was found to be far less than in the case of heating the foils in open air. Recent investigations of the thermal etching of silver (24) suggest that evaporation (kinetic model) correctly describes etching processes for silver under vacuum. In contrast, under oxygen a drive to attain equilibrium (thermodynamic model) appears to accurately explain most observations. Specifically, it was found that under ultra-high vacuum conditions hill-and-valley structures form only when free evaporation is permitted. When evaporation was suppressed, no facets formed. This modern work may also explain the earlier experimental discrepancies. It was shown that as the temperature is raised and evaporation increased, faceting disappeared even in the free evaporation case. This was attributed to a reduction in surface anisotropy and hence differences in relative evaporation rates which generally accompany increases in temperature. Experiments conducted by the same group for the thermal etching of silver in an oxygen atmosphere suggest that evaporation of metal plays little role in the thermal etching process. The thermodynamic model appears to best explain the observations. That is, identical faceted surfaces formed both in the case of suppressed and free evaporation.

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Winterbottom (25) found that in any gas, at least at very low pressures ( 3), the samples became covered with relatively large platinum particles (average diameter approximately 1 pm) (region 11). In excess fuel mixtures (02/C2H4 < 3) over the same temperature range (region IV), the sample surfaces were covered with a thick layer of carbon in which a high concentration of small platinum particles (diameter < 0.1 pm) were embedded (Fig. 17) (35, 169). At temperatures exceeding 677°C in excess oxygen (region 111), the foils were again found to be simply faceted. In excess fuel at the same high temperatures (region V), the foils were faceted but also partially covered with carbon. It was concluded that the structures that formed in regions I1 and IV could not be attributed to the sum of the etching effects of the individual gases. There was a synergism resulting from the mixing of gases. Moreover, this synergism existed only over a limited range of temperature. It was postulated that the synergism resulted from the gas-phase formation of a particular free radical, perhaps methylene (33, 170-1 72), which was formed as part of the principal mechanism for homogeneous ethylene oxidation over a limited temperature range. It was also suggested that the hydroperoxyl radical might be responsible. Similar postulates of free-radical-initiated processes have been made to explain other phenomena, such as the formation of soot, which forms only over the range of temperatures at which a particular free radical is produced (I 73-176). Another study of the same system was conducted, with the focus on thin silica-supported platinum films in an impinging jet flow (120). It was shown (Fig. 18) that in regions I1 and IV all of the platinum directly below the impinging jet was removed from the thin film. Large platinum particles (ca. 1 pm in diameter) were found downstream on top of regions of unetched platinum film several millimeters from the area of impingement of the jet. Many


399

FIG.17. The character of etching of platinum foils in ethylene/oxygen mixtures is a function of gas phase composition. Under excess fuel conditions between 500 and 700°C, carbon films containing platinum particles form on top of the foil. TEM micrograph of a carbon film formed after 40 h of treatment at 587°C in a fuel excess 0&H4 mixture (region IV). (Reprinted with permission from (35).Copyright 1985 American Chemical Society.)

platinum particles were also found as far as several centimeters from this area on top of silica in regions where no metal was initially present. In region IV, platinum etching was accompanied by the deposition of thick carbon films. In all , other regions and in the presence of each of the individual gases (N2, 0 2 C2H4, H20,C 0 2 ) , no metal transport and no large particle formation was observed. Only limited sintering, quantitatively identical to that occurring under vacuum at the same temperature, was observed. Again, particular attention was paid to the influence of oxygen, and again no evidence was found for any difference between the influence of oxygen, any other gas, or vacuum. To determine the plausibility of the postulate of hydroperoxyl radical intermediates, a study was conducted of the effect of hydrogedoxygen mixtures on the same types of films in the same impingingjet reactor (142). Detailed kinetics suggested that at low temperatures (ca. 480°C) in either ethylene/oxygen or hydrogedoxygen mixtures the hydroperoxyl radical is the dominant free radical present during combustion. Moreover, it is clear that no hydrocarbon radicals were present under reaction conditions. Thus, if etching were observed in a

400


FIG.18. Etching in ethylene/oxygen mixtures of platinum thin films on Si02 substrates in an impinging jet reactor dramatically illustrate that metal is volatilized under some conditions: Hole formation in (a) region 11 and (b) region IV (see Fig. 16). In all other regions of the phase diagram and in any single product or reactant gas no metal was removed from the films (120).

hydrogedoxygen mixture, it would seem unlikely that any hydrocarbon free radical (e.g., methylene) would be responsible for the etching observed in ethyleneloxygen mixtures.


