Advances in Catalysis, Volume 36

ADVANCES IN CATALYSIS VOLUME 36 Advisory Board M. BOUDART Stanford. Calijomia V. B. KAZANSKY Moscow, USSR. G. A. S...

Author: D. D. Eley | Herman Pines | Paul B. Weisz

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ADVANCES IN CATALYSIS VOLUME 36

Advisory Board

M. BOUDART Stanford. Calijomia

V. B. KAZANSKY Moscow, USSR.

G. A. SOMORJAI Berkeley. Calijornia

G. ERTL Berlid!hlem. F.R.G.

A. OZAKI Ilbkya Japan

W. 0. HAAG Princeton, New Jersey

W. M. H. SACHTLER Evanston, Illinoh

J. M. THOMAS London, UK.

ADVANCES IN CATALYSIS VOLUME 36

Edited by D. D. ELEY The University Nottingham. England

PAULB. WEISZ

HERMAN PINES Northwestern University Evanston, Illinois

University of &nnsylvania Philadelphia, &nnsylvania

ACADEMIC PRESS, INC. Hmourt Brace Jovnnovich, Publishers

San Diego New York Berkeley Boston London Sydney Tokyo Toronto

COPYRIGHT 0 1989 BY

ACADEMICPRESS. 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 INFORMATION STORAGE AND RETRIEVAL SYSTEM, WITHOUT PERMISSION IN WRITING FROM THE PUBLISHER.

ACADEMIC PRESS, INC. San Diego, California 92101 United Kingdom Edition published by

ACADEMIC PRESS LIMITED 24-28 Oval Road, London NW I 7DX

LIBRARYOF CONGRESS

CATALOG CARD

ISBN 0-12-007836-8 (alk. paper)

PRINTED IN THE UNITED STATES OF AMERICA 89909192

9 8 7 6 5 4 3 2 1

NUMBER: 49-7755

Contents CONTRIBUTORS ............................................................... PREFACE ....................................................................

vii ix

Studies of Model Catalysts with Well-Defined Surfaces Combining Ultrahigh Vacuum Surface Characterization with Medium- and High-pressure Kinetics

.

CHARLES T CAMPBELL

I. I1 I11. IV

. .

Introduction ....................................................... Experimental Design Considerations................................... Elucidation of Structure-Function Relationships in Catalysis ............. Conclusion ......................................................... References .........................................................

2 4 14 47 49

The influence of Particle Size on the Catalytic Properties of Supported Metals

.

MICHELCHE AND CARROLL 0 BENNETT I. I1 I11.

.

. . VI . VII. IV V

. . .

I I1. I11 IV V. VI .

Introduction ....................................................... Preparation of Supported Metal Particles .............................. Background for Evaluating the Catalytic Properties of Small Supported Particles ......................................... Characterization of Small Supported Metal Particles .................... Effect of Particle Size on lhrnover Frequency and Selectivity: Presentation and Comparison of the Data ............................. Possible Explanations of Particle Size Effects: Experiments versus Models ........................................... Conclusions and Future Directions .................................... References .........................................................

Metal-Support interaction: Group Vill Metals and Reducible Oxides GARYL. HALLER AND DANIEL E. RESASCO Introduction ....................................................... Titania-Supported Catalysts .......................................... Model Studies ...................................................... A Comparison of Rh/TiOl and Pt/TiO, ...............................

55 59 72 91 109 140 155 160

173 179 203

209 Current Explanations of the Promoting Effect of TiO, on Catalytic Activity ......................................... 214 Bonding and Charge 'Ikansfer in Group VIII-Ti Systems ................. 221 V

vi

CONTENTS

VII. VIII

.

I.

. Ill . IV. V. VI. VII . VIII . IX. 11

Metal-Support Interactions in Other Oxide Supports and Related Phenomena ............................................. Conclusions ........................................................ References .........................................................

226 229 230

Structure and Reactivity of Perovskite-Type Oxides LUISG. TEJUCA.Jost LUISG. FIERRO. AND JUANM . D. TASC6N Introduction ....................................................... Structure .......................................................... Preparation ........................................................ Nonstoichiometry ................................................... Stability in a Reducing Atmosphere ................................... Adsorption Studies .................................................. Perovskites in Catalysis .............................................. Miscellaneous ...................................................... Prospective Lines of Research ........................................ References .........................................................

237 240 244 254 258 270 279 315 318 319

New Catalytic Materials from Amorphous Metal Alloys ARPADMOLNAR.GERARD v . SMITH.

. 111. I I1.

IV.

AND

MIHALYB A R T ~ K

Introduction ....................................................... Preparation and Characterization of Amorphous Alloys ................. Amorphous Alloys in Catalytic 'Ifansformations ........................ Conclusions ........................................................ References .........................................................

329 330 336 374 317

ADDENDUM TO STRUCTURE AND REACTIVITY OF PEROVSKITE-TYPE OXIDES ............ 385 INDEX...................................................................... 387

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

MrnALY BART~K, Department of Organic Chemistry,J b m f Attila University, Szeged, Hungary (329) CARROLL 0. BENNETT,Department of Chemical Engineering, University of Connecticut, Storm, Connecticut 06268 (55) CHARLES T. CAMPBELL ChemistryDepartment, Indiana University,Bloomington, Indiana 47405 (1) MICHELCHE, Laboratoire de Rkactivitk de Surface et Structure, UA 1106, CNRS, Universitk Pierre et Marie Curie (Ihris Vl). 75252 Aris Cedex 05,

Fmnce ( 5 5 ) JOSELUISG. FIERRO, Instituto de Catdlisisy €ktmleoqulmica. CS.I.C, Sermno 119, 28006 Madrid, Spain (237) GARYL. HALLER Department of Chemical Engineering, Yale University,Nav Haven, Connecticut 06520 (173) ARPADM O L N ~ Department , of Organic Chemistry,Jbzsef Attila University, Szeged, Hungary (329) DANIELE. RESASCO,Department of Chemical Engineering, Yale University, New Haven, Connecticut 06520 (173) GERARDV. SMITH,Department of Chemistry and the Molecular Science h g m m , Southern Illinois University, Carbondale, Illinois 62901 (329) JUANM. D.TAXON,Instituto Nacional del Carbdn y sus Brivados, C S I C , Apartado 73, 33080 Oviedo, Spain (237) LUISG. TEJUCA, Instituto de Cbtdlisisy &tmleoquulmica, CSIC, Sermno 119, 28006 Madrid, Spain (237)

vii

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We live in the midst of an information explosion. Publications are mushrooming. Even journals dedicated to any one “field” like catalysis are multiplying. Furthermore, when we look at catalysis, we deal with a diversity of basic ingredients and interests which can be equally at home in many disciplines (chemistry, solid state physics, materials science, enzymology, chemical engineering, electro-chemistry, crystallography, to name but a few). How do we keep informed when the proliferation of reporting is so massive and scattered over a wide spectrum of the scientific information media? The answer is, and always will be, “with difficulty.” But progress in science is dependent on two equal partners in symbiotic coexistence: (1) the generation of knowledge and (2) its successful transfer. The latter has been the legitimate driving force behind the tradition of publishing. But printing alone does not lead to a successful implant in the body of knowledge In fact, the more that gets offered in print, the more difficult it becomes for that body to choose the essentials for its growth. Mere quantity of reported data, theories, and concepts does not contribute to the progress of science. It takes a continuing process of sieving, identifying, and formulating the probable key ingredients. In good science, the key concepts, observations, and theories are those that produce the most consistency within the broader experience of the field, as well as within the basic sciences in general. That characterizes the environment, challenge, and task of the editors and authors of a series such as Advances in Catalysis. It is an objective which is not served by the daily stream of publications. Nor is it served by books with subject titles of interest to us, but which are collections of diverse and unconnected symposium papers. Instead, we seek to fill the gap between the papers that “report” and the textbooks that “teach.” We look for the competent sketch of the current state of knowledge that will help us grasp, teach, or use the current “total picture”: its key elements, their meaning and implications, and the associated challenges and remaining inadequacies. Writing our traditional day-to-day publications has a large element of “look what I have done now,” with comments on what it does to the field of knowledge. On the other hand, our authors must first focus on the field of knowledge, with appropriate comments on individual contributions of course Such a task is not easy. There is not enough space to discuss all contributions, and in their choices and analyses authors cannot be infallible predictors of the ultimate truth. If that were the case, the final chapter would have been written, not in a series of Advances, but of Arrivals. We are fortunate to have authors who have dedicated their efforts to ix

X

PREFACE

recording progress. This volume of Advances could well be considered a textbook on the current state of knowledge concerning structure and catalysis of metals and metal oxide particles, old and new. Campbell addresses the basic and broad problems of what the catalytically relevant surface structures of metals are, where we stand in techniques capable of attacking this problem, and what the current state of knowledge is. The focus is on a long-standing, important, and central problem of general investigative methodology and strategy: the “pressure gap” is created by the fact that the best techniques of surface analysis require high-vacuum conditions, while “useful” catalysis is confined to conditions of near ambient or higher pressures. Che and Bennett review the basic question of the influence of particle size on catalytic behavior of metal particles which involves questions of the basic sciences as much as practical considerations of catalyst design and use. There are so many aspects involved that this chapter is a unique resource of information on preparatory techniques, analytical technology, and methods of characterization of these materials, as well as a discussion of the many phenomena, mechanisms, and catalytic effects that deal with the question of the dependence on particle size. Haller and Resasco provide an in-depth examination of another basic phenomenon, the existence and nature of interaction between metals and their oxide supports. Their approach focuses on an exhaustive study of titania as a support. That consideration then becomes a useful and central model from which to view the behavior of other supports. Tejuca, Fierro, and Tascbn review the perovskite class of metal oxides. Many challenging concepts have been associated with these. They called early attention to the relationship of catalytic activity to electronic properties; they have been seen as potential substitutes of platinum; and more recently the discovery of superconductivity has rekindled the interest in the “electronic connection” of oxide catalysts. Smith, Molnir, and Bartbk introduce us to amorphous metal alloys, recent newcomers to the scene of materials science, and catalysis. They offer new variants of metal structures to research and application. The authors point to work (some 90 publications) of only about the last 7 years, to remaining questions, and to potentials barely touched. Reading these topics in the same volume, one discovers existing and potential relationships between them; one experiences some sparks of recognition of further relationships and of matters one might connect to develop new knowledge or innovations. We hope that this is an additional attribute of the role and manner of reporting in Advances.

PAULB. WEISZ

ADVANCES IN CATALYSIS, VOLUME 36

Studies of Model Catalysts with Well-Defined Surfaces Combining Ultrahigh Vacuum Surface Characterization with Medium- and High-pressure Kinetics CHARLES T. CAMPBELL Chemistry Department Indianu University Bloomington, lndiunu 47405

The development over the past two decades of many powerful techniques for surface characterization promised to open the area of heterogeneous catalysis for a more fundamental understanding of the relationships between atomic-level surface structure and catalytic activity. The realization of this promise has been hampered by a “pressure-gap’’ problem: the reactions are generally performed at pressures above - 100 torr, whereas surface analysis is accomplishable only in ultrahigh vacuum. To overcome this problem, many researchers have recently developed equipment that combines ultrahigh vacuum surface analytical instrumentation (AES, XPS, LEED, TDS, ISS, etc.) with a medium- or high-pressure microreactor in a single apparatus, where the catalyst can be transferred between these pressure regimes without intermediate exposure to air or other gases. A “model catalyst” approach has generally been adopted, whereby well-defined surfaces are prepared (usually starting from a single crystal) that are very homogeneous in surface structure and chemical composition. Through comparison of the activity and selectivity of a series of such model catalysts with controlled changes in detailed surface structure, a fundamental understanding of structure-function relationships in the surface chemistry of catalysis is beginning to evolve. The purposes of this paper are to (1) review the apparatus design considerations for experiments of this type, (2) review the results that have been obtained thus far using this model catalyst approach in combined presI Copyright 0 1989 by Academic Press. Inc. All rights of reproduction in any form reserved.

2

CHARLES T. CAMPBELL

sure-vacuum experiments, and (3) discuss the future of this area of research. The results so far have given exciting but not yet completely satisfying insights into many catalytic phenomena including structural sensitivity, surface modification, poisons, alkali promoters, ensemble and electronic effects at bimetallic surfaces, and strong-metal support interactions. I. Introduction

It has long been a goal of the catalytic chemist to fundamentally understand the relationship between the atomic level structure of a catalyst surface and the materials’ performance as a catalyst (Lea,its activity and selectivity). Much beautiful work has been dedicated to this effort, especially since development of modern surface analytical techniques such as the many electron spectroscopies. A major stumbling block to progress in this direction has been the fact that, while our most powerful surface analytical techniques operate only under the pristine conditions of ultrahigh vacuum (UHV), most catalytic reactions of commercial importance occur with measurable rates only at pressures of a few torr and above. Thus, in order to inspect the surface of the working catalyst, one is faced with the problem of decreasing the pressure above the catalyst by some ten orders of magnitude while at the same time maintaining its surface chemical integrity. This is, of course, an impossible task since the weakly adsorbed species (i.e., those with short residence times) are in equilibrium with the gas phase and removed almost as rapidly. Nevertheless, much effort has been devoted to developing techniques that allow for UHV inspection of catalyst surfaces that retain, as well as possible, their high-pressure structural characteristics and thereby “bridge” this “pressure gap.” In 1974 the Somojai group published a pioneering study that combined catalytic kinetic measurements at pressures up to 1 atm with pre- and postreaction surface analysis in UHV in the same apparatus (I). The group soon developed new systems that extended the kinetic regime up to -100 atm. (2, 3). A few years later, the group of Yates et al. had developed and proved a system that allowed a catalyst sample to be transferred relatively rapidly (a few minutes) between a vessel at high-pressure reaction conditions and a UHV chamber for surface analysis (4, 5 ) . At about the same time, the groups of Bonze1 et al. (6, 7) and Polizzotti and Schwarz (8)developed similar systems, each with its own relative advantages. Since then, many applications of these and related systems have been published (see below), and numerous improvements to the original

PRESSURE-VACUUM ANALYSIS OF MODEL CATALYSTS

3

designs have appeared (9-12). It is safe to say that this combination of high-pressure reaction kinetics or sample treatment with UHV surface analyses is now a more or less routine method of the surface scientist’s trade, and it is finding its way into many application areas such as corrosion science (1.3-15) and electrochemistry (14-18). A most enlightening utilization of these high-pressure-UHV methods involves what is referred to as the “model catalyst” approach, whereby a surface derived from a clean and well-defined, single-crystal plane is used to model a particular site or set of sites expected to exist on practical highsurface-area catalysts. Commercial catalysts are exceedingly complex materials, the surface analysis of which would undoubtedly lead to clear identification of eight or more different elements and countless structurally distinct surface moieties. The catalytic surface chemist wants to know which of these elements and surface structures are responsible for catalyst activity and selectivity, and how they participate in the catalytic reaction. Therefore, model surface structures of one or several elements are synthesized, usually starting from a polished and well-oriented singlecrystal surface. In doing this, the chemist takes advantage of “preparatory surface chemistry in UHV.” Thus, a model catalyst is created, the surface of which is now extremely homogeneous in a certain structural moiety of interest. This catalyst can then be transferred to the highpressure reaction conditions for kinetic characterization, where the activity and selectivity of the model catalyst are measured. After reaction, the chemist reexamines the surface in UHV to ensure that the structural details that were tuned into the starting material have not changed as a consequence of the reaction, or, if so, to determine exactly what is the structure of the working catalyst surface that has been kinetically characterized. Because the model catalyst is based on a single-crystal surface, the whole arsenal of surface analytical techniques can be brought to bear on these problems, including those techniques that rely on long-range surface order or smoothness. The purpose of the present article is to review the existing literature with respect to model catalyst studies in which attempts have been made to obtain and compare high- or medium-pressure kinetics over well-characterized, homogeneous surfaces. (For the purposes of this article I will define the medium-to-high pressure range as -0.1 torr to -200 atm.) I start with a thorough discussion of the design criteria that should be considered in constructing an apparatus and preparing samples for such studies. I then summarize the results of such studies and show that structure-function relationships in the surface chemistry of catalysis can, indeed, be determined and understood by using this approach. Some attempt has been made here to present a relatively exhaustive bibliography

4

CHARLES T. CAMPBELL

of these types of studies, at least of those using model catalysts based on single-crystal surfaces. That is not to say that every article has been included here, and I apologize for any inadvertent omissions. I also should point out that in several cases I have omitted papers for brevity that are reviewed thoroughly in those articles that I do discuss, so that they should be relatively easy for the reader to locate. In this presentation I also attempt to discuss the articles within the more global perspective of this review format, in some cases pointing out contradictions in results or potential problems in interpretation. For most cases, only a few of the major features of a given article are presented. There is a vast amount of information in these articles that I could not cover in the limited space of this review. The reader is encouraged to seek out these articles, for they are rich in results and insightful in discussion. Finally, I conclude by attempting to critically evaluate the field in general, its impact thus far on catalytic concepts, and its potential for fundamental contributions to our understanding in the future. II. Experimental Design Considerations In general, the types of studies to be discussed here utilize a single apparatus that interfaces a microreactor for high-pressure kinetic studies with an ultrahigh vacuum (UHV)chamber for surface analysis and adsorption-desorption studies. In a typical experiment, a model catalyst is prepared by orienting, cutting, and polishing a single-crystalsurface. This sample is then mounted in the UHV chamber and the desired surface structure is prepared by using what can be called “preparatory surface chemistry.” This can include such steps as sputter cleaning with an argon ion beam, high-temperature chemical cleaning, annealing to establish surface order, vapor deposition or chemical vapor deposition of other elements (e.g., oxidation) to modify the surface composition, ion-beam surface modification, and molecular-beam epitaxy . The chemical treatments are sometimes even performed at higher pressures in the microreactor or in an atmospheric load-lock position (e.g., liquid droplets of ultrapure solutions). The surface is then characterized with an array of UHV surface analytical techniques. Typical methods of choice include Auger electron spectroscopy (AES), X-ray photoelectron spectroscopy (XPS), ionscattering spectroscopy (ISS), secondary ion mass spectroscopy (SIMS), high-resolution electron energy loss spectroscopy (EELS), and many others. Since the samples usually have smooth single-crystal surfaces, methods such as low-energy electron diffraction (LEED) and other surface crystallographic methods that rely on long-range surface order can


5

also be used here. The surface analytical methods themselves will not be discussed here since they have been excellently reviewed elsewhere (1922). Following surface characterization, the sample is enclosed in the microreactor; it is pressurized with ultrapure reactant gases, and a catalytic reaction is initiated by heating the sample to the desired reaction temperature. The rate of the reaction is determined from product buildup versus time by monitoring the composition of the gaseous reaction mixture using gas chromatography, mass spectroscopy, or other methods. After reaction, the sample is again evacuated and the surface is structurally and compositionally analyzed in UHV to determine the characteristics of the working catalyst’s surface. Extreme caution must, of course, be exercised here to ensure that the surface structures observed in postreaction surface analysis are indeed those characteristic of the working catalyst under high-pressure conditions. As noted above, this is, strictly speaking, an impossible task since weakly adsorbed species with short surface residence times (67%) from adsorbed cyclohexane if the surface is kept free of contaminants and other adsorbed species that block Pt sites necessary for hydrogen abstraction (60). The mechanism of benzene production from cyclohexane on Pt(ll1) is now fairly well understood, as is the ensemble requirement for this reaction (60). Other mechanisms may be important on Pt catalysts containing step and/or kink sites, or support surfaces. It is interesting to point out that significant inverse kinetic isotope effects (RDIRH= 1.3-3.3) for the initial rates of a number of alkane conver-


19

sion reactions have been observed by Davis et al. ( 6 0 , where the unlabeled hydrocarbon was reacted in the presence of H2 or D2 gas over a variety of Pt surfaces. They attributed this to the fact that the coverage of atomically absorbed deuterium should be lower than that of hydrogen under identical reaction conditions, and that the deuterium would block fewer Pt sites and therefore inhibit the reaction less severely than hydrogen. In the case of Dz gas, the products were extensively deuterated, which suggests that the adsorbed intermediates were also largely deuterated. Given the recent observations of strong normal isotope effects (RHI RD = 4-15) in C-H bond cleavage steps on Pt surfaces (60, 62), the inverse kinetic isotope effect in the overall catalytic reaction might instead be due to a lower extent of dehydrogenation in the adsorbed hydrocarbon layer when deuterated. Since severely dehydrogenated hydrocarbons are known to irreversibly poison Pt surfaces (52, 6.9, deuterated hydrocarbons should show a lower degree of poisoning of surface sites by carbonaceous residue. Since the overall rate of reactant conversion is positive-order with respect to H2 pressure (61),the original explanation based on inhibition of the reactions by adsorbed hydrogen (or deuterium) seems less satisfactory than a model where the main species poisoning the reaction is partially dehydrogenated hydrocarbon residue. Very high coverages of such residue are generally observed in surface analysis following hydrocarbon reactions over Pt (63). The reaction rates for hydrocarbon conversion on Pt generally decrease with reaction time, and this has been correlated with the deposition of irreversibly adsorbed carbonaceous species (55, 63). The reactions of hydrogen and cyclohexene over Pt single crystals have been studied by Davis and Somorjai (64)over ten orders of magnitude in pressure. Although benzene was the major product in vacuum conditions, when the reactions were carried out near atmospheric pressure, mainly cyclohexane was obtained. This higher pressure reaction was structurally insensitive, which was attributed to the presence of a near monolayer of carbonaceous species. Massardier et al. (65) studied the competition between hydrogenation of benzene versus toluene over several Pt crystal surfaces and compared the results to those of high-area Pt catalysts. The hydrogenation of ethylene to ethane has also been studied over Pt single crystals by the Somorjai group (66).This reaction was found to be structurally insensitive as well, a fact that again was attributed to the presence of a monolayer of carbonaceous fragments under the high-pressure reaction conditions. In this case, there is strong evidence that most of these fragments are adsorbed ethylidyne species (=C-CHd. On the basis of measurements made in the presence of D2 gas, the authors pro-

20

CHARLES T. CAMPBELL

posed a reaction mechanism whereby the ethylidyne species serve as a deuterium transfer agent between the Pt surface (which cleaves the D-D bond) and ethylene molecules that are adsorbed uboue the ethylidyne adlayer. Radiotracer studies with I4C showed that the ethylidyne species are readily hydrogenated and removed from the surface at atmospheric hydrogen pressure and temperatures above 350 K, but this reaction occurs much more slowly than does ethylene hydrogenation under similar conditions (66). Similar studies (67,68) of the hydrogenation of ethylidyne on Rh( 111) and H / D exchange over well-defined ethylidyne adlayers suggest that the reaction on rhodium surfaces is very similar. Very high coverages of strongly adsorbed carbonaceous deposits are a ubiquitous occurrence in high-pressure hydrocarbon conversion reactions over Pt. The uncovered-site concentration (measured by postreaction CO adsorption-TDS) has been correlated with the total surface carbon coverage (as measured by AES) (6.3, which, in turn, has been correlated with the catalytic activity and selectivity for several reactions (55, 63). The major roles of the deposit are as a nonselective site-blocking agent and as a pool for the storage and rapid exchange of surface hydrogen (63). The ease or reversibility with which these deposits can be removed from the surface decreases with increasing temperature of their deposition or treatment and increases with their H/C atomic ratio (63). Given the constant presence of these carbonaceous deposits, it is somewhat surprising that any structural sensitivity is ever seen in hydrocarbon catalysis over Pt. The fact that some is seen is probably related to the fact that the underlying Pt structure controls the structure of this adlayer and the concentration of the few free sites found within this adlayer, which, in turn, can control the catalytic reaction rates. It is useful in this respect to point out that reaction probabilities (per collision of hydrocarbon molecule with the surface) are many orders of magnitude lower at high-pressure reaction conditions than in UHV where the surface is partially clean (64).Therefore, only a tiny (immeasurably small) fraction of free Pt sites are necessary to explain the overall observed rates of catalysis at high pressures. Oudar et ul. (69) studied the hydrogenation of butadiene over Pt( I lo), paying special attention to the mechanism of carbon deposition that deactivates the surface. At high hydrogen pressures the rate was first-order in H2 and zero-order in butadiene, but at low HZpressures the rate became second-order in H2. For high H2, they argued that the surface was covered with one butadiene molecule per two surface Pt atoms and that the sites for butadiene adsorption are different from those for H2 dissociation. At low H 2 , the sites for H2 dissociation were said to be poisoned by an additional hydrocarbon species. Above a critical temperature, irreversible


21

deactivation by carbon occurs. This latter was attributed to a simple siteblocking effect, since the rates of both butadiene hydrogenation and H2/ D2exchange were proportional to the remaining carbon-free area, and the activation energy was unchanged.

2. Hydrocarbon Conversion over Other Metals Goodman has compared the hydrogenolysis of ethane, cyclopropane, and butane over Ni( 100) and Ni( 1 11) (70-73). At similar pressure conditions (H2 = 100 torr, C2H6 = 1 torr) the Ni(100) surface is considerably more active and has a lower overall activation energy then Ni(ll1) for methane production from ethane (25 versus 46 kcal/mol) (72).The steadystate kinetics for Ni( 1 11) compare favorably with those for SiO2-supported Ni catalysts, which suggests that Ni(ll1) facets dominate on the supported catalyst surfaces. Indeed, this is the most thermodynamically stable facet (in vacuum, at least). The hydrogenolysis of cyclopropane to methane, propane, and ethane also proceeds much faster on Ni(100) than on Ni(ll1) (73),as can be seen in Fig. 6 by the much lower temperatures required for comparable rates on Ni( 100). The hydrogenolysis of butane to methane occurs far more rapidly over Ni(100) than over Ni(l1 l), although in this case the activation energies may be very similar (70, 71).At low temperatures (400K)large amounts of propane and ethane are also produced. Interestingly, the absolute rate and apparent activation energy for hydrogenolysis of butane to methane over Ni( 100) was very similar to that of ethane hydrogenolysis and CO methanation (70).This indicated a common rate-determining step for all three methane production reactions. This step was thought to be removal of surface carbide (or a partially dehydrogenated adsorbed methane precursor) by hydrogenation with H2 (70). Also consistent with this picture is the fact that alkane dissociation proceeds on clean Ni(100) with rates that are one to two orders of magnitude faster than the steady-state rate of their hydrogenolysis (74).These catalytic reactions proceed with steady-state rates on Ni surfaces with the buildup of only submonolayer quantities of carbonaceous residue (70, 72) and were not slowly poisoned by carbon buildup as frequently occurs for Pt (see above). Engstrom et al. (75-78) have compared the reactions of several small hydrocarbons with H2 over different single-crystal surfaces of iridium. For ethane, propane, n-butane, and neopentane, the dominant reaction pathway on both Ir( 111) and Ir( 1 10)-(1 x 2) was cleavage of a single C-C bond, except for the case of n-butane or Ir(ll1) (77). The major reaction channels involved demethylization of the parent hydrocarbon, except for butane on Ir( 1lo)-( 1 X 2), which made mostly ethane (75).This

22

CHARLES T . CAMPBELL 1.o

-

A

I

u)

i k

d 6F

10’’

8 LL

10-1

1 I

I

I

I

I

I

I

I

I

I

1.7

1.8

1.9

2.0

2.1

2.2

2.9

2.4

2.5

TEMPERATURE (I/K

x lo3

FIG.6. Product formation from the reaction of cyclopropane with hydrogen over a Ni(l1l) surface (open symbols) and a Ni(100) surface (filled symbols) (total pressure = 100 tom; H2/cyclopropane = 100). (0, O), Methane; (A, A), propane; ( 0 ,H), ethane. From Ref. 73.

production of ethane was attributed to the stabilization of a special intermediate, suggested to be a metallacycle pentane, by the low-coordination C, sites present on Ir( 110) but not found on Ir(ll1). This structural sensitivity in ethane selectivity correlates well with results for supported Ir catalysts, where the selectivity increases markedly with decreasing average size in the supported Ir particles (75). At low temperatures, ethane hydrogenolysis proceeds much more rapidly on Ir( 1 1 1) than on Ir( 110)( 1 x 2), but both propane and neopentane gave comparable rates on these surfaces (77). The activation energies vary depending on the temperature and the surface structure, so distinct activity comparisons do not hold over the whole range of reaction conditions. A full analysis of the kinetic


23

dependencies on temperature and reactant partial pressures has been performed, and the mechanistic implications have been discussed (77). The reaction of cyclopropane and hydrogen over Ir gave both hydrogenation to propane and hydrogenolysis to methane and ethane, with the former dominating below 500 K (78). For both products, the Ir(ll0)(1 x 2) surface was more active by a factor of -3.5 than Ir(ll1) (78). Again, this difference was attributed to the presence of C7sites on Ir(l10). The hydrogenation of methylcyclopropane on both surfaces was dominated by production of n-butane, although a pathway for isobutate production was also competitive at low temperatures and had a much lower activation energy. The Ir( 1 lo)-( 1 x 2) surface has a greater activity for hydrogenation of propylene, and generally superior hydrogenation capability, than does Ir( 1 11). This finding was correlated with the presence of higher adsorption energy states for Ha on the former surface, which suggests a higher hydrogen coverage under reaction conditions (78). As with Ni, steady-state reaction rates for these hydrocarbon reactions were achieved over Ir single crystals with the buildup of only submonolayer quantities of carbon (& < 0.5) (77, 78). Wax et al. have studied the hydrogenolysis of ethane over W(100) (79). This surface is itself not active in this reaction until a monolayer of surface carbide is formed, after which an active and very stable model catalyst results. In general, the activity of cleaned metal single crystals for this reaction increased in the order W < Ni < Ru < Ir at the same conditions (573 K, 1 torr ethane, 100 torr H2)(72, 75, 79, 80). The cyclotrimerization of acetylene to benzene has been studied by Rucker et al. (81) over Pd( 11I), (loo), and (1 10) at pressure near 1 atm. The (1 10) surface was four-fold less active than Pd(ll1) or (loo), which contrasts with their relative selectivities during TDS under UHV conditions. The activity at high pressures was correlated with the fraction of the various surfaces that exposed clean Pd atoms, as probed by postreaction CO adsorption-desorption. In all cases, most of the surface was covered with a carbonaceous residue. The authors stated that the reaction rate is first-order in acetylene pressure for all three surfaces. The extensive data on the Pd(ll1) surface clearly indicate first-order kinetics in that case. However, the limited data presented for Pd(ll0) seem (to the present author) to be better fitted by an order of -2.5, which is closer to the value of three suggested by the overall stoichiometry. The structure of carbonaceous deposits formed from medium pressure exposure of ethylene between 325 and 800 K have been compared for initially clean surfaces of Ir, Pt, and Rh by Niemantsverdriet and van Langeveld (82) using SIMS and AES. On all surfaces, the hydrogen con-

24

CHARLES T. CAMPBELL

tent (H :C ratio) decreased with increasing reaction temperature. The hydrogen content also decreased in the order Ir > Pt > Rh. The deposits were, however, most graphitic in nature on Pt surfaces. The authors point out that the tendency of Pt to graphitize carbon may explain the relatively low hydrogenolysis activity of Pt. 3 . Ammonia Synthesis (N2

+ 3H2 --* 2NH3)

The synthesis of ammonia over both Fe and Re single-crystal surfaces has been studied by the Somorjai group and found to be extremely structural sensitive (83-85). The activity of iron increases in the order (1 10) 4 (210) = (100) 4 (221) = (1 1 1) (83,84).Because dissociative N2adsorption is thought to be the rate-determining step, this order has been correlated (83) with increasing relative activity and decreasing activation energy for dissociative N2 adsorption, as measured by the Ertl group (86, 87, 136). They attributed the higher activity of the more open crystal planes to the more aggressive behavior of low-coordinated Fe atoms. However, Strongin et al. (84) attribute the high activity of very open Fe(ll1) and (211) surfaces to their exposure of highly coordinated C7 sites at the surface for chemisorptive bond formation. As shown in Fig. 7, Re surfaces showed amazing structural sensitivity in this reaction, with relative activitiesof(0001):(10~0):(11~0):(11~1) = 1 :94:920:2820(85). That is, the specific activity increases dramatically as the atomic roughness or openness of the surfaces increases. Again, this catalytic activity correlates well with the ability of the surface for dissociative N2 chemisorption, which is thought to be the rate-determining step (88). Strongin et al. (84) argue that the surface roughness is important only to the extent that it can again expose highly coordinated surface atoms (C,, and C ~inOFig. 7) that are present in the second and third layers. They support this argument with theoretical calculations by Falicov and Somorjai (89) (which do not consider steric hindrance). 4.

Carbon Monoxide Oxidation (CO + b Reduction (NO + CO + C02 + #N2)

0 2 4

C02) and Nitric Oxide

The oxidation of CO by either O2or NO was studied by Peden et al. and Oh et al. over Rh, Pd, Pt,and Ir single crystals (90-92). The CO + 0 2 reaction was relatively insensitive to the atomic structure of the surface, and the specific activities and kinetic parameters agreed for both crystal surfaces and for alumina-supported catalysts. The Rh surfaces deactivated at high 0 2 pressures due to the formation of a near-surface oxide (91, 92). On the other hand, the CO + NO reaction was very sensitive to


TF

C6(SOLID)

I4

25

C, (SOLID)

C,o(DOT TED)

f 9

I

I

FIG.7. Structure sensitivity in the ammonia synthesis over rhenium single-crystal faces. The turnover frequencies (TF) are given as NH, molecules/cm2/s (PtOtd= 20 atm, H2 : N2 = 3 : 1, TcWstd = 870 K).Schematics of the atomic structure of each surface are given above each bar. From Ref. 85.

surface geometry (90).A mechanism and kinetic model for these reactions has been successfully developed that quantitatively fits the medium-pressure kinetic data, but that was parameterized using the results of UHV adsorption/desorption kinetic studies (92). This contrasts with an earlier kinetic study of the NO + CO reaction over Rh(100) by Hendershot and Hansen ( 9 3 , who modeled their data with a somewhat different mechanism that included adjustable kinetic parameters. It is difficult to directly compare their actual data with that of Peden et al. (90) because of differences in the reactant pressures and temperatures used.

26

CHARLES T. CAMPBELL

The CO + 0 2 reaction also appears to be structurally insensitive over Ru, since kinetic data for Ru(0001) and silica-supported Ru are fairly similar (94). The highest activity occurs when the Ru is covered with a monolayer of oxygen. Since the reaction is, instead, inhibited by an oxygen-covered surface for the case of Rh, the higher activity of Ru compared to Rh at medium-pressure conditions could be explained (94). The oscillatory kinetics of CO oxidation over Pt single crystals has been studied at atmospheric pressure by Yeates et al. (95). They present a model to explain rate oscillations that relies on the oscillatory formation of surface platinum oxide, which was observed in postreaction analysis and was related to the presence of silicon impurity.

5 . Hydrogenation of CO The methanation reaction (3H2 + CO --$ CH4 + H2O) has been thoroughly studied by Goodman and co-workers (4,5, 71,96) over Ni single crystals. Since the specific rates, activation energies, and pressure dependencies are very similar over Ni( loo), Ni( 1 1 l), and A1203-supportedNi, the reaction is structure insensitive (71,96). Transient kinetic studies at medium pressures combined with postreaction AES analysis on Ni( 100) have identified a carbidic form of adsorbed carbon as the reaction intermediate, and graphitic carbon as a poison formed at higher temperatures (71,96). The hydrogenation of CO was studied by Castner et al. over Rh( 1 1 1 ) and polycrystalline Rh foils at 6-atm pressure (97). These model catalysts produced mainly (90%) methane and showed similar specific rates, again suggesting a lack of structural sensitivity. When the Rh surface was oxidized, the initial rates were much larger, the activation energy changed, and the selectivity increased for heavier hydrocarbons and oxygenated products such as methanol, ethanol, and acetaldehyde. Logan et al. compared C O hydrogenation over Mo(100) and polycrystalline Mo foil, which gave similar rates and, therefore, apparent structural insensitivity (98). Here, the major product was methane, but ethene and propene were also significant products. The hydrogenation of CO was observed by Udovic et al. (99) and Szuromi et al. (ZOO) to proceed readily over W(llo), with methane as the dominant product at steady state. The surface was covered with a monolayer of carbidic carbon under reaction conditions, as well as some surface oxygen. The activation energy for methanation over W( 110) was only about half that seen for Ni surfaces. The synthesis of methanol from CO and H2 has been studied over Pd( 110) at medium pressures (500-1800 torr) by Berlowitz and Goodman


27

(101). The (extrapolated) specific rates and activation energy agree well with results for Pd dispersed on noninteracting supports such as Si02and some basic supports such as ZnO and MgO. Palladium on highly interactive and acidic supports such as zeolite, alumina, or lanthana give rates and selectivities (i.e., significant dimethyl ether production) substantially different from those of Pd(l10). The synthesis of methanol from CO(+CO2) and H2 was also attempted by Campbell and co-workers (102) over a Cu(ll1) surface and over welldefined Cu overlayers on a ZnO(OO0i) single-crystal surface. They could set an upper limit of < 2 x 1OI2 molecules/cm2/son the methanol production rate at temperatures up to 600 K and pressures up to 1500 torr. These limits are consistent with expectations based on (extrapolated) kinetics for high-surface-area Cu/ZnO catalysts. Their results also showed that C02 has extremely low reactivity for oxidation of the clean Cu(ll1) surface (102), and they argue that it is very unlikely that significant oxygen concentrations exist on the Cu(1 1 1)-like surfaces of practical catalysts under methanol synthesis conditions (102, 103).

6 . Other Reactions Thiophene HDS Activities have been reported by Bussell et al. over three low Miller index single-crystal surfaces of Mo and four of Re (104106). The reaction over Mo was insensitive to the surface structure. The Re(0001) surface showed about the same activity as the Mo surfaces, but Re( 1121) was twice as active, and the (1 120) and (lOi0) surfaces were approximately sixfold more active (106). While Mo surfaces were covered with a near monolayer of partially hydrogenated carbon after reaction, the Re was not moderated by a carbon overlayer. The product distribution over Mo(100) was similar to that reported for powder MoS2 catalysts, although the single-crystal surface of pure Mo was much more active (104). Measurements of the rate of hydrogenation of 35Son Mo( 100) suggest that sulfur adatoms are not intermediates in thiophene HDS (104). The water-gas shift reaction (CO + H20 + C02 + H2) has been studied in detail over a Cu( 111) surface at pressures up to 15 torr CO and 200 torr HzO by Campbell and co-workers (26). The specific activity, activation energy, and reaction orders are very similar to those extrapolated from kinetics at somewhat higher pressures over Cu/ZnO and Cu/ZnO/A1203 catalysts. Similarly, doping of the Cu(ll1) surface with ZnO, had no distinct effect on the observed activity. These results suggested that there is no specific Cu-ZnO interaction necessary for an active water-gas shift catalyst, the essential ingredient of which is metallic Cu suiface. Kinetic analysis using kinetic parameters obtained from both UHV and medium-

28

CHARLES T . CAMPBELL

pressure adsorption-desorption experiments indicated that the rate-determining step was the dissociative adsorption of water. Since there is significant structural sensitivity for this step (lo?'), the catalytic activity may depend on crystal orientation. A word of caution should be remembered concerning the general concept of structural sensitivity for catalytic reactions, since this can depend in some cases on the reaction conditions chosen. This is because, as the temperature or reactant pressures are changed, the rate-determining step can change from one that is structurally sensitive to one that is not. Similarly, optimum practical catalytic conditions are often not those where a single elementary step is rate-determining, but where a delicate balance of kinetic competition exists between steps in the mechanism. Thus, global interpretation of the above results should be used only with caution and proper qualification.

B. THEROLEOF CATALYST ADDITIVES: SURFACE MODIFICATION Another major contribution of the pressure-transfer-model-catalyst studies has been to clarify the role of catalyst additives and impurities in controlling catalytic activity and selectivity. Particularly important here has been the controlled study of the influence of surface chemical modification on the medium- or high-pressure reaction kinetics. Typically, the kinetics are followed as a function of the surface coverage of an added element or complex, and the changes in global kinetics are then correlated with the influence of the additive on the rates of individual steps in the reaction mechanism. This type of data has been found to be extremely useful in identifying the rate-determining step and in clarifying the role of surface modifiers such as poisons or promoters, or the presence of a second metal in bimetallic or alloy catalysts. (The subject of bimetallic catalysts will be discussed in a following section, as will metal-support interactions.) 1.

Electronegative and Electroneutral Elements as Additives

With a few exceptions, electronegative elements have been observed to inhibit the reaction rates of catalyzed reactions when added to well-defined transition-metal surfaces. The mechanism of this poisoning, however, can vary from a simple, steric site-blocking effect to a longer range, electronic interaction acting either through the orbitals of the metal surface or through space. Sometimes selective poisoning of one competing reaction in a branching mechanism can actually lead to a desirable increase in selectivity.


SULFUR ON "100) AT 0.25 MONOLAYERS

1

29

SULFUR ON Ni(100 AT 0.5 MONOLAYERS

\

P

10

!i

E 10'

0

0.1

0.2

0.3

0.4

i

ADATOM COVERAGE IN MONOLAYERS

FIG.8. Rate of the methanation reaction (CO + 3H2 -+ CH4 + HzO) as a function of sulfur (0) and phosphorus (0)coverage on a Ni(100) surface (pressure = 120 tom, H2: CO = 4 : 1 , reaction temperature = 600 K). From Ref. 71. Reprinted with permission from Acc. Chem. Res. 17, 194. Copyright 1984 American Chemical Society.

The influence of sulfur or phosphorous addition on the catalytic activity of a Ni(100) surface in CO methanation has been thoroughly studied by Goodman and co-workers (71, 96, 108). As shown in Fig. 8, the rate decays very nonlinearly with the coverage of S or P. The initial slope of the curve for sulfur has been interpreted to suggest that 210 Ni surface sites are deactivated by each sulfur adatom (108).Phosphorous, which is slightly larger but considerably less electronegative than sulfur, poisons the rate much less rapidly than sulfur, so the authors concluded that longrange electronic effects were playing the dominant role in catalyst poisoning here. According to their model, a phosphorous adatom sits in a fourfold hollow site and poisons only its four nearest neighbors, while the more electronegative sulfur adsorption leads to deactivation also in the second shell of neighboring Ni atoms. This model relies on LEED studies which suggests that both sulfur and phosphorous spread uniformly over

30

CHARLES T. CAMPBELL

the surface in the low-coverage range (80,109). Interestingly, sulfur addition did not alter the activation energy below 600 K (fog),which suggests that a given Ni site is either completely poisoned or not affected at all, depending on its distance from a sulfur atom. Very similar effects of sulfur were also reported on Ni/A120~catalysts (108).Nakamura et al. (109a) have shown that the rate of H2-D2 scrambling at 12 torr over Ni(100) decays linearly with sulfur coverage to zero at 0s = 0.45. Similar effects of sulfur have been observed by Goodman and coworkers for CO hydrogenation over Ru(0001) and Rh( 111) model catalysts (80) and C02 hydrogenation over Ni(100) (110).In the case of CO2 hydrogenation, the reverse water-gas shift reaction (C02 + H2 CO + H20) apparently establishes a rapid equilibrium, and then methane is produced from the resulting CO (110).The hydrogenation of CO over sulfur-dosed Mo(100) was studied by Logan et al. (981,who found that whereas methane production was poisoned, the effect was much less dramatic than on the surfaces outlined above. In addition, the amount of ethylene product became comparable to that of methane upon sulfur poisoning. The effect of adsorbed sulfur upon the rate of CO methanation over W(110) has been studied by Szuromi et al. (100, 111). The rate decayed much more slowly with sulfur than on Ni surfaces, and the activation energy was unchanged. At high coverages, the rate and the saturation hydrogen coverage decayed in a similar way. These authors also showed that low levels of 0 2 torr) completely poison the surface with oxygen. The rate of hydrogenolysis of ethane to methane over Ru(0001) is much more slowly poisoned by sulfur than is methanation (80).The hydrogenolysis of cyclopropane to methane, ethane, and propane has been studied by Goodman over sulfur-dosed Ni( 111) and Ni( 100) (7.3). The production of methane and ethane on Ni( 11 I ) and methane on Ni( 100) is dramatically poisoned by low sulfur coverages. The yield of propane plus propylene decreases much more slowly. In contrast to the clean Ni(lI1) surface, ethylene production increases, becoming a significant product, after sulfur addition. Overall, the addition of sulfur causes a reduction in the ability of the Ni surface to cleave C-C bonds. Oudar et al. (112) have studied the influence of sulfur on the hydrogenation of 1,3-butadiene and H2-Dz equilibration over Pt(ll0). The rates decayed linearly with sulfur coverage, so that each sulfur atom poisons one dissociation site for hydrogen without influencing the activation energies or mechanism. The authors established the first isotherm for sulfur adsorption under actual catalytic reaction conditions. The adsorbed hydrocarbons influenced the equilibrium coverage of sulfur on the Pt surface. The thermodynamics of adsorbed sulfur on several metal single crystal surfaces have been presented by Bernard et al. (114).

