Advances in Catalysis, Volume 23

ADVANCES I N CATALYSIS VOLUME 23 Advisory Board G. K. BORESKOV Novosibirsk, U.S.S.R. P. H. EMMETT Baltimore, Maryland...

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ADVANCES I N CATALYSIS VOLUME 23

Advisory Board G. K. BORESKOV Novosibirsk, U.S.S.R.

P. H. EMMETT Baltimore, Maryland

M. BOUDART Stanford, California

J. HORIUTI Sapporo, J a p a n

G. NATTA

E. K. RIDEAL

Milan, Ztaly

London, England

H. S. TAYLOR Princeton, New Jersey

M. CALVIN Berkeley, California

W. JOST Gottingen, Germany

P. W. SELWOOD Santa Barbara, California

ADVANCES IN CATALYSIS VOLUME 23

Edited by D. D. ELEY The University iV ottingham, England

HERMAN PINES Northwestern University Euanston, Illinois

PAULB. WEISZ Mobil Research and Development Corporation Princeton, New Jersey

1973

ACADEMIC PRESS

NEW YORK AND LONDON

COPYRIGHT 0 1973, BY ACADEMIC PRESS, INC. ALL RIGHTS RESERVED. N O 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, WI T H O U T PERMISSION IN WRITING FROM THE PUBLISHER.

ACADEMIC PRESS, INC.

111 Fifth Avenue, New York, New York 10003

Unired Kingdom Edifion published by ACADEMIC PRESS, INC. (LONDON) LTD. 24/28 Oval Road. London N W l

LIBRARY OF CONGRESS CATALOG CARDNUMBER:49-7755

PRINTED IN TH E UNITED STATES O F AMERICA

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

vii ix

Metal Catalyzed Skeletal Reactions of Hydrocarbons J. R. ANDERSON

I. 11. 111. IV. V. VI.

Introduction.. . . . . Catalyst Structure

....,........... 1 .......................... ......... 2 16 ....................... Isomerization and Dehydrocyclization 25 Hydrogenolysis on Metals. . . . . . , . . . . . . . . . . . . . . . _ . . . . . . . . . . . _ . . _ .62 Reactions over Chromium Oxide Catal 81 References. . . , , . . . . . , . . . . , . , . , , , , . . . . . . . . . . . . . . . . . . . . . . . . . 84

Specificity in Catalytic Hydrogenolysis by Metals J. H. SINFELT

I. 11. 111. IV. V.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91 General Discussion on Hydrogenolysis Reactions. . . . . . . . . . . . . . . . . . . . . . . 92 Comparison of Metals as Hydrogenolysis Catalysts. . . . . . . . . . . . . . . . . . . . . 97 Contrast between Ethane Hydrogenolysis and Other Reactions, . . . . . . . . . 106 116 Conclusion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 116 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

The Chemisorption of Benzene R. B. MOYESAND P. B. WELLS

I. 11. 111. IV. V.

Introduction, . . . ....................... Chemisorption , , . . , , , , . , . , , , , , , , , . , , , . , , , . . . . . . . . . . . . . . . . . . . . . . . . . .

121 122 133 . . . . . . . . . . . . . . . . . 148 Some Aspects of Benzene Hydrogenation. . , . . 152 ......._........... Conclusions. . . . . . . . . . . . . . . . . . . . . . 154 References. . .

The Electronic Theory of Photocatalytic Reactions on Semiconductors TH.WOLKENSTEIN Introduction, . , . . . . . . . , . , . , , , , , . , , . , . , . , . . , . . . . , . , . . . . . . . . . . . . . . . . . 157 I. The Mechanism of the Influence of Illumination on the Adsorption and Catalytic Properties of a Surface. . . . . . . . , . . . . . . . . . . . . . . . . . . . . . . . . . . . . 158 V

vi I1. 111. IV . V. VI .

CONTENTS

The Photoadsorptive Effect . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . The Reaction of Hydrogen-Deuterium Exchange . . . . . . . . . . . . . . . . . . . . . . . The Reaction of Oxidation of Carbon Monoxide . . . . . . . . . . . . . . . . . . . . . . . The Reaction of Synthesis of Hydrogen Peroxide . . . . . . . . . . . . . . . . . . . . . . . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

170 179 189 197 203 206

Cycloamyloses as Catalysts DAVIDW . GRIFFITHSAND MYRONL. BENDER

I. I1. I11. IV . V. VI .

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Physical Properties of the Cycloamyloses . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Reactions in Which the Cycloamyloses Participate Covalently . . . . . . . . . . . Noncovalent Catalysis by the Cycloamyloses . . . . . . . . . Catalytic Properties of Modified Cycloamyloses . . . . . . . . . . . . . . . . . . . . . . . . Concluding Remarks . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

209 210 222 242 249 258 259

Pi and Sigma Transition Metal Carbon Compounds as Catalysts for the Polymerization of Vinyl Monomers and Olefins D . G . H . BALLARD

I . Introduction., . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 263 I1. Soluble Transition Metal Alkyl Compounds as Polymerization Catalysts . . 266 111. Ligand Replacement in Transition Metal Alkyl Compounds and Polymeri...................................

288

erization Catalysts Derived from Tran Alkyl Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . V. Stereoregular Polymerization with Transition Metal Alkyls . . . . . . . . . . . . . . VI . Mechanism of Polymerization. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . VII . Conclusion ...................................... ... References .................................................

293 298 304 323 324

327 AUTHORINDEX ............................... SUBJECTINDEX ............................... . . . . . . . . . . . . 337 347 CONTENTS OF PREVIOUS VOLUMES ..........................................

Contri b u to rs Numbers in parentheses indicate the pages on which the authors’ contributions begin.

J. R. ANDERSON,CSIRO Division of Tribophysics, University of Melbourne, Parkville, Australia (1)

D. G. H. BALLARD,Imperial Chemical Industries Limited, Corporate Laboratory, The Heath, Runcorn, Cheshire, England (263) MYRONL. BENDER,Department of Chemistry, Northwestern University, Evanston, Illinois (209) DAVIDW. GRIFFITHS,Department of Chemistry, Northwestern University, Evanston, Illinois (209) R. B. MOYES,Department of Chemistry, The University, Hull, England (121) J. H. SINFELT,Corporate Research Laboratories, Esso Research and Engineering Co., Linden, New Jersey (91) P. B. WELLS,Department of Chemistry, The University, Hull, England (121) TH. WOLKENSTEIN, Institute of Physical Chemistry, Academy of Sciences, Moscow, U S S R (157)

vii

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Preface Three months after the Fifth International Congress on Catalysis I still find my mind turning back to V. Haensel’s introductory remarks in which he pointed out the gap between the industrial art and the academic science of catalysis. At the present time, enzymologists are pulling ahead, and more real knowledge is available about the structures and mechanisms of macromolecules such as lysozyme, ribonuclease, hemoglobin than of any industrial catalysts. It is true that while chemical kinetics was the only weapon the enzyme field moved slowly and that the advance started with computcr-aided crystallographic studies (Perutz and Kendrew a t Cambridge, England). It is also true that the equivalent structural techniques for surfaces are only just emerging, i.e., Auger, ESCA, LEED, etc., and that a t present new techniques tend to be concentrated on single crystal surfaces. However, this is a hopeful sign as is the way in which the mechanisms of organometallic homogeneous catalysts are being decided and their possible relevance for heterogeneous catalysis investigated. So turning to our present volume, we refer first to the last article in the volume by D . G. H. Ballard. He concludes that our knowledge of catalytic mechanisms is limited “because the majority of useful catalysts for practical reasons are heterogeneous and therefore unsuitable for mechanistic studies.” Ballard’s article well illustrates the fact that where all the techniques are available to establish structure (as they are in homogeneous organometallics), kinetic studies take on a new depth and progress is rapid. The articles by J. R. Anderson, J. H. Sinfelt, and R. B. Moyes and P. B. Wells, on the other hand, deal with a classical field, namely hydrocarbons on metals. The pattern of modern work here still very much reflects the important role in the academic studies of deuterium exchange reactions and the mechanisms advanced by pioneers like Horiuti and Polanyi, the Farkas brothers, Rideal, Twigg, H. S. Taylor, and Turkevich. Using this method, Anderson takes ultrathin metal films with their separated crystallites as idealized models for supported metal catalysts. Sinfelt is concerned with hydrogenolysis on supported metals and relates the activity to the percentage d character of the metallic bond. Moyes and Wells deal with the modes of chemisorption of benzene, drawing on the results of physical techniques and the ideas of the organometallic chemists in their discussions. Th. Wolkenstein’s article provides mechanisms for certain light-accelerated catalytic reactions on solids. This is a field where very explicit ix

X

PREFACE

models may be constructed in relation to the band theory of semiconductors and where detailed mathematical treatments have been made and compared with experiment. Finally, we come to enzyme models. D. W. Griffiths and M. L. Bender describe the remarkable catalytic property of certain cycloamyloses which act through formation of inclusion complexes, and in this respect recall the clefts containing the active sites in enzymes such as lysozyme and papain. I believe these articles show that there is a general move forward over a wide field in catalysis and that in the future we may reasonably hope that the academic-industrial gap in the heterogeneous field will start to close.