40 1

Studies of etching in hydrogedoxygen mixtures appeared to support the freeradical-induced etching postulate. Again it was discovered that etching occurred only in the presence of reaction mixtures and only over a limited range of temperatures and gas compositions. Specifically, etching took place only at temperatures greater than about 450°C and in mixtures in which the oxygen/ hydrogen molar ratio was greater than about 2/l. The nature of the etching was very similar to that observed in ethylene/oxygen mixtures. Metal was removed only in the area of impingement of the jet, and many large particles of platinum were found downstream. The only significant difference was that a band of large particles (cu. 5 pm in diameter) formed at the edge of the region from which all metal had been removed (Fig. 19a). A band of particle clusters was found just beyond the large-particle band. These clusters were clearly composed of smaller particles about 1 pm in diameter (Fig. 19b). The similar nature of etching in ethyleneloxygen and in hydrogedoxygen mixtures suggested that the same free radical was responsible for etching in both cases. Thus all carbon-containing free radicals were eliminated as candidates. Moreover, recent investigations of the combustion of hydrogen (I 77-183) and of ethylene (183) suggest that at low temperatures the dominant reaction mechanism involves the hydroperoxyl (H02) radical. Earlier suggestions that the methylene radical dominates the oxidation process at low temperatures are no longer considered valid. To explain the particles that formed in both the ethylene/oxygen and hydrogen/oxygen mixtures, it was postulated that they form in the gas phase and that the overall etching process takes place in three steps. First, free radicals are formed homogeneously in a boundary layer adjacent to the surface. Second, these radicals interact with metal atoms in the surface. This interaction results in the formation of volatile intermediates. Third, the metastable, volatile intermediates interact in the gas phase so that metal particles are formed and stable product molecules released. Individual metastable species presumably interact with each other and also with particles formed from multiple collisions. The larger particles interact with each other as well. The third step in the model of particle growth in the gas phase is very similar to that describing the growth of soot (I 73-1 76). The general mathematical analysis of this type of particle growth was first developed by Smoluchowsky (184-186) and is also used to describe other processes, such as particle sintering. Next it was demonstrated that the nature of the sintering of platinum particles in Pt/Si02 is different in ethylene/oxygen mixtures than in any single-component gas (82). Specifically, it was found that sintering was dominated by Ostwald ripening (single-atom transport) under etching conditions and by coalescence (whole-particle transport) growth under all other conditions. Some (187, 188) have also suggested that particle sintering is greatly accelerated under

FIG. 19. Particle structure on catalytically etched Pt thin films exposed in H2/02 mixture is a function of position relative to the impinging jet. (a) Just beyond the area from which platinum is fully removed a ring of large faceted single crystals is found. (b) Beyond the ring of large single crystals lies a ring of particle-agglomerate structures. It is suggested that both structures are initially composed of small particles which formed in the gas phase and precipitated onto the surface. Particles in the outer ring are agglomerates of these small particles. Particles in the inner ring are re-etched agglomerates. (Reprinted with permission from (142). Copyright 1988 American Chemical Society.)


403

reaction conditions. It is possible that these are also examples of free-radicalenhanced transport. X-Ray diffraction was used to investigate silica-supported platinum particles after treatments in either single-component gases or after treatment in reaction regimes known to result in etching. The X-ray Pt( 111) peak was deconvoluted after various treatments to yield the platinum particle size distribution. It was found that in any single-component gas or under vacuum, the particle size distribution changed as anticipated for coalescence growth. In short, it appears that the smaller particles in the distribution coalesced to form larger particles. The large particles, because of stronger interactions with the support, diffuse much more slowly and are rarely involved in coalescence. In contrast, under reaction conditions known to result in etching, only the larger particles grew. The number of smaller particles decreased, but there was no indication of growth. This type of behavior is consistent with Ostwald ripening. During Ostwald ripening there is atom-by-atom transport from smaller particles to larger ones, driven by the net difference in vapor pressure as a hnction of particle size. These results support the hypothesis of etching by free radicals. Only under reaction conditions is it expected that there will be a mechanism for atomic transport resulting in Ostwald ripening, whereas in other environments only coalescence growth is anticipated. The results of other studies add support to the free-radical hypothesis. For example, the etching of Pt/Rh (904 0) gauzes in hydrogerdoxygen mixtures occurred only over the same range of reaction conditions under which pure platinum is etched (51). Etching was demonstrated by two different phenomena. First, in a simple flow reactor, operated so that rapid mass transfer away from the metal surface was favored, the gauzes underwent dramatic weight loss, but only under the known etching conditions. No weight loss was measurable in any pure gas, or under reaction conditions which were previously shown not to etch pure platinum. Second, in an impinging jet reactor, which does not favor mass transport away from the metal, there was massive restructuring of the surface, but only under conditions also known to etch pure platinum. Very little weight loss was detected. The observed reconstruction of Pt/Rh (90/10) gauzes in hydrogetdoxygen mixtures provides additional insight into the etching process. In the flow reactor in which significant weight loss was detected there was no observable modification of the surface morphology. The surface was striated and lightly faceted both before and after treatment. In the impinging jet reactor the reconstruction took the form of a layer of particles forming on the gauze surface. As shown in Fig. 20, directly under the impinging jet the particles were relatively small (

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