*


31

The general subject of sulfur poisoning of metal catalysts has been reviewed by Bartholomew et al. (115), who discussed results mostly on high-surface-area catalysts. The hydrodesulfurization of thiophane over Mo( 100) is deactivated by sulfur adsorption, although the rate for nearly a monolayer of sulfur is still almost half that measured in the absence of sulfur (104). Excessive sulfur exposure leads to formation of a MoS2 layer and severe deactivation of Mo(100)(104). Campbell and Koel have studied the deposition kinetics of sulfur from H2S onto Cu(ll1) and the resulting influence of sulfur adatoms on the kinetics of the water-gas shift reaction (H20 + CO + H2 + COZ)(24). Although H2S deposits sulfur adatoms on the surface and liberates H2 gas (H2S+ Hz + S,) with a reaction probability of -0.03 per H2S collision with the surface at 500 K, the reverse reaction proceeds only very slowly, with a reaction probability of 5 per H2collision. This kinetic discrepancy suggests that a few ppm of H2S will be sufficient to deposit almost a full monolayer of sulfur onto Cu surfaces under catalytic reaction conditions. The rate of the water-gas shift reaction decreases linearly to zero with sulfur coverage over the coverage range between zero and a saturated, close-packed monolayer. Thus, it appears that in this case each sulfur adatom simply poisons a fraction of the Cu surface equivalent to that which it can sterically block (-2.6 Cu atoms), and that only one or maybe two Cu atoms are required for the rate-determining step. The reaction is positive-order with respect to H2O adsorption (26), a fact that does not change on sulfur addition. This suggests that only one or two Cu surface atoms are needed to cleave the 0-H bond in H20. The stability of adsorbed sulfur and its influence on the water-gas shift kinetics helps to explain the well-known sulfur intolerance of Cu-based catalysts for this reaction (24). The influence of adsorbed Si, P, S, and C1 on the medium-pressure cyclotrimerization of acetylene to benzene over Pd(l1 l), (loo), and (1 10) has been studied by Logan et al. (113). Whereas both sulfur and chlorine decrease the activity, silicon increases the activity. The effect of phosphorous depends on the crystal face. According to their work function measurements, sulfur withdraws electron density from the Pd surface (as is also expected for chlorine), whereas Si donates electron density, and P has the least effect on the work function. Thus, the qualitative influences on catalytic activities correlate with the influences of the additives on the electronic character of the surface. In addition, Si decreases the carbon coverage seen in postreaction AES from -82 to -70% of a monolayer, whereas sulfur and chlorine increase the amount of carbonaceous residue. The authors interpreted these results by suggesting that the electrondonating ligands keep the Pd surface cleaner for the desirable reaction by

32

CHARLES T. CAMPBELL

preventing surface polymerization. Inconsistent with this model, potassium, which should have the greatest extent of charge transfer of all the ligands studied, caused a decrease in the catalytic activity of Pd(llO), although it did double the rate on the other two surfaces. Garfunkel et al. (116) studied the influence of preoxidation of Re and Fe polycrystalline surfaces on their activity in CO hydrogenation. Oxidation usually led to higher methane selectivity and less carbon buildup on the surface. The activation energy for methanation was changed by oxidation of Fe(poly), which suggests either an electronic effect or a change in the rate-determining step. In contrast, Castner et al. (97) found that Rh surfaces produced a greater fraction of higher-molecular-weight hydrocarbons from CO + H2 mixtures when preoxidized, and that the selectivity for oxygenated products was substantial on oxidized Rh but insignificant on the metal alone. Hasenberg and Schmidt (119) studied the synthesis of HCN from CH, + NH3 mixtures at 0.01 to 10 tom over a polycrystalline Pt surface that was initially clean. Under reaction conditions the surface was covered with nearly a monolayer of carbon. Addition of 0 2 to the reaction mixture reduced the carbon coverage, decreasing the HCN production rate at the expense of NO production. A kinetic model was developed that fit the data well. Campbell and co-workers (48, 117, 118) have studied the influence of chlorine adatoms on the selective oxidation of ethylene to ethylene epoxide over Ag( 110) and Ag( 1 1 1) surfaces. As also occurs when trace levels of chlorinated hydrocarbon are added to the reactant feed over aluminasupported Ag catalysts, the deposition of chlorine adatoms on Ag single crystals led to a decrease in the steady-state rates of both epoxide production and full combustion (COz + H20). The latter was poisoned more rapidly than the former, and this lead to a marked improvement in selectivity, to a limiting value of 80-85%. This proved that the true promoter is atomically adsorbed chlorine, and that no support effects are required to understand (in first order) the role of the promoter. The selectivity is improved at significantly lower chlorine coverages on Ag( 111) than on Ag(I10) (48). In both cases the reaction rates decrease only at considerably larger chlorine coverages than those required to markedly suppress dissociative 0 2 adsorption, even under conditions where the reaction rate is limited by 0 2 supply. This was originally interpreted to indicate that molecularly, rather than dissociatively, adsorbed 0 2 was the true oxidizing agent (48, 117). The relative importance of molecularly versus dissociatively adsorbed oxygen in the reaction mechanism is, however, still a subject of some controversy (46). Figure 9 summarizes the results of a series of experiments in which the


c

,,

0.9

0

0.1 I

02

0.3

04

I

I

I

05 I

06 1

33

1

FIG.9. Influence of chlorine coverage on the kinetic parameters for selective ethylene oxidation over a Ag( 110) surface. Parameters for both the production of ethylene epoxide (EtO, circles) and the undesired side reaction to full combustion (COz, squares) are presented. Steady-state reaction orders in Ps and PE,and activation energies E, versus chlorine coverage near 563 K, PE,= 20 torr, and Po2 = 150 torr. From Ref. 118.

influence of chlorine coverage on the steady-state kinetics of ethylene epoxidation was determined over a range of temperature and partial pressure conditions on Ag(l10) (fZ8).The ability demonstrated here to monitor directly the changes in kinetic parameters as a function of modifier coverage highlights the strength of the surface science approach in addressing the influence of surface additives. The most obvious feature here is that the kinetic parameters for full combustion (COz) essentially track those for epoxidation (EtO), which strongly suggests that a common adsorbed intermediate is formed in the rate-determining step of both path-

34

CHARLES T. CAMPBELL

ways, and that the selectivity is determined only subsequent to this step by the branching ratio of two relatively rapid competing reactions available to this intermediate. Campbell and co-workers have analyzed these chlorine-induced changes in the kinetic parameters using the results of UHV adsorption-desorption experiments (kinetics and energetics) measured for ethylene and oxygen as a function of chlorine coverage (118). For example, low chlorine coverages increase the heat of adsorption (and therefore steady-state coverage) of ethylene. As a consequence, under conditions in which the reaction rate is positive-order in ethylene pressure, chlorine addition leads to an increase in epoxidation rate and a decrease in the reaction order with respect to ethylene pressure. Tan et al. (120-122) have also studied the influence of chlorine on the epoxidation of ethylene over Ag( 111) in detail. The Lambert group (123126) has also performed numerous other mechanistic investigations of this reaction over Ag(ll1). These and other mechanistic studies of this reaction are excellently reviewed in a recent article by van Santen and Kuipers (46).

2. Alkali Promoters Catalysts of various types are frequently treated with solutions of alkali salts in order to improve their subsequent activity or selectivity in specific applications. A fundamental understanding of the role of alkali promoters has long been a goal of the catalysis community. In 1982, Campbell and Goodman (127) investigated the role of potassium promoters in nickel catalysts for CO hydrogenation by using model catalysts synthesized by the vapor deposition of well-defined K overlayers on a clean Ni(100) surface. As with K impregnation of supported Ni catalysts, the steady-state rate of catalytic methane production decreased and that of higher hydrocarbons and olefins increased (see Fig. 10). This result showed that the dominant role of the alkali in this case is unrelated to interactions with the catalyst support material. The activation barrier for CO dissociation on Ni(100) was shown to be greatly reduced by K addition; therefore, the steady-state carbide coverage under reaction conditions increased markedly. Within the accepted “carbide” mechanism, this was expected to cause an increase in the C :H ratio in the surface layer and, therefore, a decrease in the probability for hydrogenation of carbon and an increase in the probability for chain growth (C-C bond formation) in the adsorbed intermediates. Thus, less CH4 and more olefins and higher hydrocarbons were seen in the products, and the graphitization of the surface occurs at a lower temperature. Since that study, the coadsorption of CO and alkalis on transition-metal surfaces has been


METHANE

ETHYLENE

35

ETHANE

FIG. 10. A comparison of the product distributions (weight percent) over a Ni(100) catalyst observed during CO hydrogenation over clean (solid bars) and K-doped (crosshatched bars) catalysts at T = 500 K, H2: CO = 4 : 1, and a total pressure of 120 torr. Potassium coverage = 0.10 monolayers (ML).From Ref. 127.

extensively studied, and numerous short- and long-range electronic interactions have been shown to operate in the coadsorbed layer, which can facilitate C-0 bond cleavage (see Ref. 128 and refs. cited therein). It should be noted, however, that the K on Ni(100) was not atomically absorbed under reaction conditions, but instead stabilized by some oxygen-containing species, perhaps in a CO, H20, or OH complex (127). Similar results have been obtained for K-promoted Fe foils and surfacecharacterized Fe powders by Bonze1 and co-workers (129-132) and by Dwyer and Hardenbergh (133, 134). Many of these results over Fe have been discussed in detail in a recent paper by Wesner et al. (132). On Fe foil, the chain-growth factor (Y in the Fischer-Tropsch reaction increased by -30% on K addition. Dwyer and Hardenbergh (134) observed a new surface-carbon species (neither carbidic nor graphitic) in postreaction XPS analysis of K-promoted Fe powders and suggested this species to be “polymethylene-like.” This same K-induced species was later seen on Fe foils by Wesner et al. (132) and shown to be kinetically active in hydrogenation, and its presence was well correlated with the enhancement in chain growth. In general, postreaction surface analysis indicates much higher carbon coverages on Fe compared to Ni surfaces, both with and

36

CHARLES T. CAMPBELL

without K. As on Ni, the carbon graphitizes at high temperatures. The graphitization temperature is considerably higher on Fe powders than on Fe foil (133). Similar structural and kinetic effects were observed for vapor-deposited and solution-deposited K on polycrystalline Fe foil (132), suggesting that the catalytic reaction conditions rather than the method of deposition controlled the structure of the alkali on the working catalyst’s surface. Evidence from XPS measurements suggested a surface KOH complex for this structure (131, 132). Similar influences of Na have been reported in CO hydrogenation over Re foils (116), where the selectivity for longer-chain hydrocarbons increased and the overall rate declined. Wesner et al. (135) studied K promotion of Co foils in the FischerTropsch reaction. Again, quite similar results were obtained. Like Ni, Co showed much less surface carbon following reaction than on Fe, although on Co the carbon is mostly graphitic in nature. However, after K addition, the carbon is much more carbidic. The reactivity of carbidic carbon with H2 on both Fe and Co is reduced in the presence of K. On Mo(loo), somewhat similar results were also obtained in CO hydrogenation by Logan et al. (98),where the selectivity for higher hydrocarbons and olefins increased on K addition. In this case, the overall rate of reaction also increased, which may be related to the unusual positive CO pressure dependence for this surface. The hydrogenation of C02 over K-promoted Ni(100) was studied by Peebles et al. (110), who showed a strong enhancement in both Co and methane production rates but no changes in activation energies. They also showed that K could be used as an antidote to low-level sulfur poisoning, where the combined effects of K + S was almost a coverage-weighted average of their individual effects. The addition of potassium to industrial Fe catalysts leads to an increase in activity for ammonia synthesis (Nz + 3H2 --.* 3NH3)(136).This promotion effect has been the subject of considerable attention from the surface science community, particularly with regard to the coadsorption of K or K + 0 and N2 (136-139). Ertl and co-workers have shown that potassium addition to single-crystal Fe surfaces can lead to a 10- to 100-fold enhancement in the rate of dissociative N2 adsorption, which is thought to be the rate-determining step in NH3 synthesis (136-139). However, Bare et al. (140) were unable to promote the activity of Fe( 1 1 I ) , (loo), or (1 10) surfaces for this reaction at 20-atm pressure with either K, K + 0, or K + AIO, addition. They interpreted this result to indicate that the promotional role of K in industrial catalysis may be cooperation with other promoters, such as the support material, to cause structural rather than electronic promotion. These results were for very low conversions, however, so that the product (NHj) partial pressure was low. Strongin and

PRESSURE-VACUUM

ANALYSIS OF MODEL CATALYSTS

37

Somojai (183) later found that potassium does promote the activity of Fe( 1 11) and (100) at higher conversions, and that the effect increases with increasing conversion. They interpreted this in terms of the poisoning of surface sites by adsorbed ammonia and a decrease in the coverage of NH3 due to the presence of potassium. They also found that potassium decreases the heat of NH3 adsorption. Asscher er al. (85) also saw at best only a very slight increase in the NH3 production rate over Re singlecrystal surfaces after potassium addition. For both Fe and Re, surface oxygen was required to stabilize the surface alkali under reaction conditions. Stolze and Norskov (181, 182) have modeled reasonably accurately the kinetics of high-pressure ammonia synthesis over Fe/AI2O3and commercial K/Fe-based catalysts using kinetic parameters for elementary reaction steps taken from UHV studies over clean and K-dosed Fe single crystals (86, 87, 136-139). Zaera and Somojai (141) studied the influence of potassium on the reactions of n-hexane and H2 over Pt( 1 11) near atmospheric pressure. The overall reaction rate decreased dramatically with as little as 2% of a monolayer of K, and the coverage of carbonaceous residue also increased slightly in this K coverage range. At higher K coverages, the rates continued to decrease but the carbon coverage also decreased. The results were explained in terms of inhibition by K of the dehydrogenation of adsorbed hydrocarbons. Campbell and Koel (142, 143) have studied cesium promotion of the medium-pressure water-gas shift reaction (H2O + CO + H2 + C02) over a well-defined Cu(ll1) surface under conditions in which the rate is limited by dissociative H2O adsorption. As shown in Fig. 1 1, they observed a 15-fold enhancement in catalytic activity at 8cs = 0.15, with no major change in reaction order with respect to H20 and only a slight increase in the overall activation energy (17 to 20 kcal/mol). Under reaction conditions, the Cs was stabilized on the surface in an oxidic or hydroxidic complex of approximate ratio Cs : 0 = 1 : 1, which formed, at high coverages, islands of ap(2 x 2) overlayer structure. The maximum promotional effect was achieved at a Cs coverage just below that required to nucleate these p(2 X 2) islands. They attribute the promotional effect to direct participation of the surface CsOH complex in H20 bond cleavage. Again, coadsorption of S poison and Cs promoter led to nearly directly additive individual effects. Campbell (144, 145) has also studied Cs promotion of the selective oxidation of ethylene over Ag( 11 1) at low conversions. Under mediumpressure reaction conditions, the Cs is stabilized as a surface cesium oxide of approximate stoichiometry Cs : 0 = 1 : 3, where bonding of the oxygen atoms both to the surface Ag atoms and to at least one Cs atom is

38

C H A R L E S T. C A M P B E L L

1.6 1.4

3

1.2

u)

s; 1.0 W

4

i , -1

o*8

0 0.6

I

v

0.4

Cs/Cu

AES RATIO / ( l / l O O )

FIG.1 I . Variation of the water-gas shift reaction rate with cesium coverage on Cu( I 1 1) at 612 K, 26-torr CO, and 10-tom H20. The open data point represents dosing of pure water (no CsOH). From Ref. 142.

suggested both by this stoichiometry and by XPS,AES, and TDS results. This C s / O complex coalesces into islands of (2 fi x 2 f i ) R 3 0 " structure according to postreaction LEED analysis. This same cesium structure results either by running the reaction over a pure Cs overlayer on atomically clean Ag( 11I), or by dosing Cs from ultrapure aqueous solutions of cesium hydroxide, carbonate, or nitrate onto Ag( 11 1) using procedures very similar to those used in preparing Cs-doped industrial catalysts. This result lays to rest fears that the vapor-deposition method chosen for alkali promotion by surface scientists yields a final alkali structure any different from what would be expected from the more industrial approach of impregnation from solution. Small coverages of Cs led to a decrease in both rate and selectivity, whereas very high coverages continued to poison the rate but led to only a very small improvement in selectivity compared to alkali-free Ag(ll1). Although these minor effects cannot be neglected, the major promotional effect in industrial catalysis appears to be related to other factors such as interactions with the support or other promoters and/or further oxidation of the product ethylene epoxide.

PRESSURE-VACUUM


39

Lambert and co-workers (121-123, 125) have addressed this issue of further oxidation and other features of Cs promotion using model catalysts based on Cs/Ag(11 1). Their results indicate that the simultaneous presence of Cs and COz lead to marked enhancements in selectivity (125). In contrast to Campbell’s results, they report that Cs enhances the rate of epoxidation at low conversion (122). They also show that, although Ag( 11 1) is active in the isomerization of the product (ethylene epoxide) to acetaldehyde, its activity for this side reaction can be suppressed by surface Cs addition (122, 123). Since this isomerization may lead to further oxidation of the epoxide, this may help to explain the enhancement in selectivity derived from Cs addition in industrial catalysis. C. BIMETALLIC A N D ALLOYCATALYSTS The subject of bimetallic or alloy catalysis has long been of deep interest to the catalytic chemist. Addition of a second metal to a moderately good catalyst can frequently improve either the activity, the selectivity, or the lifetime of the catalyst. This improvement can arise from any combination of effects, including those that are basically electronic in nature, whereby the metal-metal bonding changes their electronic structure, or those that are basically geometric in nature, where the size, shape, and availability of the ensembles or groups of active metal sites can be altered by the presence of a second, frequently more inert metal. These subjects have been excellently reviewed in a number of recent works by Sachtler and van Santen (146, 143, Ponec (39), and Sinfelt and Cusumano (148,149).The goal of the surface science experiment here has been to carefully distinguish these two effects in certain reaction systems, and even to quantify the surface electronic, structural, or ensemble characteristics required for an effective catalyst. This hope is, of course, spurred by a knowledge that these properties can be carefully controlled and quantitatively characterized by use of the modern spectroscopic and structural tools of surface science combined with model catalysts based on single crystals. Here, the second metal is either present as an alloy in the bulk crystal itself or is vapor-deposited (or chemically deposited) onto the surface under the pristine conditions of UHV. In the latter case it has been repeatedly demonstrated that well-ordered overlayers or surface alloys of the second metal can often be generated that present the entire range of surface compositions of the two metals (see Refs. 150-154 and refs. cited therein). Although UHV adsorption-desorption studies on these types of model catalyst have been pursued extensively for some

40

CHARLES T. CAMPBELL

time, actual catalytic kinetics at realistic pressures has been only a recent development. In 1983 Sachtler and Somorjai (155) studied the reactions of n-hexane + H2 at -230 torr over a Pt( 1 1 1) single-crystal surface that had been carefully modified with vapor-deposited Au overlayers and surface alloys generated by annealing these overlayers. Although Au itself is very inert for these reactions, the surface alloys were found to be more active than pure Pt( 1 1 1). Large increases in the isomerization rate and simultaneous exponential decreases in the rates of hydrogenolysis and aromatization with increasing Au concentration resulted in high isomerization selectivity. The authors attributed this result to an ensemble effect, where the probability for finding large groups of Pt sites, required for hydrogenolysis, aromatization, and possibly deactivation by carbon buildup, decreases as Au dilutes the Pt alloy. These authors soon extended this study to include cyclohexane dehydrogenation (156), where they found a large enhancement in the rate of benzene production, which maximized at a surface composition of -50% Au and declined thereafter. Gold addition also induced cyclohexene production, which occurred most readily at 90% Au surface concentration. Again, the results were explained by an ensemble effect, where a decrease in the sizes of available Pt sites reduced the extent of site poisoning by the benzene product. More recently, the authors have further extended this work to include a comparison of the Au-Pt( 1 1 1) versus Au-Pt( 100) alloy surfaces in nhexane reactions (157). As shown in Fig. 12, the same general decreases in reaction rates and increases in isomerization selectivity with Au addition were seen for both Pt surface orientations, except for the fact that, in contrast to Pt( loo), the rate of methylpentane production actually increases with Au addition. Other more subtle differences were also observed, demonstrating clearly the existence of structural sensitivity in alloy catalysis. Again, ensemble effects were used to explain the results. Unfortunately, the sizes and shapes of available Pt atom ensembles could not be determined in any of the preceding studies even for the starting alloys. This is due to the difficulty of determining the lateral distribution of elements in bimetallic surfaces, even when they give good LEED patterns as in these cases. This prevents any quantitative assessment of ensemble requirements for the various reactions and clearly highlights a glaring need for the development of surface analytical tools that determine the lateral distributions of elements on multimetallic surfaces. Recently, Campbell and co-workers (158) have demonstrated new methodology that may allow for semi-quantitative assessment of ensemble effects. This method relies on the inertness of bismuth adatoms and their

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potential as nearly ideal site-blocking agents that have well-defined lateral distributions on transition metal surfaces. The bimetallic Cu/Ru system does not alloy, but supported Cu/Ru catalysts have been well studied because Cu addition very selectively suppresses hydrogenolysis reactions compared to simple dehydrogenation over Ru (148, 149). Recently, Peden and Goodman (159-163) have investigated this system using well-defined adlayers of Cu on a Ru(0001) surface. In this case, Cu adatoms clusters into two-dimensional islands on the surface. The addition of submonolayer Cu causes a dramatic (eightfold) increase in the rate of cyclohexane dehydrogenation to benzene at 100 torr (159, 160, 163) and a linear decrease in the rate of ethane hydrogenolysis to methane (159, 161). Ultimately this leads to the same marked

42

CHARLES T. CAMPBELL

increase in selectivity for dehydrogenation over hydrogenolysis as seen for supported Cu/Ru (159). However, the enhanced activity for dehydrogenation has not been observed on the supported catalysts. In Fig. 13, Peden and Goodman compare the supported and single-crystal catalysts in terms of their specific activity (now on a per surface Ru atom basis) for these two reactions (159). In the case of the unsupported catalysts, Peden and Goodman (see Ref. 159 and refs. cited therein) have seen substantial spillover of hydrogen chemisorption from Ru to Cu sites on their model catalysts. Because the Ru surface atom density in the experiments with supported catalysts was measured by HZchemisorption, they conclude that there may have been an overestimation of the Ru surface concentration in those experiments (159). This correction would bring the data of Fig. 13 for the two types of catalyst at least into qualitative agreement. This subject, however, remains controversial. Peden and Goodman attribute the enhancement in cyclohexane dehydrogenation to the altered

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43

geometric and electronic properties of the surface sites, or perhaps to a cooperative effect such as they see for hydrogen spillover from Ru to Cu. Ensemble effects may certainly be at work here since carbonaceous residue and Cu both suppress hydrogenolysis on the Ru surface in a similar way (163). Goodman and co-workers have also studied the kinetics of methanation and cyclohexane hydrogenolysis on their Cu/Ru(0001)model catalysts. In these cases, as with ethane hydrogenolysis, Cu merely serves as an inactive diluent, suppressing the rates on a simple one-to-one site-blocking basis (161, 162). Greenlief et al. (164) have studied CO methanation and ethane hydrogenolysis over submonolayer Ni films on W(110) and W(100). For these reactions, W alone is very inactive. They have compared the kinetics with that seen for pure Ni crystals. When compared on a per surface Ni atom basis, the turnover frequencies and activation energies for CO methanation are very similar for pure Ni crystals and for the whole range of coverages of Ni on W surfaces. This highlights the extreme structural insensitivity of this reaction on Ni. Although the activation energies for ethane hydrogenolysis over Ni/W( 110) and Ni/W(100) were very similar, the rate (per Ni atom) on Ni/W(l 10) decreased with increasing Ni coverage.

D. METAL-SUPPORT INTERACTIONS The oxidic support material (e.g., SiOz, Al203, TiO2, ZnO, ...) for a catalytically active metal can favorably or adversely influence the performance of that metal in a particular catalytic process. This has been a subject of considerable study for many years, and the general picture has evolved that the catalytic kinetics can be affected by electronic influences on the active metal atoms due to their specific bonding to the support, structural effects where the dispersion of the metal phase or its atomic geometry can be controlled by the underlying support, and cooperative effects were adsorbed intermediates bond to both the metal and the oxidic surface sites, or where certain necessary or undesirable steps in the reaction mechanism take place on the surface of the oxidic support. However, the specific details of these effects have been clarified in only a few cases. Recently, surface science methods have been applied to the study of these interactions, which are referred to as “SMSI” effects (strong metalsupport interaction) when particularly noticeable. The general subject of SMSI effects has recently been reviewed in detail by Tauster (165), who discusses results for both practical and model catalysts.

44

CHARLES T. CAMPBELL

In 1982 Kao et al. (166) reported on the CO hydrogenation activity of nickel overlayers on the TiOz(100) surface as a model SMSI catalyst. They found that an average Ni thickness of -5 8, gave optimum activity, displaying a rate that was 3.3-3.7 times that for an equivalent size sample of pure Ni( 11 I ) , with an accompanying increase in the selectivity for higher hydrocarbons. As shown in Fig. 14, the activation energy for this model SMSI catalyst was unchanged from Ni(ll1). Later, Chung and coworkers (167) studied the activity of a Ni(ll1) surface containing controlled amounts of reduced titania. This catalyst was optimally promoted at a titania coverage of -0.1 monolayer, displaying an activity and product distribution similar to Ni/Ti02(100) and essentially identical to the practical Ni/TiO2 catalysts. The model Ni/TiO2 catalysts prepared by different methods showed similar surface Ti3+ concentrations. These results, together with their companion study (168) of the surface composition of model Ni/Ti02catalysts that had been given the high-temperature SMSI reduction, suggested that partially reduced titania species migrate up from the support material to partially cover the Ni surface. Since this reaction is known to be structurally insensitive (see above), these species must improve the catalysts by some favorable electronic influence on the Ni surface atoms or else by direct participation in the reaction mechanism. Demmin et al. (169-1 72) studied methanation and CO and H2 chemisorption on model SMSI catalysts prepared by deposition of titania onto T(C)

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FIG. 15. Methanation rates for 100-torr CO and 400-torr H2 on a Pt foil. The turnover is in units of molecules/site/second and is based on 1 x 1OIs sites/cm2for the number (hiCH4) foil. The dashed lines indicate rates reported by Vannice and Twu (185) for Pt/TiO, (--4, Pt/Al2O3(---), and Pt/SiO, (-----). The rates for the niobia- and titania-covered foils were measured with an oxide coverage sufficient to completely suppress CO and H2adsorption. (x), Clean Pt; (O),TiOJPt; (O), NbOJPt. From Ref. 172.

Pt, Pd, and Rh foils and niobia onto a Pt foil. In all cases, a partial oxide formed uniform layers on the metal surfaces, suppressing chemisorption of H2 and CO. Nevertheless, the rate of methanation increased substantially and its activation energy decreased, as can be seen in Fig. 15. These three results are almost identical to those reported for the well-known “SMSI state” of high-surface-area niobia- and titania-supported metals. On the basis of these results, Demmin and co-workers attribute the SMSI effect to the migration of a titania or niobia species onto the metal crystallites. Demmin and Gorte (170) also showed that, while clean Pt foil showed no surface carbon in postreaction AES, the titania-dosed foil showed significant quantities of reactive carbon following reactions at low H2/CO pressure ratios, suggesting that CO dissociation on Pt is enhanced by the presence of titania species at the surface. Levin et al. (173) studied the surface structure and chemisorption properties of Ti0,-dosed Rh foil following high-temperature reduction in medium pressures of H2. Again, CO adsorption was inhibited. Later, they

46

CHARLES T. CAMPBELL

showed that an optimum coverage of TiO, of -0.15 monolayers led to a threefold enhancement in activity for CO hydrogenation at atmospheric pressure, and a higher selectivity for olefins, a lower activation energy, and higher reaction orders with respect to H2 and CO (174, 186). These results are very similar to those for Ni and Pt (above). The authors argued that surface Ti3+participates in CO dissociation. In contrast to titania and niobia, alumina is a support that does not display the classic SMSI behavior outlined above, at least for practical catalysts (173, 175). Bischke et al. (176) studied CO methanation over model Ni/A1203 catalysts prepared by evaporation of Al onto a W foil, subsequent oxidation of the Al, and finally vapor deposition of Ni onto the thin A1203film. Specific reaction rates (per Ni surface atom) and the activation energy were similar to those found for pure Ni crystals and practical Ni/A1203 catalysts. This agrees with the classical picture of A1203 as a more or less inert support for this catalyst and the known structural insensitivity of this reaction over Ni (see above). Submonolayer quantities of oxidized aluminum were shown by Levin et al. (186) to only decrease the rate of methanation over a Rh foil, in proportion to the fraction of Rh sites covered by the oxide film. Lee et al. (17) prepared model Fe/MgO catalysts by deposition of Fe onto MgO(100). These surfaces were then characterized in detail following Hz + CO treatments at -0.1 torr by application of a powerful array of surface analytical methods. For Fe layers less than four atoms thick, the Fe was oxidized after reaction with CO + H2. For higher coverages, the Fe was carburized. Carbon deposition was seen at all Fe thicknesses. They saw at least two carbon species, with C-Fe distances of 1.78 and 2.06 A. A strong synergy between Cu and ZnO has been reported for highsurface-area Cu/ZnO catalysts in methanol synthesis (CO + H2 + CH20H) by Klier (see Refs. 178 and 179 and refs. cited therein). To study the pertinent Cu/ZnO interactions, Campbell et al. (102) prepared and characterized model catalysts consisting of well-defined Cu overlayers on the oxygen-terminated ZnO(0M) single-crystal surface and others consisting of ZnO, films on Cu( 111). They attempted to produce methanol on these catalysts at temperatures of 500-600 K and H2 + CO(+COz) total pressures of up to 1500 torr. Although they were unable to observe any methanol production, they were able to set an upper limit on the activities of these model catalysts at < 2 X molecules/site/s. This limit is consistent with the rates expected for high-area Cu/ZnO catalysts (extrapolated from a somewhat higher pressure regime). Campbell and Daube (26) also studied the kinetics of the water-gas shift reaction over a


47

Cu( 11 1) surface and saw no major changes in the rate due to the addition of ZnO, to the surface. Strongin et al. (180) studied metal-support effects in Fe-catalyzed ammonia synthesis by depositing aluminum oxide onto the Fe(l1 l), (loo), and (1 10) surfaces and observing the resulting influences on the kinetics of NH3 production from N2 + Hz.Treatment of the model catalysts in 20torr H 2 0 caused a lasting restructuring of the alumina-dosed Fe(l10) and (100) surfaces. This restructuring resulted in a marked increase in catalytic activity for these Fe surfaces, to rates near those for the (already very active) Fe(ll1) plane (see above). This was explained in terms of an alumina-induced geometric rearrangement of the Fe surface atoms to a local geometry more favorable for participation in NH3 production. IV. Conclusion It is safe to say that the combination of UHV surface analyses with medium- and high-pressure kinetics over well-defined planar surfaces has already added considerably to our understanding of structure-function relationships in catalysis. This is certainly nor to say that experiments using this approach are all that will be required to develop a thorough understanding of the relationships between catalyst performance and atomic-level surface structure. Indeed, experiments with high-surfacearea catalysts will play an equally important role. Studies of either type are certainly most meaningful when they can be directly compared with results using the other approach. We are slowly approaching a happy symbiosis between careful experiments over practical, high-surface-area catalysts and model, single-crystal catalysts. Because of the inherent surface heterogeneity of practical catalysts, many types of experiment are preferably performed over the homogeneous surfaces of model catalysts such as those reviewed here. However, present technology limits many of the kinds of experiment that can be performed on well-defined planar surfaces. For example, the amount of gaseous product is small and difficult to analyze accurately. Therefore, transient kinetics and isotope-labeling experiments designed to elucidate the reaction mechanism are currently most conveniently performed with powdered or supported catalysts. It is notable that even state-state kinetics over single-crystal metal oxides have never been reported, partly because the specific rates in oxide catalysis are very low. Similarly, the surface area of planar model catalysts is small, so that competing side reactions over undefined reactor surfaces can lead to significant problems

48

CHARLES T. CAMPBELL

that are difficult to detect and avoid. In addition, recipes for preparing truly homogeneous, well-defined planar samples presently exist for only a very limited range of materials. The preparation of well-defined surfaces of more complex materials remains as one of the major challenges of “UHV preparatory surface chemistry. ” Catalysis over biphasic materials, such as is thought to be the case in many redox catalysts for selective oxidation (e.g., copper and copper oxide mixtures), represents a class of reaction where study over a single homogeneous surface may well be impossible at steady state. However, transient kinetics of elementary steps over homogeneous surfaces should provide useful insight into the relative roles of the separate phases even in this case. In any case, many types of experiments will continue to be most preferably performed over powdered and supported high-surface-area catalysts. Indeed, these catalysts certainly are much closer in structure to those of commercial import, and many methods of surface analysis can be performed only with higharea samples. It should be pointed out that this field of catalysis on well-defined, planar surfaces is still in its infancy. As in the early days of vacuum surface physics, the results should generally be taken “with a grain of salt” until they are corroborated over another sample by another group. Complications in surface preparation and cleanliness are very difficult to avoid and are more common than expected by the “high-surface-area” catalysis community. The quality of kinetic results over homogeneous surfaces is presently limited by the quality of sample preparation and mounting. There is a tendency to see the sophisticated instrumentation involved and trust implicitly any results from so impressive an apparatus. Scientifically speaking, this can be as dangerous as the reactionary response of discarding all kinetic results from the surface science approach after a few (or several) contradictions are discovered. At the same time, surface scientists must make a dedicated effort to familiarize themselves with the literature of high-area catalysis, as a knowledge of the experimental complications in that field can help them to understand and appreciate seemingly contradictory results found there. Clearly, some highsurface-area catalysts, however heterogeneous, have been better characterized than others. In the future we will see the development of more and more sophisticated surface analytical techniques capable of analyzing a variety of surface types under high-pressure reaction conditions. Eventually this will eliminate the complications of sample transfer into UHV described here. At the same time, more sensitive detection schemes for gaseous products will appear, as will reproducible recipes for preparation of homogeneous surfaces of complex materials. All these developments will facilitate the


49

study of accurate transient and steady-state catalytic kinetics over welldefined, planar surfaces. They will also aid in the quality of structural and kinetic characterization of high-surface-area catalysts resembling those used industrially. The present author is optimistic that investigations using these two different approaches, when coupled with proper theoretical input, will eventually lead to a relatively complete understanding of structure-function relationships in the surface chemistry of catalysis. ACKNOWLEDGMENTS

The author gratefully acknowledges support for this research by the Department of Energy, Office of Basic Energy Sciences, Chemical Science Division. I thank D. W. Goodman and G. A . Somojai for helpful discussions, for communication of results prior to publication, and for contribution of figures. I am also indebted to R. J. Gorte and Y.-W. Chung for contributions of figures and Elizabeth McGaw for typing of the manuscript.

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5 . Goodman, D. W., Kelley, R. D., Madey, T. E., and Yates, J. T., J . Catal. 63, 226 ( 1980).

6 . Krebs, H. J., Bonze], H. P., and Gafner, G., Surf. Sci. 88, 269 (1979). 7. Bonze], H. P., and Krebs, H. J., Surf. Sci. 91,499 (1980). 8. Polizzotti, R. S . , and Schwarz, J. A., J . Vac. Sci. Techno/. 17, 655 (1980). 9. Cabrera, A. L., Spencer, N. D., Kozak, E., Davies, P. W., and Somojai, G . A.. Rev. Sci. Instrum. 53, 1888 (1982). 10. Campbell, C. T., and Paffett, M. T., Surf. Sci. 139, 396 (1984). 11. Ichikawa. S . , and Wilson, M. S . , Rev. Sci. Instrum. 58, 317 (1987). 12. Rucker, T. G . . Frank, K., Colomb, D., Logan, M. A., and Somojai, G. A . , Rev. Sci. Instrum. 58, 2292 (1987). 13. Castle, J. E., Surf. Interface Anal. 9, 345 (1986). 14. Peuckert, M., Coehen, F. P., and Bonze], H. P.. Surf. Sci. 144, 515 (1985). I S . Peuckert, M., J . Electroanul. Chem. 185, 379 (1985). 16. Ross, P. N., and Wagner, F. T., Adu. Electrochem. Electrochem. Eng. 13,69 (1985). 17. Hubbard, A. T., Acc. Chem. Res. 13, 177 (1980). 18. Schardt, G . C., Stickney, J. L., Stern, D. A . , Wieckowski, A . , Zapien, D. C., and Hubbard, A . T.. Lungmuir 3, 239 (1987). 19. Ertl, G . , and Kiippers, J., “Low Energy Electrons and Surface Chemistry.” Verlag Chemie, Weinheim, 1985. 20. Woodruff, D. P., and Delchar, T. A,, “Modern Techniques of Surface Science.” Cambridge Univ. Press, New York, 1986. 21. Feldman, L. C., and Mayer. J. W., “Fundamentals of Surface and Thin Film Analysis.’’ North-Holland Pub]., New York, 1986.

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22. Clarke, L. J., “An Introduction to Low Energy Electron Diffraction.” Wiley, New York, 1985. 23. Davis, S. M., Zaera, F., and Somojai. G . A., J . A m . Chem. SOC. 104, 7453 (1982). 24. Campbell. C. T.,and Koel, B. E.,Surf. Sci. 183, 100 (1987). 25. Campbell, C. T., and Paffett, M. T., Appl. Surf. Sci. 19, 28 (1984). 26. Campbell, C. T . , and Daube, K. A., J. Caral. 104, 109 (1987). 27. Campbell, C. T.,J. Vac. Sci. Technol. A 6, I108 (1988). 28. Herglotz, H. K., and Lynch, D. R.. J . Coll. Interface Sci. 105, 447 (1985). 29. Hoffman, F. M., and Robbins, J. L., J. Electron Spec. Relar. Phenom. 45,421 (1987). 30. Fischer, D. A., Dobler, U., Arvanitis, D., Wenzel, L., Babershki, K., and Stohr. J., Surf: Sci. 177, 114 (1986). 31. Blum, L., Abruna, H. D., White, J . , Gordon, J. G., Borger, G. L., Samart, M. G., and Melroy, 0. R., J . Chem. Phys. 85, 6732 (1986). 32. Zaera, F., Fischer, D. A., Shin, S., and Gland, J. L., Surf. Sci. 194, 205 (1988). 33. Redhead, P. A., Vacuum l2,203 (1962). 34. Falconer, J. L., and Madix, R. J., Surf. Sci. 48, 393 (1975). 35. Chan, C. M., Ark, R., and Weinberg, W. H., Appl. Surc Sci. 1, 360, 377 (1978). 36. Robertshaw, G., and Estrup, P. J., J. Vac. Sci. Technol. 15, 554 (1978). 37. Boudart, M., Proc. Inr. Congr. Card., 6rh, London 1, 1 (1976). 38. Katzer. J. R., J . C u r d 32, 166 (1974). 39. Ponec, V., Adu. Card. 32, 149 (1983). 40. Boudart, M., C H E M T E C H Nov., P-688 (1986). 41. Kummer, J. T.,J . Phys. Chem. 60, 666 (1956). 42. Wilson, J . N.. Voge, H. H., Stevenson, D. P., Smith, A. E., and Atkins. L. T., J . Phys. Chem. 63, 463 (1959). 43. Campbell, C. T.,J. Card. 94, 436 (1985). 44. Campbell, C. T.,J . Vac. Sci. Technol., A 2, 1024 (1984). 45. Sajkowski, D. J., and Boudart, M., Caral. Rev.-Sci. Eng. 29, 325 (1987). 46. van Santen, R. A., and Kuipers, H. P. C. E., Adu. Card. 35, 265 (1987). 47. Delgass, N., personal communication. 48. Campbell, C. T., J. Card. 99, 28 (1986). 49. Campbell, C. T.,and Paffett, M. T.,Sucf. Sci. 143, 517 (1984). 50. Somojai, G . A., Adu. Card. 26, I (1977). S l a . Somorjai, G . A., Science 227, 902 (1985). 516. Somojai, G. A., “Chemistry in Two Dimensions.” Cornell Univ. Press, New York, 1981. 52. Somojai, G. A., Chem. Soc. Rev. 13, 321 (1984). 53. Somojai, G . A., and Yeates, R. C., J. Electrochem. SOC. l31, 228C (1984). 54. Davis, S. M.. Zaera, F., and Somojai, G. A., J. Caral. 85, 206 (1984). 55. Gillespie, W. D., Herz, R. K.. Petersen, E. E., and Somojai, G. A., J . Carol. 70, 147 (1981). 56. Zaera, F., Godbey, D., and Somojai. G. A., J . C a d . 101, 73 (1986). 57. Dauscher, A., Garin, F., and Maire, G., J . C a r d . 105, 233 (1987). 58. Garin, F., Aeiyach, S . , Legare, P., and Maire, G., J . Card. 77, 323 (1982). 59. Herz, R. K., Gillespie, W. D., Petersen, E. E., and Somorjai, G. A., J . C a d . 67,371 (1981).

60. Rodriguez, J. A., and Campbell, C. T.,J . Phys. Chem. (in press) (1989). 61. Davis, S. M., Gillespie, W. D., and Somojai, G. A., J . Caral. 83, 131 (1983). 62. Seimanides, S. M., Campbell, J . M., and Campbell, C. T., J. Phys. Chem. (in press) ( I 989).

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51

63. Davis. S . M., Zaera, F., and Somojai, G. A., J . Caral. 77,439 (1982). 64. Davis, S. M., and Somojai, G. A., J . Catal. 65, 78 (1980). 65. Massardier, J., Bertolini, J. C., Tri, T. M., Gallezot, P., and Imelik, B., Bull. Soc. Chim. Fr. 3, 333 (1985). 66. Zaera, F., and Somojai, G. A., J . Am. Chem. SOC. 106, 2288 (1984). 67. Davis, S. M., Zaera, F., Gordon, B. E., and Somojai, G. A., J. Card. 92,240 (1985). 68. Koel, B. E., Bent, B. E., and Somojai, G. A., Sutf Sci. 146, 211 (1984). 69. Oudar, J., Pinol, S., and Berthier, Y., J . C a r d 107, 434 (1987). 70. Goodman, D. W., Proc. inr. Congr. Caral., 8rh, Berlin 4, 3 (1984). 71. Goodman, D. W., Acc. Chem. Res. 17, 194 (1984). 72. Goodman, D. W., Surf Sci. l23, L679 (1982). 73. Goodman, D. W., J. Vac. Sci. Technol., A 2, 873 (1984). 74. Sault, A. G., and Goodman, D. W., J. Chern. Phys. 88, 7232 (1988). 75. Engstrom, J. R., Goodman, D. W., and Weinberg, W. H., J . Am. Chem. Soc. 108, 4653 (1986). 76. Engstrom, J. R., Goodman, D. W., and Weinberg, W. H., J. Vac. Sci. Technol., A 5 , 825 (1987). 77. Engstrom, J. R.. Goodman, D. W., and Weinberg, W. H., J . Am. Chem. SOC. 110, 8305 (1988). 78. Engstrom, J. R., Goodman, D. W., and Weinberg, W.H., J . Am. Chem. Soc. (in

preparation). 79. Wax, M. J., Kelley, R. D., and Madey, T. E., J. Card. 98,487 (1986). 80. Peden, C. H. F., and Goodman, D. W., ACS Symp. Ser. No. 288, 185 (1985). 81. Rucker, T. G . , Logan, M. A., Gentle, T. M.. Muetterties, E. L.,and Somotjai, G. A., J. Phys. Chem. 90,2703 (1986). 82. Niemantsverdriet, J. W., and van Langeveld, A. D., in "Studies in Surface Science and Catalysis: Catalysis 1987" (J. W. Ward, ed.), Vol. 38, p. 769. Elsevier, Amsterdam, 1988. 83. Spencer, N. D., Schoonmaker, R. C., and Somojai, G. A., J. Carol. 74, 129 (1982). 84. Strongin, D. R., Carrazza, J., Bare, S. R., and Somojai, G. A., J. C a d . 103, 213 ( 1987). 85. Asscher, M., Carrazza, J., Khan, M. M., Lewis, K. B., and Somojai. G . A., J. Catal. 98, 277 (1986). 86. Bozso, F., Ertl, G., Grunze, M., and Weiss, M., J. C a d . 49, 18 (1977). 87. Bozso, F., Ertl, G., and Weiss, M., J. Carol. 50, 519 (1977). 88. Haase, G., and Asscher, M., Surf. Sci. (in press). 89. Falicov, L. M., and Somojai, G. A., Proc. Narl. Acad. Sci. USA 82,2207 (1985). 90. Peden, C. H. F., Goodman, D. W.,Blair, D. S., Berlowitz, P. J., Fisher, G. B., and Oh, S. H., J. Phys. Chem. 92, 1563 and 5213 (1988); Goodman, D. W., and Peden, C. H. F., J. Phys. Chem. 90,4839 (1986). 91. Oh, S. H.. Fisher, G. B., Carpenter, J. E., and Goodman, D. W., J. Carol. 100, 360 (1986). 92. Peden, C. H. F., Berlowitz, P. J., and Goodman, D. W., Proc. Inr. Congr. Catal., 9th p. 1214 (1988). 93. Hendershot, R. E., and Hansen, R. S., J . Card. 98, 150 (1986). 94. Goodman, D. W., and Peden, C. H. F., J . Phys. Chem. 90,4839 (1986); correction, J. Phys. Chem. 91, 5832 (1987). 95. Yeates, R. C., Turner, J. E., Gellman, A. J., and Somojai, G. A., Surf. Sci. 149, 175 ( 1985). 96. Goodman, D. W., and Houston, J. E., Science 236, 403 (1987).