D. D. ELEY

Metal Catalyzed Skeletal Reactions of Hydrocarbons J. R. ANDERSON CSIRO Division of Tribophysics University of Melbourne Parkville, Australia

I. Introduction. .................................. 11. Catalyst Structure. ............................. A. Evaporated Metal Films.. ....................... B. Supported Catalysts. ........................ C. Gas Adsorption Behavior. . . . . . . ........ 111. Experimental Techniques. ............................... A. Evaporated Metal Film ........................... B. Supported Catalysts.. ............... C. Use of Hydrocarbons rbon . . . . . . . . . . . . . . . IV. Isomerisation and Dehydrocyclisation Reactions on Metals. . . . . . . . . A. Reactions on Platinum. .....................................

V. Hydrogenolysis on Metals.. VI. Reactions over Chromium References. . . .

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

14 16 20 25 26 62

1. Introduction It is the purpose of this article to review the mechanisms of reactions undergone by the carbon skeletons of aliphatic and alicyclic hydrocarbons in the presence of metallic catalysts. As well as skeletal isomerization reactions, we shall deal with cyclization, ring opening, and hydrogenolysis reactions, while recognizing that this last group necessarily also requires the net removal or addition of hydrogen atoms. Indeed, although our central concern is with the fate of the carbon skeleton, it is impossible to avoid a detailed consideration of the state of hydrogenation of the reacting species, particularly when discussing reaction pathways in terms of likely intermediates, and when discussing reactions which result in the formation of aromatic cyclization products. Catalytic reactions of this type have a long history, and a wide range of catalyst types have been used. However, we do not intend this review to be 1

2

J.

R. ANDERSON

exhaustive for all catalyst types. Our purpose is to emphasize those hydrobarbon reactions that are laregely confined to the metal. We do not propose to review catalysts for which dual-function activity is important. Such reviews are already available elsewhere (e.g., 1-6). Technical operating catalysts all consist, of course, of metal dispersed on a support. However, mechanistic studies have used both unsupported and supported metals, and we therefore will discuss in some detail the structure of various unsupported and supported catalysts upon which these reactions have been studied.

II. Catalyst Structure A. EVAPORATED METALFILMS Evaporated films are, of course, not “practical” catalysts. Their use as model catalysts is however justified by the insight which such work may give toward an understanding of catalytic reaction mechanisms. An initial state of high surface purity may be achieved with evaporated films using relatively straightforward techniques, and it is the elimination of initial surface contamination as a significant experimental variable which makes evaporated films desirable as model catalysts compared to bulk supported catalysts.

1. Continuous Films

It is often found that the ratio R (measured, for instance, by gas adsorption methods) of actual metal surface area accessible to the gas phase, to the geometric film area, exceeds unity. This arises from nonplanarity of the outermost film surface both on an atomic and a more macroscopic scale, and from porosity of the film due to gaps between the crystals. These gags are typically up to about 20 A wide. However, for film thicknesses >500 A, this gap structure is never such as completely to isolate metal crystals one from the other, and almost all of the substrate is, in fact, covered by metal. In practice, catalytic work mostly uses thick films in the thickness range 500-2000 A, and it is easily shown (7) that intercrystal gaps in these films will not influence catalytic reaction kinetics provided the half-life of the reaction exceeds about 10-20 sec, which will usually be the case. It is difficult to assess with high precision the crystal planes exposed to the gas phase in low temperature ( O O C ) polycrystalline films. The assumption has sometimes been made [e.g., Brennan, Haywood, and Trapnell (8)],that for fcc metals the surface consists of an equal exposure of ( l l l ) , (loo), and (110) planes, with a similar assumption for bcc metals with

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

3

regard to (1lo), (loo), and (211) planes. However, for low-temperature polycrystalline transition metal films in the thickness range 500-2000 11, high index planes are undoubtedly present to an appreciable extent, and this is the more serious the more refractory the metal. This conclusion is clearly implicit from photoelectric work-function data (9-16) and gas adsorption data (7, 16, 17) for various evaporated films. Nevertheless, it must also be said that in low-temperature films grown to great thickness and thus consisting of very wide crystals, low index planes are probably dominant. For instance, in a polycrystalline nickel film deposited on glass at 20°C to a thickness of 3.1 pm and with an average crystal width in the region of 0.2 pm, surface replicas clearly show regular faceted crystal shapes with dominant (111) and (100) surface planes (It?), as expected from both thermodynamic (19) and kinetic arguments (9). Polycrystalline films deposited on amorphous substrates are of lower crystallographic surface heterogeneity the higher the temperature of annealing subsequent taodeposition or the higher the substrate temperature during deposition; again, photoelectric work-function data serve to empha size the point (9, 10,12,IS). The question arises of the extent to which, in polycrystalline films reactant gas has access to the substrate. It is clear that in high-temperature films the total absence of intercrystal gaps means that such access of gas is completely absent. In the case of films deposited a t O'C, one may estimate from the measured roughness factor and from transmission electron microscopic evidence that, of the total substrate area, more than 90% is in direct contact with metal; in any case, the substrate at the base of a gap is almost certainly covered with a thin layer of metal. Thus, even in this case the gas cannot have more than trivial access to the substrate. Deposition on glass or other amorphous substrate at higher temperatures may result in some degree of preferred crystal orientation (20,21). The tendency toward preferred orientation tends to be greater at larger film thicknesses, although it can undoubtedly occur in the initial stages of film growth (22).In general, however, the occurrence and extent of preferred orientation on glass is of poor reproducibility, and when preferred crystal orientation is deliberately required, glass is not the best choice as a substrate. Film deposition on a single crystal substrate can, in principle, lead to the formation of an epitaxed single crystal film. However, relatively limited use has been made of well-epitaxed single crystal films for catalysts for two reasons: In many instances the single crystal substrate is only available with very limited dimensions so that the film catalyst is also correspondingly restricted in its area; second, in most cases it is either inconvenient or impossible to design the single crystal substrate and the evaporation source

4

J. R. ANDERSON

so that evaporated metal falls only on the single crystal substrate. Failure to achieve the latter means that if the reaction is to be confined to the epitaxed film, some method has to be found for transferring the epitaxed specimen from the preparation chamber to reaction vessel without breaking the vacuum. One way of achieving this is with a UHV-compatible winch. Only two crystalline substrates have had appreciable use for the prepttration of the metal film catalysts. These are mica and rocksalt. Mica is a convenient substrate for film growth. A cleaved mica surface is extremely flat and it therefore obviates one uncertainty inherent in the use of glass: Although a freshly fire-polished glass surface has a high degree of smoothness, it is subject t o corrosion in aqueous media, particularly if acidic. Decoration of a cleaved mica surface shows the presence of only an extremely low concentration of surface imperfections, and the surface is mainly featureless. Such imperfections as do occur are so relatively infrequent as to be of negligible effect on the degree of surface perfection of a thick metal film. Mica has the added desirable property of being flexible in thin sheets, so that it is not difficult to arrange a cylindrical substrate geometry so that most, if not all, the evaporated metal is deposited on the mica. With some metals it is possible to obtain a high degree of single crystal epitaxy on mica [e.g., silver ( 2 3 ) ] . However, single crystal film catalysts have not been prepared on mica from transition metals such as those we are mainly concerned with in this chapter, no doubt because of temperature limitations imposed by glass apparatus. With these fcc metals, deposition on mica at 35O"40O0C in HV or UHV leads to polycrystalline deposits in which each crystal is oriented with a (111 ) axis normal to the substrate, but with the crystals oriented with rotational disorder about this axis (cf. 2.4). In some cases, rotational disorder is not completely random. Complete preferred crystal orientation is not obtained with total reproducibility. The microscopic results show that the exposed surface of each crystal is overall relatively flat, so that the whole exposed film surface must be close to (111). Presumably, some higher order planes are exposed to a relatively small extent in the immediate vicinity of the grain boundaries; nevertheless, the proportion of (111) surface exposed is estimated to be not less than 90% and this is very similar t o the estimate for the proportion of (111) surface exposed in completely epitaxed silver films ( 2 3 ) . This estimate is also in agreement with that obtained from a patch-model analysis of photoelectric work-function data for nickel films deposited on mica a t 320°C (12). Although the surface of such a high-temperature film may appear relatively flat and featureless to shadowed replication, decoration shows clearly