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97. Castner, D. G., Blackadar, R. L., and Somojai, G. A., J . Card. 66, 257 (1980). 98. Logan, M., Gellman, A., and Somojai, G. A., J. Caral. 94, 60 (1985). 99. Udovic, T. J., Kelley, R. D., and Madey. T. E., Surf. Sci. 150, L71 (1985). 100. Szuromi, P. D., Kelley, R. D., and Madey, T. E., J. Phys. Chem. 90, 6499 (1986). 101. Berlowitz, P. J., and Goodman, D. W., J. Catal. 108, 364 (1987). 102. Campbell, C. T., Daube, K. A., and White, J. M., Surf. Sci. 182, 458 (1987). 103. Campbell, C. T., Appl. Card. 32, 367 (1987). 104. Gellman, A. J., Neiman. D., and Somojai, G. A., J. Catal. 107, 92, 103 (1987). 105. Bussell, M. E., and Somojai, G. A., J. Card. 106, 93 (1987). 106. Bussell, M. E., Gellman, A. J., and Somojai, G. A., J . Catal. 110, 423 (1988). 107. Spitzer, A., and Luth, H., Surf. Sci. 160, 353 (1985). 108. Goodman, D. W., Appl. Surf, Sci. 19, 1 (1984). 109. Goodman, D. W., and Kiskinova, M., Surf. Sci. 105, L265 (1981). 109a. Nakamura, J., Yamada, T., and Tanaka, K., Surf. Sci. 185, L515 (1987). 110. Peebles, D. E., Goodman, D. W., and White, J. M., J . Phys. Chem. 87,4378 (1983). 111. Szuromi, P. D., Kelley, R. D., and Madey, T. E., J . Vac. Sci. Technol., A 5, 867 (1987). 112. Oudar, J . , Pinol, S.. Pradier, C. M., and Berthier, Y., J. Catal. 107, 445 (1987). 113. Logan, M. A., Rucker, T. G . , Gentle, T. M., Muetterties, E. L., and Somojai, G. A., J . Phys. Chem. 90, 2709 (1986). 114. Bernard, J., Oudar, J., Barbouth, N., Margot, E.. and Berthier, Y . , Surf. Sci. 88, L35 (1979). 115. Bartholomew, C. H., Agrawal, P. K., and Katzer, J. R., Adu. C a d . 31, 135 (1982). 116. Garfunkel, E. L., Parmenter, J., Naasz, B. M., and Somojai, G. A., Langmuir 2, 105 (1986). 117. Campbell, C. T., and Paffett, M. T., Appl. Surf, Sci. 19, 28 (1984). 118. Campbell, C. T., and Koel, B. E., J . C a r d 92, 272 (1985). 119. Hasenberg, D., and Schmidt, L. D., J . Catal. 104, 441 (1987). 120. Tan, S. A., Grant, R. B., and Lamberg, R. M., J . Caral. 100, 383 (1986). 121. Tan, S. A., Grant, R. B., and Lambert, R. M., J. Catal. 106, 54 (1987). 122. Grant, R. B., Harbach. C., Lambert, R. M., andTan, S. A., J. Chem. SOC.,Faraday Trans 1 , 83, 2305 (1987). 123. Grant, R. B., and Lambert, R. M., in “Catalysis on the Energy Scene” (S. Kaliaguine and A. Mahay, eds.), p. 251. Elsevier, Amsterdam, 1984. 124. Grant, R. B., and Lambert, R. M., J . C a d . 92, 364 (1985). 125. Tan, S. A., Grant, R. B., and Lambert, R. M., Appl. Caral. 31, 159 (1987). 126. Tan, S. A.. Grant, R. B., and Lambert, R. M., J. C a d . 104, 156 (1987). 127. Campbell, C. T., and Goodman, D. W., Surf. Sci. l23, 413 (1982). 128. Rodriguez, J . A., and Campbell, C. T., J . Phys. Chem. 91, 2161 (1987). 129. Bonzel, H. P., and Krebs, H. J., Sutf Sci. 109, L527 (1981). 130. Bonzel, H. P., and Krebs, H. J., Surf. Sci. 117, 639 (1982). 131. Bonzel, H. P., Broden, G., and Krebs, H. J., Appl. Surf. Sci. 16, 373 (1983). 132. Wesner, D. A., Coenen, F. P., and Bonzel, H. P., Langmuir 1, 478 (1985). 133. Dwyer, D. J., and Hardenbergh, J. H., J. Caral. 87, 66 (1984). 134. Dwyer, D. J., and Hardenbergh, J. H., Appl. Surf. Sci. 19, 14 (1984). 135. Wesner, D. A., Linden, G., and Bonzel, H. P.. Appl. Surf. Sci. 26, 335 (1986). 136. Ertl, G . , Catul. Rev.-Sci. Eng. 21, 201 (1980). 137. Lee, S. B., Weiss, M., and Ertle, G., Surf Sci. 108, 357 (1981). 138. Paal, Z., Ertl, G., and Lee, S. B., Appl. Surf. Sci. 8, 231 (1981). 139. Ertl, G., Weiss, M., and Lee, S. B., Chem. Phys. Lett. 60, 391 (1979).


53

Bare, S . R., Strongin, D. R., and Somojai, G. A., J. Phys. Chem. 90,4726 (1986). Zaera, F., and Somojai, G. A., J. Curd. 84, 375 (1983). Campbell, C T , and Koel, B. E., Surf. Sci. 186, 393 (1987). Campbell, C. T., Koel, B. E., and Daube, K. A,, J . Vuc. Sci. Techno/., A 5, 810 ( 1987). 144. Campbell, C. T., J. Phys. Chem. 89, 5789 (1985). 145. Campbell, C. T., and Daube, K. A., J. Curd. 106, 301 (1987). 146. Sachtler, W. M. H., CHEMTECH July, p. 434 (1983). 147. Sachtler, W. M. H., and van Santen, R. A.. Adu. Curd. 26,69 (1977). 148. Sinfelt, J. H., Sci. A m . 253(3), 90 (1985). 149. Sinfelt. J. H., and Cusumano. J. A., “Advanced Materials in Catalysis,” p. 1. Academic Press, New York, 1977. 150. Bauer, E . , in “Chemical Physics of Solid Surfaces and Heterogeneous Catalysis” (D. A. King and D. P. Woodruff, eds.), Vol. 3B, p. I . Elsevier, Amsterdam, 1984. 151. Biberian, J. P., and Somojai, G. A., J. Vuc. Sci. Techno/. 16, 2073 (1979). 152. Weissman-Wenocur, D. L., and Spicer. W. E., Surf. Sci. 133, 499 (1983). 153. Campbell, C. T., Surf. Sci. 167, L181 (1986). 154. Campbell, C. T., Paffett, M. T., and Voter, A. F., J . Vuc. Sci. Techno/.,A 4, 1342 (1986). 155. Sachtler, J. W. A., and Somojai, G. A., J. Curd. 81, 77 (1983). 156. Sachtler, J . W. A., and Somojai, G. A., J. Curd. 89, 35 (1984). 157. Sachtler, J. W. A., and Somojai, G. A., J . Card. 103, 208 (1987). 158. Campbell, C. T., Campbell, J. M., Dalton, P. J., Henn, F. C., Rodriguez, J. A., and Seimanides, S. G., J . Phys. Chem. (in press) (1989). 159. Peden, C . H. F., and Goodman, D. W., J . Curd. 104, 347 (1987). 160. Peden, C. H. F., and Goodman, D. W., J . Card. 100, 520 (1986). 161. Goodman, D. W., and Peden, C. H. F., 1nd. Eng. Chem. Fundum. 25, 58 (1986). 162. Goodman, D. W., Houston, J. E., and Peden, C. H. F., J. Vuc. Sci. Techno/., A 5 , 823 (1987). 163. Goodman, D. W., and Peden, C. H. F., J.C.S. Furuduy 1 8 3 , 1967 (1987). 164. Greenlief, C. M., Berlowitz. P. J., Goodman, D. W., and White, J. M., J . Phys. Chem. 91, 6669 (1987). 165. Tauster, S . J., A r c . Chem. Res. 20, 389 (1987). 166. Kao, C.-C., Tsai, S.-C., and Chung, Y.-W., J. Card. 73, 136 (1982). 167. Chung, Y.-W., Xiong, G., and Kao, C.-C., J. Curd. 85, 237 (1984). 168. Takatani, S., and Chung, Y.-W., Appl. Surf. Sci. 19, 341 (1984). 169. Demmin, R. A., KO, C. S., and Gorte, R. J., ACS Symp. Ser. No. 298, 48 (1986). 170. Demmin, R. A., and Gorte, R. J., J . Curd. 105, 373 (1987). 171. Demmin, R. A., KO, C. S., and Gorte, R. J., J. Phys. Chem. 89, 1151 (1985). 172. Demmin, R. A., and Gorte, R. J., J . Curd. 98, 577 (1986). 173. Levin, M. E., Salmeron, M., Bell, A. T., and Somojai, G. A., Surf. Sci. 169, 123 ( 1986). 174. Levin, M. E., Salmeron, M., Bell, A. T., and Somojai, G. A,, J. Card. 106, 401 140. 141. 142. 143.

(mn.

175. 176. 177. 178. 179. 180.

Tauster, S . J., Fung, S. C., and Garten, R. L., J . Am. Chem. Soc. 100, 170 (1978). Bischke, S . D., Goodman, D. W., and Falconer, J. L., Surf. Sci. 150, 351 (1985). Lee, Y. C., Tong, P., and Montano, P. A., Surf. Sci. 181, 559 (1987). Klier, K . , Adu. Curd 31, 243 (1982). Klier, K., Appl. Surf. Sci. 19, 267 (1984). Strongin, D. R., Bare, S. R., and Somojai, G. A., J . Cutal. 103, 289 (1987).

54 181. 182. 183. 184. 185. 186.

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Stoke, P., Phys. Scr. 36, 824 (1987). Stoke, P., and Norskov, J. K., J . Card. 110, 1 (1988). Strongin. D. R., and Somorjai, G . A., J . Caral. 109, 51 (1988). Sinfelt, J . H., J . Cural. 29, 308 (1973). Vannice, M. A.. and Twu. C. C., J . Catal. 82, 213 (1983). Levin, M. E.,Salmeron, M., Bell, A. T., and Somojai, G . A,, J.C.S. Faraduy I 83, 2061 (1987).


The Influence of Particle Size on the Catalytic Properties of Supported Metals MICHEL CHE Laboratoire de RPactivitP de Surface et Structure, UA 1106, CNRS Universiti Pierre et Marie Curie (Paris V I ) 75252 Paris Cedex 05, France

AND

CARROLL 0. BENNETT Department of Chemical Engineering University of Connecticut Storrs, Connecticut 06268

1.

Introduction

Metals constitute a wide class of catalysts, and because catalysis occurs on the surface, there is an economic incentive, especially for precious metals, to obtain catalysts in the form of small metal particles. This, however, raises two main problems. One is fundamental in nature and addresses the question as to below which particle size the metallic properties are lost. The other is more practical and concerns the preparation and characterization of very small particles and their catalytic activity. Let us consider a large metal particle that has electronic levels so close that they actually form bands. The spacing between adjacent levels is approximately expressed as (l a , b ): 6

z

&FIN

where eF is the Fermi level energy and N is the number of atoms in the particle. As the spacing between the levels becomes larger than the thermal energy kT, the levels begin to behave individually and the particle may lose its metallic properties. The critical size of the particle can be calculated by the above formula. At room temperature 6 = 2.5 x eV, 55 Copyright 0 1989 by Academic Press, Inc. All rights of reproduction in any form reserved.

56

MICHEL CHE A N D CARROLL 0. BENNETT

and with eFof the order of 10 eV, N is calculated to be approximately 400, which corresponds to a diameter of about 2 nm. This estimation is, of course, very approximate, but a lot of data confirm that the physical and catalytic properties start to change near this size and that catalysis is one of the most sensitive means to probe the surface of such particles. The other problem concerns the practical ways of producing metals with particle sizes around the critical value calculated above. The easiest and most practical way to achieve this goal is to deposit the metal in low concentration on a high-surface-area support. The main purpose of using a carrier is, of course, to achieve a high dispersion of the metal and to stabilize it against sintering. In a number of reactions, however, the metal support is not inert and the overall process is actually a combination of two functions: that of the metal and that of the catalytically active support. Thus, if an intrinsic size effect on catalytic properties is to be evidenced, much care should be taken to choose as much as possible a neutral nonactive support and to avoid using conditions that could possibly induce any metal-support interaction. These precautions are often not sufficient, and the preparation methods, as we shall see, are also likely to affect the final properties of the particles obtained. In the last decade, remarkable progress has been made in the preparation methods of supported metal catalysts. This has stimulated interest in the study of the relation between the catalytic behavior of metal particles and their mode of preparation with the aim of achieving the goal indicated above. Many workers have investigated the dependence of reaction rate on particle size. Early attempts to approach this problem were made by Kobosev and co-workers in the 1930s from the viewpoint of atomic dispersion and active ensembles. They studied the behavior of catalysts containing very small amounts of supported metal and were able to derive the number of atoms within the ensembles that were active for specific reactions (one atom for SO2 oxidation, two atoms for benzene hydrogenation, three atoms for ammonia synthesis, four atoms for acetylene oligomerization) (2a-c). These results as well as later ones have been reviewed by Gil’debrand (3). Boreskov et al. (4, 5 ) were the first to complete a systematic investigation of the relationship between particle size and catalytic activity, after their development of a technique for measuring the surface area of platinum catalysts by means of the selective chemisorption of hydrogen (6). They showed that the specific activity of platinum in the oxidation of sulfur dioxide ( 4 ) and of hydrogen (5) varied by less than one order of magnitude for catalyst samples differing in platinum surface area by four orders of magnitude. A few years later, Kobosev’s ideas were further

METALS: PARTICLE SIZE AND CATALYTIC PROPERTIES

57

investigated by Poltorak and Boronin on highly dispersed metals prepared by adsorption techniques (7u-d). Their studies led to many of the conclusions reached independently by Boudart et al. (8) for the hydrogenolysis of cyclopropane over supported Pt catalysts. One of the most important conclusions was that the rate per unit surface area of metal seemed for certain catalytic reactions very insensitive to the degree of dispersion of the metal. Boudart called these reactions facile (8)or structure-insensitive (9). The term structure refers to the coordination number of surface atoms that can be varied by exposing different crystallographic planes and making them imperfect by means of steps and kinks or by varying the particle size as can be done with supported metals (10). In the early work of Boreskov et ul. an apparent exception to this lack of sensitivity of activity to platinum particle size was observed for the H2DZexchange at low temperature (11). Other examples of the sensitivity of the specific activity to particle size were reported later on. Thus hydrogenolysis of ethane was found to depend on particle size for the Ni/AI2O3Si02 (12) and Rh/SiOz (13) systems. In the former case, the catalytic activity for unit surface area decreased for increasing particle size, whereas in the latter, an optimum particle size was associated with a maximum specific activity. These reactions were called “demanding” (8) or “structure-sensitive” (9) by Boudart, who drew attention to the difficulty of differentiating between metal-support interactions, the role of impurities, and mass-transfer limitations. This author proposed to select a molecule that can react along two parallel paths and to follow the change of selectivity with the particle size or the mode of preparation of the metal. The reaction of neopentane in the presence of hydrogen is a typical example (14). Reactions can now be classified into four categories (15) depending on how the turnover frequency TOF (the rate of the reaction expressed in moles per exposed metal atom and per unit time) varies as a function of the particle size d or of the fraction exposed (Fig. 1): the TOF of structure-insensitive or facile reactions does not depend on the particle size (curve 1). The TOF of structure-sensitive or demanding reactions may vary in two opposite ways: it may decrease when the particle size decreases, i.e., larger particles are more active than small ones; this is termed a negative particle size effect or antipathetic structure sensitivity (curve 2); or it may increase for decreasing particle sizes, i.e., smaller particles are more active than larger ones; this is called a positive particle size effect or sympathetic structure sensitivity (curve 4). The TOF may go through a maximum, for instance, if small particles exhibit a negative effect and larger ones a positive effect. In this case, those of an intermediate size will have maximum specific activity (curve 3). We have chosen

58

MICHEL C H E A N D CARROLL 0. BENNETT

d,nm

0

3

c I

Ln

LL

0

I-

9 ; 1

c

c

\

\

2

\

0

FE

1

FIG.1. Some of the ways in which turnover frequency has been found to vary with fraction exposed and particle size (see text for a discussion of the curves).

this numbering because it appears that only curves 1-3 can be reasonably understood as explained at the end of this review. The number of papers dealing with metal particles is extremely large: a computer survey based on volumes 66-101 of Chemical Abstracts gave the following numbers of references for the Group VIII metals. Fe, 213,818; Cu, 167,896;Ni, 140,550; Co, 83,827; Ag, 59,960; Pt, 39,268; Au, 36,444; Pd, 30,150; Rh, 14,761; Ru, 13,015; Ir, 7390; and Os, 4976. The same search restricted to particle size effects still gave 3,541 references. Our coverage had to be more limited and has been restricted to those papers in which curves of the types described above were presented. Unfortunately, the catalytic results are not always fully comparable. In fact, most works do not distinguish initial rate results (on clean surfaces) from steady-state rates (on semipoisoned surfaces), even though most refer to low conversions. Further, if a true particle size effect is to be evidenced, the metal particles must correspond as much as possible to the following criteria: they must be ( 1 ) monodisperse, (2) fully reduced, (3) unpoisoned (by liquid solvents or adventitious species), (4) unperturbed (by the carrier), and (5) unpromoted (by Ca, K, etc.). This goal is most clearly approached by methods based on vaporization of a metal by laser pulses, leading to a cloud of clean clusters of bare metal containing as few as one atom. Catalytic activity of clusters as a function of number of atoms can be obtained by time-of-flight mass spectrometry (16-18). Reviews have appeared in the past on particle size effects (3, 9, 10, 15, 19-24), but in the present one we attempt to give a comprehensive and


59

comparative study of particle size effects, particularly with regard to catalytic properties. Although this review is limited to the intrinsic particle size effect, as exhibited by either model or practical systems, it is clear that activity or selectivity in a useful catalyst may be altered favorably (or unfavorably) by support interaction (25a), addition of promoters (25b,c), and bifunctionality, matters that were discussed earlier. Particle size effects are also important in related fields such as electrocatalysis (25d-31). For example, the reduction of dioxygen over Pt/C shows antipathetic structure sensitivity, leading to an optimal particle size of about 4 nm for a working electrode (31). However, work on electrochemistry is not included in this review. II. Preparation of Supported Metal Particles

A. INTRODUCTION Recent years have witnessed marked progress in the preparation of small metal particles. This has been achieved by the choice of a suitable support, the selection of the appropriate preparation method, or the combination of both. The selection of the carrier is relatively simple. It may be imposed by the type of reaction to be promoted. For instance, if the latter requires a bifunctional catalyst (metal + acid functions), acidic supports such as silica-aluminas, zeolites, or chlorinated aluminas, will be used. On the other hand, if the reaction occurs only on the metal, a more inert support such as silica will be used. In certain cases, other requirements (shock resistance, thermal conductivity, crush resistance, and flow characteristics) may dominate and structural supports (monoliths) have to be used. For the purpose of obtaining small metal particles, the use of zeolites has turned out to be an effective means to control their size. However, the problem of accessibility and acidity appearing on reduction may mask the evidence of the effect of metal particle size on the catalytic properties. The choice of the method is more complicated and depends on whether an industrial or a model catalyst is prepared. The most common industrial preparation methods involve several steps: (1) the distribution of a precursor compound over a carrier, (2) the drying and calcining of the resulting solid, and (3) the reduction of the precursor compound, which may, after steps 1 and/or 2, no longer be identical to the initial precursor, to obtain the active final metallic state. Rather than describe the numerous methods available in the literature, we shall in the following emphasize the advantages and disadvantages of each method with respect to the goal

60


to be achieved (i.e., the preparation of monodispersed metal particles). The design and preparation of supported catalysts have been the subject of several recent reviews (Ib, 3, 32-39, 40a-c).

B. THEVARIOUSPREPARATION METHODSFOR DEPOSITING THE METALPRECURSOR OR THE METAL This review begins with the more complex methods and moves on to the simpler ones. In this context, complexity is defined as the number of species involved before finally obtaining the highly dispersed metal particles. Table I summarizes the principal features of the different methods used to prepare the metal particles on a support, which is the common feature of all the methods and thus is not included in the table. They are subdivided into groups when the principles are similar. We could have ordered the preparation methods differently, for instance, using a chemical approach starting from the oxidation state and nuclearity of the metal in the precursor compound, but this would not have served our purpose, specifically, to show the parameters that can possibly influence the behavior of the final metallic state. 1.

Coprecipitation

Most methods deal with the formation of metal particles on a support that is preformed since this leads to simpler preparation processes. There is an important route, however, typically used for metal-SiOz and metalA1203 catalysts, which involves (Table I) the coprecipitation in a precursor form (hydroxides, nitrates, carbonates, silicates, etc.) of both the support and the active phase from a solution (37a,b, 38, 41). The advantage is to produce an intimate mixing of metal precursor and support. The precipitate leads on calcination to a support with the active component dispersed throughout the bulk as well as at the surface. After reduction to the final catalyst, it is difficult to obtain metal crystallites of uniform size (42,43) because of the presence of both the oxides (of the support and of the active metal) and other intermediate compounds [e.g., nickel aluminate or silicate for the Ni/A1203 (42) and Ni/Si02 (43) systems, respectively] that have different reducibilities. The coprecipitation can be improved in various ways so as to avoid any heterogeneity during the process. This is achieved in the superhomogeneous coprecipitation method, which consists of two steps. In the first, layers of the salt solutions are superposed and then mixed into a supersaturated solution instantaneously. The second step consists of forming a

TABLE I The Various Preparation Methods of Supported Metal Particles Further activation steps" Washing, hydrolysis, Solvent

Metal precursor

Inorganic

Salt

Inorganic Inorganic Inorganic

Solvated cation Complex anion Complex cation

Inorganic

Complex cation

Inorganic

Complex anion

organic

Metal cluster compound

Organic Organic

-

-

Metal cluster compound (solid or vapor) Metal colloid

Other species present

or Example

evacuation

Support precursor Precipitating agent Counteranion Countercation Counteranion Counteranion Competing cation Countercation Competing anion

-

Pd colloids/Si02

Atoms (metal vapor) Ions Atoms (metal vapor)

+ or - indicates whether or not the corresponding activation step has been performed. The method number, as referred to in the text, is indicated in parenthesis. TMI, transition metal ion.

Drying

Calcination

Reduction

Methodb

+

+

Coprecipitation (1)

+ + +

+

+ +

Impregnation (2)

+

+

+

+

Simple ion exchange (3) Competitive ion exchange (3)

+

+

-

+

+

-

+

-

-

-

-

Decomposition of metal cluster compound (4)

Chemical deposition ( 5 )

Ion implantation (6) Vapor-phase deposition (7)

62


homogeneous precipitate from this supersaturated solution ( 3 7 ~ )Some . authors produce the coprecipitation of the components in an organic solvent that is then evacuated under supercritical conditions. This leads to particularly high-surface-area materials (44, 45). The disadvantages of the coprecipitation method are that the precursor ions are also distributed within the bulk of the support oxide and that the pore structure of the final catalyst is more difficult to control than when one starts from a separately produced carrier. The next type of preparation method avoids these difficulties. 2. Impregnation

This method is begun by impregnating a preformed support with a solution of a metal salt. The solvent is removed by evaporation during drying (40a,b). When the amount of solution corresponds to the pore volume of the support, the method is more precisely referred to as the incipient-wetness impregnation. This method is used when there is little or no interaction between the precursor and the support. The advantages are technical simplicity, low cost, and reproducible metal loadings. If the metal content is limited by the solubility of the metal compound, multiple impregnation steps can be used. When the pores contain only air, the impregnation is referred to as capillary, whereas when they are filled by a solvent, the impregnation is called diffusional (46a). In some cases, precipitation of an active precursor onto a separately prepared carrier is also used, and this is known as the deposition-precipitation method (46b). Here the active precursor is precipitated onto the support by means of a chemical reaction, whereas in the classical impregnation method, it is “physically” deposited onto the carrier through removal of the solvent as described above. Recently, a new type of “dry” impregnation has been presented (46c,d),which consists of heating a mechanical mixture of the active precursor and the carrier. This leads to the spontaneous dispersion of the active precursor onto the support. This solid-solid adsorption is a spontaneous entropydriven process since, in contrast to what happens in gas-solid or liquidsolid adsorption, there is a change from an ordered (three-dimensional)to a less ordered (two-dimensional) structure of the active precursor, and this is a process of increasing entropy (46e). This type of solid-solid adsorption will be the subject of a forthcoming review (46f). In the case of impregnation, the subsequent step of drying (40a,b)is an important operation in the distribution of the active phase at the surface of the support. When a pellet with a uniform pore system is heated at a fast rate, a temperature gradient is established between its external surface

METALS: PARTICLE SIZE A N D CATALYTIC PROPERTIES

63

and its interior. Evaporation starts at the external surface, and the gasliquid interface moves toward the interior. The precursor concentration increases at the menisci, and the active phase is deposited on the pore walls. After reduction, the metal particles will be rather homogeneous in size. By contrast, if the drying rate is low, the active phase, which interacts little with the support, remains in the liquid phase so that the concentration in the active phase steadily increases. When the drying is finished, the active phase is found at the bottom of the pores. After reduction, the metal particles obtained will, therefore, depend on the volumes of the pores (47). If the latter are not uniform in size, the situation is more complex and there will be a correspondingly broad distribution of particle sizes in the finished catalysts. It should be noted that there is no washing step after impregnation (Table I) because this method involves little or no interaction between the precursor and the support. Thus, any washing would eliminate the precursor phase. During the calcination step, occurring between drying and reduction, two phenomena generally occur: (1) the precursor oxide is formed by decomposition of the impregnated salt, and (2) some chemical bonding is established between the precursor oxide and the support (48). As seen in Table I, the impregnation method may involve a precursor in the form of either a complex cation or a complex anion, with the disadvantage that counterions are always present. During the calcination step, such ions may lead to uncontrolled side reactions producing strongly held species, such as nitrogen-containing species (NO, N02, NO;) for samples prepared from nitrates (49) or ammonium compounds (50). The next preparation method avoids this problem to a large extent.

3. Zon Exchange This technique is used when the precursor interacts with the support. The interactions are basically controlled (52) by (1) the type of support and state of the surface (number and nature of the functional groups, their acid-base properties) and (2) the impregnating solution (pH, type and concentration of the metal precursor, and presence of competing ions). In this method, also called dipping impregnation, wet soaking impregnation, impregnation with adsorption, or adsorption from solution, the support is immersed in a solution of the metal compound. The slurry is stirred for a given time, filtered, and possibly washed. The resulting product is then dried. The concentration of the precursor solution may be readjusted and recycled. This technique is widely used in the preparation of laboratory and industrial catalysts. Depending on the strength and conditions of adsorption of the precursor species, the concentration of the

64


active phase may be varied. This method thus leads to a greater control over the dispersion and distribution of the active species within a catalyst pellet. Because the precursor interacts with the support, it is not eliminated easily by washing. The metal loading is governed by the concentration of adsorption sites on the support oxide. As a consequence, the weight of the active component that can be incorporated is limited. One can increase the metal loading by successive impregnations, but this is not recommended if a careful control of the physical parameters is to be achieved. Another disadvantage is the high cost of the process because of lower productivity and complex precursor solution recycling in the case of costly precursors. a. Surface Charge of Oxides in Suspension. Most oxides are amphoteric and when immersed in an aqueous solution present surface charges that are controlled by the equilibria S-OH;

tH+ -OH7 S-OH = +OH- S-0-

+ H20

where S-OH represents a surface adsorption site. The addition of an acid shifts the equilibrium to the left, increasing the number of positive charges on the surface: the oxide behaves as an anion exchanger. Conversely, the addition of a base shifts the equilibrium to the right, increasing the number of negative charges: the oxide behaves as a cation exchanger. When the number of positive and negative charges at the oxide surface is equal, the corresponding pH is the isoelectric point of the solid (IEPS) or the zero point of charge (ZPC). The isoelectric point can be determined by electrophoresis measurements. The, values of the IEPS, which are known for a number of oxides, reflect the chemical composition of the solid and the electrolyte in which it is immersed (52, 53). Structural or adsorbed anionic impurities shift the IEPS to more acidic values, and cationic impurities shift it to more basic values or toward the IEPS characteristic of the impurity oxide (53). Above the IEPS value, the oxides will behave as cation exchangers and as anion exchangers below that value. The further one gets from this value, the larger the surface charge and the exchange capacity (54). b. Simple Zon Exchange. Among the supports, zeolites, Si02-AI203, and Si02 adsorb cations; A1203,Ti02,and Cr2O3, which are amphoteric, adsorb anions in acidic and cations in basic solutions; MgO, La203,and ZnO adsorb only anions (36, 41, 53). The adsorption equilibria can be represented by the following equations (36): S-OH

+ C+ =S-0-C+ + H+

cation adsorption


S-OH

+ A- + H+ =S-0H;A-

65

anion adsorption

A very interesting feature of the ion-exchange method is that the counterion can be easily removed by washing (Table I) because it is repelled from the surface that possesses charges of the same sign. This is an important advantage over the impregnation method, where washing cannot be used because of the weak interaction between the precursor and the surface. In the ion-exchange method, the cations C+ or the anions A- tend to saturate the first adsorption sites so that most of the compound adsorbs near the pore mouth, and a large concentration gradient develops within the pellet pore. There are several ways to obtain a more uniform distribution: (1) a large supply of compound is used so as to saturate every adsorption site, (2) the support may be left for a long time in contact with the solution, and (3) competing ions are added to the solution, which will adsorb on the same sites as the ions of the metal precursor. This is the subject of the method discussed next. c . Competitive Zon Exchange. This method, also called chromarographic adsorption, involves the competition between two types of ion for the same adsorption sites (Table I). The competitive exchange has long been known for anions (55), whereas the competitive cation exchange has been developed more recently (56). The competitive ion-exchange equilibria can be represented by the following: S-0-C+

S-0H;A-

+ D+ =S-O-D+ + C+ + B- =S-OH;B- + A-

competitive cation exchange competitive anion exchange

A typical example is illustrated by the following reactions: Si-O--H+ 2Si-0-NH:

+ NH3

+ [Pt(NH3)4]2+

Si-0-NH; (Si-0-)2[Pt(NH3)4]2+

+ 2NH:

where the NH: and [Pt(NH3)4]2+ ions are competing for the same surface Si-0- sites. By varying the concentration of NH: competing ions, it is possible to shift the equilibrium of the second reaction presented in the typical example above and to control the deposition of the precursor metal species onto the oxide surface to obtain a homogeneous distribution. The adsorption of the precursor metal species by oxide surfaces can be accompanied by a ligand reaction (57a) between a surface site and a ligand on the metal complex, with the result that the surface becomes a mono- or polydentate ligand within the metal ion coordination sphere (57b-e):

66


(Si-O-)2[Ni(NH,),]2+

=[(Si-O)2Ni(NH3)4] + 2NH3

This stronger interaction between the metal precursor and the surface and its homogeneous distribution leads, after reduction, to highly dispersed metal particles. Very often, though, because of their strong interaction with the surface, the precursor ions are difficult to reduce at the temperatures where metal aggregation is not important. The final state thus consists of highly dispersed metal particles anchored onto the surface by means of a “chemical glue” composed of partially reduced or unreduced ions (58). In order to avoid this problem of incomplete reduction, which becomes all the more important as the particle size decreases (59), a number of attempts have been made to bypass the reduction step and to deposit the metal directly in its zero-valent state. There are different methods for this, and they are discussed separately below. 4.

Decomposition of Metal Cluster Compounds

Ever since Parkyns (60) prepared nickel particles on oxide (A1203,Si02) surfaces by decomposition of nickel carbonyl, metal cluster compounds (alkoxides, carbonyls, organometallics, etc.) have been used increasingly for the production of laboratory metal catalysts (Table I), and several reviews have appeared on this subject (61-65a,b). Metal cluster compounds were initially employed for the two following purposes. The first concerned the formation of heterogeneous analogs of homogeneous active catalysts, and the second dealt with the production of small metal particles. The major advantage of the latter approach is that the temperatures usually required to decompose the metal cluster compounds and produce small metal particles are substantially lower than those observed for the reduction of metal salts (nitrates, chlorides) involved in conventional methods so that the number of atoms initially present in the complex can be preserved, e.g., Ni3 if one starts from trinuclear nickel clusters (66). However, the results using this approach have not been always as promising as originally believed (64). Metal cluster compounds are usually anchored onto an oxide surface via the two following routes: 1.

reaction involving surface hydroxyl groups S-OH: nS-OH

+ MX,

[S-O].MX,-,

+ nHX

The metal content is controlled by the concentration of surface S-OH groups, the stoichiometry of the reaction, and the number of metal atoms within the cluster.


67

2. reaction first to produce functional anchoring groups at the oxide support surface and then to anchor the metal cluster compound [A number of functions can be selected as well as their further reactions, and this subject has been extensively reviewed (62)l.

From recent literature data, there are many cases in which the resulting surface products, after anchoring, are not in a zero-valent state but rather form a mixture of species so that the metal is both in zero-valent and partially oxidized states. Most studies on functionalized supports have been made on hydroxylated oxide surfaces. The metal particles produced on such supports are not of an unusually high degree of dispersion. The nature of the decomposition process on dehydroxylated supports has not been as thoroughly studied as on hydroxylated supports. Preliminary results, however, indicate that formation of highly dispersed zero-valent metal particles may be obtained (64). The other disadvantage of the metal cluster compound decomposition is that the ligands, which have to be removed to produce the metal particle, and the organic solvent used to impregnate the support by the precursor compound may remain on the surface and thus interfere later in the catalytic reaction. The problem of the solvent can sometimes be avoided by using volatile metal cluster compounds (6%). The next method based on depositing colloidal metal particles onto a support removes the problem of incomplete reduction, i.e., the presence of the mixture of reduced and unreduced species, but does not remove that of the solvent. 5 . Chemical Deposition from a Metal Colloid Dispersion

In this method, the preformed metal is adsorbed as individual particles onto a carrier, both suspended in a solvent (Table I) (67,68).The size and structure of the colloids are not influenced by the support (69), but the liquid medium may contain species that are by-products of the preparation and that may complicate the situation unless removed by dialysis. In various attempts to use this method the metal particle size, which depends on the metal and the preparation conditions, is usually larger than that obtained by the preceding methods. A more sophisticated technique, the so-called solvated metal atom dispersion technique, has been developed by Klabunde and Tanaka (70). It involves the evaporation of metal atoms in the presence of complexing solvents to yield “solvated” metal atoms. The solution is brought into contact with the support at low temperatures and the slurry obtained is warmed to room temperature. The excess solvent is removed by further

68

MlCHEL CHE A N D CARROLL 0. B E N N E T T

evacuation. Although attractive, this method presents the disadvantage that the small crystallites are surrounded by organic fragments. Those fragments may affect the catalytic and magnetic properties of the metal particles (71). Several solutions have been proposed to avoid these organic fragments, and we shall review two of those, which can, however, produce only model catalysts. In these methods, only the metal is deposited on the support; one concerns the deposition of the metal in its oxidized state and the other, in its zero-valent state. 6 . Ion Implantation

The principle of this method is to select by mass spectroscopy a positively charged ion, for example l9W+,which is implanted by ion bombardment of a given support (Table I) (72). The ion-implantation technique is simpler when single crystals are used, for instance, Mg(100) (73). Since electrons are readily available in the ion implanter, the ion is quickly neutralized and the neutralized form of the element implanted can be finally found either as a gas (e.g., Xe) or as a metal (e.g., Pt) in the outer layers of the oxide exposed to ion implantation (74). The advantage of this method is that, in contrast to all previous methods, it involves only the metal and the support and that metal-support interactions concern only a single face of the support (75). The disadvantage is that the metal is not located only at the oxide surface but also within the oxide and the concentration profile is critically dependent on the energy of the implanted ions. The other problem is that a number of defects are produced during the bombardment that may interfere later in catalysis. However, this effect can be evaluated (75) if ions of similar weight but that are catalytically inactive (e.g., Xe') are implanted under the same conditions (energy, angle, and time of ion implantation). The next method avoids the problems caused by the implantation profile and the presence of defects.

7 . Metal Vapor Deposition This method (Table I), like the preceding one, eliminates the problem of any spurious species (counterion, ligand, solvent, etc.) and has the advantage that it does not alter physically the support surface. There are a number of other advantages, which can be listed as follows: 1. When they come to land on the surface, the atoms are in their zerovalent state, except for those atoms that may be reoxidized by surface OH groups by reaction such as (61) Ni" + OH =Ni2+ + 0'-+ 1H2

METALS: PARTICLE SIZE

A N D CATALYTIC PROPERTIES

69

2. There is thus no thermal reduction step involved, and this suppresses both any sintering process and the formation of some strongly adsorbed hydrogen phase, which, for small particles, can decrease or even eliminate the catalytic activity (58, 76, 77). The “reduction” degree is thus much larger than that obtained for similar sizes with conventional methods. 3. The dispersion can be controlled by monitoring only the time of evaporation (78). For small times of evaporation, atomic dispersion can be obtained all the more easily because the system does not need to be heated as in other methods for thermal reduction. The surface atom diffusion is thus drastically slowed down. 4. Finally, heterogeneous nucleation theory can be applied (79), and it affords an independent way of estimating the particle size, particularly difficult if not impossible to measure close to atomic dispersion by conventional methods because of the low metal content of such model catalysts.

There are, of course, some limitations. An obvious one is that this method cannot be applied yet to the preparation of industrial catalysts, another one is its cost because ultra-high-vacuum (UHV) equipment is required. This drawback explains why this method is usually coupled to surface techniques such as XPS, UPS, RHEED, and AES, which also require UHV. The last disadvantage is that the best suited supports are those that are flat, i.e., oxide single-crystal faces, oxides produced by oxidation of a metal single crystal, or compressed powder oxides. There have been several examples where the preparation chamber also serves as sample chamber for surface techniques and is coupled to a catalytic reactor. Whereas there are a number of works using this approach for bulk metals (80),there are, by contrast, few studies dealing with metals supported on either single crystals (81) or polycrystalline supports (78,79,82, 83). The latter type of system appears to be the model catalysts closest to the real catalyst.

C. FURTHER ACTIVATION STEPS The preceding paragraphs have introduced the various preparation methods leading eventually to supported metal particles. These methods fall obviously into two categories depending on whether the metal is basically in its zerovalent state (decomposition of metal cluster compound, chemical deposition, ion implantation, and vapor-phase deposition) or in an oxidized state (coprecipitation, impregnation, and ion exchange)(Table I).

70


For the latter methods, it is necessary to reduce the ions to obtain the metal. In some cases, a calcination is performed before the reduction. We shall now concentrate on these two stages.

I. Calcination Very often, calcination is carried out in air or oxygen and leads to the following transformations: 1. formation of the oxide system on decomposition of the precursor compound and subsequent bonding of the formed oxide with the support (typically observed for catalysts prepared by impregnation), 2. removal of some of the elements introduced during the preparation by formation of volatile compounds ( S 0 2 , NO2, COZ, etc.), 3. decomposition of the precursor ionic complex and further ligand exchange reactions between surface groups (OH, 0 2 -and ) ligands bound to the metal ion, resulting in stronger precursor-support interactions (typical of catalysts prepared by ion exchange), 4. elimination of carbonaceous impurities possibly introduced during the preparation, and 5 . sintering of the precursor compound or of the formed oxide system.

Consequently, calcination may have a pronounced effect on the reducibility, dispersion, and distribution of the metal in the final catalyst. For instance, high-temperature treatment may lead to very stable solid solutions or to large oxide crystals, particularly for catalysts prepared by impregnation. In the case of solid solutions, the metal is difficult to reduce because of strong precursor-support interactions. This can be overcome by higher reduction temperatures. Because of these strong precursorsupport interactions, high dispersion of the metal can be achieved in the final catalysts (84). By contrast, large oxide crystals will lead, after reduction, to low dispersions of the metal. For catalysts prepared by impregnation, it is often observed that the final particle size distribution reflects the different types of precursor-support bonding obtained after calcination, and this depends on the metal loading. Thus, for Ir-Ti02 systems prepared by impregnation (83,it has been observed that two types of particles were formed after calcination in air and reduction by hydrogen. One corresponded to small particles issued from precursors ion exchanged on the support surface and in amount never in excess of the monolayer; the other corresponded to large particles issued from precursors impregnating the support and observed for metal loading exceeding the monolayer. There have been relatively few studies on the influence of treatment


71

conditions on metal dispersion, and most concern platinum. Kubo et ul. (86) reported that the temperature of calcination in air before hydrogen reduction was the key parameter determining the platinum dispersion in RNaY and RNH4Y zeolites. The best dispersion was obtained for air calcination at 300°C. In many cases, the calcination in air is used to carefully control the decomposition of ion complexes used as precursors and to develop strong interactions with the support so that after reduction, small metal particles are formed (58, 87u,b).

2. Reduction The conditions of reduction are particularly important since they can strongly affect the final distribution of the metal particles (88). In studies of nucleation and film growth from metal vapor, it is generally observed that the crystallite size is determined by the saturation density of nuclei, which is found to increase with decreasing substrate temperature due to shorter adatom migration distances (89).In the case of nucleation in metal oxide reduction, an opposite trend is found. Higher temperatures will favor more rapid initial reduction of the metal oxide and higher density of nuclei but also coalescence or sintering of particles. An intermediate temperature must be found that leads to the maximum nuclei concentration without any appreciable coalescence. The parameters probably effecting this compromise are the presence of nucleation sites that may correspond to surface defects (90) and the presence of water vapor, which usually leads to an increased mass transfer (92).It is thus anticipated (88)that the smaller particles should be obtained (1) by using starting materials that will maximize the defect concentration in the final support, (2) thoroughly drying the supports at the lower temperatures to avoid changing the defect concentration, and (3) using a high flow rate of hydrogen in the reduction process while keeping the temperature low. This high flow rate of hydrogen lowers the partial pressure of water vapor and increases the rate of reduction at lower temperatures (92).Once nucleation has started, hydrogen can dissociatively chemisorb on the metal nuclei that are formed. The hydrogen atoms are mobile and can migrate to the metal particle-metal oxide interface and be used in reduction of the oxide. This causes a much more rapid reduction since hydrogen atoms react with either the metal ion (93) or the oxide ion 02-(94) more readily than does molecular hydrogen at the same temperature. The technique of hydrogen atom beams produced by a microwave discharge has been used to produce small metal particles (58, 93). The microwave discharge technique has also been used to decompose metal carbonyl compounds and small metal particles have been obtained (95).

72 111.


Background for Evaluating the Catalytic Properties of Small Supported Particles

A. DEFINITIONS In his interesting article on particle size effects, Bond (21) gives the results of model calculations based on spheres and on cubes with five faces exposed, for Ni, Pt, and Pd. For these simple systems, the fraction of total metal atoms exposed, FE, is related to the diameter or length of a side d (in the figures, d is written with a bar above it to highlight the fact that it is an average value) by the approximate formula Bld,

FE

= 11.0,

d > 1 nm d < 1 nm

where d is in nanometers and B is about 1 nm. In discussing the literature on structure sensitivity we shall use B = 1 .O nm in Eq. ( I ) , realizing that B varies from one metal to another or as the morphology changes with d for a given metal. We define the atomic rate of reaction AR as the rate of production of a given product, with the units S - I , or moles of product per mole of total metal atoms per second. This is the rate that we strive to maximize for economic reasons, as far as catalyst development is concerned. However, it is convenient to decompose the atomic rate AR into two factors of a more fundamental nature:

AR = (TOF)(FE) (2) where TOF is the turnover frequency, s - I (per second), a rate expressed as moles of product per mole of surface metal atoms per second. In general, TOF is a function of d ; such systems are structure-sensitive. For certain systems TOF is not a function of d , and these cases are called structure-insensitive. In this work we review the behavior of TOF(d) for various metals and supports and for various reactions. The essence of this behavior was presented elegantly by Boudart (10). However, it is now clear that other factors in addition to d commonly affect the turnover frequency. We prefer the term fraction-exposed (96) to dispersion, since the latter conveys an idea of distribution in space that is not directly related to particle size and shape. The quantity FE is commonly measured by hydrogen chemisorption at room temperature, using the equation H/M FE = H/M,

(3)


73

The quantity H/M, i.e., H atoms irreversibly adsorbed divided by total metal atoms, is measured experimentally. H/M,, where M, refers to surface metal atoms, may be known on the basis of surface science studies [e.g., low-energy electron diffraction (LEED)] and is usually assumed equal to unity. However, we shall discuss some papers that show that H/M, is also a function of d and other variables. In addition, reversibly adsorbed hydrogen may also be important (97,98). For interpreting most data, nevertheless, we can only assume that FE = HIM. It seems reasonable to use surface metal atoms as a basis, and not surface sites. The number of surface atoms required to form a site is in many cases a subject of debate, whereas H/M is measured experimentally. Turnover frequency is preferred in the IUPAC report (96) to turnover number TON, for it has the correct units. A number implies a dimensionless quantity, whose magnitude is independent of the system of units. TON also seems to suggest an aura of a physical constant, referring to a property of a particular catalyst system. Since TOF is a rate, a function of temperature, pressure, and fluid composition in addition to d, etc., turnover rate had been suggested as a more appropriate name (99). These ideas are clearly discussed by Boudart and Djega-Mariadassou (100, 101), where rate is expressed as vitesse de rotation or turnover rate. We have defined TOF in terms of moles of a certain compound, relating to the idea of a yield. Thus, for methanation, we have TOFco and TOFcH,, etc. Ichikawa et al. ( 1 0 2 ~have ) used the term site-timeyield (STY), but we do not propose such a departure from ingrained habits. Often rates are expressed per unit mass of supported catalyst, and we can define the activity A, moles of product per second per kilogram of supported catalyst, as A = (AR)(C)/AW

(4)

where C is the catalyst loading, kilogram (metal)/kilogram (metal + support), and AW is the atomic weight of the metal. In line with the usual meaning of “specific,” A might be called a specific rate, but this term has, unfortunately, become associated with a rate per unit area of metal. This term seems unnecessary, because for well-defined surfaces TOF is interpreted as molecules reacted per surface metal atoms per second. Because of these difficulties, the latest version of the IUPAC manual (102b) expresses activity as mass-specific rate, rate per unit area as area-specific rate, and volumetric rate (see below) as volume-specific rate. For reactor design, the usual volumetric rate is convenient, written as r = ApB

(5)

74


where r is moles per second per cubic meter of bed and p~ is the bed density expressed in kilograms of catalyst per cubic meter of bed. We have taken pains to present these definitions because it seems that the organization and understanding of catalytic results is hampered by lack of agreement on these simple details. Rates may be expressed as turnover frequency, atomic rate, activity, or volumetric rate, provided the authors determine also H/M, C, and p ~ and , defend their choice of H/M,. Of course, d should be measured independently if possible, as should the morphology of the metal particles. We emphasize the simplicity of Eqs. (1-5). It would seem desirable to express all kinetic results in terms of these rates. From a fundamental point of view, results on TOF(d) furnish important data for testing models of catalysis. We shall discuss how such results may be explained in terms of the geometric and the electronic properties of supported metals. From a practical point of view, the form of Eq. (2) shows that there may be a value of d that maximizes the value of AR, depending on how TOF varies with d.

B. POSSIBLE EXPLANATIONS OF PARTICLE SIZEEFFECTS 1. Electronic Effects

As the number of atoms in an isolated metal particle is reduced, the differences in energy levels of the valence electrons will eventually become appreciable with respect to kT. In the limit the discrete levels associated with an isolated atom would prevail. Thus, we expect a variation in the catalytic properties of small particles because of this electronic effect, which applies to all the atoms of the particle. However, the proportion of the total atoms present on the surface and thus in various states of coordination with respect to the other metal atoms will also change. The electronic configuration of these surface atoms is changed, leading to changes in the catalytic properties. This second electronic effect may be viewed as having a geometric origin, so that the usual division of size effects into electronic and geometric is somewhat arbitrary. Some attempts have been made to measure the electronic properties of small particles by X-ray photoelectron spectroscopy (XPS). The preparation of samples of isolated small metal particles is not easy. The most successful methods are either vapor deposition of noble metals (Pt or Pd) on carbon or silica, or ion exchange used to prepare metals in Y zeolite. For the noble metals and inert supports used, it is assumed that the metal particles are isolated from each other. The results of several studies are presented in Fig. 2, in which the


1 I

3.0

10

i.nm L O 2.0

1.0

7s

0.L

ZD > al

Y

m

a in

FIG. 2. Binding energy shifts ABE of metal core levels measured-by photoelectron spectroscopy (XPS) versus fraction exposed F E and mean particle size d for various supported metal systems (see Table I1 for details of the studies).

TABLE I1 Details of the Supported Metal Systems Studies Presented in Fig. 2 Curve (Fig. 2)