5

that the surface is not completely smooth on a quasi-atomic scale, due to the presence of surface steps (cf. Fig. Id, p. 4 of ref. 7 ) ,and this would be the more important the more refractory the metal under similar temperature conditions of preparation. When deposited on mica at O"C, surface replicas show surface roughness comparable to that of films on a glass substrate, and the degree of preferred crystal orientation is also usually negligible. A wide range of metals has been grown epitaxially on a (100) rocksalt face, including (fcc) gold, silver, aluminum, nickel, copper; and (bcc) chromium, iron ( 2 3 , 2 5 ) .All can readily given an orientation of (100) metal planes parallel to rocksalt (loo), but gold, silver, copper, and aluminum can also give (111) metal planes parallel to rocksalt (100) depending on the conditions during metal deposition and during rocksalt cleavage, and this also affects the quality of epitaxy. As shown by decoration, a rocksalt cleavage face is far from absolutely smooth. Because of the convenience of using, for catalytic purposes, a film deposited on a relatively large area of substrate, a technique has been developed (24) for producing an evaporated layer of rocksalt as a substrate for subsequent film deposition. The evaporated rocksalt layer is of course microcrystalline, but consists largely of crystals exposing (100) faces upon which films of metals may then be deposited at elevated temperatures. By decoration, the presence of growth steps on the (100) surfaces is clearly revealed. To avoid problems due to sintering, thermal etching, and incipient evaporation of the rocksalt layer, and to maintain adequate vacuum conditions for surface cleanliness of the metal film, the substrate temperature is limited to about 250°C during metal deposition or subsequent annealing. Although a substrate temperature of 250°C will produce reasonably well epitaxed single crystal films of silver, with metals of higher melting point and greater cohesive energy, epitaxy is much more difficult (26, 2 4 ) . With this type of film, the exposed surface is far from perfect, due both to the microcrystallinity of the rocksalt substrate and to the imperfections in the surface of the epitaxed film. For a nickel film so prepared, the proportion of (100) surface exposed, as judged from rare gas adsorption data (11) is no more than 70% and is probably rather less than this with platinum. 2. Ultrathin Films

For the prescnt purpose, we take the term "ultrathin" to refer to an evaporated metal film where the concentration of metal on the substrate is low enough for the film to consist of small isolated metal crystals. If the average concentration of metal atoms on the substrate is of the order of a monolayer or less, the metal crystals are small enough for ultrathin films to serve as models for highly dispersed metal catalysts, but where surface cleanliness and catalyst structure can be better controlled.

6

J. R. ANDERSON

The degree of dispersion, i.e., the average crystallite size of the metal in supported catalysts, is important not only in controlling the surface area per unit weight of metal, but there is also the question of whether the nature of the catalytic process is dependent on metal crystallite size. Most interest here centers on extremely small crystallites, for instance 1 Torr) and temperatures ( >20°C) and, in practice, these are the conditions usually used to measure the total surface areas of metallic nickel in supported catalysts (e.g., 61-63) . The effect of other surface impurities may be more severe than that of oxygen. For instance, adsorbed sulfur strongly inhibits hydrogen adsorption on nickel ( 5 8 ) ,while chlorine adsorbed on nickel is also likely to be a tenaciously held surface contaminant. The general comments made above concerning the character of hydrogen adsorption in relation to surface cleanliness apply to platinum. Again, and for similar reasons to those given for nickel, total hydrogen uptakes on supported platinum are generally measured a t >1 Torr and >20"C, often about 100 Torr and 200°C (e.g., 64-66). Under these conditions adsorbed oxygen reacts with hydrogen according to (65, 67). (@)

Pt-O(,)

+ N Hz

+

Pt-Hw

+ HzO

( 1)

the water being taken up by the support. On clean platinum surfaces (e.g., for a surface evaporated films), the surface stoichiometry H(,,/Pt saturated with adsorbed hydrogen is unity a t 0"-20°C (68,69) ; at - 195°C the ratio rises somewhat due to further adsorption of more weakly bound molecular hydrogen ( 7 0 ) . At >20"C and > 1 Tow, one would expect a supported platinum specimen saturated with hydrogen to have a H ( e ) /

16

J. R . ANDERSON

Pt ( 8 ) ratio of about unity, and a recent critical survey of the situation by Wilson and Hall (48) indicates that this is, in fact, the case. Reaction (1) is the basis of the technique for the titration of chemisorbed oxygen by hydrogen (67). Hydrogen chemisorption a t 20°C has been used with ultrathin platinum films (29,SO). I n general it is found that the hydrogen uptake is rather larger (by about 20%) than can be reasonably accounted for by the particle size as determined by the electron microscope. It is found that on this type of surface, hydrogen adsorption isotherms are very similar in character to those observed with known clean platinum (e.g., thick films), thus confirming their surface cleanliness. Pure silica appears to be inert to hydrogen, but a t >1 Torr adsorption may occur on alumina and carbon and various workers have reported proportions of the total hydrogen uptake attributable to the support in the following ranges : alumina O-25%, and carbon 50% or more. I n addition to adsorption occurring directly onto the support, there is also the possibility that when metal is present, adsorption on the support may be augmented by the transfer of dissociatively chemisorbed hydrogen from metal to the support. This possibility was envisaged by Spenadel and Boudart (64) who concluded however, that it was unimportant, a t least with platinum/silica catalysts. However, there is a substantial body of evidence, based on hydrogen chemisorption studies (71-74) as well as on catalytic reactions over supported metals (for a summary of references see Sancier (75) to show that this sort of hydrogen transfer can occur with silica, alumina, and carbon supports, particularly a t highish temperatures ( >300"C). The effect is particularly severe with carbon. It also seems likely to occur with ultrathin platinum films. With highly dispersed platinum catalysts this behavior is not unexpected in view of the activation energy for surface mobility of H on platinum of 4.5 kcal mole-' (69), so that a t 20°C the migration of H through a distance of (say) 50 could occur very quickly ( isobutane > n-butane, and this agrees with experimental results. The type of intermediate shown in structure (B) has also been supported by Muller and Gault (119) who showed that in the reaction of 1,l-dimethylcyclopropane with deuterium over a series of thick evaporated metal film catalysts, it was only on platinum that 1 , l ,3-da-neopentane (and 1,1,3,3-d*-neopentane) were dominant products. On palladium, iron, rhodium, nickel, and cobalt the major product was 1,5-dz-neopentane. Anderson and Avery's bond shift mechanism has the consequence of predicting that a quaternary carbon atom cannot be generated in the hydrocarbon product. In fact, Anderson and Avery (24) showed that in the isomerization of isopentane over platinum films, only a very small amount ( < 1%) of ncopentane was produced (although the equilibrium constant for isopentane neopentane is 0.16 a t 27SOC). Furthermore,

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBONS

35

in a large number of experiments with n-hexane, 2-, and 3-methylpentane over various types of platinum film catalysts (28, 29),neohexane has never been observed in the reaction products. On the other hand, other workers have found small but significant amounts of neohexane in the isomerization products from the hexanes; for instance, about 0.1% from 2-methylpentane ( 8 4 ) , and nf-R-/L Pt

Pt

Pt

Pt

Pt

(9)

The small extent to which this reaction apparently occurs is in agreement with the high energy of the C1ring. Reaction (9) is closely related to (13) which, as we shall see, is a possible route for dehydrocyclization to form Cg ring compounds. The importance of this reaction path appears to be highly variable and is presumably dependent on the nature of the catalyst surface. It is of interest that no neohexane was observed from reactions of the hexanes over platinum/alumina catalysts (0.2 and 10% platinum) ( 8 4 ) . Since Anderson, Macdonald and Shimoyama (28, 29) failed to observe any neohexane on UHV platinum film catalysts ranging from highly sintered, thick films to ultrathin films consisting of discrete, very small metal crystals, it seems likely that the feature required to generate neohexane may well not be the structure of the platinum surface per se, but may be an impurity-generated site. The relative rate of isobutane isomerization has been shown by Anderson and Avery (24) to be markedly increased by using a (111) platinum film surface. On the other hand, this did not occur with n-butane, nor did it occur with either iso- or n-butane over a (100) platinum surface (cf. Table 11).A triangular array of adjacent sites on a (111) platinum surface can be readily fitted by an adsorbed isohydrocarbon, and this structure also fits to allow the carbon orbitals to be directed normally to the surface. On simple geometric grounds, this adsorbed structure is specific to the (111)/

36

J. R. ANDERSON

isohydrocarbon system. The extra residue reactivity from this structure can be understood in terms of the bond shift mechanism already discussed. Thus, if double bonding with the surface occurs a t two out of the three carbon atoms, by comparison with the two-point adsorbed structure, e.g., (B) , st'ructure (E) can provide for an increase in isomerization rate on

CfC\(

IIPt

Pt

II

Pt

purely statistical grounds by a factor of two. However, a more important factor may well br the extent to which (E) allows electron removal from the r-system thus lowering the energy of the bridged structure. AS discussed in previous sections, the extent to which polycrystalline evaporated metal films expose (111) planes in the surface will be heavily dependent on the conditions of preparation. I n particular, deposition or sintering a t high substrate temperatures will lead to an increased proportion of low index planes. This is also true for supported metal catalysts, and has recently becn explored by Boudart et al. (121) who have correlated a n increasing facility for neopentane isomerization with the increasing temperature of platinum catalyst pretreatment (425"-900°C). We have so far tacitly assumed that an adsorbed hydrocarbon molecule undergoes only a single bond shift process in one period of residence on the surface. However, this is not necessarily the case. I n fact it was shown (24) that about 8% of the total isomerization product from neoprntane was n-pentane, the balance being isopcntane (cf. Table 11).Since we are considering only initial reaction products, the chance of n-pentane being formed by the readsorption of isopentanc already produced, is very small, and the only reasonable conclusion is that two consecutive bond shifts occur during one residence period. In fact, as we shall see from the subsequent discussion, consecutive reactions during a single residence period form a very important and general feature of platinum catalyzed isomerization reactions. A reacting molecule has, as an alternative to isomerization, reaction by hydrogenolysis in which lower molecular weight products are formed. This latter process will be discussed in detail in a subsequent section. However, we note hcre that the relative importance of isomerization versus hydrogenolysis decreases as the partial pressure of hydrogen in the reaction mixture increases. This has been demonstrated by Kikuchi et al. (123) for