Metal and electron

BEu for bulk (eV) 335.0 335.0 335.0 71.5 314.5 314.0 71.3 314.6

~~~

support

Referenceb

C Y-Zeolite Si02 Y-Zeolite C A1203

Si02 A1203

~~

BE = binding energy; ABE = BE(observed) - BE(bu1k). References: (a) Takasu, Y., Akimaru, T., Kasahara, K., and Matsuda, Y., J . Am. Chem. SOC. 104,5249 (1982); (b) Vedrine, J. C., Dufaux, M., Naccache, C., and Imelik, B., J.C.S. Faraday f74,440 (1978); (c) Takasu, Y., Unwin, R., Tesche, B., Bradshaw, A. M., and Grunze, M., Surf. Sci. 77,219 (1978); (d) Mason, M. G., Gerenser, L. J., and Lee, S.-T., Phys. Rev. Lett. 39, 288 (1977); (e) Huizinga, T., van? Blick, H. F. J., Vis. J. C., and Prins, R., Surf. Sci. 135, 580 (1983); (f) Masson, A,, Bellamy, B., Colomer, G., M'bedi, M., Rabette, P., and Che, M., Proc. f n t . Congr. Catal., 8th 4, 333 (1984); (g) Masson, A., Bellarny, B., Hadj Romdhane, Y., Che, M., Roulet, H., and Dufour, G., Su$. Sci. 173,479 a

(1986).

76


appropriate electron energies are shown to change with metal particle size. It appears that for particles of diameter greater than about 5 nm, the energy levels are those of the bulk metal. For particles of 5 nm, the fraction exposed is about 0.2, and the number of atoms varies from 3000 to 6000, depending on the metal and its morphology. For Pd, the energies stop changing below about 1.O nm, where all the atoms are exposed. For Pt, the changes continue until even smaller particles are reached. The three energies of Fig. 2 in the small-particle limit for Pd show considerable differences. Is this so because of the three different supports used? This explanation would imply that the particles are not really isolated from their supports. 2. Geometric Effects

a . Mathematical Models Based on Regular Small Crystallites. van Hardeveld and Hartog (20) have published the results of many calculations describing the arrangements of the atoms in small metal particles. Many graphs and sketches of small crystallites are presented. As an example, we show in Fig. 3a an octahedron made up of atoms in the facecentered cubic (fcc) system. The three types of surface atom are indicated by the symbols C9, face atoms; C7, edge atoms; and C4,corner atoms. A bulk atom is Clz. The subscripts on C indicate the coordination of a particular atom. The superscripts of Cj refer to a system for identifying which of the 8 atoms are missing in C4, for example. The original work should be consulted for details. For the various possible small octahedrons, the ratio N(C,)IN, has been calculated (20), where N(Cj) is the number of atoms of coordination Cj and N, is the number of surface atoms. These quantities vary with the total number of atoms, which, in turn, can be related to d for a given metal. Figure 3b shows, as a function of d for Ni, the fraction of the surface atoms that are made up of each of the three typical atoms of Fig. 3a. The fraction of face atoms C9 in Fig. 3b becomes about 0.9 at about 6 nm and falls to zero at about 0.6 nm, where all the atoms are at edges and corners. The effect of surface atom coordination makes itself felt in general for particles smaller than 4 nm. These changes are, of course, linked to changes in electronic characteristics also. Supposedly the catalytic properties of the metal are linked to N(Cj)IN, of Fig. 3b, and it is implied that one-atom sites are involved. For example, if catalysis is favored by atoms of low coordination, small particles would give higher rates. If catalysis is favored by face atoms, large particles

77

b

I

1

10

4.0- 2.0 1.0

0.L

d,nrn(Ni)

FIG.3. (a) Face-centered cubic (fcc) octahedron, m = 9 ( m is the number of atoms alon_g an edge). From Ref. 20. (b) Coordination of surface atoms versus the mean particle size d for nickel, assuming that the particle is a face-centered cubic (fcc) octahedron. The symbols cprt,. refer to atoms identified in (a).

would give higher rates (10). The former case corresponds to sympathetic structure sensitivity and the latter, to antipathetic structure sensitivity. Since ordinary supported catalysts exhibit a distribution of crystal sizes, it is clear that the effects may be more diffuse than those suggested by Fig. 3b. This distribution can be measured and controlled in model catalysts (79).

78


As a model of a small crystallite is built up by adding atoms one by one (103), a complete octahedron or other regular shape is reached only periodically ( 1 0 4 , for instance, at particles with a number of atoms of 6, 19,

44,etc. (20) for an octahedron. Adding the 20th, 21st, etc. atom starts a new layer and creates atoms of different coordination. At the edges of added layers, a particularly interesting ensemble of five atoms is formed, called a Bs site (20). Crystallites are considered to be fcc octahedra with extra layers of atoms on each face. If the extra layers are added in such a way as to produce the maximum total number of BS sites, the metallic particle illustrated in Fig. 4a is obtained. There are now some atoms in coordination states that did not exist in Fig. 3a. Figure 4b is based on calculations for a shape similar to that in Fig. 4a (20). For this more complicated arrangement, the fraction of face atoms CS falls to 0.9 at about 20 nm and goes to zero at about d = 2.5 nm. Geometric effects thus persist to higher d than in Fig. 3b. It is probable that catalysis for certain reactions requires a site made up of a particular ensemble of surface atoms. The number of adjacent atoms’ required may increase from 2 for the dissociative adsorption of oxygen or hydrogen to something like 12 for ethane hydrogenolysis (105). The Bs sites seem to have particular importance in catalysis, and it is claimed that they are necessary for the appearance of infrared-active adsorption of nitrogen (106). The concentration of BS sites is very sensitive to the shape of the crystallite. For various crystals containing 683 atoms, van Hardeveld and Hartog (20) calculate a number of Bs sites ranging from 76 to 0; for example, 36 for a sphere, 76 for a cubooctahedron, 13 for a cube, and 0 for many other possible arrangements. For a nickel crystal of octahedral shape van Hardeveld and van Montfoort (106) calculate that there are no BS sites for particles of d below 1.2 nm. Then the curve for the fraction of Bs sites on the surface passes through a distinct maximum of 0.17 at 1.8 nm and then falls so that it is 0.072 at 4.4 nm (4009 atoms). Sites made up of two atoms or more are thought to be important in various catalytic systems. For example, Dalmon and Martin (105) proposed that 12 atoms are needed for a site for the hydrogenolysis of ethane. At the present moment, the BS sites serve as a popular explanation of structure sensitivity, for curves of TOF versus d for a number of reactions pass through a maximum in the range of around 2-3 nm (107-112). Further discussion of this matter is reserved until after a presentation of the results from the literature. From the foregoing, it is clear that geometric effects may be important out to much larger particles than would be deduced from Fig. 2b. One lesson to be learned from these geometric considerations, which have


I

100

,

#

I

1

1

79

,

LO 2 0 -10 4.0 21) 1.0 d,nm(Ni)

FIG.4. (a) Face-centered cubic (fcc) octahedron-max B5.From Ref. 20. The shaded atoms represent one of the possible B5 sites. (b) Coordination of surface atoms versus the mean particle size a for nickel, assuming that the particle has a face-centered cubic (fcc) octahedron-max B5 structure. The symbols Cip.+‘,.-refer to atoms identified in (a).

been available for some years, is that fundamental studies on structure sensitivity must include data on the crystallite shapes. Modern computeraided electron microscopy should make this possible (113, 114). As time passes it should also become more practical to check the plausibility of models by quantum-mechanical calculations. Other calculations and experimental work (104, 114) have shown that the situation is probably in reality more complicated than that described

80


above. Most small fcc metal particles ( d = 10 nm) crystallize in a cubooctahedral arrangement (14 faces) illustrated by Fig. 5a. For even smaller crystallites, an icosahedral non-fcc arrangement may become more stable (22, 114-117); such a shape is shown in Fig. 5b and c. The differences in energy between the smaller icosahedrons and cubooctahedrons are small, so that the d below which the former are stable varies considerably from one metal to another. For more details and references, the reviews by Anderson (22), Gillet (116), and Hoare and Pal (117) are useful. It seems plausible that the catalytic activity of small metal particles would be influenced by the crystal structure. Yacaman et al. (If@ have studied pentane hydrogenolysis over Rhly-AlzO3. Rh/Si02, Rh/C, Rh/Ti02, and Rh/MgO. The support and preparation method, all for particles of d < 5 nrn, determined whether cubooctahedrons or icosahedrons were formed, but the catalytic properties depended more on d than on crystal type. b. Alloying Effects. The behavior of binary metal alloys of supported metals has been investigated by many researchers. Examples of reviews on this subject are those of Sinfelt (1 19), Sachtler and van Santen (120),

FIG.5. (a) Face-centered cubic (fcc) cubooctahedron (20). (b,c) Icosahedra. From Ref. 117.


81

b

FIG.5 . (continued)

Clarke and Creaner (121), and Ponec (122). Alloys are sometimes used as practical catalysts because of their selectivity or resistance to poisoning. Here we mention this kind of research because of the similarity of the effect of alloying a relatively inert Group Ib metal with a Group VIII metal

82


to the effect of decreasing the metal particle size. This similarity arises because in the present view the valence electrons of each metal in these alloys retain their identities, so that relatively inert Cu can be considered a diluent in an active Ni particle, for instance. The physical basis for this notion that each constituent in an alloy retains its own atomic character was first obtained through photoelectron spectroscopy in the period 19681970, notably by Seib and Spicer ( 1 2 3 ~ )Evidence . based on catalytic results led to similar conclusions (120). A geometric interpretation has been applied to kinetic measurements on the hydrogenolysis of cyclopentane which indicates that a mixed PtRe ensemble is much more active than either pure metal (123b). Thus in our discussion of the role of the multiatom sites (ensembles) usually considered important in antipathetic structure sensitivity we shall have occasion to refer to related results on catalysis over alloys. For example, the hydrogenolysis of ethane over Ni is inhibited as Cu is added and the necessary ensembles are broken up. c. Experimental Models via Stepped Single Crystals. The understanding of adsorption and catalysis has been advanced by studies on well-defined surfaces, to which it is possible to apply a variety of electron spectroscopic and electron diffraction methods. Often it is found that the TOF varies from one face to another of a single crystal. Thus, changes in TOF for supported particles may be related to changes in the relative surfaces of the faces exposed as a function of d , method of preparation, or catalyst history. In order to apply the powerful methods of surface science to atoms of low coordination, such as those encountered on small particles, the use of stepped surfaces of single crystals has been developed. Davis and Somorjai (124) have reviewed this method. If a single crystal is cut at a slight angle from that required to produce a principal face, steps or kinked steps are formed. The heights and orientations can be changed, as expressed for example by the Miller indices shown in Fig. 6. Atoms at the edges of the steps and at the corners of the kinks can be assimilated to the low-coordination C7 and C4 atoms of Fig. 3. Thus we have a geometric model for some of the effects produced by small supported particles. Of course, the low-coordination atoms are still surrounded by many other metal atoms. Somorjai and his collaborators have studied a number of such systems [Blakely and Somorjai (125); Smith et al. (126); Davis and Somorjai (127)l. The data shown in Fig. 6 for cyclohexane dehydrogenation imply that this reaction should be structure-sensitive. Some results from the older literature [e.g., Cusumano et al. (128)] seem to indicate that this is a structure-insensitive reaction over Pt/A120, and Pt/silica-alumina. A

METALS: PARTICLE SIZE A N D CATALYTIC PROPERTIES 10

1

1

I

I

1

1

I

83

I

15 torr Cyclohexone 100 torr Hvdroaen

Time (minutes)

FIG.6 . Dehydrogenationof cyclohexane. Benzene accumulation in the gas phase at 573 K and I atm in a closed (batch) reactor containing a Pt crystal of known exposed surface. The crystals expose the flat Pt(l11) (M), stepped Pt(557) (V),kinked Pt (25, 10, 7) (O),and Pt(l0,8,7)(A) surfaces. Initial rates are proportional to the slopes at zero time. From Herz, R. K., Gillespie, W. D., Petersen, E. E., and Somojai, G . A.. J . C u d . 67, 371 (1981).

number of possible reasons for such differences will be discussed in later sections. C.

MODELCATALYSTS V I A CONTROLLED DEPOSITION OF SMALL PARTICLES

In discussing electronic effects and Fig. 2, a number of references have been made to model catalysts formed by deposition of metal from the vapor phase onto an inert support. An attempt is made to simulate the surface of a real catalyst so that it can be better studied by XPS and electron microscopy. In addition, Masson et al. (78, 79) have developed techniques to control and make almost uniform the size of the particles. Rutherford backscattering is used to find the total number of atoms deposited, and the number of nuclei that grows into discrete metal particles is found by electron microscopy. From these measurements the average number of atoms per particle is obtained. Longer deposition times give larger particles. With such a model catalyst (79)it is possible to measure both electronic and geometric characteristics of the metal particles, along with their reactivity, as the particle size is changed. It is of interest to apply this method not only to particles smaller than 4.0 nm, for which electronic effects may

84


be important, but also to particles up to 50 nm. For the latter, catalytic effects must be related to the identities of the exposed faces of the metal crystallites, and these are particularly accessible because of the method of sample preparation. In this way the structure sensitivity shown by many systems above d = 4 nm may be better understood. D. EXPERIMENTAL DIFFICULTIES In the preceding sections we have discussed how observations on the effect of d on catalysis might be explained by electronic and geometric factors. Usually it has been assumed that one is dealing with isolated crystallites of completely reduced metal. Often such may not be the case, and we now consider some common situations in which complicating factors may arise. 1. Metal Bonding with the Support

The atoms of a metal that are in contact with the support may undergo some electron transfer with the support, leading to some ionic or covalent bonding. Such effects would influence only a small fraction of the atoms of a large particle, but the fraction of atoms affected can be expected to increase as the particle size decreases. Thus measured variations in catalytic behavior may be due to a combination of intrinsic size effects and support effects. Strong metal support effects are often observed even for particles larger than 4 nm, after reduction at temperatures above 450°C. These results on supports such as Ti02 have been explained as arising from a large increase in metal-support contact made possible by the decoration of the metal by a reduced form of the oxide support [Resasco and Haller (129); Tatarchuk and Dumesic (130); Sadeghi and Henrich (131a)l. Clearly the decoration also permits a geometric interpretation (129) similar to that invoked to explain the effects of alloying. In any event, most studies of the effect of particle size do not involve the consequences of high-temperature reduction on reducible oxides.

2. Incomplete Reduction of Metal What may seem to be an intrinsic decrease in TOF with particle size may sometimes be caused by the decreasing extent of reduction of smaller particles. Thus, for nickel, small particles are not completely reduced, and the degree of reduction must be considered in evaluating the effect of particle size on TOF (59). Attempts to produce small clusters of iron

METALS: PARTICLE SIZE A N D CATALYTIC YKOPERTIES

85

supported on alumina by the decomposition of iron carbonyls led mostly to iron oxides (64). 3 . Changes in Crystal Structure Figures 3 and 4 are based on the maintenance of the same crystal shape as the size varies. For small particles, needles or even two-dimensional shapes may predominate, as discussed by Bond (21). More subtle changes may occur, such as the replacement of cubooctahedrons by icosahedrons for crystals of d < 2 nm. For large crystals kinetic effects during the preparation may cause deviations from the expected equilibrium distribution of faces, often assumed to be equal portions of the (1 1I), (loo), and (1 10) planes. In their measurement of the TPD of CO on small Pt particles on A1203, Altman and Gorte (13Zb) explain their results by proposing that (1 1 1)-type facets are preferred on the larger particles. 4.

Changes in Deactivation with Particle Size

Deactivation may be structure-sensitive. Thus Lankhorst et al. (132) have shown that variations of TOF with d for hexane reforming on Pt were caused more by increases in carbon deposition with d than by decreases in the intrinsic rate of reaction with d. For model stepped-surface single crystals of platinum, Somorjai and Blakely (133) showed that carbon deposition was favored on terraces over edges and corners, so that the reforming reactions favored by the latter sites are poisoned less rapidly than those occurring mostly on the flat surfaces. Thus small particles are expected to be poisoned less rapidly than large ones, for this example.

5 . Metal Reconstruction It has long been suspected that chemisorption and catalysis may actually cause major changes in the shapes of small crystallites from those observed in the absence of adsorption. The reality of such changes has been demonstrated dramatically by van’t Blik et al. (134) for the chemisorption of CO on well-reduced crystallites of 6-10 Rh atoms supported on alumina. On adsorption the particles are broken up to form Rh(C0): bonded to the oxygens of the support. In the absence of CO, a raft structure for such Rh particles has been observed ( 1 3 5 ~ )However, . in the presence of CO this structure is destroyed and the Rh atoms oxidized, on alumina. For nickel, a dramatic change in particle size may occur by another process. When kinetically and thermodynamically favored, gaseous

86


Ni(C0)4 may be formed in the presence of CO so that there is transport of Ni from small particles to large ones (135b). Under certain conditions, usually to be avoided, all the nickel may be removed from the support and deposited on cold surfaces downstream of the reactor. OF STRUCTURE SENSITIVITY E. SOMECONSEQUENCES

Before considering some of the details of structure sensitivity in catalysis, it is of interest to present a few examples. In this way some of the practical consequences of structure sensitivity are illustrated, although it must be kept in mind that there are many steps between the laboratory results we show and the development of an industrial process.

I . Structure Sensitivity and Reaction Rates Results that illustrate three examples of structure sensitivity are presented in Fig. 7. Curve 1 represents the TOF for the structure-insensitive reaction of benzene hydrogenation on Pt/SiOz (136). The TOF is about 0.22 s - I and is not a function of particle size over the range studied. Also plotted for this system is the atomic rate AR, calculated from TOF by Eq. (2). As already mentioned, the atomic rate is an approximate measure of

u) l[O

Y

i,nm

2,5

1,0

AR

2;5

1p

0;C

I

FIG.7. Turnover frejuency (rate) TOF and atomic rate AR versus fraction exposed FE and mean particle size d for a selection of reactions and metal systems (see Table Ill for details of the studies).

87


TABLE 111 Details of the Structure Sensitivity Studies of Metal Systems Presented in Fig. 7 Curve (Fig. 7)

Metal

1

Pt

Si02

2

Rh

Si02

3

Ru

AI2O3,SOz, etc.

Support

Temperature, pressure

Reaction

Reference

50°C, 1 atm, Hz/C6H6 = 12.75 253"C, I atm, H*/C2H6 = 6.67 250°C I atm, HJCO = 2

C & ,-k 3H2 --* C6Htz C2H6 f H2-* 2CH4 co -F 3H2+ CH4 + H20

136 13 137

the rate of product produced per unit cost of metal in the catalyst. Thus, it is of interest to maximize the atomic rate, and for the present example this reasoning leads to using a catalyst of the highest FE that is practical. The dotted parts of the curves are extrapolated to FE = 1, where TOF = AR. Naturally, processing costs may become prohibitive as FE increases, and this factor is not included in the above reasoning. Curve 2 of Fig. 7 shows TOF and AR for a sympathetic structuresensitive system, the hydrogenolysis of ethane over Rh/Si02 (13). The TOF rises with FE and then falls for very small particles. The authors expressed some reservations about this behavior, but in our review of the literature we shall see that such behavior is often observed. The curve for AR goes through a maximum in the region of d around 1.5 nm. Thus, for this system there is an optimal value of d, close to the smallest sizes that are practical. Finally, curve 3 of Fig. 7 shows some results for an antipathetic structure-sensitive system, the hydrogenation of carbon monoxide over supported ruthenium (137). For such systems there is usually a maximum in the curve of AR versus d, and this occurs at about 5 nm for the present case. For particles smaller than 5 nm, the decrease of TOF is the predominating effect, and for particles larger than 5 nm, the decrease in FE is the predominating effect. We shall see that CO hydrogenation over almost all systems studied shows an antipathetic structure sensitivity, so that there is no economic incentive to decrease the crystallite size below a fairly large size, 5 nm for the data shown. This reasoning is based on a consideration of the reaction rate, but the selectivity is also important, and we discuss this matter presently. It should be noted that in constructing Fig. 7 and the other figures in this article we use Eq. (1) with B = 1 nm. Actually, B may vary between 0.7 and 1.0 nrn, depending on the metal and the shape of the crystallites.

88


However, in view of the precision of the data and the facility of presentation with only one value of B, we have retained B = 1 nm for all figures. It should also be clear that the relation between d and FE loses its meaning as FE approaches I .O. From the geometric point of view, the properties of the surface atoms of a small cubooctahedron are different from those of the same atoms arranged as a one-dimensional raft. From Figs. 3a and 3b, it is deduced that structure sensitivity should appear only for particles with d less than about 5 nm. It is convenient to call systems that actually exhibit this behavior systems with limited structure sensitivity. On the other hand, if geometric effects are governed by ensembles of several atoms, then structure sensitivity should extend to diameters considerably greater than 5 nm. This behavior we call extended structure sensitivity. In other words, with limited structure sensitivity, the system is structure insensitive ford > 5 nm. A glance at the figures in the rest of this review shows that many actual systems present extended structure sensitivity. 2. Structure Sensitivity and Selectivity

The key to the practical exploitation of a catalyst is often its selectivity, so it behooves us to see how the selectivity may vary with d . Curve 1 in Fig. 8 sketches the variation of the molecular weight of the product of the Fischer-Tropsch synthesis over supported cobalt as a function of FE (1.38). Large particles favor a higher production of hydro-

1

100 - 1.0 1

25

i,nm 10

25

10 L

0 L

al

>

a

FE FIG.8. Selectivity ad ! average carbon number of product versus fraction exposed FE and mean particle size d for a selection of reactions and metal systems (see Table IV for details of the studies).

TABLE IV Details of the Reactions and Conditions of the Metal Systems Studies Presented in Fig. 8 Curve (Fig. 8)

Metal

Support

1

co

2

Pt

A1203, Ti02, Si02, C Si02

3

Ag

Si02

Temperature, pressure 225"C, 1 atm, H-JCO = 2 240-375"C, 1 atm, HJC5 = 10.5 220°C, 1 atm, 02/C2H4 = 1.0

Reaction

CO + H2hydrocarbons nCsH,2 + H2 + isomerization, hydrogenoly sis C2H4 + 0 2 -, C2H40. c02

Reference

Comments Fischer-Tropsch synthesis Selectivity = (isomerization)/ (hydrogenolysis isomerization) Selectivity = (C2H4OV ( 0 2 + C2H4O)

138 140

+ 141

90


carbons in the gasoline range, one of the original goals of the process. Since the atomic rate is also favored by relatively large particles, the use of such particles is doubly indicated. This is in accord with the old literature, which indicated that highly dispersed supported catalysts were not desirable. Precipitated or fused (iron) catalysts were preferred (139). Curve 2 of Fig. 8 concerns the competitive isomerization and hydrogenolysis of normal pentane as a function of particle size over Pt/Si02 catalysts (140). Isomerization is favored over large particles and hydrogenolysis over small particles. It is clear again that the best catalyst probably corresponds to an intermediate particle size. For the partial oxidation of ethylene over Ag/SiOz catalysts (141), curve 3 shows that the selectivity toward ethylene oxide (versus COz) is favored by the largest crystallites studied. However, both the production of CzH40 and of COz show sympathetic structure sensitivity. Thus, there must exist an optimal crystallite size for the most effective catalyst. It should be noted here that the rates and selectivities usually reported are initial rates, found in a differential reactor. It is clear that there exists a complicated optimization problem in the design of the industrial reactor and catalyst. The best values of crystallite size, conversion, and other operating conditions must be sought. Often these various searches are carried out by separate groups, at separate times, so that the attainmknt of the optimal process is difficult to achieve. It is worth noting here that the metal loading on the catalyst [C in Eq. (4)] is an important variable. Crystallite size can be manipulated by variation of C and of the pretreatment (sintering) temperature and gas-phase composition. For instance, a given d may be obtained from a low loading and a high sintering temperature, or from a high loading and a low sintering temperature. Most authors find that the TOF is about the same for two or more such catalysts of the same d (or FE) and different combinations of loading and sintering temperature. Thus we express results with d as independent variable rather than C. An early review actually used metal loading as the independent variable, reporting catalytic activities as a function of metal loading without reference to FE or d (3). For Ti02 and other reducible supports, a high reducing temperature may induce a strong metal-support interaction (SMSI). In this case TOF becomes a function of sintering (reduction) temperature at a given d. SMSI is absent in most of the literature we shall discuss. However, in the observed variation of TOF with d, the explanation of the observations may be sought as arising from electronic effects, ensemble effects, support effects, reconstruction effects, etc. We shall see that which explanation, or combination thereof, is valid is not evident for many of the literature studies.


91

3 . Structure Sensitivity and Chemisorption In our discussion of the influence of structure on the turnover rate our understanding is frequently hampered by lack of information on the ratedetermining step and the most abundant surface intermediate. It would be logical to consider the structure sensitivity of the rate of an elementary step, such as the desorption of a chemisorbed gas. Results on temperature programmed desorption as a function of particle size might be simpler to interpret than those of global reactions consisting of a sequence of steps. However, few such data are available. There are, however, a number of papers that treat the equilibrium relations in chemisorption, and we shall discuss some of these. Such measurements can lead to information on the structure and binding of adsorbed species, and also to the characterization of metal catalysts. For example, CO adsorbed on PtO shows a lower stretching frequency in the infrared than when adsorbed on oxidized platinum (142). Another reason to consider chemisorption in this review is that it is often used as the basis for estimating d or FE and thus also influences TOF via Eq. (2). As already mentioned, hydrogen is usually used for measuring FE and in most cases the best value of H/M, seems to be close to 1.0. However there are a number of complications, and we discuss them in the following section on characterization.

IV. Characterization of Small Supported Metal Particles

A. INTRODUCTION In order to obtain data on structure sensitivity, a reliable measure of the metal fraction exposed is needed. The simplest method of measuring FE is via chemisorption of hydrogen; carbon monoxide and oxygen are also sometimes used. After a discussion of chemisorption, we consider later in this section other methods of characterization. Chemisorption leads to an estimate of the average particle size; by electron microscopy and electron microdiffraction, information can be obtained on the particle size distribution and in favorable circumstances on the morphology of the small particles. It is also important to know the chemical state (degree of reduction) of the particles, and to this end electron paramagnetic resonance (EPR), X-ray photoelectron spectroscopy (XPS), ultraviolet photoelectron spectroscopy (UPS),Auger electron spectroscopy (AES), and infrared spectroscopy (IR) are useful. These and other methods will be discussed in what follows. Our goal in the discussion is not to review the methods

92

MICHEL CHE AND CARROLL 0 . BENNETT

considered but to show how they are used to further the study of small particles. In order to determine the effect of d or FE on TOF,evidence is required that the metal particles are indeed approximately monodisperse, fully reduced, unpoisoned, little influenced by the support, and unpromoted. These matters have been reviewed in the previous section under the heading “Experimental Difficulties.” Thus we shall discuss how FE can be measured and then consider other methods of characterization that will furnish more details about the geometric and chemical states of the small particles. We shall, however, not consider metal-support interactions since this aspect is covered in a separate review (25a).

B. PARTICLE SIZEA N D PARTICLE SIZEDISTRIBUTION 1. DeJinitions

Although the expression “particle size” is very often used in the literature, it is very seldom clearly defined, mostly because many definitions are possible and none is unambiguous except if crystallites are all-cubic or all-spherical. This, in practice, never happens and it is necessary to define arbitrarily some quantity as the particle size. Matyi et al. ( M a ) propose that the simplest way to do this is to equate “size” with “diameter,” where the diameter is defined as any straight line passing through the center of mass of the particle and terminating at the particle boundary. The authors present other definitions and show that the most suitable choice depends on both the experimental technique used and the nature of the system being examined. As discussed in Section V, we do not need be too concerned about this aspect, as important as it might be, since there are other more important difficulties to resolve before any rationalization of the data becomes possible. The same type of problem is encountered in the definition of particle size distribution since the starting assumption is that all particles have the same shape. In this case, the particle size distribution is defined as the relationship between a given particle size and the frequency, or number of particles, with that certain diameter or size. Matyi et al. ( 1 4 3 ~discuss ) in detail for a given particle size distribution the several features that can be interpreted as the average size. We refer the reader to this review article.

2. Chemisorption a . Hydrogen. The use of the chemisorption of hydrogen to characterize metal catalysts has been discussed in a useful review by Scholten et al. (14%). Usually an isotherm is measured at room temperature, and mono-


93

layer coverage (H/M, = 1.0) is assumed to correspond to irreversibly adsorbed hydrogen. The latter is determined by some arbitrary definition. In the best cases, the actual average particle size and morphology have been determined by some reliable physical methods, so that the chemisorption method is calibrated. Following a certain recipe of waiting and pumping, the amount of H2 adsorbed at 25°C and a given pressure are said to correspond to a certain H/M, ratio, preferably unity. Sometimes the appropriate amount of H2 is taken as that found by extrapolating the highpressure part of the isotherm (supposedly linear) to zero pressure. In some cases the isotherm of total adsorption is obtained, the reversibly adsorbed H2 is removed according to a recipe, and then the reversibly adsorbed isotherm is measured. The irreversibly adsorbed H2 isotherm is obtained by difference, and then the monolayer is deduced from this isotherm. For details we refer the reader to Scholten et al. (143b) and appropriate textbooks (e.g., 1b). We next discuss a number of points to keep in mind concerning the use of H2 chemisorption as a method of characterization. 1. Chemical adsorption on metals is usually considered as nonactivated. For H2 on Ni/A1203, however, the isobars pass through a maximum at about 100°C (144), meaning that at 25°C kinetic factors (adsorption time) have an important influence on the amount of H2 adsorbed. For Rh/SiOz (144) the amount of H2 adsorbed is found to depend very little on time; the isobar descends smoothly as the temperature is increased. Thus each metal has its own behavior, which must be taken into account. The thermodynamics and kinetics of adsorption and desorption are discussed in the literature (145, 146). 2. As discussed in the interesting article of Crucq et al. (145),for most chemisorptions the AH of adsorption (exothermic) decreases owing to lateral repulsive interactions as coverage increases. Often the relation is approximately linear, leading to a Temkin isotherm. The change in AH may arise because of a distribution of metallic sites and/or because of increasing interactions between adsorbed H a s the coverage increases. Thus the fraction of adsorbed gas that can be removed depends on temperature, time, and arrangement of the equipment; the activation energy for desorption obviously also decreases with increasing coverage. Thus the term reversible adsorption has only an artificial meaning and might better be replaced by weak (and strong) adsorption (147). Sayari et al. (98) have studied the weak adsorption of H2 on various supported ruthenium catalysts of 0.9 < d < 12.4 nm. They express their results in terms of the ratio HwIHT,where the subscripts W and T, stand for weak and total adsorption. At 0.9 nm, HwIHTis about 0.1, passes

94


through a maximum of 0.32 at 1.5 nm, and drops to zero above 2.5 nm. The authors (98) associate the weak adsorption with special sites, invoking in particular the Bs sites. Although these sites do reach a maximum concentration at about the right d, their decline for higher d is gradual (106) and extends far above 2.5 nm. It is interesting to recall that Aben et al. (97) have found that the hydrogenation of benzene is structure-insensitive if the TOF is based on the FE on the basis of the adsorption of weakly bound hydrogen only. In Fig. 13 (Section V,B,5), the TOF (curve 4) is based on strongly adsorbed hydrogen and shows a maximum at about 2.0 nm. 3. Although a calibration of H2 chemisorption may be desirable, there must exist another reliable measure of particle size, and such has not been the case for d < 3 nm. In recent years, however, improved techniques in electron microscopy have permitted the measurement of particle sizes down into the range 0.3-1.0 nm. Even though the resolution of modern instruments may be of the order of 0.3 nm, it is often difficult to obtain enough contrast to distinguish the metal from the support. By careful sample preparation McVicker et al. (148) were able to see particles of Ir on A1203of about 0.6 nm. They found that for particles in this range HIM, is about 2, whereas it drops to 1.O for larger particles (strongly adsorbed H2). The contrast in electron micrographs can be remarkably enhanced by computer-aided image processing. An example is the recent work of Fuentes et al. (149) for Rh/Ti02,but they did not report results for hydrogen chemisorption. It is also possible to use extended X-ray absorption fine-structure (EXAFS) spectroscopy to deduce particle sizes, and this method has been used, along with hydrogen chemisorption, to study Rh, Pt, and Ir supported on Si02 and A1203 (147, 150). The average coordination number of the metal atoms is calculated from the EXAFS results. A model implemented by computer is then used to estimate the particle size from the coordination number. Metal support interaction can be included in the model. Kip et al. (147) have used total adsorbed hydrogen in reporting chemisorption results. For Pt they find 0.6 C H/M, C 1.2; for Rh (FE always > 0.8), 1.5 < HIM, < 2.0; and for Ir, 2.5 < HIM, < 3.0 for 0.2 < FE < 0.9. In these data the higher H/M, refers to FE close to 1 .O and the lower one, to the samples of lowest dispersion measured. Although there is some scatter in the results, it is clear that H/M, is a function of FE for a given metal, and that its range changes from one metal to another. Roughly, at FE = 1.0, it seems that H/Pt, = 1, H/Rh, = 2, and H/Irs = 3, where H/M has been obtained by extrapolating HT/M to zero pressure.


95

It should be noted that hydrogen held by refractory supports such as SiO2 and A1203 usually has a negligible effect on the results, either via physisorption or spillover (147, 148, 149a) at 25°C. Clearly, metal-support interactions are important, for instance, in the formation of rafts of rhodium or the location of groups of only one or two atoms (149b).However, on the flat Rh particles found on a Ti02 support, there should be at least two H atoms adsorbed per Rh atom (147, 148). For hydrogen adsorption on reducible oxides, there may be an apparent slow adsorption that actually arises from the reduction of the support. For Rh/TiOz it has been found that at 22°C there is rapid chemisorption of Hz corresponding to H on Rh followed by further slow disappearance of H2 corresponding to the gradual creation of Ti3+by spilled-over hydrogen (151). The resulting lines representing Hz adsorbed versus In (time) have slopes and intercepts that are functions of reduction temperature (SMSI) and subsequent oxygen treatments. Oxygen adsorption is also measured and a coherent picture of the various processes is presented (151). It should be noted that the preparation of catalysts of low metal loading (7) to achieve a cation exchange. Analogously, we have also effected anion exchange by working at low pH values (lower than the isoelectric point of Ti02, i.e., a pH of 5-6) and using metal chloride solutions. Under these conditions, anion complexes (A-) of the type [Mn+Clx](x-")-are present in solution. They can be adsorbed onto the support by the following mechanism: S-OH

t A- t H+

S-OHiA-

(1)

On the other hand, the impregnation process involves a combination of both adsorption of metal complex ions and deposition of the solute as the solvent is evaporated. Thus, the resulting metal dispersion depends on the fraction of metal ions undergoing exchange with support species relative

183

METAL-SUPPORT INTERACTION

to the fraction of metal precursors deposited on the support as occluded solute. We prepared a Rh/TiOz catalyst series by varying the pH of the Rh(N03)3 impregnating solution from 3 to 11 (46). As shown in Fig. 3, as the pH was increased, the cation-exchange capacity of the support increased, resulting in a higher dispersion. At excessively high pH values, precipitation of rhodium can occur, resulting in a lower dispersion. In an electron microscopy study of a cation exchanged Rh/TiOz catalyst, Fuentes et al. (52) used image processing techniques to filter the noise and enhance the contrast. They observed two-dimensional structures of Rh with about 1-nm cross section and one-dimensional rows of Rh along the [Ool] direction of rutile (110) planes. We have recently rationalized this structure in terms of the genesis of the catalyst (53). Considering the surface of the Ti02 powder during the impregnation as that of a fully hydroxylated TiOz(ll0) plane, we expect to find both bridging hydroxyls bound to two Ti4+cations and terminal hydroxyls bound to one. In the former case, the bridging oxygen will be more acidic and, therefore, will preferentially act as cation-exchange sites. We can then expect that during the first treatment with NHXOH) those bridging sites will preferentially hold NH:, which will subsequently be exchanged by the cationic metal complex (probably [Rh(H20)4(OH)2]+). The chemistry

\

0

/

A

E E

0.25

\

-

3

6

9

pH OF IMPREGNATING

12 SOLUTION

FIG.3. Metal dispersion (H/M) as measured by H2chemisorptionafter reduction at 473 K of a series of 2-wt%Rh/Ti02 catalysts as a function of pH of the solution used during the impregnation step in the catalyst preparation. (After Ref. 46.)

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GARY L. HALLER A N D DANIEL E. RESASCO

by which the metal precursor loses water, adsorbs H2, and reduces to the metal is difficult to imagine in detail, but formally we can write 3/2& + [Rh(H20)3(0H)2(0Ti)]3+ SH20 + [Rh(HOTi)13+ (2)

-

As written, only the Rh has been reduced and all the titanium cations remain Ti4+. This is a reasonable assumption for the low-temperature reduction (LTR; usually 473 K) case. When the reduction temperature increases, surface oxygen vacancies may be created, causing the reduction of Ti4+to Ti3+.In that case, the HTR step may be written as [Rh(HOTi)]’+ + 1/2H2 [RhTi]’+ + H 2 0 (3)

-

In both cases, the resulting structure would be one-dimensional rows of Rh with the spacing of the Ti rows, as observed experimentally (52).

B. EFFECTOF REDUCTION TEMPERATURE ON THE CHEMISORPTIVE A N D CATALYTIC PROPERTIES OF TiO2-SUPPORTED CATALYSTS Following the original work of Tauster et al. (13), every Group VIII metal has been supported on Ti02 and investigated as a SMSI catalyst. Among these catalysts, Pt/TiOz and Rh/Ti02 stand out as the most thoroughly investigated systems with regard to this phenomenon. Perhaps platinum, which is more widely used in petroleum refining catalytic processes, has received more attention than rhodium. Nonetheless, a number of particularly illustrative studies have been carried out on Rh/Ti02. Our own investigations have concentrated on the Rh/TiOz system because Rh has the advantage over Pt of having a much higher activity for some structure sensitive reactions, e.g., alkane hydrogenolysis. Thus, it is possible to measure the catalytic activity at temperatures below 473 K, which is used for the LTR. This is not as practical for either ethane hydrogenolysis or CO hydrogenation on Pt because of its relative inactivity compared to Rh. The first-row Group VIII metals have the added disadvantage that one cannot always be certain of complete reduction to the metal after a LTR. Because of the rich literature and our own interests, we will concentrate our discussion of the effects of reduction temperature on Rh/Ti02 and Pt/TiO2 catalysts. We will briefly discuss the case of Ni catalysts because this metal exhibits a behavior rather different from that of other Group VIII metals and has also been widely studied. 1.

H2

and CO Chemisorption

Tauster et al. (13) first identified what has become a distinctive feature of the so-called SMSI catalysts, the decrease in H2 and CO chemisorption


185

capacity as the reduction temperature is increased. It is now well established that the chemisorption suppression after HTR is a common characteristic of all Group VIII metals supported on TiO2, and the extent of suppression is usually taken as an indicator of the extent of metal-support interaction. Even though a direct comparison of data from different laboratories is not straightforward, in part because of the different preparation procedures and pretreatment conditions, a general trend may be ascertained. This is illustrated in Fig. 4, which includes data from several laboratories (13, 16, 22, 34, 54-63). There exists a correlation for TiOzsupported catalysts between the loss of adsorption capacity after HTR (723-773 K) and the original metal dispersion, as measured after LTR (473-523 K). This trend was first pointed out by us for the case of Rh/Ti02 catalysts (16). We explained it in terms of the electronic description prevailing at that time and speculated that the effect of an electron transfer should be more noticeable on the smaller particles. Later, we realized that this trend could also be interpreted in terms of the extent of interfacial perimeter around the metal particle, which increases as the particle size decreases. Thus, a localized interaction could be used to explain the observed correlation. As mentioned above, the idea of localized interac-

El

I

?

0.15-

0.1.

0.05.

FIG.4. Hydrogen uptake (H/M) after reduction at 773 K over a series of Group VIII metals supported on Ti02 as a function of the uptake (H/M) after reduction at 473 K. These measurements have been performed in several different laboratories (13,16,22,34,54-63).

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GARY L . HALLER A N D DANIEL E. RESASCO

tion lead us to our current TiO, decoration model. Most of the data in Fig. 4 lie approximately on the same curve. Those corresponding to Ir and Pd

catalysts appear to follow a different trend. In the case of Ir this difference may be due to a different H/M stoichiometry. It is generally believed that for Ir the adsorption stoichiometry H/M is about 2 (64). In that case, the Ir points would also lie on the lower curve which assumes a H/M = 1. It is more difficult to explain the different behavior of Pd, for which there are no reports indicating H/Pd stoichiometries greater than one. However, the larger particles of Pd can form bulk hydrides that could increase the resulting H/Pd ratio. The effect of reduction temperature on the chemisorption of CO on Rh/Ti02catalysts has been studied by IR spectroscopy (36,37).The lowtemperature reduced catalysts evidence the presence of the three typical adsorption forms of CO on Rh, i.e., linear (2070 cm-I), dicarbonyl (2030 and 2100 cm-I), and bridged (1800 cm-I). As the metal dispersion increases, the proportion of dicarbonyl form, characteristic of small particles (65) or nonzero valent Rh ions (66), increases. As shown in Fig. 5 (36), an increase in reduction temperature to 517 K causes a more pronounced decrease in the amount of multisite bridged form. At higher reduction temperatures (623 K)there is a substantial decrease of both the

WCIVENUHOERS

wvmmocns

FIG.5 . Infrared absorption spectra of CO adsorbed on a 2-wt% Rh/Ti02 catalyst after saturation at room temperature following reduction at (A) 473 K for 1 h; (B) 517 K for 2 h; (C) 623 K for 2 h; (D) 773 K for 1 h; (E) 773 K for 3 h. (After Ref. 36.)


I87

linear and the dicarbonyl forms, although the latter is clearly more affected. In the case of highly dispersed, ion-exchanged Rh/TiO2 catalysts (36,37), only the gem-dicarbonyl form is observed after LTR. However, increasing the reduction temperature in this case causes the appearance of the linear form as the dicarbonyl form disappears. Even though this conversion could be partially explained by particle growth during the HTR it must be pointed out that it is reversed by oxidation followed by LTR (36). This might indicate that a non-zero-valence species, i.e., Rhl+, exists, which is responsible for the CO dicarbonyl form and these are reduced to Rho during HTR. van’t Blik et al. (67) have ruled out this possibility for Rh/AI2O3catalysts by TPR, EPR, and EXAFS experiments. They have proposed that all the Rh is in the zero-valence state but that the CO adsorption itself causes the rupture of the Rh particle and the oxidation to the + I valence state. This concept had previously been proposed by Primet (68) from IR and XPS data. Similar effects on chemisorptive properties were observed for Pt catalysts. In a combined IR and kinetic study (69)it was demonstrated that the HTR of Pt/TiO2 may, indeed, change the ability of H2 and CO to compete for adsorption sites. It was observed that after HTR the CO IR, bands appear to be significantly reduced and at slightly higher frequency, particularly the band corresponding to the bridged form. The latter feature was considered as evidence of a weakening of the surface bond. More recent TPD studies have shown that after HTR the CO becomes more weakly adsorbed (30).Chemisorption studies indeed show that irreversible chemisorption of CO no longer occurs on Pt particles highly dispersed on Ti02after HTR (69a).In addition, under CO hydrogenation reaction conditions, the IR bands of CO could not be detected on Pt/TiO2 catalysts, but they were still present when other supports, i.e., Si02 or A1203, were used. The question is, however, whether such bands are kinetically important. A more significant finding (70)is perhaps the fact that, even at room temperature, the presence of hydrogen readily decreases the intensity of the CO bands. Unfortunately, this experiment has not been reproduced by Robbins in his IR investigation of Pt/Ti02 catalysts (71).This effect was not observed on Pt catalysts supported on other oxides. If this effect is real, the important consequence would be that hydrogen may more effectively compete with CO for the available adsorption sites on Pt when supported on Ti02 compared to less reducible oxides. The alteration of heats of adsorption by SMSI is still a controversial matter. From differential scanning calorimetry measurements Vannice ef al. (72)originally reported an increase in the heat of adsorption of H2 with Pt particle size, but not much effect of the support or reduction temperature. More recently (73)they have corrected the earlier work to conclude

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GARY L. HALLER A N D DANIEL E. RESASCO

that neither particle size nor support affects the heat of adsorption of H2 on Pt. Herrmann et al. (74)made more accurate measurements in a microcalorimeter and reported that the initial heat of adsorption decreases from 22.2 to 19.1 kcal/mol and the average heat from 14.9 to 12.9 kcal/mol for LTR and HTR, respectively. In the case of CO, the initial heat is not changed but the average heat of adsorption decreases from 25.2 to 20.6 kcal/mol for LTR and HTR reduction, respectively. One characteristic of the SMSI originally reported by Tauster et al. (13) was the restoration of normal chemisorption capacities on oxidation at 673 K followed by reduction at 473 K. Table I gives an example of this effect for the case of Rh/TiOz catalysts (46). It is observed that the chemisorption capacity is partially restored when the oxidation treatment is carried out at room temperature. An interesting feature was observed for the chemisorption of H2 and CO on Pd/Ti02 catalysts (75). The HTR does not have a significant influence on the distribution of hydrogen adsorption states, but causes important changes in CO adsorption states. The four CO desorption peaks observed by TPD were attributed to different adsorption structures. The TPD peaks appearing at 638 and 773 K for a 1.7% Pd/TiOz catalyst are ascribed to bridged bonded CO adsorbed on Pd( 100) and Pd( 11 1) planes, respectively. The low-temperature peaks (400 and 488 K)are attributed to linearly adsorbed CO. As shown in Fig. 6, the attenuation of the intensity of the peak at 638 K at increasing reduction temperatures is much more pronounced than the other three peaks. According to the assignment of TPD peaks, this trend would suggest that the bridge-bonding sites on Pd(100) planes are deactivated first. Generally speaking, geometric arguments alone, independent of the surface structure, suggest that bridged sites will be blocked before linear sites. In this case, however, there is a preferential deactivation of a particular type of site. This observation can be related to a recently reported field-emission microscopic study (76) of a TABLE I The Reversibility of SMSI in RhlTi02 by 298 and 673 K Oxidation" Treatment

COIRh

H/Rh

HTR 02 (673 K) + LTR HTR 02 (298 K) + LTR

0.01 0.85 0.01 0.18

0.02 0.76 0.01 0.15

From Ref. 46. Measured at 523 K .

Ethane hydrogenolysis rateb (molecules1Rh atom min) 0.36 0.09 nm, and rB > 0.051 nm in the case of oxides. Goldschmidt (19), on the basis of geometric considerations, defined the tolerance limits of the size of ions through a tolerance factor t = ( r A + rx)/*(r~ + r x ) , where r A , TB ,and rx are the radii of the respective ions; t would be equal to one for the ideal cubic structure (Fig. 2). In fact, the perovskite structure exists in oxides only between the limits 0.75 < t < 1.0 with t between 0.8 and 0.9 in most cases. For t > 1 the calcite and aragonite structures are prevalent, whereas for t < 0.75 the stable structure is ilmenite. Roth (20)has classified the limits of the existence of these competing structures according to the ionic radii values. Besides the ionic radii requirements, the other condition to be fulfilled is electroneutrality, i.e., that the sum of charges of A and B ions equals the total charge of X anions. This is attained in the case of oxides by means of charge distribution of the form A'+B5+03,A2+B4+03,or A 3 + B 3 + 0 3Moreover, . partial substitution of A and B ions giving rise to complex oxides is possible while keeping the perovskite structure. Figure 3, elaborated from some comprehensive compilations of data on the structure and properties of this type of compound (4,15,17,21,22), shows that almost all the stable elements have been included in the perovskite framework, many of them in both the A and B positions. In what follows we will A X H

I X Er

A

B A

B A B

A U

Po

Tb

B

B A B

B

A

A Pu

B

A Eu

B Np

B

A Sm

h

Nd

B Am

B

A

A

B Cm

A

Cf

A Er

Ho

B

B EC

B

A Dy

Tb

Gd

B Es

A Tm

B Fm

A Yb

B Md

Kr

Lu

B No

. Lw

242

LUIS G. TUUCA

et al.

limit the structural description of perovskites to the case of X = oxygen, for the reasons indicated in Section VI1,I. The ideal perovskite structure appears only in a few cases for tolerance factors very close to 1 and at high temperatures. In other conditions different distortions of the perovskite structure will appear. The compound CaTi03 was originally thought to be cubic, but the true symmetry was later shown to be orthorhombic (23). Distorted structures with orthorhombic, rhombohedral, tetragonal, monoclinic, and triclinic symmetries are known, but the last three types are very rare and poorly characterized (21), so we will only describe the orthorhombic and rhombohedral distortions. For t values in the range of 0.75 < t < 0.90 a cooperative buckling of corner-shared octahedra takes place, leading to the orthorhombic distortion. This network, sometimes typified as GdFe03 structure, has a space group Pbnm,and its relationship to the perovskite structure is as shown in Fig. 4a. It is obtained by tilting oxygen octahedra in such a way that the A atoms are displaced along (1, 1, 0) pseudocubic directions or (0, 1, 0) directions. The true orthorhombic cell is usually referred to as “0orthorhombic,” characterized by a lattice parameter ratio (cla) > fi,to be distinguished from the 0‘-orthorhombic structure, with (cla) < fi. This latter form is the result of a superimposed Jahn-Teller distortion of the perovskite structure. When there is no octahedra buckling, a small deformation from cubic to rhombohedral symmetry may take place. This occurs for tolerance factors in the range of 0.9 < t < 1 .O. The rhombohe-

ox,, (3x1 @ A FIG.4. Orthorhombic (a) and rhombohedral (b) distortions of the perovskite structure. (Reprinted by permission from Ref. IS.)

PEROVSKITE-TYPE OXIDES

243

dral distortion, sometimes referred to as “LaAI03 structure,” has the RSc-D% symmetry. Its relationship to an ideal perovskite is shown in Fig. 4b. In most cases the anions are displaced, thus requiring a large unit cell; in the most general situation the anion displacements may be decomposed into R?c and R3m components. Nonstoichiometry in perovskites can arise from either cation deficiency (in the A or B site), anion deficiency, or anion excess. This subject has been widely discussed in some excellent reviews (24, 25); thus, we will focus only on general features of the different types of nonstoichiometry. Because of the stability of the BO, groups, A cations can be missing without collapse of the perovskite network. The ReO3-type structure is thus the limiting case of A-site vacancy nonstoichiometry. The most typical example is that of tungsten bronzes, &W03. The question as to whether A-site atoms and vacancies are ordered is not fully resolved (24). Because of the large formal charge and the small size of the B cations in perovskites, B-site vacancies are not energetically favored; B-B interactions, which may be a compensating factor, are favored by hexagonal stacking of A03 layers. Accordingly, a number of hexagonal perovskites exhibiting B-site vacancies have been described (24). Normally, these vacancies are ordered between h-h layers where the B 0 6 octahedra share faces, in agreement with Pauling’s rules for the sharing of coordination polyhedra. Anion vacancy in perovskites is more common than cation vacancy. Unlike the well-known case of W 0 3 , anion-deficient nonstoichiometry is not accommodated by the crystallographic shear mechanism, but by assimilation of vacancies into the structure, resulting in supercells of the basic network. The review by Rao et al. (24) contains numerous examples of this kind of behavior. Anion excess has been described in a more limited number of systems. Structural details of this type of compounds can be found in Rao et al. (24) and Smyth (25). B. RELATEDSTRUCTURES The strong bonding between B-site ions and oxygen ions leads to retention of the BO3 grouping even when a three-dimensional network is no longer permitted by stoichiometry. The simplest case is that of the SnF4type structure, consisting of the superposition of octahedra nets, with the free peaks of octahedra of one layer lying in the holes formed between the peaks of octahedra in the neighboring layers. The peaks of octahedra are sunk so deeply into these holes that an almost planar structure is obtained. In the K2NiF4-typecompounds, slices of the perovskite structure

244

L U I S G . TWUCA

et al.

one unit cell thick are displaced relative to one another along the c axis. Two neighboring K-F layers are positioned over one another as in the rock-salt-type structure. Compounds such as La2Ni04having this type of structure are important in relation to perovskite transformation in mild reduction conditions. If the slices of the perovskite structure have the thickness of n unit cells, the composition of the crystal is An+,Bn03,,+, . The compounds Sr3Ti*O7and Sr4Ti3010 (26) are typical members of this family. This type of compound shows disordered intergrowth (27). A further variation results from a combination of the above-described structures with those of the tetragonal PbO-type structure. A typical example of the resulting class of compounds., referred to as Aurivillius phases, is BitNbzOsF, where the bismuth atoms combine with the oxygen atoms to form a highly altered polymeric network (28). Other examples of both ordered and disordered intergrowth phenomena from Aurivillius phases have been given by Rao et al. (24). 111. Preparation

A number of methods have been used in the synthesis of perovskites; the choice of a particular one depends mostly on the expected use for these oxides. Obviously, no attention has been paid to textural characteristics of samples whose uses are based on their electric or magnetic properties. However, application of perovskites in the field of catalysis requires solids with a well-developed porous network. As "the present review is concerned particularly with the surface and catalytic properties of perovskites, we will place special emphasis on preparation methods leading to a high surface : volume ratio. Also, methods yielding homogeneous solids will be discussed because of the important effect that inhomogeneities may play in heterogeneous catalysis. A primary characterization of perovskite-type oxides must include textural analysis and X-ray identification of the phase(s) present. For a more detailed characterization, structural analysis for establishing the lattice position of cations and surface analysis (by means of techniques such as XPS) for defining the surface concentration and oxidation states of cations are desirable. Consequently, information provided by these techniques will furnish the essential criteria for comparing the different preparation methods. For convenience, we will classify the methods used to date for the preparation of pure perovskite phases according to the scheme proposed by Courty and Marcilly (29) for the whole field of mixed oxides. Table I gives a survey of methods used as a function of the phenomena on which they are based.

245


TABLE1 Preparation Methods for Perovskite-Type Oxides

Method Reaction Solid-solid Liquid-solid

Physical

Chemical

-

Ceramic Crystallization Coprecipitation Single compounds Mixtures Complexation

Dry evaporation Explosion Spray-drying Free-drying

A. METHODSBASEDON SOLID-SOLID REACTIONS

Solid-solid reactions are the basis of the most frequently used procedures for preparing mixed oxides, especially when the surface areas of the resulting solids are not an important parameter. Indeed, these high-temperature methods are essential for preparing perovskites with special morphologies, such as monocrystals or thin layers. Because this kind of method is most frequently used for the preparation of ceramic materials, it is usually referred to as the “ceramic method.” The solid-solid methods offer the advantage of their simplicity. It suffices to calcine physical mixtures of the single-oxide components or other adequate precursors. Very high temperatures (usually > lO00”C) are required for a complete reaction between the single phases to take place, thus leading to a drastic loss of surface area by sintering. Consequently, surface areas of perovskites that have been prepared by this method are very low. As an example, Voorhoeve et al. (30) prepared by the ceramic method a series of 20 perovskite oxides, almost all of them having surface areas smaller than 1 m2/g. Thus, in spite of its simplicity, this method is not very attractive for preparing perovskite catalysts. A further drawback is the lack of homogeneity of the solids obtained, because of an incomplete reaction between the mixed precursors. Several approaches have been made toward gaining a better degree of mixing and, if possible, lowering the temperature necessary for complete reaction. Besides repeating grinding and firing (a usual procedure in solid-state reactions), joint grinding of precursors in a liquid medium (3I),pelleting of powder mixtures (32), and high-pressure calcination (33) have been tried, this latter method being especially efficient; moreover, use of high oxygen pressures in a belt-type equipment (33, 34) has permitted stabilization of high oxidation states of several elements in the perovskite structure.

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LUIS G. TUUCA

et al.

More recently ( 3 3 , a new method using peroxides as oxidizing agents has been reported to obviate the need of high oxygen pressures. In other cases the nature of the precursors was changed in order to generate in situ single-oxide phases with small particle sizes and highly reactive surfaces. Precursors whose anions can be easily eliminated, such as hydroxides, nitrates, carbonates, acetates, and oxalates, have been preferred for perovskite preparation. High-temperature solid-state methods deserve particular attention in the preparation of perovskite phases with special morphologies such as thin films or single crystals. Thin films of perovskites, used in microcircuitry, are seldom used in catalysis, and their preparation will not be discussed here. Indeed, this subject was reviewed by Galasso (41, who concluded that vacuum evaporation provided the best results in the case of BaTiO3 thin layers. Single crystals of perovskites have been, however, widely used for fundamental studies in catalysis. Among the various preparation techniques also reviewed by Galasso (4,growing from a flux has been the most commonly used. Voorhoeve et al. (36) prepared Lal-,Pb,Mn03 by growing from a lead borate flux heated at 1250°C. Minute flux inclusions were found to significantly alter the surface composition, but they were readily removed by leaching with diluted acid (37). Etching with a more concentrated acid has permitted elimination of residues of unreacted oxides and, in some cases, was used for increasing the surface area of perovskite phases (38,39). The main disadvantage of the flux growing technique is the high-temperature reaction with crucible material, which is then incorporated into the perovskite as an impurity. This has led to undesired presence of platinum in perovskites which, as will be discussed in Section VII, brought about misleading results of catalytic activity for these materials.

B. METHODS BASEDON LIQUID-SOLID REACTIONS The simplest physical method based on a liquid-solid reaction is dry evaporation. It has much in common with the ceramic method, as the homogeneity of the solution is not preserved during evaporation. Thus, a heterogeneous solid is formed that resembles the solid mixtures used as starting material for solid-state reactions, and consequently, inhomogeneous materials with low surface areas are obtained. Arai et al. (40) obtained LaM03 (M = Cr, Mn, Fe, Co, Ni, Cu) perovskites with surface areas in the range of 0.6-4.8 m2/g.However, Gysling et al. (41) reported that LaRhO3 prepared by dry evaporation was much more homogeneous than the same perovskite obtained by the ceramic method. The explosion


247

method proposed by Wachowski er al. (42,43)uses the explosive properties of ammonium nitrate added to a precursor solution of perovskite, which is then dry-evaporated. Surface areas in the order of 30 m2/gwere obtained by carrying out a preliminary vacuum decomposition at 300°C and then calcination in oxygen at 500°C. This method offers the advantage of using simple laboratory equipment and requiring a relatively low expenditure of work. Another possibility for preparing perovskites from a solution consists in increasing the rate of evaporation by powdering the liquid to give a mist that is subsequently spray-dried. Using the two-stage method proposed by de Lau (44),Johnson, Jr. et al. (45) prepared LaMnO3-substituted perovskites with surface areas in the range of 9-17 m2/g. Later on, Imai and Orito (46), using a three-stage method, prepared Lac003 with 12 m2/g. Imai et al. (47) have used several additives dissolved in the starting solution to increase the surface area of the final sample. The best result was obtained with NH4Cl, in which case a surface area of approximately 50 m2/g was attained. More recently, Murphy and King (48) have used spray-drying for preparing L~.sSro.zCOo.sNio.z03, which was subsequently plasma-sprayed on nickel grids (see Section II1,D) to be used as an electrocatalyst. Freeze-drying is probably the physical method that best preserves the homogeneity originally present in the solution. This results in a homogeneous precursor that needs a rather low temperature for complete transformation into perovskite. Tseung and Bevan (49) first prepared Lac003 by means of this method, with a surface area of 38 m2/g. Johnson, Jr. et al. (45) and Wachowski et al. (42,43)have also obtained perovskites having surface areas greater than 30 m2/g. However, in other cases (50,51) high calcination temperatures were needed, thus resulting in surface areas smaller than 4 m2/g. Crystallization of complexes from a liquid phase is, a priori, a method expected to produce extremely homogeneous mixed oxide precursors, although, as Courty and Marcilly pointed out (29),the difficulty in preparing complexes containing the stoichiometric ratio required for the catalyst severely restricts its applicability. Crystallization of oxalate complexes La[M(C204),] * 9H20 (M = Fe, Co) constitutes a typical application of this procedure (52, 53). However, the high temperatures necessary for complete decomposition of these complexes (640°C for Fe and 800°C for Co) reduce the surface areas of the corresponding perovskites. Nag and Roy (52,53)studied in detail the decomposition of precursors. Usha et al. (54) obtained some further insight into these thermal decomposition reactions. Coprecipitation constitutes the most widely used chemical procedure for separating a precursor from the solution. However, two different

248

LUIS G . TWUCA

et al.

cases should be distinguished. In the first one, the cations to be present in the mixed oxide are separated from the solution in the form of a single mixed compound. This method, which has many characteristics in common with that of crystallization, is limited by the stoichiometric ratio of cations in the precursor and in the final mixed oxide. A typical application to perovskites, first reported by Gallagher (55), is the precipitation of cyanide complexes such as L~[CO(CN)~] * 5H20, which are then calcined to LaCoO3. Gallagher (SS), Gallagher and Schrey (56),and Tascdn et al. (57, 58) have studied in detail the thermal decomposition of cyanides leading to the formation of CaCo03, LaFe03, and similar perovskites. Three reaction steps were found: dehydration of the pentahydrate, oxidation of the anhydrous cyanide, and final reaction to yield the mixed oxide (58). Because of the low tepperatures necessary for obtaining the perovskite phases (450°C for Lac003 ; 330°C for LaFe03), surface areas of 37.5 m21g (LaCo03) and 9.5 m2/g (LaFe03)were obtained (57). Other applications of this method to perovskites have included precipitation of Ba[Ti(OH)6](59) or barium titanium citrate as precursors for BaTi03 (60) apd strontium titanyl oxalate, as a precursor for SrTiO3 (61).No surface area data were given for these preparations. The second case refers to the joint precipitation of at least two different compounds. O~viously,this will result ie a very heterogeneous precursor; however, there are no constraints imposed by stoichiometry. For the preparation, of perovskites, coprecipitation as hydroxides, carbonates, and oxalates have been very frequently used. In the case of hydroxides, the precipitating agents were KOH (62),NH40H ( 6 3 , or tetraethylammonium hydroxide (64).Coprecipitation as carbonates was performed by means of (NH4)&03(45,62),K2C03(62),or NH4HC03+ NH40H (65). Oxalic acid (42,57)or ammonium oxalate (52) was used for the coprecipitation of oxalates. In principle, precipitation by means of oxalic acid Seems to be the most adequate method because (1) the presence of alkaline ions, which could contaminate the precipitate, is avoided and (2) the absevce of ammonia or their derivatives prevents a possible selective dissolution of one of the precipitates because of the well-known ability of transition metal ions to form ammoniacal complexes. Gallagher et al, (62) have compared the different coprecipitation methods and indicated that carbonates required higher temperatures than did hydroxides to form the perovskite phase. Oxdates and carbonates required similar temperatures as oxalates form carbonates as intermediates in their decomposition. Tasc6n et al. (57, 58) found Co304, La202CO3, and La203 as intermediate compounds in the formation of LaCo03 from mixed oxalates. Small differences were observed amongst surface areas of perovskites obtained by coprecipitation, which are generally lower

PBROVSKITE-TYPE OXIDES

249

than 10 m2/g(42,45,57,62,63,65), irrespective of the precipitating agent used. An exception was found in a sample of SrTiO3 with a surface area of 64 m2/g (64),, whose precursor had been totally decomposed at 550°C. Complexation in the form of aniorphous compounds with a vitreous structure is another chemical procedure for separating from the solution a solid precursor while preserving as much as possible the homogeneity present in the solution. A method uses the complexing properties of hydroxiacids, in particular citric acid (66,67). It was originally employed for preparing perovskites such as LaCr03(66,67), and LnA103(Ln = Y, La, Sm) (67); no surface area data are available in these cases. Later, Tasc6n et al. (57) prepared the oxides LaM03 (M = V, Cr, Mn, Fe, Co, Ni) by this method, with surface areas ranging from 3.5 m2/g (LaCr03)to 55 mz/g (LaVO,), showing the influence of the transition-metal ion on the decomposition temperature of the corresponding citrate complexes. The mechanistic features of citrate decomposition have been analyzed by Courty er al. (67), Delmon and Droguest (68) and Tasc6n et al. (57), who found three decomposition steps corresponding to dehydration, first decomposition, and pyrolysis; the latter was catalyzed by the presence of cations whose oxides were oxidation catalysts. This may explain the different behavior observed for the first-row transition series of metal ions (57). C. COMPARISON OF METHODS As we have outlined in the preceding discussion, the presence of different metal ions leads to strong differences in the final heating temperature for obtaining a single perovskite phase. Consequently, comparison of the performance of different preparation methods is possible only when data are available concerning the preparation of a specific mixed oxide by different methods. Unfortunately, scarce examples exist in the literature in this respect. In the following, some of these will be discussed. Berndt et al. (69) compared the ceramic and coprecipitation methods for the preparation of interlanthanide perovskites. Although the objective of their work was not to prepare catalysts and, consequently, no surface areas were measured, clear differences were observed among the temperatures necessary with the different methods for obtaining pure perovskite phases. The solid-state reaction of mixtures of separately precipitated hydroxides proceeded much faster than that of physical mixtures of single oxides (ceramic method). This is a clear indication of the influence of particle size in solid-solid reactions. The spray-drying, freeze-drying, and coprecipitation methods were compared in a study carried out by Johnson, Jr. et al. ( 4 3 , concerning the preparation of La,-,M,Mn03 (M = Sr,

250

LUIS G . TEJUCA

et al.

Pb, K, Ce, Co, Ni, Mg, Li) perovskites. Coprecipitation produced wellcrystallized oxides, with surface areas smaller than 10 m2/g. Spray-drying led to solids with some compositional segregation; their surface areas were below 17 m2/g. Freeze-drying permitted the synthesis of perovskites at temperatures low enough to preserve larger surface areas (14-32 m2/g). The relative order of catalytic activity per unit surface area was freezedrying > spray-drying > coprecipitation; according to the authors, this could be due, in part, to the differences in the calcination temperature (i.e., merely changing the calcination temperature would change the surface properties of these oxides). According to Delmon and De Keyzer (701,the chemical liquid-solid methods are used in the great majority of preparations of mixed oxides at both laboratory and industrial scales. Consequently, these should provide the essential guidelines for the preparation of practical mixed oxide catalysts. Tascdn et al. (57) have carried out a comparative study of coprecipitation and complexation methods for the preparation of LaMO3 (M = V, Cr, Mn, Fe, Co, Ni) catalysts. Coprecipitation as a single cyanide compound, or as a mixture of coprecipitated oxalates, as well as the citrate method, was used. Heating temperatures for obtaining a single perovskite phase were higher for oxalate decomposition (900-1000°C) than for cyanide or citrate decomposition. Surface areas varied in the ranges of 2-6 m2/g (oxalates), 10-40 m2/g (cyanides), and 4-55 m2/g (citrates), with a clear influence of the cation in position B of the structure. Wachowski et al. (42,43) have compared the surface areas of a series of eight perovskites prepared by different methods: ceramic ( 0.4. This is due to the instability of Co4+, which tends to be reduced by a simultaneous release of oxygen. The reductive nonstoichiometry,which increases with increasing x, was found to have a remarkable effect in the reducibility of these oxides (see Section V,B) and in their activity for oxidation (95, 96). Gibb et al. (98) studied the series of substituted perovskites SrFe,RU~-,O~ by- ~Mossbauer spectroscopy. The incorporation of Fe in the structure takes place exclusively as Fe3+.Substitution of Ru4+by Fe3+for x < 0.3 leads to the appearance of oxygen deficiency. For x > 0.3 there is an increasing proportion of Ru5+ and a parallel decrease in oxygen deficiency. Thus for x = 0.5 the reductive nonstoichiometry appears to be I

X

0.2 0.3

0.4 0.5 I

400

1

I

I

500

I

600

I

I

1

700

T,OC

FIG.7. Percent mass loss as a function of the temperature for the compositions Ba,La,-,CoO,. (Reprinted by permission from Ref. 93.)

258

LUIS G . TWUCA

et al.

minimal. It seems that this system is unable to tolerate more than approximately 4% oxygen deficiency and, as a result, partial oxidation of Ru4+ occurs. Other oxygen-deficientperovskites were described and characterized. Thus, rather unstable SrVO3-,, (A =z 0.1) oxides have been synthesized by Dougier et al. (99). MacChesney et al. (ZOO) prepared a series of SrFeO3-,, oxides. By decreasing the equilibration temperature from 1400 to 550°C and increasing the oxygen pressure from 2.03 x 104 to 3.41 x lo7 Pa, these authors obtained perovskites with increasing oxygen content from SrFe02.72to the stoichiometric compound. The tendency of this oxide to reductive nonstoichiometry as opposed to the behavior of the lanthanum ferrite (92) seems to be justified by the charge-compensating factor, resulting in an oxygen loss, introduced by the presence of Sr2+in the lattice. Wachowski et al. (92), using the explosion method, obtained the slightly oxygen deficient perovskite LaNi02.9B.Soderholm et al. (101) found that the X-ray diffraction lines of BaCe03 broaden and shift to lower angles on substitution of Ce4+by Dy3+,This has been attributed to the incorporation of oxygen vacancies into the lattice as charge compensators. Then this compound can be formulated as BaDy,Cel-,03-,,. Voorhoeve et al. (30) reported oxidative nonstoichiometry for the perovskites L ~ M O J(M + ~= Cr, Mn, Fe) and reductive nonstoichiometry for Lacoo,-,,. XPS measurements carried out on a similar series of LaMO3 oxides after heat treatment in air at 900°C indicated a surface nonstoichiometry that changes from oxidative (M = Cr, Mn) to reductive (M = Fe, Co, Ni, Rh) (102).This nonstoichiometry was found to be much more marked than that observed in the bulk. The more oxygen deficient perovskites were those that are more easily reducible (M = Co, Ni, Rh). Tabata et al. (103)found also significant differences between the chemical composition of the surface (determined by XPS) and of the bulk (determined by X-ray fluorescence spectroscopy) in a series of SrxLal-,Coo3 oxides. These results indicate a very different behavior of the surface with respect to the bulk in these compounds. This is an important factor to be considered when trying to correlate the composition of a perovskite with its catalytic performance. V. Stability in a Reducing Atmosphere

A. TEMPERATURE PROGRAMMED AND ISOTHERMAL REDUCTION Temperature-programmed reduction (TPR) experiments were performed mainly on LnMO3 oxides, where Ln and M are, respectively, a rare earth and a transition element. Reduction of LaRh03 with CO (4"C/


259

min) (104) resulted in a weight loss equivalent to a reduction of about 75% of Rh3+to Rho at 580°C (reduction up to 1000°C in H2 or CO affects only the metal M; reduction of the rare earth cation requires higher temperatures). After this the sample weight increased with temperature because of carbonate formation or carbon deposits. Formation of carbonaceous species has been observed in the reduction of Lal-,Sr,Fe03-A with CO (105). To avoid this effect, the reduction is carried out most frequently in a hydrogen atmosphere. Fierro et al. (106) and Tascdn et al. (107) studied the series of LaM03 oxides (M = Cr, Mn, Fe, Ni, Rh) (reduction took place in H2 at a heating rate of 4"C/min unless otherwise specified). LaC1-03 ( 1 0 6 ~exhibited ) the highest stability. At 1000°C this oxide undere- per molecule (3e- per molecule went a reduction of only 1.3 x would amount to full reduction of M3+to MO).LaMn03 (106b) showed a single reduction step of le- per molecule at 800°C (temperature where the weight loss corresponding to the indicated reduction degree is attained). Vogel et al. (10% H2-N2,2OoC/min) (81) reported, however, a first reduction step of LaMn03.],to the stoichiometric compound at 550°C and then a second reduction of le- per molecule (to Mn2+)at 1050°C. TPR data with reduction of Mn3+to Mn2+in the partially substituted perovskites LaMno.5C&.503and L ~ o . & . ~ M ~were O ~ also given by Vogel et al. (81) and Vogel and Johnson (108). LaFe03 (1074 reduced to 3e- per molecule at 1OOO"C. This temperature is practically coincident with those found by Wachowski et al. (19% H2-He, 20"C/min) (92) and Carreiro et al. (15% Hz-Ar, 0.4"C/min) (109) for full reduction of this perovskite. The total weight loss observed was higher than that expected for reduction of Fe3+ to metallic Fe (92, 1074, indicating that the lanthanum and iron oxide exhibit oxidative nonstoichiometry (LaFe03.20).In addition, Wachowski et al. (92) found a plateau for a weight loss of 4.2%, close to that that would be expected for reduction of Fe3+to Fe2+.Reduction of YFe03 (15% H2-Ar, 0.4"Clmin) (109) to 3e- per molecule took place at a substantially lower temperature (740°C). LaCo03 (D2,3"C/min) (64)and LaNi03 ( 1 0 6 ~both ) exhibited two reduction steps of le- and 3e- per molecule at 450 and 625°C (LaCo03)and 325 and 475°C (LaNi03). A similar reduction behavior has been reported by Wachowski et al. (92) and by Futai et al. (110) for LaCo03, LaNi03, and LnCo03 perovskites. Crespin et al. (111) and Levitz et al. (112) detected, however, three reduction steps for LaNi03 (H2, O.OS"C/min) of 1, 2, and 3e- per molecule at 300, 370, and 455"C, respectively. LaRh03 (107b) reduced in a single step to 3e- per molecule at 600°C. Carreiro et al. (109) and Gysling et al. (41) recorded lower temperatures for reduction (15% H2-Ar, 0.4"Clmin (109);H2, lO"C/ min (41))of LaRh03 (430"C), and YRh03(400°C). No intermediate reduction state has been detected for rhodium in these perovskites. Rh(1) has

260

LUIS G . TEJUCA

et al.

been found in the reduction by heating in vacuum of Rh-Y zeolites (113, 1 14). The ease of reduction increases, therefore, from Cr3+to Ni3+ in the series of LaM03 perovskites. The same trend has been found for the simple oxides of Fe, Co, and Ni (115). On the other hand, the mixed oxides LaM03 (M = Fe, Co, Ni) were found to be more stable in a H2 atmosphere than the simple oxides NiO, Fe203, and Co304(92); this shows the increased stability of transition-metal cations in a perovskite structure. The TPR experiments also show that the stability of perovskite oxides increases with increasing size of the A ion. Different TPR steps correspond to different reduction mechanisms. These can be studied by kinetic experiments of reduction in isothermal conditions (116). As illustration, reduction data of PrCo03(117) are given in Fig. 8. The TPR diagram (Fig. 8a)shows reduction of Co3+to Co2+ (step a) at 400°C and of Co2+to metallic Co (step b) at 525°C. The plots of (I! (reduction degree) versus t (time) for step a in Fig. 8b show that the reduction rate decreases continuously with time, indicating that the reduction occurs according to the contracting-sphere model. The process starts with a very fast nucleation, which results in a total coverage of the PrCoO3 grains by a thin layer of the reduced phase. This causes a continuous decrease in the rate of the interfacial reaction as the grains of the starting oxide are consumed in the reaction. The reduction process is described by the Mampel intermediate law ( I 18)

1 - (1 - a ) " n

=

kt

+b

(1)

where cy is the reduction degree; k and b are temperature-dependent constants; and n = 2 would correspond to a two-dimensional growth and n = 3, to a! three-dimensional growth: This equation holds over the middle range of (I! values. Kinetic runs in step b in Fig. 8c started with a very fast reduction of approximately le- per molecule, after which a slow reductioh took place, yielding sigmoidalJreductioncurves. This. indicates that reduction of Co2+ to Coo is controlled by the formation and slow growth of reduction nuclei of metallio cobalt on the surface of the reduced phase in step a (nucleation model). Initially, the reduction rate increases because of the growth of nuclei already formed and the appearance of new ones.'At a certain point the reduction nuclei start to ovedap; at the inflection point, the interface of ,the oxidized and reduced phases and the reduction rate both begin to decrease. Reduction of this type i$ described by the Avrami-Erofeev equation (1 18) 1 - (I! = e-kr' (2)

26 1


1.o

0.8 0.6 0.4

0.2 0.0

I

1

0

20

I

40

I

I

60

8U

I

too

t,min Reduction of PrCoO, : (a) temperature-programmed reduction; (b, c) isothermal reduction in the first (to le- per molecule) and second (to 3e- per molecule) steps, respectively, of diagram (a) (90cmVmin H j flow rate;4"C/min heating rate). (Reprinted by permission of the publisher, Chapman & Hall, from Ref. 117.)

262

LUIS G. THUCA

et al.

[where a,k , and n have the same meaning as in Eq. (l)], which holds over a values of 0.05-0.90. Kinetic studies were done with other perovskites. For example, LaNiO3 ( 1 0 6 ~ exhibited ) a behavior similar to that of PrCoO3; LaFe03 (107a)and LaRh03 (107b)reduced to the metal according to the contracting-sphere model. It should be noted, however, that the distinction between the two reduction mechanisms is somewhat arbitrary because the contracting-sphere model starts with a very rapid nucleation, and the nucleation mechanism ends according to a contracting-sphere model (note that reduction after the inflection point of the kinetic curves in step b in Fig. 8c is similar to reduction in step a in Fig. 8b. B. EFFECTOF CATIONS IN POSITIONS A

AND

B

Arakawa et al. (119) showed that the extent of reduction in H2 of LnCoO3 oxides increased from LaCo03to EuCo03. In the same direction the energy of the metal-oxygen bond decreased. Futai et al. (110) performed TPR experiments for a series of LnCo03oxides and found that the temperature for maximum consumption of H2 (Tmax)for the reduction steps to 1 and 3e- per molecule decreased from Lac003 to EuCo03; however, T,,, increased of remained constant for the perovskites of Gd, Tb, and Dy. Carreiro et al. (109) observed a higher stability in H2 of LaFe03 and LaRh03 than the corresponding yttrium perovskites. In concurrence with these results, Sakai et al. (120) found that the strength of the V-0-V interactions in LnVO3 oxides is weakened by the substitution of La3+with heavier lanthanide ions. On the other hand, Sugihara et al. (121) and Katsura et al. (122) calculated the Gibbs free energy of formation AG of a series of LnFe03 oxides from metallic iron, LnzO3, and oxygen at 1200-1297°C and found a linear relationship between AG and the tolerance factor t , as defined by Goldschmidt (19). This is illustrated in Fig. 9, where it is observed that AG becomes more negative and the change in entropy (AS)decreases for increasing values of t . A similar relationship holds for the Gibbs free energy of formation of LnFeO3 oxides from Fez03 and LnzO3 (12.3). However, in this case the plots AG versus t were found to be independent of the temperature from 877 to 1024°C. These data are consistent with the TPR results and suggest that the stability (or reducibility) of perovskite oxides increases (or decreases) with increasing size of the A ion. This conclusion would be consistent with the preferred occupancy of the larger Ln3+ ions in the 12- coordinated A sites. Partial substitution of the A ion by an ion of different charge may also cause significant changes in reducibility; for instance, when

263


1

I 1200oc

- -30 h

--28

?

Q)

v

--26

-55

Yb '

Dy Tb Gd Eu Sm 0,93

Nd 0,94

2

La 0,95

t FIG.9. The Gibbs free energy of formation (AG) of a series of LnFeOl oxides from Ln203,metallic iron, and oxygen (solid lines) and the change of entropy (AS, dashed lines) as a function of the tolerance factor t . (Reprinted by permission from Ref. 122.)

La3+is partially replaced by Sr2+in LaCoO,, the charge compensation is accomplished by oxidation of Co3+to Co4+and by the formation of an oxygen-deficient perovskite, Lal-,Sr,C003-~ (see Section IV). The increasing concentration of both Co4+ (unstable) and oxygen vacancies (which facilitates the diffusion of lattice oxygen from bulk to surface) accounts for the increased ease in the reduction of this oxide with increasing strontium content (95, 96). Nakamura et al. (82) studied the structural changes undergone by LaM03oxides in a reducing atmosphere at I bar pressure (1 bar = lo5 Pa) and 1000°C as a function of the O2partial pressure, Po,. Their results (Fig. 10) are qualitatively in agreement with those discussed above (Section V,A). La2C020s,La2Ni205,and LaNi02are unstable at high temperatures (111,112) and thus were not observed. Instead, the dismutation products La2M04were found. Neither has the formation of the La2Mn04 phase been observed, since it is unstable above 925°C (108). The order of stability expressed in terms of the -log Po, values appears to be LaV03 = LaCr03 > LaFe03 > LaMn03 > LaCo03 > LaNi03. This is the order of the Madelung constant except for Lac003 and LaNi03, which have higher constants. According to Nakamura et al. (82), a cause for this

264

et al.

LUIS G . TUUCA

-2

-3

0

1

1

1

1

1

1

1

1

1

10

5

1

1

1

1

15

20

- log Po*

La R O3

-3

-o\" o

-4

n

-7 -5 -6

l

,

3 -8 ' -9 3 -10

l ,+Fe

a

*

-1 2

-log Po2 FIG.10. Weight and structural changes undergone by LaMOl oxides in a reducing atmosphere at lo-' Pa and 1OOO"C. (Reprinted by permission of the publisher, Pergamon Journals, Ltd., from Ref. 82.)


265

discrepancy may be the existence of stable compounds La2M04for Co and Ni. These results indicate the same stability order found by means of TPR. It should be noted, however, that both sets of data would indicate a higher stability of LaMn03 relative to LaFeO3 if reduction of M3+to Mo is considered. Carreiro er al. (109) discussed the higher stability toward reduction of LaFe03 and YFe03 as compared with the corresponding rhodium perovskites in terms of the standard Gibbs free energy of formation of the corresponding sesquioxides Fe2O3 (- 177 kcal/mol) and Rhz03(-49 kcal/ mol): the less negative the value of AGO, the greater the tendency for dissociation to the metal and oxygen. The above stabilities could be explained considering that similar relative values in AGO can be expected for the mixed oxides LnFe03 and LnRh03. ON REDUCTIONA N D OXIDATION C. PHASETRANSFORMATIONS

Perovskites of titanium and chromium such as SrTi03,BaTi03( 6 4 , and LaCr03 (106a) remained virtually unchanged even after reduction at 1OOO"C. Other perovskites of transition metals underwent substantial changes after less severe treatments. Reduction of LaMn03 to le- per molecule by isothermal heating in H2at 720°C (this temperature is, of course, somewhat lower than that needed to attain the same reduction degree in a TPR experiment) yielded MnO and La203 (106b).By reduction of alkali-substituted LaMn03 above this temperature, Vogel and Johnson (108) reported a breakup of the perovskite structure with formation of La2Mn04.However, when the heating is carried out below 700"C, the initial structure can be preserved; e.g., LaMno,5Cuo.502,93 after reduction to LaMno.sC~.S02.26 (containing Mn2+ and Cu') showed X-ray diffraction (XRD) lines of a perovskite phase only (81). After reduction of LaFe03.1sto approximately 1.2e- per molecule, the perovskite structure was partially preserved ( 1 0 7 ~ )besides, ; lines of a-Fe and Laz03 were observed. Further reduction led to the disappearance of the perovskite. Sis et al. (124), by means of thermogravimetric, calorimetric, XRD, and magnetic measurements, showed that Lac003 does not reduce directly to Co and Laz03but rather through the formation of intermediate oxygendeficient structures. Studies carried out by Wachowski et al. (92) on reduced LaM03 (M = Fe, Co, Ni) and by Arakawa et al. (119, 125) on reduced LnCo03 oxides (Ln = La, Nd, Sm, Eu) point to the same conclusion. In the initial weight-loss region of LaCo03 (0-3.25 wt% or a reduction lower than le- per molecule) the oxygen-deficient compound LaCo,03-, (monoclinic)and metallic cobalt are formed. At a weight loss

266

LUIS G . TEJUCA

er al.

of about 3.25 wt% ( l e - per molecule) the original rhombohedra1 LaCo03 with a loss of lanthanum and cobalt appears. At 5.5-6.5 wt% loss (about 2e- per molecule), Coo and La2Co04 are formed (124). This latter compound was previously synthesized and characterized by Janecek and Wirtz (126). Crespin and Hall (64)found that the perovskite structure of LaCo03was slightly modified by reduction to le- per molecule. The XRD pattern of the sample reduced to 3e- per molecule (at 400°C) showed only La203 peaks, indicating that the metallic cobalt was highly dispersed; when the sample was heated in He at 800"C,the expected Coo lines appeared. These results are consistent with those of Sis er al. (124). The reduction step of le- per molecule in Ref. 64 would correspond to the formation of L ~ C O ,in~Ref. ~.~ 124, which is then transformed according to Eq. (3):

-

L~CO,O~.~

+ $ C03,-2 + Q La20,

f LaCoO,

(3)

According to Crespin et al. (111, 112), reduction of LaNi03 to le- per molecule at low temperatures (300°C) occurs according to the following transformation : 2LaNiO3 $. H2

-

La2Ni203+ H20

(4)

A comparison between the observed distances obtained from XRD data of the reduced sample and the distances calculated from the proposed unit cell demonstrated the existence of a single phase with the stoichiometry of LazNi~Os.Similar compounds of Co and Ni were prepared and described by Vidyasagar er al. (127). Reduction of LaNiO3 to 2e- per molecule above 300°C also yielded LazNi20~besides metallic Ni and La2O3. However, when the reduction is effected below 300"C, LaNiO2 is formed. Its structure has been described using XRD and EXAFS techniques. Both La2Ni20~and LaNi02 are stable in air at room temperature. However, after heating at 1000°C in helium a dismutation reaction occurs with formation of La2NiO4, NiO, and NiO. These results are in agreement with those of Fierro et al. ( 1 0 6 ~ 1who , found formation of La2Ni04after reduction of LaNiO3 to le- per molecule at 240°C and heating in helium at 800°C. La203 is the only phase that is detected after reduction of LaNi03 to 3e- per molecule at 300°C. Similar to that observed for LaCo03 (M), the lines of Nio appear in the XRD pattern only after heating at 800°C in an inert atmosphere (1 1 I). Perovskites may undergo reversible reduction-oxidation cycles when these are carried out at temperatures where sintering of the oxidized or reduced species does not occur. Thus, reoxidation at 400°C of LaCoO, reduced to 3e- per molecule fully restores the perovskite structure. However, reduction to 3e- per molecule, heating in He at 800°C and reoxida-


267

tion as above did not produce a perovskite phase. Instead, Co304appeared (64).This irreversibility in the redox cycle is caused mainly by the drastic increase that the particle size of the metal undergoes in the sintering process. This increase for metallic nickel (after reduction of LaNi03at 430°C and heating in He at 900°C) was found to be larger than an order of magnitude ( 1 0 6 ~ )The . reduction products (up to le- and 2e- per molecule) of LaNi03 (LaZNizOs and LaNiOz; see above) were fully reoxidized in O2at 180°C (111). Vidyasagar et al. (127) reported a somewhat higher reoxidation temperature (325°C) for LazNizOs.In the reoxidation at 180°C of the fully reduced (3e- per molecule) LaNi03 a strongly exothermal reaction was observed and a mixture of LazNi04 and NiO was obtained (111). This result is similar to that observed for reduction at 500°C of LaCo03 to Coo and La203 and reoxidation at 400°C with pure 0 2 (64). In these conditions substantial amounts of Co304 plus La203 besides perovskite were obtained. The formation of simple oxides may be avoided by carrying out the reoxidation at a slower rate, using diluted oxygen in an inert gas or air as an oxidizing agent. Reversible reduction-oxidation processes were also reported for LaRh03 (107b) and PrCo03 (117). In these cases the particle sizes of the regenerated perovskites were found to be smaller than those of the starting samples. Reller et al. (128)presented evidence by XRD, high-resolution electron microscopy, and selected area electron diffraction of CaRu03 reduction to Ruo and reoxidation at low temperatures to form a pure perovskite phase. When these redox processes were repeated, the size of the Ruo crystallites decreased beyond the detection limit of X-ray diffraction. This and other examples mentioned above [Lac003 (64),LaNiO3 (111)l show that by reduction treatments or redox cycles, the metal in position B is in a highly disperse state on a matrix composed by the oxide of the metal in position A. Considering the importance of dispersed metals in heterogeneous catalysis, the reduction or reduction-oxidation of perovskites in controlled conditions may represent a promising pathway for the preparation of highly active catalysts. D. XPS STUDIES The surface reduction of perovskites (mainly cobaltites) has been studied by examining the evolution of the XPS photolines of the metal in position B and oxygen. Lombard0 et al. (129) reported a multiplet splitting of the C02p photolines in oxidized and reduced LaCo03 of 15.4 eV (Co2p3l2at 779.9 - 779.3 eV and C02p,/~at 795.5 - 795.0 eV for Co3+ ions), which is in excellent agreement with that given previously by Ichi-

268

LUIS G . TFJUCA

f?t

Ul.

mura et al. (63) for unreduced LaCo03. Formation of Co2+on reduction can be detected by the appearance of two shake-up satellite peaks that are situated at about 7 eV upscale from the main C02p peaks (63) [8.4 eV upscale in SrCo03-, (130)].These satellites are characteristic of Co2+and can be considered as the fingerprints of this ion since the lines of the multiplets for both Co3+and Co2+ions appear very close in the spectrum. The C02p photolines for metallic cobalt are situated at 1.5-2.0 eV downscale from those of Co3+.Following the evolution of the C02p photolines, Marcos et al. (51) studied the surface reduction of pure and Sr- or Thsubstituted LaCo03as a function of the bulk reduction (Fig. 11). At a low extent of reduction these oxides are homogeneously reduced. However, for a bulk reduction of 1.5e- per molecule, the surface is reduced to approximately 3e- per molecule. At this point the maximum difference between bulk and surface reduction occurs. It is observed that the surface of the Sr-substituted perovskite is more easily reducible than that of LaCo03, in agreement with Nakamura et al.'s data for bulk reduction of Lal-,Sr,Co03-k (95, 96). Likewise, Gysling et ul. (41) observed that

"

0

I .5

3.0

3.6

Bulk reduction (e-/mol ) FIG. 1 I . Correlation between surface reduction, determined by XPS and bulk reduction, calculated from the hydrogen uptake in a gas recirculating system. (Reprinted by permission from Refs. 51 (0)and 129 ( O ) . )

269


whereas the bulk of LaRh03 underwent reduction to Lat03and Rh in H2 at 430"C, the surface reached the same extent of reduction at 300°C. These results illustrate the higher reactivity of the surface relative to the bulk in these perovskites. The XPS spectra of perovskites present two 0 1 s photolines that correspond to two different oxygen species (63,94,103,129,131).Yamazoe et al. (94), by means of studies on the evolution of the 0 1 s photolines with the outgassing temperature and also with x in the perovskite La,-,Sr,Coo3 in combination with a TPD study of oxygen, associated the 0 1 s photolines with lattice (lower binding energy B E ) and absorbed (higher B E ) oxygen, although this latter fraction should also contain adsorbed oxygen. Some representative examples of BE values of these photolines and their assignation for some perovskites (51, 63, 94, 102, 117) and simple oxides (132) are given in Table 111. Marcos et al. (51) and Fierro and Tejuca (102) reported an intensity increase of the 0 1 s line at higher BE on reduction of Lac003 and PrC003 at increasing temperatures (1500°C). This was assigned to the formation of hydroxyls in the presence of the water generated in the reduction process. After reduction of Tho.2Lao.sCo03 at 300-500°C no distinguishable OH- signal was detected. However, after reduction of Sro.4Lao.6Co03 at temperatures as low as 250°C the lattice oxygen signal is lost under the tail of the OH- signal, i.e., TABLE 111 0 1 s Photoemission Lines in Oxides ~~~

~

Compound LaCoO, LaCoOJ SrxLal-,CoOl LaMO,

PrCoO, c0104

NiO

a

Adsorbed.

* Lattice.

Absorbed.

~

~

BE (eV)

Assignation

Reference

531.9 528.3 532.2 529.7 530.2 - 531.4 528.2 531.4 - 532.6 529.1 - 530.0 530.9 528.4 531.0 529.6 53 1.4 529.4

Oxygen" Oxygenb Hydroxyl anion" Oxygenb Oxygen" Oxygenb Oxygen or hydroxyl aniona Oxygenb Oxygen or hydroxyl anion4 Oxygenb Water or hydroxyl anion" Oxygenb Water or hydroxyl anion" Oxygenb

63 51

94

102 117

132 132

270

LUIS G . TUUCA

et al.

the hydroxyl concentration increases or decreases in the reduced cobalt perovskite where La3+has been partially substituted by Sr2+or Th4+ions (51). This effect is in concurrence with those of Yamazoe et al. (94) and Tabata et al. (103), who observed an increase of the 01s signal at higher BE with increasing values of x in LaI-,Sr,Co03. The above results are consistent with the higher reducibility of the Sr-substituted perovskites as compared with Lac003 (51, 95, 96) and also with the higher basicity of Sr2+with respect to La3+or Th4+and, therefore, with its higher tendency for hydroxide formation. VI. Adsorption Studies

This section refers to adsorption studies that have been directed to the characterization of perovskites or to the determination of the role of adsorbed species on the catalytic activity. These include mainly equilibrium and kinetics of adsorption, successive or simultaneous adsorption of two gases, infrared spectroscopy (IR), and temperature-programmed desorption (TPD). Physical adsorption has been discussed in Section I11 because of the dependence of the textural characteristics of these oxides on the preparation method.

A. HYDROGEN AND OXYGEN Crespin et al. (133) found associative adsorption of hydrogen at -5°C on LaNi03 reduced to an extent of le- per molecule; the coverage reached was lower than that reported by Tejuca (134) on LaNi03 reduced at 300°C. Ichimura et al. (135) reported TPD spectra after adsorption of H2 on LaAI03, LaCrO3, and LaCoO3 at 25°C showing superposed desorption peaks. The species that desorb above 70°C were considered responsible for C2Hs hydrogenolysis and C2H4 and C2H2 hydrogenation. However, TPD spectra reported after H2 adsorption at 25°C on reduced LaM03 oxides showed one peak at 65-105"C, except for the system H2-LaNiO,, which showed also a second desorption peak at 335-340°C (134). The intensity increase of these peaks with the reduction temperature of the oxide suggests that hydrogen adsorbs on reduced transition metal ions. The interaction of oxygen with perovskites has been studied mainly because of the importance of these materials as oxidation-reduction catalysts. Data of oxygen adsorption on LaMO3 (M = Cr, Mn, Fe, Co, Ni) oxides were reported by KremeniC et al. (136). The adsorption profile at 25°C showed two maxima for Mn and Co (Fig. 12) that coincide with the maxima observed by Iwamoto et al. (137)for the respective simple oxides


27 1

FIG.12. Total (open symbols) and reversible (filled symbols) adsorption profiles of O2on LaMOl oxides on a clean surface (circles) or on a surface with preadsorbed isobutene (triangles) ( P= 2 x l(r Pa; T = 25°C). (Reprinted by permission from Ref. 136.)

of transition metals and also for the catalytic activity of oxidation of this series of perovskites (136) (see Section VI1,B). The reversibly adsorbed oxygen represents a small fraction of the total adsorption, in agreement with Nakamura et d ’ s results (95). The oxygen-adsorption heat for Lac003 (138) was found to be higher than the adsorption heat for LaCr03 ( 1 0 6 ~ that ) showed a lower oxidation activity. Shimizu (139) found a clear correlation between the oxygen adsorption and the metaloxygen binding energy in this series of LaM03 oxides. Yamazoe et al. (94)and Seiyama et al. (140) reported the appearance of two oxygen desorption peaks after oxygen adsorption on Lal-,Sr,Co03 at 800°C. The low-temperature peak (a)can be ascribed to absorbed and adsorbed oxygen, whereas the high-temperature peak ( p ) was assigned to lattice oxygen. These oxygen species were also detected by XPS. The influence of x in Lal-,SrxM03 (M = Mn, Fe, Co) on the intensity of the low-temperature peak (a)has been explained on the basis of the nonstoichiometry and defect structure of these oxides (94, 95, 140-142) (Fig. 13a). Thus, for M = Fe and Co, the Sr substitution increased oxygen desorption (a).Since A-site substitution with a divalent ion is expected to lead to the formation of oxygen vacancies in these oxides, this result suggests that the low-temperature species is associated with oxygen va-

272

LUIS G. TUUCA

et al.

i"

9

!\

-'cn

3.0-

>

b

E

i2.0-

t In

2

Temperature pC

1.0-

Temperature ,"C

FIG.13. Temperature-programmeddesorption of oxygen from La,-,Sr,MO,+, (M = Co, Fe, Mn) (a) and from LaMO, (M = Cr, Mn. Fe. Co, Ni) (b). Oxygen preadsorption, 80025"C, 1.33 x 104 Pa. (Reprinted by permission of the American Chemical Society, from Ref. 140.)

cancies. In LaMn03+A the partial substitution of La3+by Sr2+for x = 0.2 decreases both the cation vacancies and Mn4+concentration without the formation of oxygen vacancies (see Section IV). Therefore, a desorption is not promoted. For x = 0.4, oxygen vacancies are formed and a small a desorption appears (140).The high-temperature peak ( p )were found to be more specifically associated to the B cation, although it is also affected by A-site substitution. This is illustrated in Fig. 13b, where peaks of oxygen desorption from unsubstituted LaM03 perovskites are shown. Oxygen adsorption and desorption were found to be larger for the perovskites, showing a higher catalytic activity for total oxidation, that is, perovskites of manganese, cobalt, and nickel (Fig. 12) (136,140). Seiyama et al. (140) associated the onset temperature for these large desorption peaks with the decomposition temperatures of the respective cgmponent oxides. George et al. (143) found that LaCoO, did not adsorb CO or 02 separately but adsorbed both gases from their mixtures. Enhanced adsorption of oxygen after CO-02 (138,144) and isobutene-02 (136,145) successive adsorption on this perovskite has also been observed by several authors. Some of thcse results are illustrated in Fig. 12, which shows that the enhanced adsorption measured after isobutene preadsorption is greater for the oxides (LaFeO3 and LaCo03) adsorbiqg larger amounts of the


273

hydrocarbon (136). This seems to indicate that the increase in 02 adsorption is directly related to the amount of preadsorbed hydrocarbon. This effect causes the maxima for LaMn03 and LaCoO, on a clean surface to become attenuated by isobutene preadsorption. This behavior has also been described for other systems (146). According to Weller (146), enhanced adsorption of one reactant in the presence of a second reactant accounts for the negative adsorption constants, when kinetic data are analyzed by using rate equations of the Langmuir-Hinshelwood type. The opposite effect has been observed for COz preadsorption, which moderately decreased both the irreversible and reversible subsequent adsorption of 0 2 on LaCo03 (138). Since it was shown that 02 and C02 adsorb on different surface centers (138, 147) this result was ascribed to steric hindrance. Other techniques have been used to study adsorbed or lattice oxygen in perovskites. Thus, XPS permitted to follow the changes undergone by adsorbed oxygen before and after CO adsorption (148)or on outgassing at increasing temperatures (94). In addition, this technique provided evidence of large departures from stoichiometry at the perovskite surfaces (65,102). Electron spin resonance (ESR) has been used to identify oxygen species on Lal-,Ca,Mn03+k (65) and Lac003 (138).After adsorption of oxygen on this latter perovskite at 150°C and lowering the temperature to -196"C, a signal attributed to 0; was found that changed rapidly into a diamagnetic species. Conductivity changes as a function of 02 pressure were associated to a fast adsorption process followed by transformation of adsorbed oxygen into lattice oxygen (149).DTA has been used to study oxygen uptake on cobalt perovskites in order to obtain a relative measurement of the degree of nonstoichiometry (150).

B. NITRICOXIDEA N D CARBON MONOXIDE Comparison of NO adsorption at 25°C on the series of LaM03 oxides revealed a profile similar to that found for 02 showing maxima for Mn and Co (151). NO adsorption was found to be independent of temperature for some perovskites such as LaFe03 (152) and LaNi03 (151) for a wide temperature range (0-400°C). This suggests that the surface sites for NO on these oxides did not change substantially in character with the temperature. For LaCrO3 (153), LaMn03 (154), and LaRh03 (107b) a narrower interval was found, where NO adsorption changes moderately ( ~ 2 0 % ) with temperature. Coverages were of the same order of magnitude as those reported for simple oxides (155).

214

LUIS G . THUCA

et al.

The infrared spectrum obtained after adsorbing NO on LaMn03 above room temperature included bands at 1910 cm-l of dinitrosyl species; at 1610, 1485, 1135, and 1045 cm-* of bidentate and monodentate nitrates; and at 1300 cm-I of nitrite structures (154). The intensity of these bands increased with temperature, showing a progressive chemisorption character for the adsorption process. Similar species were found for LaFeO3 (152); for this oxide, however, formation of N2O at 100°C and higher temperatures was observed, suggesting that NO adsorbs in dissociative as well as molecular forms. Voorhoeve et al. (156)found also a TPD peak at 100-250°C assigned to nitrosyl groups on low-valence metal ions after NO adsorption on potassium- and ruthenium-substituted LaMn03 perovskites. These results indicate that NO interacts with both cations and anions on the surface of these perovskites. Quantitative measurements of the successive NO-CO and CO-NO adsorption on LaMO3 oxides at 25°C showed an inhibiting effect of NO on subsequent CO adsorption, which was found to be larger than the inhibiting effect of CO on NO adsorption (151, 152). Thus, NO appears to be more strongly adsorbed than CO on the perovskite surface, in agreement with previous results of Shelef et al. (155) for other transition-metal oxides. Indeed, a comparative analysis of the adsorption energetics of NO and CO on LaCr03 (153) showed higher adsorption heats and lower adsorption entropies for NO. Supporting these results, infrared bands of carbonates formed after CO adsorption at 25°C on LaMn03, which contained preadsorbed NO, were eliminated by a thermal treatment at 300°C (154). However, carbonates formed after CO adsorption on a fresh perovskite sample showed a higher stability (50). Infrared spectra recorded after simultaneous adsorption of NO + CO on LaMnO3, LaFeO3, and LaCoO, at 300 and 500°C provided evidence of the presence of N 2 0 , isocyanate species, and NO adsorbed with a donortype or coordinative bond (151). An additional band of nitrosyl groups was detected on LaCo03. These results may provide some clues for the mechanism of NO + CO reaction on these oxides, as N20 and isocyanate species have been previously suggested as intermediates for this reaction on simple oxides (157). The chemisorption of NO seems to play an important role in this reaction catalyzed on perovskites. Thus, Chien et al. (73) observed a higher NO adsorption rate for activated (reduced) than for unactivated LaCo03 and L%.~~B%.&oO~. Also, a higher adsorption rate was found for the Ba-doped oxide, which exhibited a higher degree of catalytic activity. The relative constancy of NO adsorption with temperature and the strength of its bond with perovskite surfaces have suggested the use of this molecule over CO for determining surface metallic centers (107b,


275

151-154). However, the IR evidence indicates that NO does not show any particular specificity for adsorption on metallic or oxide ions (151, 152, 154); therefore, the assumption of a 1 : 1 correspondence between adsorbed NO molecules and adsorption centers, as that reported for simple oxides (155) may not give a proper estimate of the number of surface transition-metal ions. Ulla et al. (258) used the poisoning effect of NO adsorption in ethylene hydrogenation at -20°C for the estimation of metallic centers in reduced LaCoO3. Active-site concentration was found to be lower by one order of magnitude than the theoretical concentration of metallic cobalt. This was assumed to be due to the fact that only a small fraction of the metallic sites are active for hydrogenation. The adsorption profile for CO on the surface of LaMO3 oxides at 25°C showed a maximum for LaFe03 (151) resembling that found for isobutene adsorption (136). The extent of CO adsorption was found to be substantially lower than that of NO adsorption (151). Data on the kinetics of CO adsorption on LaCr03 fitted satisfactorily the integral Elovich equation (159). Analysis of plots of time versus the reciprocal of the adsorption rate suggests the existence of a precursor state previous to the adsorption process. Experiments of CO2-CO successive adsorption on LaCoO, strongly suggest that these molecules adsorb on the same center (160). On the contrary, CO adsorption on a surface with preadsorbed oxygen was practically equal to that measured on a clean surface showing the noncompetitive character of the adsorption of these molecules (160). The IR and TPD spectra obtained after CO adsorption on LaM03 oxides in the temperature interval 25-500°C showed the presence of different types of carbonate besides linear and bridged CO (50, 151, 154, 160162). Thus CO, as NO, interacts with surface oxygen and metallic ions. Carbonate formation increased for increasing adsorption temperatures. This was particularly so for LaMnO3 (50), where large increases in the intensity of the infrared bands were observed for adsorption temperatures above 200°C. Adsorption heats determined from the TPD peak temperatures showed that CO is slightly more stable on the Rh3+ cations of the LaRh03 than on metallic rhodium (163, 164).

C. CARBON DIOXIDE, SULFUR DIOXIDE, WATER,AND PYRIDINE Studies on equilibrium and energetics of C 0 2 adsorption have been carried out on LaCr03, LaFeO3, and LaCo03 in the zone -78,500"C by Tejuca et al. (147,165-167). Surface coverages followed the general trend LaCo03 > LaFe03 > LaCrO3. The adsorption isobars on LaCr03 and LaCo03 showed activated adsorption above approximately 150"C,

276