37

TABLE 111 Proportion of Reaction of n-Pentanea by Zsomerization at 4SO"C

Catalyst

Partial pressure of hydrogen (atm)

Proportion of Reaction by isomerisation (mole percent)

5

19.9 8.0 3.9 14.3 8.2 4.2

Pt/silica

10 20 Pt/carbon

5

10 20 a

n-Pentane partial pressure 0.5 atm. From Kikuchi et al. (123).

reaction of n-pentane over 5% platinum/silica and platinum/carbon catalysts. Conversion to isopentane would be expected to occur by a bond shift process. Some typical results are shown in Table 111. 2. Larger Molecules

When the reactions of alkane molecules larger than the butanes or neopentane are studied, and in particular when the molecule is large enough to form a C5 or a Cs ring, the complexity of the reaction pathway is considerably increased and an important feature is the occurrence, in alddition to isomerization product, of important amounts of cyclic reaction products, particularly methylcyclopentane, formed by dehydrocyclization; this suggests the existence of adsorbed cyclic species. The question is whether the reaction paths for dehydrocyclization and isomerization are related. There is convincing evidence that they are. Skeletal interconversions involving n-hexane, 2- and 3-methylpentane may be represented. 3- Methylpentane

[&-] C

2-Methylpentane

n-

38

J. R. ANDERSON

If this were so, one would expect a correlation between, the ring-opening behavior of methylcyclopentane and the distribution of reaction products in the isomerization of the hexanes. A reasonable correlation of this sort has bcen shown to exist for various reactions on supported platinum and platinum film catalysts (113, 131, 132, 115, 84, 133, SO) and some typical data are shown in Table IV. Because there are complications arising from multiple reaction pathways for isomerization, one does not necessarily expect exact agreement in the values for the isomer ratios. However, the data in Table IV can leave no doubt that an adsorbed Cs cyclic intermediate provides an important reaction path. It is interesting that the agreement is poorest with thick film catalysts (massive platinum) and best with highly dispersed catalysts. This accords with the model that will emerge from the subsequent discussion as to the relative importance of an adsorbed C5cyclic intermediate as a function of catalyst structure. The use of 'V-labeled reactant molecules for mechanistic studies has recently been extended to reactions of the hexanes by Gault and his collaborators (84, 81, 115), using 0.2%, 10% platinum/alumina, as will as platinum film catalysts. The results are quite strongly dependent on the type of catalyst used. The results, which indicate the proportions of the TABLE IV Isomer Distributions in Initial Products from Hydrogenolysis of Methylcyclopentane and from Isomerizalion of Hexanes over Platinum Catalysts" Reactant hydrocarbon Catalyst

MCPb n-H 2-MP/3-MP ratio

Ultrathin Pt film, 0.02 p g om-2 2.44 (273°C)" Ultrathin Pt film, 0.13 pg cm-2 -2.7 (273°C)c 0.274 Platinum/alumina (300°C)" 2.20 10% Platinum/alumina (270°C)" 3.3 Thick fringe-free polycrystalline P t films, deposited 28O0-30O0C (273°C)" 1.90

MCP

2-MP

3-MP/n-H ratio

2.61

0.80 0.67

2.18 2.20

-0.6 0.55 0.55 3.1 4.5

-

2.05

12

5.3

MCP

3-MP

_______ 2-MP/n-H ratio

1.95

-

-1.6 1.41 1.10 1.10 10.3 16

23

5.1

Data: film catalysts (SO);platinum/alumina catalysts (113, 116, 84). MCP = methylcyclopentane; 2-MP = 2-methylpentane; 3-MP = 3-methylpentane; n-H = n-hexane. Reaction temperature. a


39

various isotopically labeled isomerization products, are contained in Tables V, VI, and VII. Referring first of all to the reactions over 0.2% platinum/alumina (Table V) the major features of the product distributions may be explained by a simple reaction via an adsorbed C5 cyclic intermediate. For instance, if reaction had proceeded entirely by this path, 2-methylpentane-2-13C would have yielded 3-methylpentane labeled 100% in the 3-position (instead of 73.4%) and would have yielded n-hexane labeled 100% in the 2-position (instead of 90.2%). Similarly, 3-methylpentane-2-13C would have yielded a 2-methylpentane labeled 50% in the methyl substituent (instead of 42.6%), and would have yielded n-hexane labeled 50% in the 1-and 3-positions (instead of 43.8 and 49% respectively). The other expectations are very easily assessed in a similar manner. On the whole, the data of Table V lead to the conclusion that some 80% or so of the reacting hydrocarbon reacts via a simple one step process via an adsorbed C5 cyclic intermediate. The departures from the distribution expected for this simple process are accounted for by the occurrence of bond shift processes. It is necessary to propose that more than one process (adsorbed Cg cyclic intermediate or bond shift) may occur within a single overall residence period on the catalyst; Gault’s analysis leads to the need for a maximum of three. The number of possible combinations is large, but limitations are imposed by the nature of the observed product distributions. If we designate a bond shift process by B, and passage via an adsorbed (3.5 cyclic intermediate by C, the required reaction paths are C (dominant reaction), B, CB, BB, BC, BCB, CBB, BBB where CB, for instance, means reaction via a C process followed by a B process without intervening desorption. These all represent parallel reaction pathways. Even so, not all possible bond shifts occur and, in particular, the ones that are absent include those leading to a 2,3-dimethylbutane skeleton (an insignificant type of reaction product), as well as those which convert between 2-methylpentane and n-hexane skeletons (a conclusion from the distribution of the labeled isomers). Unlike the behavior over 0.2% platinum/alumina, the main features of the labeled product distributions obtained over 10% platinum/alumina and over platinum film catalysts (Tables VI and VII respectively) cannot be explained in terms of a single dominant reaction pathway via an adsorbed C5 cyclic intermediate. Again, parallel, multiple-step reaction pathways are involved. The results from 2-methylpentane-2-13C have been qualitatively accounted for (84) by the pathways

C, B, CB, BC, BCB.

TABLE V Proportions of Isotopically Labeled Products from Isomerizaiion of Hexanes over 0.2% Platinum/Alumim Catalyst at B7S°Ca Initial products (mole percent) 3-methylpentanes

2-methylpentanes

n-hexanes

Reactant hydrocarbon

n/v 4.1

73.4

22.5

-

-

-

9.8

90.2

0

49.4

4.9

45.7

-

-

-

68.5

4.3

27.2

L

5.9

1.4

92.7

-

-

-

-

-

-

r

-

-

-

1.3

90.8

7.9

-

-

-

"r

-

-

-

42.6

50.0

7.4

43.8

7.2

49.0

v

-

-

-

7.5

0

92.5

-

-

-

AA

LA

a

fw

Corolleur, Corolleur, and Gault (81).


41

TABLE VII Proportions of Isotopically Labeled Producta from Zsomerizatiun of 8-Methylpentane-8-13C over Thick Polycrystalline Platinum Films at 87S"Ca Initial productsb(mole percent) ?

3-methylpentanes Reactant hydrocarbon

/L Corolleur (84).

T

n-hexanes

Y T+T

Arv

F

hr\/

w

63-64

17-18

P

3M z

0-2

57-61

3W1

20


43

A more detailed analysis of the results obtained over 10% platinum] alumina (115) leads to an extended array of parallel, multistep reaction paths, and it was concluded (for 273°C) that a n adsorbed species had a chance of reacting via an adsorbed CS cyclic intermediate of about 0.3, of reacting via a bond shift of about 0.2, and a chance of desorption of about 0.5. One would expect these probabilities to be temperature dependent, but to different extents, so that the nature of the product distributions should also be temperature dependent. An immediate conclusion from this work with the 13C-labeled hexanes is that the reaction is strongly dependent on catalyst structure. This conclusion is also evident from the product distribution data summarized in Table VIII. With 0.2 and 10% platinum/alumina catalysts [inert alumina (116)1,the former gave considerably greater proportions of cyclic reaction products than the latter. With the film catalysts, the ultrathin films gave greater proportions of cyclic products than fringe-free thick films, and within the ultrathin films the proportion of cyclic products decreased with increasing average platinum particle size. Inasmuch as one expects the average platinum particle size to be greater with 10% than with 0.2% platinum/alumina, all of these results are consistent. It was previously argued (28) that an enhanced proportion of methylcyclopentane indicated an enhanced formation of adsorbed C5 cyclic reaction intermediate, suggesting that reaction pathways via the latter are more important the smaller the average platinum particle size. This argument assumes, of course, that the ease of product desorption is not too different on the various catalysts. The proportion of cyclic product varies more strongly with particle size from n-hexane reactant than from 2- or 3-methylpentane. (The failure reported previously (28) to observe a dependence on average platinum particle size of the proportion of cyclic product from 2-methylpentane when using ultrathin films, was due to a failure to use films of sufficiently low mass per unit substrate area.) That ultrathin platinum films favor reaction via an adsorbed C5 cyclic intermediate agrees with preliminary results (134) from reaction of 2-methylpentane-2-13C on ultrathin and thick films. The same trend with particle size was reported (115) for this reaction with supported catalysts. Work with ultrathin and thick fringe-free platinum films has shown that not only does the product distribution change with catalyst structure, but the specific rate of raction (per unit platinum area) changes also (SO). The data in Fig. 12 for the reaction of 2-methylpentane and n-hexane show a decrease in the specific rate with increasing particle size. Other work has been reported in the literature on the influence of platinum particle size (in supported catalysts) on isomerization and dehydrocyclization reactions. However, the reaction conditions tend to vary widely