LUIS G . TEJUCA

et al.

whereas on LaFeO3 the coverage decreased continuously with temperature. In these systems the Freundlich model of adsorption, which assumes an exponential decrease of isosteric heat with the coverage, was obeyed. Results of adsorption of C02 and other molecules have led to propose the Freundlich equation for the estimation of surface areas of oxide catalysts (167). Experiments on successive 02-CO2 and CO-CO2 adsorption on LaFeO, (166) and Lacoo, (147) at 25°C were in agreement with previous results of 0 2 and CO adsorption; that is, they showed noncompetitive and competitive adsorption, respectively. Isobutene preadsorption did not affect in a significant way the subsequent C02 adsorption on LaCo03 (145). Kawai et al. (168) found IR bands of monodentate and bidentate carbonates on C02 adsorption on BaTiO3. Likewise, IR spectra after C02 adsorption on LaMO3 oxides in the interval 25-500°C contained bands of free carbonates and monodentate, bidentate, and bridged carbonates (50, 147,165,166). In some cases, transformation of monodentate into bidentate carbonates was observed with increasing temperature (166). The width of the IR bands, together with the fitting of the equilibrium results to the Freundlich model, was interpreted as indicating heterogeneity at the perovskite surface. TPD spectra after C02 adsorption at 25°C on oxidized and reduced LaMn03 (161) and LaCo03 (162) provided further evidence of the formation of different types of carbonate. Whereas CO and C02 desorption peaks were observed for reduced LaMnO,, no CO peaks were found for reduced LaCoO3. This may be associated to the higher catalytic activity of this latter compound for CO oxidation. Adsorption of SO2 on La,,,sSro.SMn03has been studied by Minming et al. (169), who took advantage of the poisoning effect of this molecule on the oxidation activity of the perovskite, in order to determine active center concentration on the surface. The adsorption at monolayer coverage suggested that the (1, 1, 0) face is the most frequently exposed. Infrared spectroscopic data of SO2 adsorbed on Lao.aSro.4Col -xM,O, were reported by Wan et al. (270).These are discussed in Section VII,J. Crespin and Hall (64)determined hydroxyl concentrations between 13.8 and 0.3 x lOI4 OH cm-2 on the surface of BaTi03, LaCo03, and SrTiO, perovskites previously equilibrated with water vapor at room temperature and then heated from 25 to 600°C. These concentrations were similar to those found by Fierro and Tejuca (106a) after adsorbing water vapor on LaCrO3 at 150°C and pumping from 125 to 525°C. The infrared spectra yielded by this latter adsorbent-adsorbate system showed bands at 3680 and 3550 cm-I. The adsorption and dissociation of water was assumed to take place on pairs of surface acid-base centers, anion va-

277


cancy-02-, yielding an acidic OH on an anion vacancy placed between coordinatively unsaturated La3+and a reduced transition-metal ion (band at lower wavenumber) and a basic OH on a terminal oxide ion. Infrared spectra of pyridine adsorbed at 25°C on LaM03 (M = Cr, Mn, Fe, Co) oxides previously outgassed at 500°C are given in Fig. 14. For LaMn03 and LaFe03 (Fig. 14a) (50), they include Lewis bands at 1595, 1490, and 1440 cm-I whose intensity remained constant on H2O adsorption. For LaMn03 a shoulder at 1540 cm-' was also observed. A similar spectrum was found for pyridine adsorbed on LaCr03 (165). Pyridine adsorbed on the reduced (500"C, H2) and outgassed (as above) LaM03 oxides yielded a spectrum similar to that of the unreduced samples and also a weak Bronsted band at 1540-1545 cm-I whose intensity increased on H20 adsorption. The appearance of Bronsted acidity (Fig. 14b, Q-Q) may arise from the heterolytic dissociative adsorption of H2 on co-

I

1600

I

I

1400

3 /cm-'

I

1

I

1600

I

I

1400

3/

C d

FIG.14. (a) Infrared spectra after pyridine adsorption on LaMnO] (al) and LaFeOl (a2) and after subsequent H20adsorption (bl, b2). (b) Spectra after pyridine adsorption on reduced (HI, 500°C) LaMOl [M = Co (al), Fe (a2),Mn (a3),Cr (a,)] and after subsequent H20 adsorption (bl , b2, b3, b,). (Reprinted by permission from Ref. 50.)

278

LUIS G. TEJUCA el al.

ordinatively unsaturated M3+ and 02-ions, as occurs on some simple oxides: -0-M-0

+ H2

-

H

I

0-M-OH

The increase in Bronsted acidity (Fig. 14b, 61-b4) may be accounted for by the increase in anion vacancy concentration produced in the reduction process and the dissociative adsorption of the H20 molecule as indicated above. These results show the rather low acidic character of these perovskites. AND ALCOHOLS D. HYDROCARBONS

The adsorption of C2H6 and C2H4 at -5°C on reduced LaNiO3 was studied by Crespin ef al. (133). Whereas C2H6 adsorption followed Henry’s law, C2H4 adsorption obeyed the Langmuir isotherm for dissociative adsorption. The adsorption plateau for C2H4 was found to be close to the total number of Nit+ ions on the surface. Min and Peiyan (171) identified by IR spectroscopy the species -C2H4(ads) and -C2Hs(ads) after adsorption of C2H6 on Ndo.7Sro.3MnO3 at room temperature. Ichimura ef al. (135) found TPD peaks of CzH6, C2H4, C2H2, and CH4 after adsorption of C2H6 and C2H4 on Lac003 at 27°C. However, the TPD spectra obtained after adsorbing C2H6 and C2H4 on LaAI03 and LaFe03 contained only peaks of the undissociated molecules. These results, together with XPS data, emphasize the contribution of cobalt ions in the perovskite structure for the C-C bond scission. Tascon and Tejuca (172) and KremeniC ef al. (136) reported data on propene and isobutene adsorption on LaCrO, and LaMn03. These authors found physisorption on LaMnO3 below 0°C.A wide IR band centered at 2330 cm-I of COZwas observed on adsorption of these hydrocarbons above 25°C; its intensity was higher on the oxide with the higher oxidation activity (LaMn03) (136). In the system propene-LaMn03 bands of carbonate species were also found. Isobutene adsorption on LaM03 oxides at 25°C showed a maximum for the oxide exhibiting the lowest catalytic activity for CO and hydrocarbon oxidation (LaFe03) (136). Results of successive adsorption of 02-isobutene on LaM03oxides and C02-isobutene on LaCo03 pointed to noncompetitive adsorption (136, 145).

Madhok (173) found that the sorptive capacity of LaCo03 for alcohol vapors decreased in the order methanol > ethanol > n-propanol > nbutanol. No data were reported for aromatic alcohols. However, this


279

oxide showed catalytic activity for vapor-phase oxidation of benzyl alcohol to benzaldehyde. VII. Perovskites in Catalysis

A. CO OXIDATION The oxidation of carbon monoxide over perovskite-type oxides has been widely studied. Voorhoeve et al. (30) brought forward new ideas in explaining the role of defect chemistry of perovskites such as cobaltites, manganites, chromites, and ruthenates. They suggested that two different oxidation processes should be distinguished; (1) the catalyst participates in the reaction as a reagent, being partially consumed and regenerated in a continuous cycle, and (2) the catalyst provides the atomic orbitals of the proper symmetry and energy to activate the reactant molecules. These two alternatives were termed reagent or intrafacial catalysis and template or suprafacial catalysis, respectively. The oxidation of CO in the middle temperature range ( 1OO-3OO"C) has been suggested as a suprafacial catalytic process where it is expected to observe important effects of the ferroelectric and magnetic order of surface spins and of semiconductivity on the catalytic reaction. Some of these will be reviewed below. The first study on the CO oxidation using perovskites was conducted by Parravano (2,2) (see Section VI1,I). Later on, Kawai et al. (168)found a discontinuity in the activation energy for the oxidation of CO by 0 2 or N 2 0 near the ferroelectric Curie temperature (T, = 120°C) of BaTiO3. The surface properties of these ferroelectrics are particularly different from those of the bulk. Alkaline-earth niobates and BaTi03 display charged surface layers and polarization reversal at the surface that are influenced by the charge layer. The reaction rate is very slow in the temperature range 100-200°C and under steady-state conditions was found to be limited by the rate of desorption of C 0 2 . The activation energy is 10 kcal/mol above T, and only 1.8 kcal/mol at lower temperatures. This suggested a large dipole-dipole contribution to the bond of COZas carbonate above T,. The low-rate and low-activation energy observed suggest that the oxidation reaction proceeds via surface defects. LnM03 perovskites in which the lanthanide (Ln) ions are essentially inactive in catalysis and the active transition-metal (M)ions are placed at relatively large distances (ca. 0.4 nm) from each other are excellent catalytic models for study of the interaction of CO and 0 2 on single surface sites. It must be stressed, however, that idealized correlations between catalytic activity that is confined to the surface, and a single collective

280

LUIS G . TEJUCA

et al.

parameter, (conductivity, ferromagnetism, etc.) should not be taken as conclusive. Voorhoeve et al. (14,30, 174), Shimizu (175), and Tasc6n and Tejuca (I76) have shown a suggestive correlation between the activity data, using mixtures of CO and 0 2 at atmospheric pressure, and the electronic configuration of the transition-metal ion. In Fig. 15 the catalytic activity is plotted either as the reciprocal of the reaction temperature at which the activity is lo+' mol of CO converted per square meter of catalyst per second (Fig. 15a) or as the rate of mole CO converted per unit area and unit time (Fig. 15b) versus the occupancy of the d levels for the transition M3+ion. It is known that the octahedral environment of the B ions splits the d-orbitals into two levels; the lower (tZe) contains orbitals that are repulsed less by negative point charges than are the orbitals in the higher (e,) level. It is observed that the maximum activity is attained in both --5 I

c

In n I

E

-48

P

--I -B .-c

- -8 - -9

FIG.15. Activities of first-row transition-metal oxide perovskites for CO oxidation in a 2 : I mixture of CO and O2at atmospheric pressure (a) or in a 1 : I mixture of CO and O2 at 227°C at atmospheric pressure (b). The activities of vanadates (B), chromates (O), manganates (A), ferrates (O), cobaltates (01,and nickelates (A)are plotted at the appropriate dorbital occupation corresponding to the average valence of the transition-metal ion. (Redrawn by permission from Refs. 14 and 176.)


28 1

cases (Fig. 15a,b) for an occupation of the eg levels of less than one electron being the t2plevels half-filled or totally filled. This is in agreement with the data of Sazonov et al. (177)for the homomolecular 0 2 exchange on perovskites (see Section VIIJ and also with those of Boreskov (178) for the methane and HZoxidation and homomolecular 0 2 exchange reactions on several spinels. Voorhoeve et al. (14,301 have also stressed that the catalytic activity of perovskites is influenced by their stoichiometry. A simple way of varying the oxidation state of the ion at the position B is by substitution of the A ion by a different ion with an oxidation state other than 3. This method has been used by several authors (9,62,88, 96, 179-181) to understand the role of the 3d-orbital occupancy in the LaM03 series on the catalytic oxidation of CO. For M = Co the appearance of Cot+ions by introduction of Ce4+in position A enhances the rate of oxidation of CO, whereas the presence of Co4+ ions by substitution with Sr2+reduces the rate. The explanation for this behavior has been given by assuming that CO is bonded to the transition-metal ion as a carbonyl, as occurs on metals (182), with donation of the carbon lone pair into the empty 3 4 2 orbital of M to form a a-bond accompanied by back-donation of the tzgelectrons of the metal to the antibonding .rr-orbital of CO. It should be noted that the dz2orbital is the lowest e, level for the M3+ions at the surface, and in order to have a partially empty dzzlevel, the occupation of all the e, levels must be below unity. It must be also emphasized that for either nonsubstituted LnCoO, (Ln = La, Sm, Nd, Dy) or substituted Ndl-,Ba,CoO3 perovskites, two Co ions in oxidation states differing by one, bonded by an oxygen ion, can exchange an electron through the 02p orbital (183-185): C03+-O-C04+

C04+-O-C03+

(5)

This exchange, known as Zener double exchange, accounts for an average oxidation state of Co ions (183) at least on the time scale of the reaction of CO oxidation and reaches a maximum for an angle Co-0-Co of rr-radians (184),i.e., for the cubic structures. It has been suggested that such an average oxidation state of Co facilitates the adsorption-desorption of the reactants, whereas individual Co3+ions in their discrete spin and oxidation states lead to an adsorption process that is too strong or too weak for the catalytic reaction to take place (183). The kinetics of CO oxidation on Lac003 has been studied by Tasc6n et al. (186). In the temperature range 120-155°C these authors found a strong C02 inhibition and proposed a kinetic equation by assuming the controlling step to be the surface reaction between adsorbed CO and

282

et a/.

LUIS G . THUCA

dissociatively adsorbed oxygen:

and COz, where PCO, P O , , and Pco2 are the partial pressures of CO, 02, respectively; bo, is the adsorption coefficient of oxygen and k' = k K (k is the specific rate constant, and K is the equilibrium adsorption constant). Experimental data according to the linear form of Eq. (61, ( Pq/r)1/2versus Pbt and l / r versus Pco2 for variable partial pressure of 0 2 and CO2, are given in Fig. 16. From Arrhenius plots, these authors calculated an activation energy of I5 kcal/mol similar to that found by Rao and Chakrabarty (150) and lower than that reported by Yao (282). From these results and from IR spectroscopic data (see Section VI) the following scheme for the mechanism of CO oxidation on LaCoO3 was proposed:

coladsl

-

+ 2oladsl

Co31adrl

CO2ldr) + Olads)

(9) (10)

COU,,

(11)

C03(sdn)

C021adr1-

(slow)

where equation (9) should be rate-controlling. Oxygen is adsorbed as molecular 0 2 species on Co ions of low oxidation state, which subseadsorbed on the same quently dissociates yielding atomic oxygen (0-) 6-

b

ul

\ N

-2

N

3.0-

v

a 125 11 13511

P

\

N O ' 04

'

'

0.8

'

'

12

'

'

1.6

I

0

2

6

4

1/2

P i /mmHg

pcrJz/""

Hg

FIG. 16. Linear plots of experimental data for oxidation of CO according to Eq. (6) for variable pressure of O2 (a) and C02 (b) [Pco = 6.75 x 102 Pa (a); POI = 3.37 x 102 Pa (b)]. (Reprinted by permission from Ref. 186.)


283

center. CO adsorbs on surface 02-ions, producing a labile species that interacts with adsorbed atomic oxygen, producing a carbonate that then decomposes, yielding adsorbed C02 and oxygen. Previous work strongly suggested that the activation centers for oxygen and CO are cobalt ions and oxide ions, respectively, whereas the La3+ions are catalytically inactive (138, 160). A slightly different mechanism has been proposed by Gunasekaran et al. (187) for the oxidation of CO on La2M04(M = Ni, Cu) oxides. On the basis of electrical conductivity changes in different atmospheres, these authors found that the rate of 0 2 adsorption was independent of the state of the surface (freshly evacuated or with preadsorbed CO); however, the rate of CO adsorption underwent a remarkable decrease by the presence of preadsorbed 0 2 . Because the rates of CO and 0 2 adsorption were faster than the rate of the oxidation reaction, they concluded that the surface reaction is the rate-determining step. Viswanathan and George (149) studied the oxidation of CO on LnCo03 (Ln = La, Nd, Sm, Gd) and found important gradient changes in the Arrhenius plots in the temperature range 190-220°C, where the adsorption isobars of O2and CO showed a maximum or a minimum. Tascon and Tejuca (138,160) and Chakrabarty et al. (148), using IR and XP spectroscopies, have explained this behavior by assuming that the adsorption of oxygen and CO occurs as described above. On the basis of conductivity measurements (149), a mechanism similar to that depicted by Eqs. (7)(1 1) was proposed, although other possible reactions involving formation of carbonates, namely:

-

CObds)

+ 02(ads)

Cokis)

+ 20%~)

COiiadsl

(12) (13)

C0,ads)

+ 2OL)

CO:i,dsi + 2e-

(14)

coiiadsl

have been advocated. In this reaction sequence any of the reaction steps of Eqs. (12)-( 14) between the adsorbed species may be rate-controlling. B. OXIDATION OF HYDROCARBONS AND OXYGENATED COMPOUNDS Libby ( 7 ) and Pedersen and Libby (8) were the first investigators to suggest the potential application of perovskites as oxidation catalysts. In the research work that ensued, the particular behavior of these materials was explained in terms of the relative ease with which oxygen species can be released from the catalyst surface. To confirm these ideas, work concerning the oxidation of paraffins (40,131,140-142,171,188-191), olefins (136,140,181,192),aromatics (193,194),and oxygenates (140,193,195199) have been carried out.

284

LUIS G . T U U C A ('1

d.

The oxidation of light paraffins, as methane (40, 140, 190), propane (141, 142, 189), and n-butane (140) has been frequently taken as a test reaction for perovskite oxides. Arai et al. (40) studied the catalytic combustion of methane over LaM03 (M = Cr, Mn, Fe, Co, Ni, Cu) and partially substituted Lal-,A,MO3 (M = Mn, Fe, Co; A = Sr, Ca, Ba, Ce; 0 e x =e0.4) perovskites and compared these oxides with Pt/A1203 catalysts. The Sr-substituted La-Mn, La-Fe, and La-Co oxide systems showed high oxidation activities; Lao.6Sro.4Mn03 was almost as active as Pt/A1203for a conversion level below 80%. These authors examined the dependence of the oxidation rate on the partial pressures of CH4 and 0 2 and proposed an empirical rate equation to describe their experimental data: r = k * P&,

*

P&

(15)

where r is the oxidation rate of methane and m and n are the reaction orders of methane and oxygen, respectively. The m and n values of two substituted La,,&o.2MO3 (M = Co, Mn) perovskites and a reference of 1 wt% Pt/A1203for methane conversions below 20% in the temperature range 450-650°C are collected in Table IV. The results clearly indicate a first-order kinetics with respect to the partial pressure of methane. Although Eq. (15) describes quantitatively the kinetic data, it is desirable to deduce formal rate equations on the basis of the precise knowledge of the reactivity of oxygen. Two kinds of oxygen species, adsorbed and lattice oxygen, with different bond strengths are present at the surface of perovskite-type oxides (see Section V,D). The adsorbed oxygen is believed to become active and to react with hydrocarbons at lower temperaTABLE IV Dependence of Methane Oxidation Rate ru on the Partial Pressure of Methane and Oxygenb

Catalyst L~(~.RS~O,ZCOOI Lao,&~MnOl 1 Wt% Pt/A120]

Temperature ("C) 450 550 650 450 550 650 450 650

Reaction order m

n

I .o

0.6 0.4 0.3 0.5 0.3

1.o 0.9 1.1 0.9 I .o 0.9 1.1

Oxidation rate, r = kP&, P & . Reprinted by permission from Arai et al. (40).

0.2 -0.5 -0.3

285


tures than the lattice oxygen. At low temperatures, the oxidation rate depends largely on both CH4 and 0 2 partial pressures, and it may be expressed by the Rideal-Eley (R-E) mechanism, which assumes reaction between adsorbed O2 and gaseous CH4:

However, this expression does not apply in a large temperature range since the experimental n values (Table IV) decrease with increasing temperature. To explain this behavior, two active oxygen species should be taken into account. At low temperatures, the methane combustion is governed mainly by molecularly adsorbed 0 2 ; the contribution of lattice oxygen is negligible. At higher temperatures the coverage of molecular O2 decreases and the lattice oxygen becomes reactive. The methane combustion at high temperatures over metallic oxides can be described by a redox mechanism. The rate of oxygen incorporation to the lattice is much faster than the oxygen consumption; hence the rate is zero-order with respect to the oxygen partial pressure: (17)

r = kl * PCH4

The decrease of n with increasing temperature can be attributed to the change of the oxygen species that participate in the reaction, specifically, from molecularly adsorbed to lattice oxygen. The extent of reaction of molecular oxygen largely depends on its coverage; that is, it decreases with decreasing oxygen partial pressure and with increasing temperature. In an intermediate temperature range the contribution of both oxygen species to the rate equation can be equally important. The complex temperature dependence of the CH4combustion on a Lal-,Sr,Mn03 ( x = 0 . 2 , 0.4) (Fig. 17a) points to the participation of both oxygen species. On the basis of these considerations, the overall rate of CH4 combustion in the whole temperature range can be described by a combination of Eqs. (16) and (17):

At high temperatures Eq. (18) can be simplified because the coverage of molecularly adsorbed oxygen is small [I >> (KO, PO,)^/^]:

-

r = r,

+ tj = k,

*

PcH, * (KO,* PO,)^'^

+ kl

*

PCH,

(19)

The contributions of adsorbed and lattice oxygen to the overall combustion rate can be estimated by means of plots of the experimental rate versus (Fig. 17b). The intercepts of the straight lines represent the contribution of lattice oxygen, which, as expected, increases with in-

LUIS G . TWUCA

et al.

a

4c

!OO

400

600

T("C1

800

t 0d v a t m

FIG.17. (a) Rate of methane combustion on La&ro.rMn03. Circles, experimental data; dashed line, rate expected for oxidation by lattice oxygen; pointed line, rate expected for oxidation by adsorbed molecular oxygen. (b) Rate for methane oxidation versus Pod; intercepts represent contributions of lattice oxygen. (Reprinted by permission from Ref. 40.)

creasing reaction temperature. These considerations are consistent with Voorhoeve et al. 's classification of oxidation reactions on perovskites in suprafacial and intrafacial processes that occur at low and high temperatures, respectively (14, 30). Conner et al. (188) studied the oxidation of propane on Bal,85 Bio.~Oo.os(Bi&I v3Te)06(0,vacancy). This defect oxide containing Bi and vacancies on both A and B positions showed 25% propane conversion at 4Oo0C, yielding partial oxidation and cracking products. Other nonsubstituted perovskites containing Bi in only the B site were tested and found to be inactive. Propane oxidation has been also used by Nakamura et al. (95) and Nitadori and Misono (142) to study the effects of Sr2+ and Ce4+ substitution in LaM03 (M = Fe, Co) perovskites and more recently by Nitadori et al. (289) to investigate the effects of Sr2+,Ce4+,and Hf4+ substitution in LaMn03. In general, the activity of Lal-,Sr,MnO3 varied in a way parallel to the rate of isotopic equilibration of oxygen, the reducibility of the catalyst surface, and the amount of reversibly adsorbed oxygen (measured by volumetric adsorption). Because these properties


287

pertain exclusively to the surface oxygen it is inferred that the activity is directly related to the nonstoichiometric character of the surface. The absence of any correlation between activity data and the oxygen peak that appears at high temperatures in the TPD spectra, which is known to arise from bulk lattice oxygen (Fig. 13), supports that assumption. Yao (181) investigated the combustion of C2H4 over LaCo03, BaCo03, LaxSrl-,MnO3, and L*.7Pbo.3MnO3 and found lower reaction rates for the cobaltites and manganites than for Co304. On the other hand, Jhaveri et al. (192) studied the oxidation of propylene on several cobaltites, CuC0204, and Ce02-Co304and fitted the kinetic data to the empirical rate equation r = k Pfic. At 400°C the catalytic activity increased in the order cobaltite < substituted cobaltites < Ce02-Co304< CuC0204. KremeniC et al. (136) studied the catalytic activity for the oxidation of C3 and C4 hydrocarbons on a series of LaM03 (M = Cr, Mn, Fe, Co, Ni) oxides and found remarkable differences among these compounds. The catalytic activity for total oxidation of propylene and isobutene at 300°C and for molar ratios HC : O2 = 1 : 4 and 2 : 1 are given in Fig. 18. The rates for

-

FIG. 18. Catalytic activity profiles of LaMO, oxides at 300°C in propene (open symbols) and isobutene (filled symbols) oxidation. Molar ratio HC:02= 0.25: 1 (circles) or 2 : 1 (triangles). (Reprinted by permission from Ref. 136.)

288

LUIS G. TEJUCA E f

a/.

isobutene were found to be higher than those for propylene. In both cases two pronounced maxima that are coincident with those found for oxygen adsorption (LaMn03 and LaCo03, Fig. 12) were observed. The partial oxidation of isobutene (propylene yielded only carbon oxides) followed the reverse sequence to that found for total oxidation. A similar pattern was found by Seiyama ef al. (140) for the propylene oxidation on LaM03oxides. Twin peak patterns were also found by Dowden et al. (200) and Dowden and Wells (201) in reactions involving hydrogen catalyzed by first-row transition-metal oxides and by Boreskov (178) in oxidation reactions catalyzed by a series of spinel-type oxides. This general behavior has usually been explained in terms of the change in crystal field stabilization energy (CFSE) as a result of the change in coordination of the M3+ ion, which shows two maxima when passing from dn to d9 configurations. Although the position of the experimental peaks either in the catalytic profiles or in the oxygen adsorption pattern does not exactly match with the predictions of the CFSE theory because other effects such as surface defects may be operative, these results provide further support to the ideas of Dowden and Wells (201) on the relationship between the local symmetry of M3+ions and adsorption and catalysis. On the other hand, the close parallelism between oxygen adsorption and catalytic activity for total oxidation of propene and isobutene indicates that these reactions occur through a suprafacial catalysis mechanism in which adsorbed oxygen is the dominant species participating in the reaction. A somewhat different application of perovskites is the oxidative dehydrogenation of olefins. Kehl et al. (202) reported that Lao,&ro.,jsFen,5503 can be used for butene dehydrogenation to butadiene and also for the conversion of other monoolefins containing at least four C atoms, such as the production of isoprene from isoamylenes. The oxidation of benzene on SrV03(193) in the temperature range 280400°C yielded only CO2 and small amounts of maleic anhydride at the higher temperature studied. The catalyst became completely oxidized to Sr2Vz07after one run. Because of the instability of V4+ in the perovskite structure, the highly exothermic deep oxidation caused the eventual thermal decomposition of the catalyst. The vapor-phase oxidation of toluene to benzaldehyde on LaCo03in the temperature range 350-600°C has been studied by Madhok (194).The kinetic analysis showed the oxidation to be first-order with respect to the hydrocarbon. The activity increased with the temperature below 500°C and decreased at higher temperatures. The activity increase was ascribed to removal of surface contaminants and to the generation of structural defects such as anion vacancies. The activity

289

PEROVSKITE-TYPEOXIDES

decrease above 500°C was assumed to be due to a decrease in surface disorder. Perovskites have also been studied as model compounds in the oxidation of oxygenated compounds. Arakawa et al. (197, 198) reported activity profiles for the oxidation of methanol on LnFe03 (Ln = La-Gd) and LnCo03(Ln = La-Eu) oxides. In the former series the activity decreased in the order Gd > Eu > Sm > Nd > Pr > La. From the analysis of the binding energy of the Fe2p31zphotoemission these authors found that the activity increased with decreasing covalence of the Fe-0 bond. Shimizu (199)found the oxides Lal-,Sr,Fe03 and LaMO3 (M = Mn, Fe, Co, Ni) to be active for oxidation of ethanol to acetaldehyde in the temperature range 2OO-45O"C. However, complete oxidation proceeded above 350°C and at high O2 partial pressure. The catalytic activity of Lal-,Sr,Fe03 increased with x up to x = 0.2 and decreased for higher x values. The activity of LaM03 oxides decreased in the sequence Co > Mn > Ni > Fe. C. NO REDUCTION Reduction of NO with CO or H2 was found to be an interesting example of intrafacial catalytic process (30). If this reaction is conducted over a transition-metal oxide, the reaction rate appears to be related primarily to the thermodynamic stability of the oxygen vacancies adjacent to a transition metal ion. Associative as well as dissociative adsorption of NO have been reported on perovskite oxides (14, 22, 80, 174) (see also Section VI,B); the adsorption on the reduced oxides is stronger than in the oxidic compounds. Dissociative adsorption takes place at moderate temperatures as in NO reduction over Lao.~sBao.lsCo03 at 100°C with the subsequent formation of N2 and N2O (73). To account for the observed products in the NO reduction over several perovskite oxides, the reaction scheme, which includes the participation of molecular and dissociative chemisorption of NO, has been proposed (22, 30, 80): M-0-M 2e-

+ NO + 0 + NO

-

M-0-M

+ 1 Nz

02(lattice)

+ Nad,

(20)

( 2 1a)

where 0 is an oxygen vacancy. The relation of the reduction of the catalyst with the conversion of NO was first demonstrated by Bauerle et

290

L U I S G . TWUCA

et al.

al. (203) and subsequently by Sorenson et al. (204), who found a close parallelism between the onset of NO reduction and the onset of bulk reduction of LaCo03. It was determined that the activity for NO reduction was associated with the formation of an oxygen-deficient structure having the general formula LaCo1-,03-,, where 0 C x S 0.08 and 0 < y S 0.5. Chemisorption of NO took place readily in this compound because of the presence of anion vacancies and low spin configuration of the Co3+ ions. From these considerations it is expected that perovskites that can easily release oxygen will be active for NO decomposition. Voorhoeve et al. (30, 80, 174, 179) demonstrated that manganites satisfy this requirement fairly well. In these perovskites the bond energy of the surface oxygen may be varied by introduction of La3+vacancies that change the electronic configuration of the Moo6octahedra, viz., lowering the binding energy of oxygen in Mn-0 and increasing the Mn4+content. Thus, these authors (14,80)compared the rate of reduction of NO over stoichiometric LaMn03 and L~.90’0,1Mn03 perovskites. They found the latter perovskite to be substantially more active than the former. In these compounds the binding energy of oxygen may be affected by the average oxidation state of Mn. To know whether the large improvement in the rate of reduction of NO was due to the Mn4+ions or to the La vacancies present, perovskites (AyAt1-,)Mn:?,Mnjt03 (where A and A’ represent one or more of the following entities: La, O’, Bi, Pb, Sr, Na, K, Rb; O’, cation vacancy) containing approximately constant proportions of Mn3+ and Mn4+ions but different A-0 binding energies, were tested for this reaction. In the (1 ,O,O) plane the lattice oxygen is coordinated by two Mn nearest neighbors at a distance of approximately 0.2 nm and by two A positions: A’

/ Mn-0-Mn / A

The heat of formation of an oxygen vacancy should be related to the sum of the A-0 and A‘-0 binding energies, which are given by the expression

A(A-0) =

AHf - m AHs - (n/2) DO 12 m

where AHf, AH,, and DOare the enthalpy of formation of 1 mol of A,O,, the enthalpy of sublimation of A, and the dissociation energy of oxygen, respectively. For A,A;-,Mni,?,Mn$+O3 perovskites, Eq. (22) gave values for the binding energy of oxygen decreasing in the order Bi, K > La, 0’> La, Rb = La, K = La, Na > La, Pb > La, Sr. This trend is closely related


29 I

to the rate of NO reduction (N2 + N20 yield). These data clearly reflect the importance of the binding energy of surface oxygen, which determines the number of active centers (anion vacancies). The cation vacancies provide weakly bound oxygen to the surface and, therefore, favor both the formation of anion vacancies and the NO reduction. D. POLLUTION CONTROL The removal of CO, unburned hydrocarbons (HC), and NO from automotive exhaust requires catalytic devices in which these pollutants are eliminated. The catalysts that are being used for this purpose are supported Pt, Pd, and Rh (205). These present a number of desirable characteristics such as high activity, stability, and resistance to sulfur poisoning. However, their use may be constrained by the limited availability of noble metals. The metal oxides constitute an alternative that should be further explored. Among this latter category, perovskite-type oxides of Co, Mn, and Ru have been investigated as catalysts for such devices. Two designs of converters have been prominent: ( I ) in the dual bed, NO is reduced in the first catalyst bed by excess of CO and HC, and then the unreacted CO and HC are oxidized over a second catalyst bed; and (2) the three-way design uses a single catalyst bed to convert NO, CO, and HC simultaneously. Experimental tests were carried out with engine exhaust gases as well as with synthetic laboratory gas mixtures. The reduction of NO with mixtures of CO and Hz (dual-bed system) has been studied over several compounds. The La0.900.1MnO3-deficient perovskite yielded 100% selectivity to Nz;however, the catalyst became progressively deactivated by reduction with the CO-HZ mixture. The substituted manganite (Lao.s~Bia.onKo.o,)MnO3 showed a high activity but a poor stability because of the presence of low-melting-point components (256). Cobaltites showed a high activity and selectivity; their activity was strongly dependent on the oxidation state of the catalyst (204, 206). Ruthenium supported on alumina proved to be unsuitable for exhaust gas purification since the noble metal is oxidized to the highly volatile R u 0 3 oxide and thus is eventually lost. To overcome this problem Gandhi et ul. (207) and Dalla Betta et al. (208) proposed the stabilization of Ru3+ or Ru4+ ions in the matrix of a perovskite structure. Additionally, the dispersion of Ru in the B sites renders the catalyst more resistant to losses in exposed atoms caused by sintering. Thus the perovskites SrRu03 and LaRuO, and the substituted (Lao.8K0.2)(Mno.94Ru0,~)0~ and (Lql.xSr0.2) ( C O ~ . ~ R Uoxides ~.~)O were ~ studied by using laboratory gas mixtures (NOCO-H2) or engine exhaust gases (256,206,208,209). The most relevant

292

LUIS G . TUUCA

et a / .

feature of the above preparations was a high activity and a high selectivity to N2. Especially the latter perovskite supported on a Torvex alumina honeycomb has been reported to undergo only minimal changes in the level of NO conversion for extended periods of time (206). This is illustrated in Fig. 19a, where conversions of NO and CO are given as a function of 0 2 or CO excess. For the stoichiometric composition, the NO conversion appears to be approximately 90%. In general, the selectivity to N2 of Ru perovskites was found to be comparable to that of Ru metal (209-2Zf). Sorenson et al. (204) proposed LaCo03 as a potential catalyst for the simultaneous reduction of NO and the oxidation of CO and unburned hydrocarbons (three-way design). These authors found this perovskite to be very effective with NO conversions of approximately 95% and CO and HC conversions of up to 85%. The perovskite (La,-,.&o.4)(Coo.NPto.o3R~o.o3)O3 has been reported to be particularly suited for this purpose. For the selective reduction of NO to N2 at the reducing side and an efficient oxidation of CO and HC at the oxidizing side of the stoichio-

I

I

L-1

O

Excess of O2(O/0)

Excess of CO(%)

;

I

I

I

0

1

2

Excess of (&(YO) Excess of CO (YO)

FIG.19. Performance of noble-metal-substituted perovskites in the treatment of exhaust from single-cylinder engines. (a) (La,,8Sro.2)( C O ~ , ~ RO3~ supported ~.~) on a Torvex alumina honeycomb after 1000-h exposure to exhaust from leaded fuel in the first bed of a dual-bed catalytic system. (b) ( L h 6Sro.4)( C O ~ . ~ ~ P O3 ~ . on ~ ~a R Torvex ~ ~ , alumina ~ ~ ) support after 800-h exposure to exhaust from unleaded fuel in a three-way catalytic system. (Redrawn by permission from Ref. 206.)


293

metric composition, this multicomponent oxide appears to provide an operating window in which the removal of all three pollutants was about 80% or higher (Fig. 19b). Other perovskites have been also used for CO and HC oxidation in both the dual-bed or the three-way systems. Cobaltites were investigated with laboratory gas mixtures (174, 181, 192) and with the exhaust gases of a single cylinder engine (204), whereas manganites were studied only with synthetic gas mixtures (36, 179, 181). The conversion of CO on LaCo03(204) was higher than those of HC in engine exhaust gases with excess air and substantially higher than that of propylene on La,,,Pbo.3Mn03 in synthetic automobile exhaust mixtures (212). The activity of the catalyst was found to be highly dependent on the oxidation state of the surface and, particularly at high temperatures, on the ratio of the partial pressures of oxidants and HC. At low temperatures and high partial pressures of CO and 02,LnCoO3 (Ln = La, Pr) and La,,7Pbo.3Mn03 and its homologues with Ba, Sr, Pr, and Nd at the A site showed an activity for CO oxidation similar to that of Pt. The activity of cobaltites, and especially of manganites, decreased with decreasing partial pressure of CO. On the other hand, platinum substitution was shown to have an important effect in the performance of perovskites. Thus, Lauder (206) reported that L~.sSro.~Coo.9pto.103 supported on monolithic alumina, besides displaying a high activity, showed a high resistance to Pb poisoning in single engine and automotive tests. The durability of these materials enhanced when used with near stoichiometric mixtures. The problem of sulfur poisoning has not been considered here. In Section VI1,J some aspects related with the interactions of sulfur dioxide with perovskite oxides as the adsorption of SO2 on B and oxygen sites, the effect of SO2 on Pt-doped oxides and in CO and hydrocarbon oxidation will be reviewed. In the light of these studies, some considerations on the present state of research on this topic are made.

E. HYDROGENATION AND HYDROGENOLYSIS OF HYDROCARBONS Libby (7) and Pedersen and Libby (8)were the first authors to report on the activity of LnCo03 perovskites (Ln = La, Nd, Dy) for cis-Zbutene hydrogenation; hydrogenolysis became important above 200°C. Later on, Lombard0 et al. (77,158,213-217) studied the hydrogenation and hydrogenolysis of several hydrocarbons on unsupported or LazO3-supported LaCoO,. These authors found a sharp activity peak in ethene hydrogenation at -20°C for a reduction of the catalyst of 1 . 5 - per molecule. Ad-

294

LUIS G . THUCA

et al.

sorbed CO presents a peak that coincides with that of hydrogenation (213-215). These results are illustrated in Fig. 20, where activity and

adsorption data are plotted as a function of the reduction degree of the oxide. Cyclopropane conversion (hydrogenation plus hydrogenolysis) at 200°C also reaches a maximum for Lac003 reduced between le- and 1.5e- per molecule (158). The selectivity for hydrogenolysis is constant up to a reduction degree of 1.5e- per molecule and then drops sharply. These results, together with data on deactivation, self-hydrogenation of ethene in the presence of the reduced perovskite, and tracer experiments with deuterium (the ethene-deuterium reaction yields multiplyexchanged ethene and ethane) (214, 215), seem to indicate that metallic cobalt is responsible for both hydrogenation and hydrogenolysis activity. The increase in catalytic activity with the oxide reduction and the observed product distribution in the hydrogenation of 1,3-butadiene on Lac003 and LaCo03 deposited on La203 support that assumption (77, 217). However, other oxidation states of cobalt, such as Co2+,may also participate in these catalytic reactions. Thus, Crespin et al. (133) showed that LaNi03, a perovskite very similar to LaCo03, is active for ethene hydrogenation after reduction to le- (La2Ni205)and 3e- (Nio/La203)per molecule. On the other hand, the adsorption plateau of C2H4 on La2NizOs appears to be close to the total number of Ni2+cations on the surface. This finding also indicates the involvement of NiZ+ions in the adsorption and presumably in the hydrogenation process. Ichimura et al. (63, 135, 218) on the basis of kinetic, tracer, and XPS studies, concluded that Co3+plays an important role in the rupture of the C-C bond whereas La3+and 02-ions contribute mainly to the dissociative adsorption of hydrogen in the hydrogenation and hydrogenolysis of C2-C5alkenes and alkanes. However, it must be noted that in the experimental conditions used (25-350°C and excess H2) some reduction of Co3+ to Co2+and Coois to be expected (64, 129). These reduced cobalt species may, therefore, be responsible for the activity. However, the possibility of the lanthanum and oxygen ion pairs acting as hydrogen atoms suppliers in hydrogenation and hydrogenolysis should not be discarded. This is supported by the fact that La203 catalyzes the H2-D2 equilibration as well as the ethene hydrogenation, thus showing the ability of this oxide for hydrogen dissociation (63). Other perovskite oxides containing Re(VI1) and cation vacancies in position B, such as BazBl,30~3Re06(B = Y,Sm), were found to exhibit hydrogenation activity of ethyl acetate to ethanol (2Z9). On the basis of the increased conversion to ethanol and the change of the catalyst color to black with time of operation, it is assumed that the activity is due to reduced Re in a highly disperse state.

z.

a

w

0

c c S T P of CO/g

d

i

O r + - -

m mol of C2H,/m2 min

lo3

296

LUIS G. TWUCA et al.

F. CO AND C02 HYDROGENATION Recently, special interest has arisen for the production of oxygenated chemicals from syngas (CO + H2) and C02 + H2. Supported rhodium appears to be one of the most promising catalysts for this purpose (220); however, it leads to very different product distributions. To understand the origin of these differences among nominally similar catalysts, model perovskite compounds have been used. Somorjai and co-workers (163, 221) have studied the activity of clean and oxidized Rh foils and of LaRhO, at different temperatures, under H2 :CO mixtures, and thus at different levels of reduction. Rh2O3 produces large concentrations of CZ and C3 products. Furthermore, addition of C2H4 to the CO + HZ feed yielded propionaldehyde, showing the carbonylation ability of Rh2O3. Under similar conditions over Rho, C2H4 was quantitatively hydrogenated to C2H6.Thus, higher oxidation states of rhodium seem to be necessary to produce the oxygenated organic molecules. The product distribution over LaRh03is highly temperature dependent and can be accounted for by a mechanism that involves associative adsorption of CO to form methanol and a dissociative mechanism for other products (163,221). The variation of selectivity with temperature is due to competing processes of hydrogenation and carbonylation and variable concentrations of molecularly and dissociatively adsorbed CO and hydrogen on the surface. Thus, below 225°C methanol production through hydrogenation of chemisorbed CO predominates. Between 225 and 350°C dissociative adsorption of CO occurs, yielding CH, species that can then undergo either CO insertion to produce oxygenates or hydrogenation to methane. At these temperatures CO insertion competes favorably with hydrogenation, producing mainly acetaldehyde and ethanol. Above 350°C CO adsorbs dissociatively; therefore, methanol production is severely depressed whereas methane is the main product. The active catalyst has been reported to contain rhodium as Rhl+and a small fraction as reduced Rho. The product distribution (hydrocarbons plus oxygenates) of CO + H2 reaction over lanthanum rhodate at temperatures of 200-350°C has been analyzed by Gysling et al. (41) and Monnier and Apai (222) by means of Schulz-Flory plots (Fig. 21). The observed monotonic decrease in the rates of hydrocarbon formation indicates an essentially common mechanism for the C-C bond formation in c2-C~hydrocarbons. However, the rate of formation of C2-oxygenates in the interval 300-350°C is greater than that of CH3OH, indicating that these compounds are formed through two different mechanisms. This finding is consistent with the results of Watson and Somorjai (163) and Somorjai and Davis (221).


297

c2

Hydrocarbons

c4

Oxygeno tes

FIG.21. Fischer-Tropsch activity of LaRh03at different temperatures, as Schulz-Flory plots. (Reprinted by permission from Ref. 41.)

Gysling et al. (41) have clearly indicated the need for careful attention to the preparation method of LaRh03 perovskites, which are to be evaluated as catalysts for the CO hydrogenation and the usefulness of XPS for monitoring the surface of such materials. They showed that under typical hydrogen reduction conditions the surface of this oxide is fully reduced to Rho. If the catalyst is contacted with a 2 : 1 H2 :CO mixture at 1 atm and 300°C the Rh3d levels can be deconvoluted into Rho and Rh3+contributions, with no Rhf species detected. On the other hand, the trends observed in the Schulz-Flory plots of Fig. 21 on LaRh03 are essentially equivalent to those observed on Rh/Si02, suggesting the same catalytically active Rh species for both catalysts. It appears, therefore, that the formation of hydrocarbons and oxygenates occurs on Rho centers. This conclusion would be in agreement with gravimetric results, indicating that the lanthanum rhodate reduces in H2 directly to metallic rhodium (107b). Conversion of syngas to oxygenate organic chemicals over perovskitetype compounds has been also studied by Broussard and Wade (223). These authors examined several perovskites with La, Ce, Nd, Sr, and Ba

298

LUIS G. TUUCA

et al.

in position A and transition metals in position B. The selectivities for oxygenates obtained with LaM03 (M = Co, Ni, Fe) oxides were in the range 24-33%. These LaMO3 perovskites, partially substituted with Ru, yielded much lower selectivities (3-14%). This is an expected result since Ru has been used in the Fischer-Tropsch synthesis for the production of higher hydrocarbons (224). Copper-substitutedperovskites such as LaCuo.5Ti0.503and L ~ C U ~ , ~ M Q yielded , . ~ Omethanol ~ selectivities of 3738%. This high selectivity suggests that copper, presumably in the reduced form, is the catalytic agent. To the authors' knowledge the influence of the nonstoichiometric (oxygen-deficient) character of these perovskites (see Section IV) has not been studied in this type of reaction. It would be interesting to investigate how the concentration of oxygen vacancies in these oxides affects the selectivity for oxygenated compounds. The highest selectivities (almost exclusively methanol) were found for noble metal perovskites as BaRh03 and BaPtO, (62 and 54%, respectively) (223). The catalytic activity for C02 hydrogenation over LaCo03and partially substituted Lal-,M,Co03 (M = Sr, Th) perovskites was recently studied by Ulla et al. (225). These authors used a recirculation system with a ratio H2 :C02 = 4 : I at 280°C and a total pressure of 2.13 x 104 Pa. The C02 conversion to oxygenates plus hydrocarbons over LaCo03 is little affected by the extent of reduction of the catalyst; however, the rates of formation of hydrocarbons were highly sensitive to this parameter. Reduced Lao.sTh0.2CoO3displayed the highest initial conversion of C02 and a high methane selectivity. L~.6Sro,4Co03 prereduced at 300°C showed a sharp maximum in activity and a high selectivity to methane; however, at higher reduction temperatures the initial conversion of COz decreased whereas the selectivity to higher hydrocarbons increased. Metallic Coo sites are assumed to be the active sites for CO2 hydrogenation; their specific. activity and selectivity are strongly influenced by the presence of Sr and Th promoters. It seems that the matrix effect of the perovskite, which permits a high dispersion of Co in the reduced materials, plays an important role in obtaining a high selectivity toward higher hydrocarbons. G. SO2 REDUCTION The reduction of sulfur dioxide by CO to elemental sulfur can be effected over perovskites of transition metals (22, 226-228). However, in addition to the desired reaction

2co + so2

-

2c02 + 4s

(23)


299

a second reaction occurs between CO and the elemental sulfur produced during the course of the reaction:

co + 4s

-

cos

(24)

Carbonyl sulfide production usually proceeds to a substantial extent over conventional iron oxide (229) and supported-copper catalysts (230-232). This compound, which is highly toxic and undesirable as an effluent contaminant, seems to be produced on the sulfide species of these catalysts. Happel et al. (22, 228) reported that COS production can be greatly decreased with a catalyst consisting essentially of the perovskite LaTi03by means of which high conversion of SO2 at space velocities of up to 40.000 h-' were obtained. This improved performance is thought to be related to the high resistance of LaTiO3 to form sulfides with the elemental sulfur produced by the reaction given by Eq. (23). The oxide-versus-sulfide stability has been measured by the difference between the heats of formation of simple oxides and sulfides for several metals. This difference was found to be a maximum for titanium (22). Kinetic data obtained in a flow reactor with several CO + SO2 mixtures and LaTi03 as catalyst indicated that the reaction is first-order with respect to SO2 partial pressure and zero-order for CO partial pressure. These results were, therefore, correlated by the rate expression r = k .Psa. The reduction of SO2 with CO on Lac003 has been investigated by Bazes et al. (226) in the temperature range 284-465°C and CO concentrations between one and three times the stoichiometric amount for complete reduction. For temperatures above 380°C and a CO concentration of 2.5 times the stoichiometric amount a reduction of SO2 of 90% or higher was obtained. However, substantial amounts of COS were produced at high temperatures and large excess of CO. The formation of COS was lessened by decreasing CO : SO2 initial concentrations to the stoichiometric ratio. These authors also'found that the kinetic data correlated with the power rate law equation: r = k Pio0,* P$o. From Arrhenius-type plots (In r vs. 1 / T ) as well as from effectiveness factor values they concluded that LaCo03 exhibited important pore diffusion effects, particularly at high temperatures. Hibbert and Tseung (227) investigated the effect of water and oxygen on the reduction of SO2 over La&ro.&oO3 as well as the changes in chemical composition undergone by the catalyst during the reduction reaction. Unlike alumina-supported metal catalysts, L ~ . 5 S r o . ~ C tolero03 ated high levels of oxygen in the gas stream, provided sufficient SO2 was present to react with the oxygen. Water vapor ( ~ 2 %did ) not adversely affect the reaction; however, hydrogen sulfide was detected, although its concentration was at least an order of magnitude lower than that of COS.

300

L U I S G . TUUCA ef at.

Tracer studies with I4C and 35S also provided information about the mechanism of the overall reaction given by Eq. (23) (233). Experiments employing reaction mixtures containing all four species or only carbon or sulfur compounds, with steam as a diluent, were carried out with labeled 14C02and %. It was possible to measure the reverse 14Ctransfer when sulfur compounds were removed from the reaction chamber. Thus the exchange reaction

+ co

1 4 ~ 0 ~

-

cot + 14co

(25)

furnished evidence of transfer of I4C from C02 to CO. Carbon-14 exchange still occurred in the presence of sulfur, although at a slightly lower rate. On the other hand, the exchange reaction

so2+ 3 5 s

-

+s

3 5 ~ 0 ~

(26)

took place at a lesser extent in the absence of CO and C 0 2 than when these compounds were present. The low rates of I4C exchange, coupled with the zero-order reaction in CO concentration, suggested that the rate of C02 desorption may be important in controlling the rate of the overall reaction. The formation of carbonate structures and their stability under reaction conditions could, therefore, be an important factor influencing the catalyst performance.

H. ELECTROCATALYSIS AND PHOTOCATALYSIS The cathodic reduction of oxygen is one of the most important processes in energy conversion systems. Metallic oxides such as Pt-doped tungsten bronzes (234-238), LaNi03 (239-242), and partially substituted Lnl-,Sr,O3 (Ln = lanthanide element, mainly La) perovskites (6, 242247) were reported to have properties suitable for this reaction. The catalytic performance of these materials for the electroiytic evolution of oxygen has also been investigated (246, 247). The Tafel slopes of oxygen evolution were dependent on the perovskite-type electrode; however, a common mechanism seemed to be operative, i.e., the electrochemical adsorption of OH- followed by the electrochemical desorption of OH (rate-determining step) yielding H202as an intermediate that then undergoes catalytic decomposition to 02. The reduction of oxygen has been investigated over either unsubstituted or partially substituted perovskites (6, 2.34-247). Matsumoto et al. (240) used a rotating ring-disk electrode that consisted of a LaNiO3 disc and a Pt-ring electrode. These authors proposed the following reactional


scheme: 0 2

-

02fnd.)

e- + O Z W ~ ) Hz0 e-

+ Olw

Oldr)

-

HOzfadr) + OH-

+ HO~adr,

HOiads,

30 1

(27) (28) (29) (30)

The rate-controlling step was found to be that represented by Eq. (30) when the Tafel relation holds (low polarization region) or by Eq. (27) for the limiting current (high polarization region) (241). The subsequent steps to (30) must conform to the experimental result of nondetection of HO; ion in the electrolyte. The rearrangement of adsorbed HOT on the electrode surface with formation of OH- in a manner such as represented by Eq. (31) has been proposed (240): 0

O-h!l-O-OH

-

PH.

0-M-b

(31)

where 0 denotes an 02-vacancy on the perovskite surface. Equation (3 1) should be fast since no dissolved HOT was detected at the ring electrode. Such a mechanism with an electron transfer from the electrode to the cr* orbital of the adsorbed HOT should involve the break of the cr bond of this species with the subsequent formation of OH-. The surface composition of the La and Ni perovskite may be represented as LaNi03-h,where A = 0.2-0.3 in the oxygen-reduction potential. Bockris and McHardy (234) found that NaxW03bronzes containing traces of Pt exhibited a catalytic activity for oxygen reduction that was higher than that of Pt metal, although other studies (236, 237) failed to confirm this finding. Na,W03 and other tungsten bronzes also showed activity for the electrochemical reduction of H202. For this reaction, Randin (248,249) found an inflection in the current-potential curves that was attributed to an inhibition phenomenon. The fact that the relative position of these curves undergoes a regular shift as the H202 concentration changes suggests that at a given potential the same fraction of surface sites are blocked independently of the Hz02 concentration. This situation can arise from a parallel reaction with incorporation of hydrogen to the bronze NaxW03t nH+ t ne-

-

Na,HnW03

(32)

This reaction has been shown to occur in the potential range where the inhibition region starts, i.e, about 0.2 V. Because the activity of Na,H,W03 for the reduction of H202 is lower than that of Na,W03, the number of active sites should decrease as the reaction progresses.

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LUIS G . TEJUCA

et al.

The oxygen evolution has been chosen by several authors for kinetic studies that involve rate-determining step evaluation. This reaction has been studied on metals as well as on oxide electrodes. Within this latter category, RuO2 has been most frequently investigated. On this oxide, the reaction has been suggested to occur with participation of lattice oxygen (250). Oxygen evolution on perovskites has been studied in alkaline solution (251-253). Bockris and Otagawa (246,247) carried out a careful and systematic investigation on several substituted perovskites in an attempt to correlate their electrocatalytic properties with their structure, laying particular emphasis on the bonding of oxygenated intermediates at the surface. These authors showed that the most likely mechanism involves a rate determining M-OH desorption (where M is the transition element in LnMO3). A careful analysis showed that differing M-OH bond strengths among the perovskites give rise to different adsorption isotherms in the liquid-solid interface. The mechanism found most consistent with the kinetic data (253) is Mz + OHMz - OH (H@Z)(phyr)

(H202hphys)

+

+ OH+ OH-

(HOThphys)

=Mz - OH + e-

-

M L H 2 O 2 + e-

(HoThphys)

+ H20

H20 + OH-

0 2

(33) (34) (35) (36)

A model based on the MO theory has been proposed for the active (0,0,1) plane of the perovskite surface, viz., the MOj cluster (Fig. 22a), in which the 3d (eg, rzg) levels will split further because of the lower symmetry at the surface. Wolfram and co-workers’ (254) calculations for SrTi03 showed that the eg levels of the M06 cluster (bulk) split into the d,~,2and d,l levels, with the latter laying below the former at the surface. Also, the t2g levels split into the double degenerated d,, and d,, states and the singlet dJystate. These energy levels are schematically represented in Fig. 22b. Assuming that these calculations also hold for other perovskites it is possible to assign the d-electron configuration of the M3+ ion at the surface as shown in Fig. 22b. The 4 2 orbital will be occupied by an electron only in M = Ni and Co (high spin configuration). On the other hand, the electronic structure of OH may be represented as follows: Is2 2s2 (2p,+ 1s ) 2pz ~ 2p:, in which the bonding state is an 02p, orbital o-bonded to a HI s state. In the interaction between an M3+ ion at the surface of the perovskite and an OH species, the d22 orbital will overlap with the (2p, + 1s) MO of an OH forming c+-typeorbitals whereas the d,,(d,,) orbital will interact with the 2px(2py)of an OH, giving r-type orbitals. On the basis of the d-electron configuration, MO diagrams for the MZOH bonding at the surface of LnM03 oxides (M = Mn, Ni) were con-