TABLE VIII Distribution of CsReaction Products from Hexanes over Platinum Catalysts Proportion of cyclic product (mole C H + B percent) Reference

Reaction temperature ("C)

2-MPb

3-MP

n-Hexane 0.8% Pt/silica 0.2% Pt/alumina 10% Pt/alumina ultrathin Pt film, 0.02 pg cm-2 ( < i5A)d ultrathin Pt film, 0.13 pg cm-2 (20A)d ultrathin Pt film, 0.5 pg cm-2 (38b)d thick polycrystalline Pt 6lms,c deposited 0°C thick fringe-free polycrystalline Pt' films deposited 275"-30O0C

295 303 298 273 273 273 273

46c 55 62 4.7 8.7 14.5 28

20 31 1.8 4.0 5.0 13

25 7.3 77 79 75 46

8.3 5.5 13

25 7.3 93 87 81 59

273

33

17

30

20

50

2-Methyl- 0.8% Pt/silica pentane 0.2% Pt/alumina 10% Pt/alumina ultrathin Pt film,0.02 pg cm-2 ( 95%) led to hydrogenolysis. Boudart and Ptak (122) have reported that among all the metals of group VIII plus copper and gold, which they tested for neopentane isomerization, only iridium ( 18O0-2OO"C) and gold (440"-480°C) were active in addition to platinum. They studied this reaction in a flow system, using gold as an unsupported powder, and a 10% iridium/silica supported catalyst. The selectivity on iridium and gold for bond shift found in this work was, although a good deal lower than on platinum, still very appreciable (-15% for Ir, -18% for Au, and -9Oo/, for P t ) . It is interesting that Anderson and Avery (24) could not detect any isomerization products from reaction over a thick iridium film. As Boudart and Ptak suggested, this difference may be a result of the use of a static reaction system with evaporated films compared with a flow system with the supported catalyst. However, we believe a more likely reason is an actual difference in the nature of the catalyst surfaces. From the comments made in earlier sections concerning the nature of the surface of supported catalysts, one would expect the surface of the supported iridium to carry a greater concentration of adsorbed contaminant than the surface of the evaporated film, and it is very probable that the effect of this would have been to diminish the rate of hydrogenolysis (v.i.) . The activity of gold for both the hydrogenolysis and isomerization of neopentane reported by Boudart and Ptak is remarkable because gold (as evaporated films) is inactive for exchange between aliphatic hydrocarbons and

60

J. R. ANDERSON

deuterium, and no chemisorption of aliphatic hydrocarbons on the surface of clean gold has been detected. At the moment we are unable to offer a a conclusive explanation for this apparent conflct in the behavior of gold surfaces. Nevertheless it is worth considering the possibility that the activity observed by Boudart and Ptak had its origin in surface impurity. The gold powder used by these workers was 99.99yo pure, and was extensively reduced at 500°C in hydrogen before use. However, it is possible that the activity might have resulted from the segregation of some active transition metal impurity at the surface of the gold particles. Since it is known (156) that the lack of chemisorptive activity of gold as a barrier to catalytic activity can be overcome if molecular species are dissociated first, prior to their presentation to the gold surface, it is possible that the impurity on the gold surface may serve as a site for dissociative adsorption, the species then being transferred to the gold surface. Kikuchi et al. (123) failed to detect isomerization products from n-pentane over silicasupported iron, cobalt, nickel, ruthenium, rhodium, and iridium, or over carbon-supported ruthenium and rhodium. Of these metals, other work (122, 157) has shown that rhodium, ruthenium, and iridium possess some isomerization activity. It is reasonable to suppose that this result by Kikuchi et al. stems from the relatively high hydrogen partial pressures used, which were generally > 5 atm (cf. Table 111). The activity of metals other than platinum for skeletal reactions of larger molecules is not well documented, particularly in a mechanistic sense. Carter, Cusumano, and Sinfelt (157) have recently studied the reaction of n-heptane on a series of group VIII metals in the form of hydrogen-reduced (300°C) metal powders. The nature of the reaction pathways is summarized in Table IX. Although many metals have been TABLE IX Reaction of n-Heptane over Reduced Metal Powdersa Percentage of reaction

Catalyst Platinum Palladium Rhodium Ruthenium Iridium a


Hydrogenolysis Isomeriaation

275 300 113 88 125

Carter, Cusumano, and Sinfelt (167).

37 91 93 92.5 87

47

6 7 7.5 13

Dehydrocycliaation

16

3

-


61

used in dual-function catalysts, there are relatively few cases studied in detail where the support is known with some confidence to be nonacidic, and where the reactions can be ascribed to the metal alone. Shuikin's review (124) mentions a number of silica-supported catalysts which have activity for skeletal isomerization and for which, by analogy with the behavior of platinum, one would expect a carbocyclic reaction path to be important. I n these instances, cyclic reaction products, including aromatics, are produced. These catalysts include the use of palladium, rhodium, and ruthenium. From Shuikin's account it appears that, as would be expected by comparison with Table IX, for rhodium and ruthenium hydrogenolysis is a very important reaction. Sinfelt, Carter, and Y a k s (157a) have studied the dehydrogenation of cyclohexane to benzene over a range of unsupported nickel/copper catalysts a t 316°C. There was a relatively small increase in specific activity on increasing the copper content from zero to about 5 at. %; following this, the activity remained constant for copper contents in the range 5-80 at. %, but it fell by a factor of about lo2on going from 80 to 100 at. % copper. It was concluded from the nature of hydrogen adsorption isotherm data that a t low copper contents the surface was considerably richer in copper than was the bulk, so that on changing the bulk concentration from zero to about 10 at. % copper, the surface composition changed from zero to about 50 at. % copper: further, from about 10 to 80 at. % copper in the bulk, the surface composition remained roughly constant a t about 50 at. copper. Clearly, the dehydrogenation activity does not parallel the nickel content of the surface, and single surface nickel atoms alone do not apparently constitute the only catalytically active dehydrogenation sites. Solymosi (158) has studied the reaction of cyclohexane over a wide range of supported nickel catalysts, and has particularly examined the influence of additives such as oxides of zinc, cadmium, and titanium, on the nature of the reaction. In the absence of additive it was reported that a t 400"-5OO"C, reaction proceeded essentially to 100% methane over catalysts variously using magnesia, alumina, or silica supports. However, by the addition of zinc oxide to any of these catalysts, hydrogenolysis was strongly suppressed, and the reaction product was reported as consisting of 8&98% benzene. The addition of cadmium oxide or titanium dioxide had a similar, but less marked, effect. Solymosi has interpreted these effects as resulting from the behavior of the additive as a dopant on the semiconducting properties of the oxide support, and that this in turn controls the electron concentration in the nickel particles which reside on the support. We are skeptical about this suggestion. For one thing, the amount of additive is enormous by comparison with the limits where the theory of impurity doping of semiconductors can be expected to be valid, while

62

J. R. ANDERSON

the expected change in electron concentration in the metal would be minute (cf. 158a) unless the particle was a cluster of only a few atoms, so the loss of only a single electron to a trap could significantly alter its electronic properties. A change in the chemical composition of the metal surface is also possible.

V. Hydrogenolysis on Metals The simplest hydrocarbon hydrogenolysis reaction is that with ethane which can yield methane as the sole reaction product. On the other hand, with larger molecules a range of reaction products is possible since the adsorbed reactant may fragment a t more than one C-C bond and, furthermore, even if only one such bond is broken in the reaction of a given molecule, the reactant will often contain more than one stereochemically distinguishable type of C-C bond. The nature of the hydrogenolysis process as revealed by the distribution of reaction products is much dependent on both the metal from which the catalyst has been prepared, and upon its structure and history. Thick metal films have been used as model catalysts for the vehavior of massive metal under conditions where adventitious surface contamination is of negligible importance. The two extremes of behavior consist, on the one hand, of total fragmentation of an adsorbed molecule to give methane as the only product, and, on the other hand, the rupture of only a single C-C bond in a reacting molecule. A comparison of the tendency for various metals, as thick film catalysts, to promote total fragmentation to methane, is available from the data of Muller (130) for the hydrogenolysis of 1 , 1 , 3 trimethylcyclopentane. This tendency is highest for cobalt (at 300°C) over which 97% of the initially reacting parent yielded methane ; with rhodium, nickel, iron (all a t 300"C), and tungsten (at 250"C), the proportion reacting to methane decreased in the listed order from 55 to 43y0; with palladium (at 320°C) and platinum (at 300°C) the proportions were 11 and Oyo,respectively. The behavior of iron was very strongly temperature dependent, and it was found that at 200°C the proportion reacting to methane was only 0.7%. Reactions over nickel films have been extensively studied, and Table X shows some typical product distributions in which the predominance of methane is evident. These data also demonstrate a tendency for the proportion of methane to increase with increasing reaction temperature, and the latter is a trend which is also apparent from the data of Kikuchi et al. (12s) for n-pentane hydrogenolysis on silica-supported iron, cobalt, and nickel catalysts. The implication of extensive fragmentation to methane is that there is little differentiation bctwecn thc reactivity of various types of C-C bonds. It will be seen from Table X that on

63


TABLE X Initial Products from Hydrogenolysis on Thick Polycrystalline Nickel F i l m

Products (mole percent)

P r ~ p a n e223°C ,~ pro pane,^, * 273°C Neopentane,' 210°C Neopentane,a 250°C n-Hexane,b 273°C Methylcyclopentane," 200°C Neohexane,a 180°c Neohexane,a 200°C

88.0 12.0 99.4 0 . 6

-

-

87.5

9.0

3.5

-

98.7

1.3

-

-

-

-

-

94.3

3.3

1.0

0.7

0.1

0.5

0.1

-

-

80

10

10

-

32

13

5

-

16

34

80

4

3

-

5

8

Anderson and Baker (68). Shimoyama (SO).

nickel film catalysts, neohexane is the only molecule where an appreciable reactivity differentiation exists, and then only if the reaction temperature is low enough; in this case, rupture of the Csec-Cprimbond is considerably more facile than any of the other bonds in the molecule which all have a quarternary carbon atom in them. The same is true for neohexane hydrogenolysis over evaporated films of rhodium (68). Platinum is an important example of a metal where, even on an uncontaminated surface such as is offered by an evaporated film, there is a strong tendency for only one C-C bond t o be ruptured in any particular reacting molecule. On this basis, one may express the distribution of reaction products in terms of relative C-C bond rupture probabilities. Some data of this sort are contained in Table XI for thick and ultrathin film catalysts, and for comparison there are included some data for reactions on a silicasupported catalyst containing 0.8% platinum. These data all refer to reactions carried out in the presence of a large excess of hydrogen, although the results of Kikuchi et al. (12s) indicate that on platinum catalysts the position of C-C bond rupture (in n-pentane) is very little dependent on hydrogen pressure. The data in Table XI show that, on the whole, the 0.8% platinum/silica catalyst used by Matsumoto et al. (110) was inter-

TABLE XI Relative C-C 1

2

3

c-c-c-c

Bond Rupture Probabilities on Platinum Catalysts 4

1

2

3

c-c-c-c

4

1

2

3

4

c-c-c-c-c-c

I

5

6

C 6

1

2

3

c-c-c-c-c

4

5

1

2

3

c-c-c-c-c

4

5

I

I

i? 1 2 s I c-c-c-c

:

2 1

2

3

5

4

I

:

F

*

4

3M

c-c-c-c I

$:

I

E Z

Hydrocarbon n-Butane Isopentane n-Hexane

Catalyst thick polycrystalline Pt filmn thick polycrystalline Pt film. thick polycrystalline Pt film" thick fringe-free polycrystalline Pt filmb ultrathin Pt film, 0.06 pg cm-% 0.8% Pt/silica


1-2

2-3

2-5

270 280 310 273

0.33 0.38 0.34 0.24

0.33 0.38 0.14 0.12

-

0.33 0.12 0.12 0.26 0.26 0.28 -

273 295

0.13 0.16 0.12 0.17

-

0.40 0.42

Relative C-C

-

bond rupture probabilities

3 4

3-5

3 4

4-5

4-6

-

-

-

0.12

-

0.16 0.17

-

-

-

-

-

5-6 Reference

-

0.24

SO

-

0.13 0.12

110

-

24 24

34 SO

2-Methylpentane

thick fringe-free polycrystalline Pt filmb thick fringe-free polycrystalline Pt filmb ultrathin Pt film, 0.08 pg om+ 0.8% Pt/silica %Methylthick fringe-free polycrystalline pentane Pt film" ultrathin Pt film, 0.13 pg cm-2 0.8y0Pt/silica 2,3-Dimethyl- 0.8% Pt/silica butane Neohexane thick polycrystalline Pt film" 0.8% Pt/silica

b

273

0.10

0.41

-

0.43

-

-

0.03 0.03

30

325

0.08 0.51

-

0.37

-

-

0.02 0.02

30

273 295 273

0.14 0.33 0.20 0.26 0.17 0.29

-

0.25 0.32 0.29

-

- 0.14 0.14 - 0.11 0.11 0.08 0.17 -

110 90

273 295 295

0.30 0.11 - 0.11 0.22 0.21 - 0.21 0.12 0.52 0.12 0.12

0.18 0.30 0.14 0.22 0.12 -

-

110 110

286 295

0.11

-

-

68 110

0

0.89 0.76

-

-

-

0 0 0 0.08 0.08 0.08

so

SO

Deposited 0°C. Deposited 270°C.

0

66

J. R. ANDERSON

mediate in its behavior between that of massive platinum (thick films) and ultrathin films although, as would be expected, the resemblence was greater between the ultrathin film and the supported catalysts. Matsumoto et al. also examined these reactions over a 0.4% platinum/alumina catalyst a t 285°C: again the behavior fell between that for massive platinum and ultra-thin film platinum catalysts, although with n-hexane the behavior more closely resembled that of massive platinum, while with 2- and 3-methylpentane the behavior more closely resembled that of ultrathin films. In comparison with these catalysts, the behavior of platinum/ carbon is appreciably different. Thus, Matsumoto et al. (159) also studied hydrogenolysis of the hexanes over a 5% platinum/carbon catalyst a t 386"C, and there was a marked tendency to favor rupture of those C-C bonds containing a primary carbon atom (C-Cprim bonds). The only qualification here is that with neohexane, reaction was confined specifically to Cquart-Cprim bonds. A generally similar trend is also apparent from the data due to Kikuchi et al. (123) for the hydrogenolysis of n-pentane. The latter work also showed that increasing reaction temperature tends relatively to diminish fragmentation a t C-Cprim bonds, and this leads to the conclusion that the trend observed by Matsumoto et al. is not due to the difference in reaction temperature. It is difficult to escape the conclusion that the surface of a platinum/carbon catalyst is qualitatively distinguished from that of other platinum catalysts, possibly as a result of a heavy concentration of carbon on the metal surface. This is a conclusion which coincides with that already reached in a previous section devoted to catalyst structure. Bond rupture probabilities have also been reported by Myers and Munns (160) for hydrogenolysis reactions over a number of supported catalysts containing platinum in the range O.l-l%. The reactions were carried out in the region of 350'-480°C. Provided one confines the comparison to nonacidic supports, these results are in tolerable agreement with the data in Table XI. It is both interesting and important to note that, judged from the behavior of neohexane, hydrogenolysis over platinum, in contrast to nickel, strongly favors rupture of C-Cquart bonds. The mechanistic implications of this will be discussed subsequently. The work of Kikuchi et al. (123) with silica-supported catalysts also shows the high tendency of iron (370"-400"C), cobalt (330"-360OC) and nickel (330"-370°C) to catalyze fragmentation (of n-pentane) to methane. This work also showed that with cobalt and nickel, the extent of methane formation tended to decrease with increasing hydrogen partial pressure. Some data are listed in Table XII. As indicated from Table X, nickel film catalysts favor extensive frag-


67

TABLE XI1 Variation with Hydrogen Pressure of Proportion of Methane in Products from n-pentane Hydrogenolysisa


Partial pressure of hydrogen (atm)

Proportion of methane in hydrogenolysis products (mole percent)

5%Co/silica

350

5%Ni/silica

350

5 30 5 15 30

91.1 76.1 86.6 62.5 50.0

Catalyst

Kikuchi et al. (123).

mentation to methane. However, with supported nickel catalysts other behavior is possible, and is well documented. I n particular, there is a trend with supported catalysts for successive degradation to methane plus a hydrocarbon containing one carbon atom less than the starting material (161-1635). It was shown (162) that on this type of catalyst, ease of bond rupture decreases in the sequence Csec-Cprim

> Ctert-Cprim > Cqusrt-Cprim

and indeed, any nickel catalyst is relatively inefficient for the rupture of a C-C bond containing a quarternary carbon atom. Again we note that with conventional supported nickel catalysts, it is very difficult to disentangle effects due to variations in the chemical composition of the surface. However, recent work with nickel films (SO) has shown that nickel particle size does influencc the nature of the hydrogenolysis product distribution, but only to a relatively small extent; the tendency is for extensive fragmentation to methane to be maximized with large nickel crystals (massive metal). Some comparative data are cont,ained in Table XIII. The ease of rupture of C-C bonds in the ring of cycloalkanes depends on the size of the ring and on the presence of ring substituents, as well as on the nature of the catalyst. I n general, a Ca ring reacts a t a temperature some 100" lower than a C4ring, which in turn reacts some 50°C lower than a Cs ring (cf. Table XV). These differences are undoubtedly due to the strain energy in the smaller rings. C-C bonds in Cg and CSrings approach the alkanes in their resistance to hydrogenolysis. The influence of alkyl substituents on the ease of ring hydrogenolysis depends on ring size. Thus, over a platinum/pumice catalyst a t 100°c,

TABLE XI11 Product Distributions from Hydrogenolysisof Propane and n-Hexane over Nickel Film Catalystsa Initial product distribution (mole percent)

Reactant hydrocarbon and catalyst Propane Thick fringe-free polycrystalline Ni film, deposited 275°C Thick polycrystalline Ni film, deposited O°Cb Ultrathin Ni film, 0.17 pg cm-* n-Hexane Thick fringe-free polycrystalline Ni film,deposited 275°C Ultrathin Ni film, 0.17 pg omp2


CH,

CzHs

C3Hs

273

99.4

0.6

273

99

273

n-CrHio

i-CdHlo

WCgHiz

i-CsHlz

Other total >CS

-

-

-

-

-

-

1

-

-

-

-

-

-

93.5

6.5

-

-

-

-

-

-

250

94.3

3.3

1.0

0.7

0.1

0.5

0.1

-

250

73.9

4.2

5.6

5.8

0.3

8.3

0.4

1.5

All data from Shimoyama (SO), except bAnderson and Baker (68).