OHSolution

OH-

OH-

303

y-

OH- OH- OH- OH OH-

0 Transition metol Mz

8

:oOXni::-ion

FIG.22. Schematic details of the perovskite surface in electrocatalysis:(a) model for the active surface; (b) d-electron configuration of M3+ ions; (c) MO diagrams for the M*+-OH (M = Mn, Ni) bonding. (Reprinted by permission of the publisher, The Electrochemical Society, Inc., from Ref. 247.)

structed (Fig. 22c). The following observations can be derived from these diagrams: (1) electrons in the tiz* orbital, e.g., Ni3+and high spin Co3+ ions, may play an important role in achieving high reaction rates, since they tend to occupy (+*-antibondingorbitals of MZ-OH, resulting in a weaker bond; (2) since OH is a saturated ligand, the contribution of rbondings, i.e. back-bonding to MZ-OH, would be negligible, especially in Co3+and Ni3+;and (3) the poor catalytic activity of V3+and Cr3+ions may be due to a higher oxidation state easily formed on those perovskites. Matsumoto et al. (252,253,255) have made an interpretation of the relative reaction rates in terms of geometrical factors of the surface. They consider the overlap integrals between the sp, orbital of 0 and the eg orbital of the transition-metal ion that are assumed to be related to the o * character of the bonds in the perovskite. In the same line of electrocatalytic applications, some perovskites have been shown to be suitable cathode materials in high-temperaturefuel cells (256, 257). Successful and economical operation of a high-temperature zirconia fuel cell has been achieved by means of porous PrCo03cathodes. With hydrogen as the fuel and oxygen as the oxidant, power densities of

304

LUIS G . TEJUCA

et ul.

about 0.3 kWlcm2 at 1000°C for periods in excess of 5000 h have been generated (256). Such cells, however, do not survive thermal cycling because the thermal expansion coefficient of &Coo3 is 2.5-3.0 times larger than that of zirconia, which leads to interfacial stresses during cooling. Perovskite-type oxides have also found application in several photoassisted reactions. Among these, the photolysis of water over BaTiO3 electrodes (258), SrTi03 powders (259), and NiO-coated SrTiO3 powders (260), the ammonia synthesis over metal oxide-coated MTiO3 (M = Sr, Ba) powders (261), the oxidation of CO over LaCo03, and MTiO3 (M = Sr, Ba) (262). have been investigated. Kennedy and Frese (258) have measured photocurrent efficiencies and the distribution of potential in semiconducting BaTi03used for the water-splitting. By applying the standard techniques for studying single-crystal electrodes to polycrystalline sintered BaTiOs, these authors have demonstrated a temperature effect of the photocurrent efficiency and used the semiconductor depletion layer theory to interpret the efficiency data in terms of solid-state properties. Later on, Lehn et al. (263) observed that SrTiO3 particles alone cannot catalyze the photodecomposition of water; rhodium coating was essential for the production of H2 and 0 2 , with Rh(II1) as the only active species. The preparation of the same type of catalyst from ultrafine powder SrTi03 (259) did not enhance the photocatalytic activity in the water-splitting reaction. Other surface modifications such as NiO coating where Ni metal is said to exist at the NiO-SrTiO3 interface have been attempted (260). The existence of Ni metal at this interface suggested the transfer of electrons between both materials. The photocatalytic synthesis of ammonia from N2 and H2O was studied by Li et al. (262) over SrTi03-and BaTi03-basedcatalysts. These authors demonstrated that the thermodynamically unfavored reaction can proceed at 50°C using band-gap illumination of a RuO2-NiO doubly coated MTi03 (M = Sr, Ba) catalyst covered with liquid water. The catalytic effect was interpreted in terms of the ability of Ru for N2 activation and of Ni for H2evolution. The reverse reaction (photodecomposition of ammonia), which was coupled with photodecomposition of water, became important when ammonia was accumulated in the reactor. A strong photocatalytic effect was observed by van Damme and Hall (262) for CO oxidation on MTiOJ (M = Sr, Ba). In addition, they found a very similar behavior related to the ferroelectric properties of the catalyst for CO oxidation over BaTiO3 in the dark. The effect of the light was discussed in terms of band-to-band transitions followed by hole capture, whereas ferroelectricity would arise from the effect of polarized surface layers on the shape of the bands.


305

I. OTHERREACTIONS

-

The homomolecular exchange of oxygen (HEO) I60

+

I80

16Ol.90

(37)

on LaM03 (M = Cr, Mn, Fe, Co, Ni) and LnCo03 (Ln = Nd, Sm) oxides has been investigated, in a broad range of temperatures, by Sazonov et al. (277).Kinetic results for the H E 0 and for the electrochemical reduction of oxygen over a partially substituted La&ho.~Co03perovskite were also reported by Hibbert and Tseung (264). The rate of H E 0 on samples that have been equilibrated with an isotopic (I602 + I8O2) gas mixture is given by C" - co In C" - C' = kt where Co,C", and Cf are the initial, equilibrium, and current concentrations, respectively, of l6O1*0molecules and k is the rate constant for the H E 0 reaction (265). Values of the rate constant k were determined from Eq. (38) at several temperatures, and the number of oxygen molecules reacting per second and per unit of surface area Z was calculated by

where NR is the total number of oxygen molecules in the reactor and S is the surface area of the catalyst. The activation energies of the exchange of oxygen for the LaM03 compounds in the high-temperature region (>300"C) were calculated from Arrhenius plots log Z versus 1/T. These, together with the activation energies for reaction (37) and the energy of the oxygen bond (M-0) for the simple oxides of the transition elements, are depicted in Fig. 23. As can be seen, the activation energy of the H E 0 for LaM03 and for the corresponding simple oxides follow almost the same trend. This behavior has been explained considering that the catalytic activity of the LaM03 oxides in the H E 0 is determined by the M3+ ions with no significant contribution of the La3+ions in the compounds. Such an explanation is supported by the important differences observed in the activation energies (AE,) and rates of the HE0 between nickel compounds where the transition metal is in + 3 or +2 oxidation states, e.g., LaNiO3 and NiO. These differences were found to be 20 kcal/mol for AE, (Fig. 23) and more than three orders of magnitude for the rate of exchange, respectively. However, the similarity of the temperature dependence of the exchange for LnM03and the corresponding LnzO3(Ln =

306

LUIS G . TEJUCA

et af.

50 t C903 Mn%

F e 3 Co304 NIO

P

40

LaCr4

I

LaFe4

LaMnO3

I

LaNi4

Lac003

FIG.23. Comparison of the activation energy of oxygen exchange for h M 0 3 oxides (squares) with the activation energy of exchange (circles) and the bond energy of oxygen (triangles) for the corresponding simple oxides. (Reprinted by permission from Ref. 177.)

La, Nd, Sm) oxides in the low-temperature region (Bnin hydrogenolysis (0,ethane; 0, cyclopropane). (Reprinted from Ref. 23 with permission of The Royal Society of Chemistry.)

The effect of stepwise heat treatments on the activity of NiszB38 was studied in the hydrogenolysis of ethane and cyclopropane (23, 41). The reaction rates of the hydrogenolysis of both compounds show maxima on samples heat-treated at 623 K and sudden decreases on samples heattreated above 623 K (Fig. 14). The explanation for the changing activities has already been given in connection with the hydrogenation of olefins (see Section III,C, 1). For hydrogenolysis the formation of more electrondeficient nickel species was found to be responsible for the increasing activity. However, the decreasing activity occurring on catalysts heattreated above 623 K was correlated with the decreasing number of active species as shown by hydrogen chemisorption, SEM,and ESCA. Additionally, the authors point out that the relative catalytic activity of hydrogenolysis to hydrogenation shows a large increase at the higher heattreatment temperatures (Fig. 15); that is, hydrogenolysis is less affected by heat treatment.

m

673 773 Temperature ( K I

FIG.IS. Change of relative rates of hydrogenolysis of ethane (0)and cyclopropane (0) against the rate of hydrogenation of ethene as a function of thermal treatment of Ni6?B3*. (Reprinted from Ref. 23 with permission of The Royal Society of Chemistry.)

372

ARPAD MOLNAR et al.

This latter was accepted as evidence for possible aggregation, since some literature data indicate that hydrogenolysis, being a structure-sensitive reaction, involves multiple adsorption (158-161). However, two recent review articles on particle size effects on metal catalysts (162, 163) emphasize the inconsistency of the rather scarce data on nickel and warn of the difficulties connected with this problem. These data, although scarce, support the conclusion that the structure in the precrystallization state is the most favorable for catalytic activity, but detailed understanding of the phenomenon requires further clarification. 2. Transformations of NO

In the decomposition of NO Nigo-,Fe,Pzo amorphous alloy ribbons showed a high initial activity followed by a decrease and establishment of a steady state (264). The process turned out to be an oxidative dissociation with the rate-limiting step being the oxidation of the amorphous surface. During NO decomposition the BET surface area of FegoPzo increased from 0.2 to 0.8 m2 g-'. With increasing temperatures NiaFezoPzo showed a sudden decrease of activity at 636 K, which is the glass transition temperature of the alloy. A sample crystallized in helium at 773 K exhibited activites similar to the alloy crystallized in situ. The data do not permit one to decide whether the low activity of the crystalline materials is due to the slower diffusion of oxygen or the lower activity of the surface for inducing dissociation. The same catalysts were studied in the reduction by hydrogen of NO to Nz and NH3 as well (125, 164). The most active NigoPzo alloy showed an activity 10 times higher than the crystalline catalyst and exhibited interesting changes in activity. The initial high activity first decreased markedly, then increased, and finally, after about 2.5 h at 523 K, reached a steady-state activity as high as the initial activity. Although a satisfactory explanation is not given, the presence of a relatively large amount of N 2 0 in the product mixture during the period of activity change might indicate partial oxidation of the surface. Similar product compositions (90% Nz plus N20 and NH3) with high (>80%) conversions were attained on Ni-Fe-Zr alloys. Pt70Z1-30was also shown to be active in the reduction of NO with CO at 523 K, giving a 9 : I N2-N20 mixture at 93% conversion (55). 3 . Ammonia Synthesis The alloys studied in the ammonia synthesis are Ni-Zr and FeglZr9(22, 49,112, 252). With these catalysts, marked increase in BET surface areas

N E W CATALYTIC MATERIALS

373

and continuous activity increases were observed during the reaction (22, 49). This behavior also is characteristic of zirconium-containing amorphous metal alloys in the hydrogenation of CO (see Section III,B,3). In the case of the ammonia synthesis the activation process continued until about 1000 h on stream. In contrast, the crystalline FeglZr9alloy showed a continuous decrease of activity, reaching a steady-state activity similar to that of the amorphous material (112) whereas a steady deactivation was observed on the crystalline Ni-Zr alloys (22). The results of long experiments (several thousand hours) indicate that both the amorphous and crystalline Fe-Zr alloys transform to a similar active surface state under the influence of the reacting gas mixture. The turnover rates of the catalysts prepared from the amorphous Fe-Zr precursor after pretreatments with different hydrogen-nitrogen mixtures at 653-723 K were one order of magnitude higher than those of a polycrystalline iron ribbon prepared by melt-spinning. XRD patterns taken at different reaction times revealed the simultaneous formation of both large and small iron particles during the activation period (49). SEM showed the segregation of irou on the surface in the form of small iron particles that were found to be related to high activity. This segregated iron resulted in an increase in BET surface area and was considered to be stabilized by the zirconium oxide present. Gradual increase of crystallinity was also observed on amorphous Ni24Zr76 and Ni64Zr36(22). The catalyst bed after reaction contained two distinct zones. Large Ni crystallites (XRD) and surface Zr02(XPS) were detected in a black, inactive bed entrance zone, but no zirconium nitride and almost no metallic zirconium were found. In contrast, an active, golden zone showed only traces of crystallinity (XRD) and had a ragged surface (SEM). It contained small, strongly disordered Ni particles, NiO, zirconium oxide, and zirconium nitride. Nonstoichiometric Z I O - ~was prevalent on the surface, whereas the subsurface region contained metallic Zr as well. The amount of ZrN increased with depth. The high activity was attributed to the presence of the small Ni particles embedded in the zirconium oxide-zirconium nitride matrix. Nitrogen adsorption measurements indicated the coexistence of molecularly and dissociatively adsorbed nitrogen, the latter reacting with added hydrogen to form ammonia. These observations, together with kinetic data, prove that dissociatively adsorbed nitrogen is the most important reaction intermediate in ammonia synthesis. A remarkable new observation is also attributed to the presence of the amorphous structure in connection with the ammonia synthesis (165). In situ X-ray diffractograms showed the presence of a-iron in catalysts prepared from unpromoted Fe304after treatment with hydrogen at 723 K. In contrast, however, promoted industrial (BASF and ICI) ammonia synthe-

374

AWAD MOLNAR et al.

sis catalyst precursors produce almost fully reduced catalysts with very broad diffraction peaks as a result of the actior, of either hydrogen or the reacting gas mixture at about 700 K. The XRD patterns are very different from those of finely divided and paracrystalline a-iron but show remarkable similarity to those of certain iron-containing glassy alloys. These data, as well as other unpublished data (Mossbauer spectroscopy, adsorption and kinetic studies, electron microscopy), indicate that the catalytically active phase of these reduced, promoted catalysts is amorphous iron, interspersed with microcrystalline promoter phases (e.g., FeA103). However, further study will be required to elucidate the exact nature of noncrystallinity in the active catalysts. 4. Decomposition of Formic Acid

Because of the milder reaction conditions (lower temperature) the decomposition of formic acid is found to be more sensitive to catalyst structure than the hydrogenation of carbon monoxide (20). Although increasing cooling rate (increasing disk velocity) has almost no effect on catalyst performance, opposite selectivites were observed on amorphous and crystalline FewBzoalloys (27). On the amorphous catalyst the main transformation is dehydrogenation of formic acid to form C 0 2 , whereas on crystallized samples excellent selectivities are exhibited for dehydrogenation. (At 458 K, CO is formed with 100% selectivity.) But during this latter process it is not clear whether HzO is also detected, a feature that seems characteristic only of Ni (166). XPS revealed substantial differences between the structure of the amorphous and the crystalline Fe-B alloys. The amorphous Fe-B, in which small iron ensembles are separated and stabilized by iron and boron oxides, can be considered an oxide-supported, highly dispersed iron catalyst. In contrast, on the crystalline alloy large iron particles exist. At present there is no explanation for the structure-selectivity correlation; also, in a similar manner, the variables influencing the formation of C02 and CO on Fe(100) were not determined in earlier studies (167). IV. Conclusions

A. MAIN ACHIEVEMENTS Despite the rather short period since the attention of catalytic chemists turned to metallic glasses as catalysts, there have already been a remarkable number of publications, and there are already some rather well-studied reactions and catalysts. These studies led to the identification of the probable active phases of certain transformations, disclosed the way they

NEW CATALYTIC MATERIALS

375

form, and permitted explanations of how they act in the catalytic process. Certain correlations could also be established between activity, selectivity, and stability, and in some reactions, the superiority of the amorphous state over the crystalline catalysts could be explained. In electrocatalysis the use of metal-zirconium alloys in water electrolysis is a good example. It seems to be well established that the very high electrocatalytic activity of the HF-treated amorphous Ni and Cu electrodes is related to the presence of a rough, Raney-type surface structure. The beneficial effect of the amorphous structure is that it ensures the ready formation of a thick, stable, active surface layer. Another thoroughly and widely studied field is the hydrogenation of carbon monoxide. Although a large number of investigations have been performed on many different catalyst compositions, only the zirconiumcontaining alloys and the Fe-(Ni)-B system yield definite and general conclusions. The continuous activity increase of the zirconium-containing alloys during reaction is correlated with the formation of a porous, highly active surface. The active phase for the CO hydrogenation reaction proved to be fine metal particles that are embedded in the ZrOz matrix. Results on the Fe-(Ni)-B system also indicate that the amorphous alloy as a precursor ensures the formation of quasi-crystalline (near-crystalline) structures. These otherwise nonexisting phases are stabilized by the amorphous matrix and may have very specific catalytic properties. A similar phenomenon turned out to be the case in the synthesis of ammonia catalysed by Ni-Zr and Fe-Zr. The advantage of these catalysts is that the amorphous structure is the precursor for generating a controlled, uniform dispersion of active species in high concentration in a chemically homogeneous environment. Progress has been made in overcoming some of the structural and stability problems inherent in metallic glasses. The low-surface-area problem has been attached by several chemical and physical roughening techniques. The problem of shaping the glass has been circumvented by adhering a ribbon of the crystalline alloy to a Ni plate and then converting the alloy surface to glass by the laser melting-self-quenching technique. Moreover, the stability problem is receding as more glasses are examined under reaction conditions and some are found to be remarkably stable. Disconcertingly, others are unexpectedly unstable under reaction conditions even though their T, are relatively high. B. FUTURE PROSPECTS One drawback of amorphous catalysts is their metastable character, which can hinder their high-temperature use. There are methods, how-

376

ARPAD MOLNAR et al.

ever, which are suitable for improving the thermal stability of amorphous materials. Alloying of properly selected components can result in amorphous alloys with either increased T, values or increased activity, which permit low-temperature applications. Overcoming the stability problem can increase the possibilities for the use of these materials as catalysts. Clearly, this is an important subject for future research. Another subject for future exploration is the surface area of metallic glasses. Ribbons produced by rapid quenching have surface areas of little larger than the geometric surface; that is, the specific surface (surface per weight) is very low for industrial applications. This problem may be overcome by some amorphous alloys prepared by new methods developed recently. In turn, different atomization techniques (gas, liquid, centrifugal, electrohydrodynamic atomizations) and comminution methods (milling and chopping) produce amorphous powders with higher specific surfaces. Powders or films prepared by either chemical deposition or electrodeposition also deserve higher attention. Although results on chemically deposited Ni-B and Co-B amorphous alloys are available (168-171), clarification of the exact nature of their structure and further examination of their catalytic activity are needed to gain insight into the structure-activity relationship and to compare the effect of preparation methods of amorphous alloys on their catalytic properties. Chemical deposition of alloys onto suitable carriers might produce supported amorphous alloys similar to the traditional, high-surface-area supported catalysts. The development of amorphous alloys with compositions specially designed for catalytic applications is another field for future work. Most amorphous alloys studied in catalysis have been those that were prepared because of their outstanding physical properties for which they might even have had technical applications as such. However, a carefully planned and executed project could result in amorphous catalysts with the proper combination of alloying elements in suitable ratios to ensure high thermal and chemical stabilities as well as high catalytic activity and selectivity for a given transformation. A good illustration of this strategy is the development of the multicomponent Pd-Ir-Ti-Rh(Ru)-P alloys, electrode materials for the electrolysis of sodium chloride. Investigations started with Pd-P; then, with the addition of other components, gradual improvements of different properties were attained, eventually reaching compositions with excellent corrosion resistance and an activity surpassing that of the traditional RuOz/Ti electrode. Continuing studies using surface characterization techniques are necessary for better understanding of amorphous catalysis. The comparison of amorphous and heat-treated catalysts can reveal what changes other than


377

crystallization (changes of nominal composition, phase separation, segregation, chemical transformations) can possibly give rise to the large differences in activities. Analysis of the catalyst surface before and after reaction can not only give information about the stability of the amorphous state but can also detect the exact nature of changes taking place on the action of the reacting chemicals. The interpretation of the continuous activation or deactivation of certain alloys during reaction is naturally intriguing, but a search for minute alterations of the surface structure and the recognition of their effect on catalytic properties are essential for understanding such unstable systems as amorphous alloys. Similar studies on catalysts that require activation can disclose the formation of the activated state and can reveal the nature of active sites formed. Besides the modern physical methods, the use of sterically complex organic probe molecules is also an excellent method for characterizing amorphous alloys, and, in turn, the amorphous structures are likely to be useful for modeling catalytic surfaces with high concentrations of unorganized coordinatively unsaturated sites. For example, the idea of active sites consisting of clusters of atoms (172-177)may be examined on metallic glasses. First it was Kobozev who suggested that the catalytically active centers are atomic in nature. The active centers consist of amorphous ensembles of a few metal atoms that do not form a crystal lattice. It follows from this observation that the ensemble theory is in serious contrast with those theories that ascribe the catalysis to the crystal phase. Similar to the Kobozev adsorption catalysts and alloy catalysts, in which the active metal is present in very low concentration, metallic glasses seem to contain the active centers in ensemble form; they are, therefore, intermediate between the molecular and the crystalline systems, that is, between homogeneous and heterogeneous catalysts in the classical sense. Finally, it seems reasonable that amorphous alloys will find applications in a broader range of catalytic reactions. In fact, their unique activities may render them suitable catalysts for reactions heretofore not ordinarily considered in the realm of catalysis. ACKNOWLEDGMENTS We gratefully acknowledge National Science Foundation Grant INT-8403357 and Hungarian Academy of Sciences grant 319/82/1.3 for support of our collaboration. REFERENCES

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Addendum to Structure and Reactivity of Perovskite-Type Oxides LUIS G. TEJUCA AND JOSE LUIS G. FIERRO lnstituto de Catdlisis y Petroleoquimica C.S.I.C. Serrano 119, 28006 Madrid, Spain

AND

JUAN M. D. TASC6N Instituto Nacional del Carbdn y sus Derivados C.S.I.C. Apartado 73, 33080 Oviedo, Spain

Recently, some interesting publications appeared on catalytic properties of perovskite oxides, including perovskite superconductors, in oxidation and reduction processes. Hibbert and Campbell (311) studied the interactions of SO2 and CO with Lal-,Sr,Co03 and used these perovskites for the catalytic removal of SO2 by CO in the interval 500-600°C. In the used catalyst, sulphides of each metal predominate, with no perovskite oxides or perovskite sulphides being detected. The catalyst with x = 0.3 gave the optimum removal of sulphur with no COS formation. France et al. (312) used a series of Lal-,A,Mn03 (A, an alkaline metal or a vacancy) oxides at 820°C as catalysts for the oxidative coupling of methane. Higher binding energies for oxygen were correlated with higher selectivities for C2 hydrocarbons. Nagamoto ef al. (313) studied AB03 oxides (A, alkaline earth metal; B = Ti, Zr, Ce) at 750°C for this transformation. The catalytic activity for C2 products was found to decrease in the order Ba > Sr > Ca and increase with increasing values of a parameter (Ad) which is a measure of the deviation of cations and anions from the ideal equilibrium distances. YBa2Cu30, (x = 6.88) was reported to retain large amounts of NO at 300°C (314). This effect was explained by absorption into the solid and/or 385

Copyright 8 1989 by Academic Press. Inc. All rights of reproduction in any form reserved.

386

LUIS G . TEJUCA

et al.

reaction with the bulk to form metal nitrates. In addition, this superconductor catalyzes the NO decomposition in the presence of oxygen (315). The oxygen deficiency of the sample seemed to influence strongly its catalytic activity. A similar superconductor (6.5 Ix I6.9) was used by Mizuno et al. (316) for the reduction of NO by CO at 300°C. Hansen et al. (317) reported data for toluene oxidation in the presence of oxygen and ammonia on YBazCu@6+, at 400°C. At low fa the catalyst is very active for mild oxidation to benzonitrile, whereas at high Po, deep oxidation to COz is preferentially catalyzed. This transition in product selectivity is reversible and occurs at Pa = 8.65 x lo3Pa. At transition conditions, the catalyst composition was found to be close to YBazCu306 (x = 0). An increase in the content of lattice oxygen (x > 0) makes the catalyst selective for COZ formation. REFERENCES 311. Hibbert, D. B., and Campbell, R. H., Appl. Card. 41, 273, 289 (1988). 312. France, J. E., Shamsi, A., and Ahsan, M. Q., Energy Fuels 2,235 (1988). 313. Nagamoto. H., Amanuma, K., Nobutomo, H., and Inoue, H., Chem. Lett. p. 237 (1988). 314. Tabata, K., Fukuda, H., Kohiki, S., Mizuno, N., and Misono, M., Chem. Letr. p. 799 ( 1988). 315. Tabata, K., and J. Mater, Sci. Lett. 7, 147 (1988). 316. Mizuno, N., Yamato, M., and Misono, M., Chem. Commun. p. 887 (1988). 317. Hansen, S., Otamiri, J., Bovin, J.-O., and Anderson, A., Nature (London) 334, 143 ( 1988).

A

ammonia synthesis, 372-374 characterization methods, 332-333 composition, 329-331 crystallization and activity, temperature effects, 334-336 electrocatalysis as electrodes in fuel cells, 342-344 sodium electrolysis, 339-342 water electrolysis, 336-339 formic acid decomposition, 374 future prospects, 375-377 hydrogenation catalysis Of alkynes, 368-369 of (+)-apopinene, 358-362 of 1.3-butadiens 365, 367 of carbon monoxide dissociative chemisorption, 344-345 over nickel(iron)-phosphorus(b0ron) catalysts, 346-351 over zirconiumcontaining catalysts,

Acetons hydrogenation over Rnney-type nickel catalysts, 370 Acetylene conversion to benzene over palladium, 23 hydrogenation over platinum, particle size and, 111 Acrolein, hydrogenation over amorphous nickel-titanium, 367 Actinides, incorporation into perovskites, 315-316

AES,see Auger electron spectroscopy Alcohols, conversion over perovskites adsorption, 278-279 dehydrogenation, 309-311 oxidation, 289 Alkylens. hydrogenation over amorphous and crystalline alloys, 368-369 Alloys amorphous, see Amorphous metal alloys singlecrystal surface cyclohaane dehydrogenation over copper-ruthenium, 41-43 n-hexane hydrogenation over goldplatinum, 40-41 metal particle size effects, 80-82 properties. 39-40 Alumim-~~pported catalysts, 46,228 Ammonia, oxidation over perovskites, 307 Ammonia synthesis over amorphous metal alloys, 372-374 over iron, 24-25 on alumina support, 47 antipathetic behavior, 150, 152 particle size and, 131-132 promotion by potassium, 36-37 photocatalysis over perovskites, 304 over rhenium, 24-25 promotion by potassium, 37 Amorphous metal alloys, see a h specifi alloys activation methods. 333-334

333, 352-357, 375

of cfs-cyclododecene, 358 of 1-huene, 358, 359 of olefins, 362-367 Raney-type nickel catalyst activities, 369-370

hydrogenolysis catalysis, heat treatment and, 370-372 preparation methods, 331 (+)Apopinenenq hydrogenation over amorphous palladium alloys catalyst preparation methods and, 358-360 catalyst purification methods and, 360-361 terminal alkyne effects. 361-362 AR, see Atomic rate of reaction Atomic rate (AR) of reaction definition, 72-73 structure sensitivity and, 86-87 Auger electron spectroscopy (AES) small metal particles. 108, 109 titania-supported catalysts, 203-205, 208-209

UHV analysis and, I, 4, 11 387

388

INDEX

B Benzene hydrogenation metal particle size and, 114-116 TPD on rhodium/titania and platinum/titania. 210 oxidation over perovskite, 288 production from cyclohexane over platinum, 18 Bimetallic catalysts, see Alloys Boron in iron-boron amorphous alloys, 344, 349-350

in nickel-boron amorphous alloys, 365-366

Brunauer-Emmet-Bller (BET)surface, amorphous alloy catalysts, 332 ammonia synthesis and, 372-373 carbon monoxide hydrogenation and, 352-354

nitric oxide decomposition and, 372 1,3dutadiene, hydrogenation over amorphous alloys copper-zirconium activity, 367 nickel-boron activity, 365 over platinum, 20-21 nlutane, hydrogenolysis, metal particle size effect, 120 1-Butene, hydrogenation over palladium/alumina, TOF, 113-114 1-Butyne, hydrogenation over palladium/alumina, TOF, 113 Butyronitrile, hydrogenation over b e y type nickel catalysts, 370 C

Calcination, in precursor-support interaction, 70-71 Carbon dioxide, adsorption on perovskites, 273, 275-276

Carbon monoxide adsorption on perovskites, 273-375 chemisorption on supported metals, 95 IR spectroscopy, 103 structure sensitivity, 138-140 titania coverage studies, 206-209 TPD studies, titania-supported metals, 186-190, 209-212, 224-225

hydrogenation. see also Fischer-ltopsch reactions, methanation reaction

over amorphous metal alloys dissociative chemisorption, 344-345 nickel(iron)-phosphorus(boron) activities, 346-351 zirconium-containing alloy activities, 333, 352-357, 375

over pemskites, 2%-298 over single crystals, 26-27 sulfur addition and, 30 over supported metals chloride effect, 201-202 ensemble effect, 149-150 metal-support interaction, 44-46 nickel-copper catalysts, 197-198 particle size and, 128-131 rhodiudtitania, reduction temperature effect, 196-200 ruthenium-based catalysts, 87 titania coverage effect, explanations, 214-220

oxidation over perovskites catalytic activity, 279-281 kinetics on lanthanum cobaltate, 281-283

mechanism of, 282-283 pollution control and, 291-293 sulfur dioxide poisoning effect, 313-315 over single crystals, 24-26 over supported metals, particle size and, 133-136 sulfur dioxide reduction over perovskites, 298-300

Carbonyl sulfide, production over pemskites, 299 Ceramic method, perovskite preparation, 245-246

Ceric oxide-supported catalysts, 226, 229 Cesium, metal catalyst promotion, 37-39 Chemisorption, on supported metals of carbon monoxide, 95 structure sensitivity, 138-140 of hydrogen, structure sensitivity. 91-95 Chlorides, effects on titania-supported catalysis, 201-202 Chlorine electrolytic evolution on amorphous metal alloys, 339-342 metal catalyst activity and, 31-34 Chromium, effect on amorphous and crystalline alloys. 368

389

INDEX Cobalt catalysts Fischer-Tropsch reaction, potassium effect, 36 supported, selectivity. 88 Copper-based amorphous alloys formaldehyde oxidation in fuel cells,

Dextrose, hydrogenation over Raney-type nickel catalysts. 370 Diffuse reflectance spectroscopy. unreduced ion detection in metal catalysts, 103 Dry evaporation, perovskite preparation. 246-247

343-344

zirconium-containing, hydrogenation catalysis, 366-367 Copper catalysts deposition on zinc oxide, SMSI, 46-47 effect on nickel activity, 197-198 effect on ruthenium activity, 41-43 single crystals, 27 poisoning by sulfur, 31 promotion by cesium, 37-38 Copper/zinc oxide catalysts, SMSI. 46-47 Coprecipitation. perovskite preparation, 247-250

Cordierite-supportedperovskites. preparation impregnation, 251-253 plasma spraying, 253 Cubooctahedron, small supported particle model, 80-81 replacement by icosahedron, 85 Cis-Cyclododecene, hydrogenation-deuteriumation over amorphous alloys, 358 Cyclohexane, dehydrogenation benzene accumulation over platinum, 18 stepped surfaces and. 82-83 metal particle size and, 115. 117 over rhodium/titania. reduction temperature and, 191-193 over ruthenium, copper addition and, 41-43 Cyclopentane, hydrogenolysis, metal particle size and. 121-124 Cyclopropane conversion over perovskites, 294 hydrogenation over metal systems, particle size and, 111-112 hydrogenolysis over nickel, 21-22 over nickel-boron amorphous alloys, heat treatment, 371 over iridium, 23

D Deuteriumation, of hydrocarbons over metal alloys. 358 over platinum catalysts, 19-20

E Electrocatalysis amorphous metal alloy activities as electrodes in fuel cells, 342-344 sodium chloride electrolysis, 339-342 water electrolysis, 336-339 perovskite activities oxygen cathodic reduction, 300-301 oxygen evolution, 302-303 Electron microscopy, small metal particle assay, 95-96 Electron paramagnetic resonance (EPR) metal particle size distribution, 99-100, 104 oxygen adsorption on perovskites, 273 poisons on metal surface, 108 Electron spin resonance (ESR), see Electron paramagnetic resonance Ensemble effect carbon monoxide hydrogenation and, 149-150

definition, 146-147 hydrogenolysis of small hydrocarbons and, 150 structure insensitivity and, 153-154 EPR. see Electron paramagnetic resonance ESR, see Electron spin resonance Ethane adsorption on perovskites, 278 hydrogen01ysis over iridium, 21-22 metal particle size effect, 117-120 over nickel catalysts boron effect in amorphous alloys, heat treatment and, 371 TOF maximum versus FE, 152-153 on tungsten support, 43 over rhodium/silica, 153 over rhodium/titania. SMSI, 194 over tungsten, 23 Ethene, see Ethylene Ethylene adsorption on perovskites, 278 combustion over perovskites, 287

390

INDEX

hydrogenation over amorphous metal alloys copper-zirconium activity, 366-367 nickel-phosphorus activity, 362-364 over perovskites, 293-295 over platinum/silica, particle size and, 110 oxidation over silver catalysts, 15-16 poisoning by chlorine, 32-34 promotion by cesium, 37-39 support effects, 90, 136-138 Ethylidine, hydrogenation over rhodium, 20 EXAFS, see Extended X-ray absorption fine structure Exhaust gas, purification by perovskites, 291-293

Explosion, perovskite preparation, 250 Extended X-ray absorption fine structure (EXAFS) analysis, 14 rhodium-titanium interaction, reduction temperature effect, 221,223 small metal particle assay size calculation from, 94 Structure, 108-109

F FE, see Fraction exposed, of total metal atoms Ferromagnetic resonance (FMR), metal particle size distribution, 99-100, 104 Fischer-ltopsch reaction over lanthanum rhodate, temperature effects, 297-298 potassium-promoted over cobalt catalysts, 36 over iron catalysts, 35 over supported cobalt, selectivity, 88 FMR, see Ferromagnetic monance Formaldehyde, oxidation on amorphous alloy electrodes, 343-344 Formic acid, decompoiition over amorphous metal alloys, 374 Fraction exposed (FE), of total metal atoms definition, 72-73 effects on turnover frequency in hydrocarbon conversion hydrogenation, 110-115 hydrogenolysis, 118-123 higher than 0.5, in metal clusters, 142-146 lower than 0.5, sympathetic behavior and, 151-152

selectivity as function of, 88, 124-127

wide range, antipathetic behavior and, 146-151 Freeze-drying, perovskite preparation, 247, 249-251 G Gas chromatography (OC), in UHV surface analds, 9-13 OC, see Gas chromatography

Oibbs free energy, perovskite formation and, 262-263, 265 Oold, platinum activity changes by. 40-41

H n-Hcxane, conversion over platinum aromatization, 16-18 gold addition and, 40-41 particle sue and, 85 1-Haanol, hydrogenolysis over amorphous alloys, 370 I-Harene, hydrogenation over amorphous and crystalline alloys, 358, 359 High-temperature reduction (HTR), effects on titania-supportedmetals, 176177, 180 carbon monoxide hydrogenation, 196-200 chemisorption of carbon monoxide and hydrogen, 184-190, 224-225 hydracarbon hydrogenation, 190-196 titania migration onto metal, 203-206 Homomolecular exchange of oxygen, over perovskit-, 305-306 HTR, see High-temperature reduction Hydrocarbons, see o h specific compoundr adsorption on perovskites, 278 conversion over single crystals iridium and, 21-23 nickel and, 21-22 platinum and, 16-21 rhodium and, 20 silver and, 15-16 tungsten and, 23 hydrogenation metal particle size effects, 110-117 over perovskites, 293-295 hydrogen01ysis ensemble effect, 150 metal particle size effects, 117-128 over perovskites, 293-294 OVCT rhodium/titania, after HTR, 190-196 oxidation over perovskites, 283-289

391

INDEX Hydrocracking reactions, over perovskites, 311 Hydrogen adsorption on perovskites, TPD and, 270 chemisorption on iron clusters, rate constants. 144-145 as method of catalyst characterization, 91-95

on titania-supported catalysts after HTR. 184-186, 194 TPD assay, 209-211, 214-217 consumption by metals, TPR and, 102 effect on SMSI, 220-221 electrolytic evolution on amorphous metal alloys, 336-338 oxidation m d particle size and, 132-133 over perovskites, 306-307 Hydrogen chloride etching, iron-boron amorphous activities and, 351 Hydrogen fluoride, amorphous alloy treatment, 333-334 formaldehyde oxidation in fuel cells, 343-344 water electrolysis and, 336-338

I Impregnation metals on supports, 62-63 perovskites on supports, 251-253 titania-supported catalyst preparation, 181-183

Infrared (IR) spectroscopy perovskitc adsorption studies, 274-278 sulfur dioxide effects, 312-313 poison detection on metal surface, 107-108

support-metal interaction, 102-103 Ion exchange, in supported metal preparation, 63-66 competitive. 65-66 oxide surface charge in suspension and, 64 simple, 64-65 titania as support, 182-184 Ionization potential, metal clusters, 144 Iridium/alumina catalysts, hydrogen chemisorption, 94 Iridium catalysts, hydrocarbon conversion, 21-24

Iridiudtitania catalysts, preparation, 70 Irodalumina catalysts, SMSI, 47

Iron-boron amorphous alloys carbon monoxide hydrogenation, 349-350 formic acid decomposition, 374 Iron catalysts ammonia synthesis, 24 antipathetic behavior, 150, 152 particle size and, 131-132 promotion by potassium, 36-37 clusters, hydrogen chemisorption, 144-145 deposition on magnesium oxide, SMSI, 46 foils, potassium-promoted, 35-36 preoxidation effect, 32 Iron/magnesia catalysts, SMSI, 46 Ironltitania catalysts, HTR, 177 Isobutene, oxidation over perovskites, 287-288

L Lanthanum cobaltate catalysts carbon monoxide oxidation, kinetics, 281-283

exhaust gas purification, 292-293 hydrocarbon hydrogenation, 294-295 reduction-oxidation, 266-267 XPS Studies. 267-270 toluene oxidation, 288-289 Lanthanum manganite catalysts, oxidative nonstoichiometry, 254-256, 258 Lanthanum oxide, in cordierite-supported perovskites, 252-253 Lanthanum rhodate catalysts, carbon monoxide hydrogenation, 296-297 Fischer-Ttopsch reaction, temperature effect, 297-298 Lead, perovskite activities and, 313, 314 Line-broadening analysis, metal particle size distribution, 97-98 Low-temperature reduction (LTR), effects on titania-supported metals benzene dehydrogenation, 210 carbon monoxide dehydrogenation. 1%. 198 chemisorption of carbon monoxide and hydrogen, 185, 187-188, 190 hydrocarbon dehydrogenation, 191, 194 LTR, see Low-temperature reduction

M Metal clusters decomposition, small particle preparation, 66-67 definition, 141

392

INDEX

supported, FE > 0, 5 electronic effects, 145 morphology, 146 structure sensitivity, 146 unsupported, FE > 0, 5 chemical properties, 144-145 electronic properties, 143 morphology, 144 size distribution, 142-143 Metallic glasses, see Amorphous metal alloys Metal-support interaction, see also

spec@ catalysts apparent, 174 electronic effects, 174-175 HTR and, see High-temperature reduction real, 175 strong, see Strong metal-support interaction TPD and, see Temperature-programed desorption Methanation reaction over amorphous alloys, 352, 354. 357 antipatheticsmcture sensitivity, 128-131,139 over nickel catalysts, 26 phosphorus effect, 29 sulfur effect, 29-30 tungsten support effect, 43 SMSI and, 44-45 over tungsten catalysts, 26 Methane, oxidation over perovskites, 284-286

Methanol oxidation on amorphous alloy electrodes, 342-343

sensing by perovskites, 316 Methylcyclopentane aromatization over platinum, 17-18 hydrogenation, metal particle size and, 127 hydrogen01ysis metal particle size and, 121-122 over rhodium/titania. reduction temperature and, 193, 195 selectivity as function of FE, 124-127 Microreactor, in UHV analysis compact design, 8 GC analysis, 9-13 sample mounting, 4-8 small-volume design, 9 Model catalysts, see aLr0 specflc catalysts alloy catalysts, kinetics, 39-43 definition, 1-2

high-pressure-UHV methods, 2-4 metal-support interactions, 43-47 structural sensitivity, 15-28 ammonia synthesis, 24 carbon monoxide hydrogenation, 26-27 carbon monoxide oxidation, 24-26 hydrocarbon conversions, 16-24 surface modification by additives alkali promoters, 34-39 electronegative and electroneutral elements, 28-34 UHV surface analysis, apparatus designs, 4-14; see also Ultrahigh vacuum surface analysis Molybdenum catalysts, 27 poisoning by sulfur, 31 Massbauer spectroscopy, metal catalyst assay, 101. 106, 177

N Nickel/alumina catalysts, SMSI, 46 Nickel-boron amorphous alloys, hydrogenolysis catalytic activity, 371 Nickel catalysts cluster compound decomposition. 66 cyclohexane dehydrogenation, particle size and, 117 deposition on alumina, lack of SMSI, 46 ethane hydrogenolysis particle size effect, 119-120 turnover frequency maximum versus

FE, 152-153 hydrocarbon conversion, 21, 22 methanation reactions, 26 phosphorus addition and, 29 sulfur addition, 29-30 potassium effects, 34-36 propane hydrogenolysis, particle size and, 119-120

Raney-type, organic compound hydrogenation, 369-370 Nickel-copper catalysts, hydrocarbon hydrogenation. HTR and, 191, 193 Nickel-iron amorphous alloys, with phosphorus or boron, carbon monoxide hydrogenation, 346-347 Nickel-metalloid amorphous alloys, hydrogenation catalysis boron and phosphorus role comparison, 365-366

393

INDEX of I,J-butadiene, 365 of olefins. heating effect, 362-365 Nickel-phosphorus amorphous alloys, carbon monoxide hydrogenation, 348-349

Nickel/titania catalysts carbon monoxide hydrogenation, 129 copper effect, 1%-197 SMSI, 44, 1% chemisorption of carbon monoxide and hydrogen, 215-217 reduction temperature effects, 190, 1% selectivity, 176 surface composition, 206 Nickel-titanium amorphous alloys, acrolein hydrogenation, 367 Niobia-supported catalysts, 227-228 Nitric oxide adsorption on perovskites, 273-275 decomposition over amorphous metal alloys, 372 reaction with carbon monoxide over rhodium catalysts, 24-25 reduction over perovskites, 289-291 pollution control and, 291-292 Nitrobenzene, hydrogenation over h e y type nickel catalysts, 370 Nitrous oxide, decomposition over peravskites, 307-309 NMR, see Nuclear magnetic resonance Nuclear magnetic resonance (NMR), metal particle size distribution, 101 0

Octahedron, small supported particle model, 76-79 I-Octyne, hydrogenation over amorphous alloys. 368-369 4-Octyng hydrogenation over amorphous

alloys. 368-369 Oxygen adsorption on perovskites. TPD and, 270-272

cathodic reduction over perovskites, 300-301

electrolytic evolution on amorphous metal alloys. 339 on perovskites, 302-303 homomolecular exchange, over perovskites, 305-306

P Palladium-based amorphous alloys hydrogenation catalysts of (+)-apopinene, 358-362 of carbon monoxide, 354-357 of ck-cyclododecene, 358 of I-hexene, 358. 359 phosphorus-containing, methanol oxidation in fuel cells, 342-343 sodium chloride electrolysis and, 339-342 Palladium catalysts acetylene conversion to benzene, 23 hydrocarbon hydrogenation, particle size and. 111. 113, 114 methanol synthesis, 26-27 poisoning by potassium, 32 Palladiudtitania catalysts, chernisorption of carbon monoxide and hydrogen, 188-189

Pentane, hydrogenolysis over platinum/silica, 90 Perovskite-type oxides, see also s p @ c lanthanum-based caralysa actinide storage in radioactive waste, 315-316

adsorption studies, temperature effects, 270-279

ammonia oxidation, 307 deposition on supports impregnation, 251-253 plasma spraying. 253 electrocatalysis, 300-303 history, 237-239 homomolecular exchange of oxygen, 305-306

hydrocarbon conversion, 293-295 hydrocracking reactions, 311 hydrogen oxidation, 306-307 methanol sensing, 316 nitric oxide reduction, 289-291 nitrous oxide decomposition, 307-309 nonstoichiometry oxidative, 254-256 oxygen deficiency and, 257-258 reductive, 254, 257-258 oxidation catalysis of alcohols, 289 of carbon monoxide. 279-283 of hydrocarbons, 283-289 photocatalysis, 304

394

INDEX

poisoning by lead, 313-314 by sulfur dioxide, 312-315 pollution control and, 291-293 preparation methods, 244-254 liquid-solid reactions coprecipitation, 247-250 dry evaporation, 246-247 explosion, 250 freeze-drying, 247, 249-251 Spraydrying, 247, 249-250 solid-solid reactions, 245-246, 249, 251 2-propanol conversion dehydration, 311 dehydrogenation, 309-311 prospective research, 318-319 reduction isothermal, 260-2452 phase transformation, 265-266 reversible reduction-oxidation cycles, 265-267

temperature-programed, 258-261 StNCtUd Changes and, 262-265 XPS Studies, 267-270 structure cubic, 240 electroneutrality, 241 nonstoichiometry, 243, 254-258 orthorhombic distortion, 242 of related compounds, 243-244 sulfur dioxide reduction by carbon monoxide, 298-300 superconductivity at high temperature, 317-318

water-gas shift reaction, 311-312 Phenylacetylene, hydrogenation over amorphous alloys. 368-369 Phosphorus mctal catalyst poisoning, 29, 31 in perovskite.catalyst, 346-349, 251 Photocatalysis, over perovskites, 304 Plasma spraying. pe~wskitcson supports, 253 Platinum catalysts particle size, effect on hydrocarbon conversion activities hydrogenation and, 110-117 hydrogenolysis and, 119-120 skeletal reactions and, 122-123 singlecrystal surface carbon monoxide oxidation, 26 gold addition effect, 40-41

hydrocarbon conversion, 16-21 potassium effect, 37 sulfur poisoning, 30 supported, hydrogen chcmisorption, 94 Platinum/ceric oxide catalysts, SMSI,226 Platinumlsilica catalysts, pentane hydrogenolysis, particle size and, 90 Platinumltitania catalysts bonding and charge transfer, 224-225 HTR, 176-177 potassium effects, 200-201 silver effects, 202-203

SMSI,44-45 comparisonwith rhodium/titania, 20!3-214 surface composition, 206-208 Poisons, detection on metal surface EPR, 108 IR spectroscopy, 107-108 secondary-ion mass spectrometry, 107 TPD technique, 106-107 Pollution control, with perovskite catalysts, 291-293

Potassium, effect on catalysts iron, 35-37 nickel, 34-36 palladium, 32 platinum, 37 rhenium, 37 titania-supported, 200-201 Praseodymium cobaltate, reduction, temperature effect, 260, 261 Propane hydrogenolysis, metal particle size and, 120 oxidation over perovskites, 286 2-Propanol, conversion over perovskites dehydration, 311 dehydrogenation, 309-31 1 Propene, oxidation over perovskites, 287-288 Propylme, oxidation metal particle size and, 136 over perovskites, 287-288 Pyridine, adsorption on perovskites, 277

R Ranw-tw nickel catalysts organic compound hydrogenation, 369-370

Reduction of catalyst precursors, methods, 101-106 isothermal, of perovskites, 260-267

395

INDEX metal-support particle preparation and, 70 temperature-programed, see 'Rmperatureprogramed reduction Rhenium catalysts, ammonia synthesis, 24-25 promotion by potassium, 37 Rhodium/alumina catalysts, carbon monoxide chemisorption, 85 Rhodium catalysts carbon monoxide conversion hydrogenation, 26 pnoxidation and. 32 sulfur addition and, 30 oxidation, 24-26 ethane hydrogenolysin, particle size and, 118 ethylidine hydrogenation, 20, 23-24 supported, hydrogen chemisorption, 94-95 Rhodium/niobia catalysts, SMSI, 227-228 Rhodium/silica catalysts ethane hydrogenolysis, 153 titania-promoted activity, 200 Rhodiumltitania catalysts bonding and charge transfer, 221-223 carbon monoxide hydrogenation. 217-218 potassium effects, 201 preparation. 181-184 in UHV apparatus, 203 reduction temperature effects carbon monoxide hydrogenation, 196-200 chemisorption of carbon monoxide and hydrogen, 184-187 EXAFS analysis, 221,223 hydrocarbon hydrogenolysis, 190-196 XANES study, 222-223 silver effects, 202 SMSI, 45-46 comparison with platinum/titania, 209-214

surface composition, 204-208 Rhodium/vanadia catalysts, SMSI, 226-227 Ruthenium catalysts carbon monoxide oxidation, 26 copper addition and, 41-43 sulfur effects, 30 supported hydrogen chemisorption. 93 structure sensitivity, 87 S

Schulz-Flory plots, lanthanum rhodate catalytic activity, 297-298

Secondary-ion mass spectroscopy, poison detection on metal surface, 107 Selectivity alkyne hydrogenation over palladium alloy catalysts, 368-369 ethylene oxidation over supported silver. 90, 136-137

methylcyclopropane hydrogenolysis over metals, 124-127 structure sensitivity and, 88-90 Silicon. metal catalyst activity and, 31 Silver catalysts effect on titania-supported catalysts, 202-203

ethylene oxidation poisoning by chlorine, 32-34 promotion by cesium, 37-39 singlecrystal surface orientation and, 15-16

over supported catalysts, particle sue and, 90, 136-138 Singlacrystal surface catalysts, see Model catalysts Small-angle X-ray scattering, metal particle size distribution, 97-98 SMSI, see Strong metal-support interaction Sodium chloride, electrolysis by amorphous alloys, 339-342 Spray-drying, perovskite preparation, 247. 249-250

Static magnetic methods, metal particle size distribution, 98-99 Stepped single crystals, in supported metal assay, 82-83 Strong metal-support interaction (SMSI) in alumina-supported catalysts, 228 carbon monoxide hydrogenation and, 196-200

chemisorption of carbon monoxide and hydrogen, 184, 187-188, 190 chemisorption loss after HTR and, 228 comparison between rhodium/titania and platinum/titania, 209-210, 213-214 hydrocarbon hydrogenolysis and, 190-196 hydrogen effect, 220-221 in model catalysts, 43-47 in niobia-supported catalysts, 227-228 in platinum/ceric oxide catalysts, 226 vanadia-supported catalysts, 226-227 Structure insensitivity ensemble effects and, 153-154

INDEX experimental examples, 141-142 overlayer effects, 154 Structure sensitivity antipathetic behavior, 147, 149, 150, 152, 155-156

of carbon monoxide chemisorption, 138-140

of carbon monoxide oxidation, 134-135 definition, 57 of hydrogen chemisorption. 91-95 limited and extended, 88 in metal clusters, supported, 146 secondary, 128, 149, 155 selectivity and, 88-90 single-crystal surface orientation effect in platinum catalysts, 16-21 in silver catalysts, 15-16 sympathetic behavior, 151-152, 155-156 TOF assay, see 'hnover frequency Sulfur metal catalyst poisoning, 29-30 in oxide-supported catalysts, chemisorption loss after HTR and, 228 Sulfur dioxide effects on perovskites adsorption, IR study, 312-313 carbon monoxide oxidation poisoning, 313-315

reduction by carbon monoxide over perovskites, 298-300 Superconductivity,yttrium role in perovskites, 317-3 18 Supported metals, small particles. see o h specflc catalysts activation by calcination and reduction. 70-71

alloying effects, 80-82 AR, definition, 72-73; see also Atomic rate of reaction bonding with support, 84 characteristics, 157 (table) crystal structure change, 85 electronic properties, 74-76. 109 FE. definition, 72-73; see ulso Fraction exposed, of total metal atoms future studies, 157-160 mathematical models, 76-81 model catalysts and, 83-84 particle size catalytic activity and. 56-59 critical, calculation, 55-56

distribution chemisorption and. 91-95 electron microscopy, 95-96 magnetic methods, 98-100 X-ray diffraction, 97-98 poisons, detection methods, 106-108 preparation methods, 61 (table) carrier selection, 59 chemical deposition, from metal colloid dispersion, 67-78 coprecipitation, 60, 62 decomposition of metal cluster compounds, 66-67 impregnation. 62-63 ion exchange, 63-66 ion implantation, 68 vapor phase deposition, 68-69, 83 reconstruction, 85-86 reduced ions, detection methods, 102-106 stepped single crystals. 82-83 structure, methods of assay, 108-109 structure sensitivity, 86-91 TOF, definition, 57-58. 72-74, 154-155; see ulso 'hrnover frequency

T Temperature-programeddesorption (TPD) benzene hydrogenation on titaniasupported metals, 210 carbon monoxide chemisorption on platinum/titania, 187-189, 224-225 comparison with rhodiudtitania, 209-212

titania coverage and, 208-209 carbon monoxide from nickel surface, 215-217

hydrogen chemisorption on titaniasupported metals. 209-211, 213-214 poison detection on metal surface, 106-107

Temperature-programedreduction (TPR) hydrogen consumption by metals and, 102 of perovskites, 258-261 comparison with isothermal reduction, 260-262

oxygen adsorption, 270-272 structural changes during, 262-265 Thermal desorption mass spectroscopy, 1, 14 Timeof-flight mass spectrometer, metal cluster assay, 142, 143

397

INDEX Titania comparative effects on platinum and rhodium activities, 209-214 coverage on metal, carbon monoxide chemisorption and, 206-209 migration onto metal during HTR, 189, 192, 1%. 201-207 nickel deposition on, SMSI. 44 platinum deposition on, SMSI,44-45 promoting effect on catalytic activity explanations, 214-221 on rhodium/silica, 200 rhodium deposition on, SMSI,45-46 Titania-supported catalysts, see also spec@ catalysts

bonding and charge transfer, 221-225 HTR. see High-temperature reduction hydrogen effect, 220-221 metal-support interactions, 173, 175-179 preparation impregnation, 181-183 ion exchange, 182-184 reduction temperature effects carbon monoxide chemisorption, 186-190 carbon monoxide hydrogenation. 196-200 hydrocarbon hydrogenation, 190-196 hydrogen chemisorption, 184-186 silver effect, 202-203 "OF, see 'hrnover frequency Toluene hydrogenation over Raney-type nickel catalysts, 370 oxidation to benzaldehyde over lanthanum cobdtate, 288-289 TPD, see 'Rmperature-programed desorption TPR, see 'Rmperature-programed reduction 'hngsten catalysts carbon monoxide hydrogenation, 26 sulfur effect, 30 ethane hydrogenolysis, 23 nickel submonolayer fiim effect, 43 F) lbrnover frequency O ammonia synthesis, 131 antipathetic behavior and, 147, 149-150, 152, 156 carbon monoxide hydrogenation, 128-130 definition, 57-58, 72-74, 84-87, 154-155

hydrocarbon hydrogenation, 110-117 hydrocarbon hydrogenolysis, 118-123 maximum versus FE, 151-153, 155-156 oxidation reactions, 132-138 structure sensitivity and, 86-87, 90-92, 94 sympathetic behavior and, 151-153, 156 unchanged with FE, 141-142

U Ultrahigh vacuum surface analysis (UHV) apparatus, 9-14 high pressure cell and, 9-11, 13 microreactor, 4-14 presssure cup design, 10-12 rhodium/titania preparation, 203-204 transfer rod design, 12-13 future assays, 48-49 -high pressure transfer, 2, 3 cleanliness requirements, 6 impurity problems, 5 homogenous surface preparation, 47-48 metal vapor deposition, 69

V Vanadia-supported catalysts, 226-227

W Water adsorption on perovskites, 276-278 electrolysis using amorphous metal alloys, 336-338. 375 photolysis over perovskites. 304 Water-gas shift reaction over copper, 27 poisoning by sulfur, 31 promotion by cesium. 37-38 over perovskites, 311-312 Wolfram, see Tungsten X XANES. see X-ray absorption near-edge

structure XPS, see X-ray photoelectron spectroscopy X-ray absorption near-edge structure (XANES), metal-support interaction. 174 in rhodium/titania catalysts, 221-223 X-ray diffraction (XRD) amorphous metal alloys. 332 during ammonia synthesis, 373-374

398

INDEX

during carbon monoxide hydrogenation nickel-iron alloys, 346 zirconium-containing alloys, 352-354,

UHV analysis and, 1.4 XRD, see X-ray diffraction

356-357

metal particle size distribution, 97-98 perovskite reduction-oxidation, 265-267 X-ray photoelectron spectroscopy (XPS) amorphous metal alloys carbon monoxide hydrogenation and,

Y Yttrium, in perovskites. superconductivity and. 317-318

347, 348, 351

hydrogen fluoride effect. 337-338 during sodium chloride evolution. 340 hydrogen role in SMSI,220-221 metal-support interaction, 174 in titania-supported catalysts, 203, 206,208

perovskite studies oxygen adsorption, 273 reduction, 267-270 platinum spectra. as function of FE,110-111 small supported particle assay, 74-75, 83 particle size distribution, 101, 104-105, 109

Z

Zinc oxide, copper deposition on, SMSI, 46-47

Zirconium-containing amorphous metal alloys ammonia synthesis, 372-373 hydrogenation catalysis of 1,3-butadiene, 367 of carbon monoxide, 333, 352-357. 375 of olefins, 366-367 hydrogen evolution, 338 nickel-based, carbon monoxide chemisorption, 344-345

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