Y

*

3tM 52

z

TABLE XIV Relative Bond Rupture Probabilities for Methylcyclopentane Hydrogenolysis on Platinum Catalysts 6

I

Relative C-C

Reaction temperature Catalyst Thick fringe-free polycrystalline Pt filma Thick polycrystalline Pt filmb Thick polycrystalline Pt filmb Ultrathin Pt film: 0.08 pg omp2 O.g(r, Ptlsilica 0.2% Ptlalumina Pt/carbon 6-20% Pt/alumina 0

b

Deposited 270°C. Deposited 0°C.

(“C) 273 266-285 272 273 295 250-310 21@260 315

bond rupture probabilities

F

3

1-2

2-3

3 4

4-5

5-1

1-6

Reference

0.01 0.05 0.06 0.15 0.10 0.21 0.04 0.06

0.32 0.30 0.29 0.25 0.28 0.23 0.33 0.30

0.34 0.30 0.30 0.17 0.23 0.12 0.27 0.28

0.32 0.30 0.29 0.25 0.28 0.23 0.33 0.30

0.01 0.05 0.06 0.15 0.10 0.21 0.04 0.06

0 0 0 0.03 0 0 0

SO, 135

0

68 113 SO,135 110 112 166 167, 168

$

#MP 0

2

0

5 oy

4 U

70

J. R. ANDERSON

methylcyclopropane reacts about 20% more rapidly than cyclopropane (164). On the other hand, over a platinum/carbon catalyst at about 3OO0C, the rate of hydrogenolysis in a series of cyclopentanes decreases with increasing methyl substitution, the relative rates being cyclopentane 1.0, methylcyclopentane 0.6, 1,3-dimethylcyclopentane 0.1, 1,2,3-trimethylcyclopentane 0.02 (165). There is, however, quite a big range possible for different isomers; thus, 1,l-dimethylcyclopentaneis of comparable reactivity to methycyclopentane. The cyclobutanes stand in an intermediate position in that the rate of ring hydrogenolysis does not much depend on alkyl substituents. There is a general trend, particularly evident on platinum catalysts, that irrespective of ring size, ring opening tends to be disfavored for C-C bonds adjacent to alkyl substituents, and particularly adjacent to gem dialkyl substituents. However, the extent to which this occurs is very much dependent on catalyst structure. Some typical data for relative bond rupture probabilities are given in Table XIV for a variety of platinum catalysts. On highly dispersed catalysts such as ultrathin films or silica and alumina-supported platinum containing < 1% metal, the chance of ring opening adjacent to the methyl group is comparable with other ring positions. However, on catalysts containing larger platinum crystallites such as those containing a larger proportion of platinum, or on thick films, the chance of ring opening adjacent to the substituent is quite low. With methylcyclopentane (166, 153) as with methylcyclobutane (112), there is a leveling effect with increasing temperature so that the ring opening probabilities become more nearly equal. As the data in Table XIV indicate, over platinum demethylation of a ring is slow compared to C-C bond rupture within a ring. On the other hand, it is well established [e.g., Kochloefl and Bazant (161)] that if one uses a supported nickel catalyst which is known to favor stepwise alkane degradation, reaction with an alkylcycloalkane is largely confined to the alkyl group(s) which are degraded in a stepwise fashion and are finally removed entirely from the ring. I n discussing the way in which hydrogenolysis occurs, it needs to be recognized a t the outset that more than one reaction pathway is possible, and their relative importance depends both on hydrocarbon structure and on the nature of the catalyst. Table XV summarizes kinetic parameters for hydrogenolysis reactions of alkanes and cycloalkanes over film catalysts and over supported catalysts for which i t can be reasonably assumed that reaction is confined to the metallic phase. These kinetic parameters refer to the overall reaction, i.e., to the rate of disappearance of the parent molecule. It will be evident from Table XV that the catalytic activity of a given metal with a given

TABLE XV Kinetic Parameters for Hydrogmlysis Reactions" Rate = Hydrocarbon Ethane

Propane

Catalyst

Eb

Ni film Ni/kieselguhr Ni on various support

58 52 28-42

W film Pt film Pt/silica Pd film Pd/silica Re/silica Ru/silica Os/silica R h /silica Ir/silica Co on various silica supports &/carbon Cu/silica Ni film W film

27 57 54 50 58 31 32 35 42 36 29-31 18 21 31 18

log,,Ac

X

35.8

-

-

0.7 0.9 to 1.0

P~H~PYHZ

Y -1.2 -1.7 to -2.4

26.2 34.2 31.8 31.9 33.6 26.3 28.1

-

-

N1.0 0.9 -1.0 0.9 0.5 0.8

-

31.8 28.7 -25.5

0.6 0.8 0.7 1.0

-

1.0 1.0

-

26.4 21.7

-

-2.5

-2.5 +O. 3 -1.3 -1.2 -2.2 -1.6 -1 t o 0

-0.6 -0.4

-

Temperature range ("C) 254-273 181-214 175-335

173-1 83 274-340 344-385 273-358 377-392 229-265 177-257 127-162 192-242 177-2 12 219-288 272-317 288-330 217-267 180-190

Reference 68 169 62, 63

170 I71 143 68 24 172

24 172

46 172 173 172 172 63, 174 176 174 63 68 68

Continued

TABLE XV-Continued Rate Hydrocarbon

Catalyst

P

logloAc

a

P’H,P~H~ Y

X

Temperature range (“0

Referenee

~

n-Butane

Isobutane

n-Pentane

Neopentane

Ni film W film Pt film P d film Ni film Pt film (111) Pt film P d film Ni/silica Pt/silica Pt/carbon Pd/silica Ru/silica Ru/carbon Rh/silica Rh/carbon Ir/silica Co/silica Fe/silica Ni film Wfilm Pt film Pt/carbon Ru/silica Os/silica

34 7 21 -38 30 21 19 21 31 28 34 49 29 37 30 31 32 31 23 32 11 21 59 36 32

28.5 16.4 21.1 -27 26.8 21.0 19.5 21 27.3 26.9 27.5 31.4 29.3 32.0 27.9 27.7 27.2 27.1 22.7 26.3 17.5 21.2 32.8 29.8 28.4

0.7 -0.3

0.5 -0.2 0.9 0.7 0.8 0.9 0.9 1.0 1.0 1.0 1.0 0.9 0.5

+1.4 -0

+1.4

-

+o.

1 -1.6 -1.4 -1.4 -1.4 -1.6 -1.5 -1.3 -1.3 -1.5 -1.5 -0.6 -

188-209 144-164 256300 276-310 201-221 265-299 294-305 270-311 315-360 360-435 36Ck-435 385-435 240-285 260-315 275-335 305-375 290-360 315-375 370-430 222-265 202-2 19 239-290 305-370 150-180 130-175

?

P

r

I

12s

Neopentane

Neohexane n-Heptane

Cyclopropane

Methylcyclobutane

Cyclopentane

Rh/silica Ir/silica Au powder Ni film Pt powder P d powder R u powder Rh powder I r powder Ni film Ni/silica Ni/pumice Pt film Pt/pumice P d film Pd/pumice Ir/pumice F e film Ni film Pt film P d film Pt/charcoal Pt/charcoal Pd/charcoal

53 46 51 >25 31 31 37 40 33 7.5

36.8 32.8 25.8

-

24.5 23.3 33.6 33.9 30.2 19.3

-

-

-

-

10.6

-

11

9 14.5 8 10 23 8 16.5 19.5 20 35 40

24.6

-

25.3

-

24.7

-

-

-

-

-

-

0.6 0.8 1

0.2 1 0.1 1 1 1 1 1 1 +ve -

-

-

-

-0.1 -0.1 0 -0.2 0 -0.9 0 0

-

1 -ve -ve -

170-200 180-200 44w80 181-200 255-305 255-355 75-100 1w200 125-200 -46 to 0 2742 130-200 -78 to -23 50-200 -46 to -8 70-200 50-150 150-170 85-200 50-150 13e210 vicinity of of 300°C

1

122 12.3 122 68

E4

157

d

k

126 176 177 126 177 126 177 177 126 178 153 179 153

I n all cases, film catalysts refer to thick polycrystalline films deposited at O"C,except for (111) Pt and (100)Pt which were deposited at 300°C and 250°C on mica and evaporated rocksalt substrates, respectively. b E, activation energy, kcal mol-l. Frequency factor, molecules cm-* sec-l. Q

5

k 3M

U

m B M

s r3

k

P2 d

0

5 0 q

4tl 6

E0 5 w

74

J. R. ANDERSON

hydrocarbon can be strongly influenced by the nature of the catalyst. The range of activation energies for ethane hydrogenolysis over various nickel catalysts is particularly large. Thus, on thick nickel films and on the nickel/kieselguhr catalyst (15% Ni), the activation energies are 52-58 kcal mole-'. Moreover the lower activation energies for supported nickel recorded in Table XV are associated with lower nickel contents (6'2, 143). On the other hand, for each of the metals platinum and palladium, the activation energies for ethane hydrogenolysis are quite similar for thick film and for highly dispersed catalysts supported on silica. The influence of the support is particularly evident with carbon supports for a number of the listed metals; the hydrogenolysis activation energy is, in most cases, considerably different from the value obtained for the corresponding silicasupported catalysts or film catalysts. Again it seems reasonable to suppose that this is due to the presence of carbon on the metal surface. An analogous comment with regard to the likely variability of the chemical composition of the metal surface is relevant to the changes which have been observed in the specific activity for ethane hydrogenolysis over nickel and cobalt catalysts, depending on the nature of the support, namely, silica, alumina, silica-alumina (and carbon) (170, 174, 180). We refer to the discussion of the nature of supported catalysts given in previous sections. Sinfelt and co-workers have presented evidence to show that with 10% nickel/silica-alumina and with various rhodium/silica catalysts, the specific activity for ethane hydrogenolysis varied with metal crystallite size. In the case of nickel/silica-alumina (143), the specific activity decreased by a factor of about 20 for an increa,se in average nickel particle size from 29 to 888, while with rhodium/silica 6181),increasing the average rhodium particle size from < 128 to about 40 A resulted in an increase by a factor of about 3-4 in specific activity, and a further increase in size to 127 8 resulted in a decrease in specific activity by a factor of about 11 relative to the value for the catalyst with rhodium particles of average size < 12 A. Nevertheless, we remain unconvinced that these results necessarily reflect real particle size effects, rather than effects due to variations in the chemical composition of the metal surface. The variability of surface composition has been emphasized in earlier sections of this review which deal with catalyst structure. I n this circumstance, we also find it difficult to be convinced that the range of activation energies listed in Table XV for ethane hydrogenolysis over nickel is not mainly due to changing chemical composition of the nickel surface. Nevertheless, the comparisons presented between the behavior of ultrathin and thick film catalysts (cf. Tables XI, XIII, XIV) where adventitious surface contamination is insignificant, makcs it clear that metal particle size can bc a real determinant to the course of these reactions, and


75

we shall allude to this factor in more detail subsequently. It is useful to note that whereas most metal film catalysts show appreciable and sometimes serve evidence for self-poisoning in hydrogenolysis reactions, supported catalysts are generally much less affected in this way. This selfpoisoning is a feature of the behavior of clean metal surfaces. In discussing the reaction pathways, we believe that the general evidence leads to the conclusion that hydrogenolysis proceeds via adsorbed hydrocarbon species formed by the loss of more than one hydrogen atom from from the parent molecule, and that in these adsorbed species more than one carbon atom is, in some way, involved in bonding to the catalyst surface. In the case of ethane, this adsorption criterion is met via a 1-2 mode or a r-olefin mode. Mechanistically it is difficult to see how the latter could be involved in C-C bond rupture in ethane. With molecules larger than ethane, other reaction paths are possible: One is via adsorption into the 1-3 mode, and another involves adsorption as a r-allylic species. We shall first consider reactions occurring on platinum. The salient points to be considered are as follows. The activation energy for ethane hydrogenolysis is much larger than that for larger hydrocarbons. Hydrogenolysis occurs on thick film catalysts with about the same activation energy with neopentane as with Cs and C4 aliphatics, and in neopentane a 1-2 adsorption mode is impossible. I n neohexane where there can be completition between hydrogenolysis by 1-2 and 1-3 adsorption modes, reaction is limited almost entirely to within the neopentyl group (cf. Table X I ) , that is, the 1-2 mode is of a relatively very minor significance. Thus, at least with aliphatic hydrocarbons up to Cg, it is difficult to avoid the conclusion that, except for ethane, an important hydrogenolysis pathway is via 1-3 adsorption, and that this process is mechanistically related to isomerization by bond shift in that both these processes involve conversion of the 1-3 adsorbed species into a bridged intermediate. If this bridged intermediate is attacked by hydrogen, hydrogenolysis results, rather than skeletal rearrangement. The detailed manner in which this hydrogen attack occurs is uncertain, but overall we may write

L'

!

Pt

lji

in which a hydrogen has been added a t Cs and CZ,and this is followed by further reaction with hydrogen to give product desorption. This process could occur by attack by Hz either from the gas phase or physically ad-

76

J. R. ANDERSON

sorbed on an adjacent site. It is also possible to formulate sequential attack by two adjacent chemisorbed H atoms. The reaction of neohexane over thick tungsten film catalysts (68) a t about 160°C shows that there is litt,le specific C-C bond rupture within the ethyl group, so again one concludes that reaction occurs preferentially via the 1-3 adsorption mode. On the other hand, neohexane hydrogenolysis over thick nickel and rhodium film catalysts a t sufficiently low reaction temperatures clearly proceeds by C-C bond rupture within both the ethyl and neopentyl groups with roughly comparable facility ; under these conditions one concludes that reactions via 1-2 and 1-3 adsorption modes are of comparable importance. We shall refer to these two processes as 1-2 and 1-3 hydrogenolysis, respectively. On this model, the catalytic selectivity for 1-3 hydrogenolysis and isomerization by bond shift depends on a complex matrix of factors: First by the ability of the catalyst to promote 1-3 adsorption (or its triadsorbed alternative, cf. structure E ) , second by the catalyst’s ability to facilitate conversion to the bridged intermediate, and third by the concentration of adsorbed hydrogen which determines whether the bridged intermediate will yield isomcrization or bond rupture. Nevertheless, a s with isomerization by bond shift, one expects that 1-3 hydrogenolysis will occur preferentially on low index surface planes. Considerations of molecular geometry show that a reaction path additional to those via 1-2 and 1-3 adsorption modes is required. Thus, cyclopentane and cyclohexane undergo hydrogenolysis on platinum catalysts with a reactivity much greater than ethane, but comparable to the other aliphatics; hence, 1-2 hydrogenolysis is improbable, and adsorption into a 1-3 mode is sterically impossible. Simple ring opening from cyclopentane or cyclohexane can be formulated as the reverse of ring closure, and mechanisms for this have already been discussed [reactions (11)-( 13)]. These involve the use of r-ally1 and/or s-olefin adsorbed species with the catalytically active site consisting of a single platinum atom. However, there is no reason to believe that a s-ally1 or s-olefin mechanism for C-C bond rupture is limited to ring opening. Rather it should be a general process occurring in straight and branch-chain aliphatics as well. For the purpose of illustration, we write reactions (22) and (23) as generalizations of the reverse of (11) and (13), respectively. We shall refer to mechanisms such as (22) and (23) s-olefin/allyl hydrogenolysis

The species on the right-hand side of (22) and (23) are assumed to be


77

hydrogenated off the surface to complete the reaction. One must also consider the possibility of the reverse of reactions (14) and (15) for hydrogenolysis, as recently emphasized by Muller and Gault ( 1 8 1 ~ ) . With catalysts such as nickel and rhodium for which it has been shown that 1-2 hydrogenolysis is seriously competitive with 1-3 hydrogenolysis, there is no need to assume that a-olefin/allyl hydrogenolysis occurs (but neither can it with certainty be excluded). This conclusion is likely to be true for other catalysts such as cobalt and iron which also favor complete hydrocarbon fragmentation to methane. Anderson and Shimoyama (135) have recently observed the variation in specific rate and in selectivity of hydrogenolysis of methylcyclopentane, 2-methylpentane, and n-hexane for changing average particle size in ultrathin (and thick) platinum film catalysts. The general trend is indicated by some typical data given in Table XVI and it is clear that, with this sort of catalyst, both the specific rate and selectivity decrease with increasing average particle diameter in the range 500°C) extensive olefin formation occurs. In the following discussion we shall, in the main, be concerned only with skeletally distinguished products. Information about reaction pathways has been obtained by a study of the reaction product distribution from unlabeled (e.g. 89, 5, 118, 184-186, 58, 187) as well as from 14C-labeledreactants (89, 87, 88, 91-95, 98, 188, 189). The main mechanistic conclusions may be summarized. Although some skeletal isomerization occurs, chromium oxide catalysts are, on the whole, less efficient for skeletal isomerization than are platinum catalysts. Cyclic C5 products are of never more than very minor impor-

82

J. R. ANDERSON

tance. Skeletal isomerization and aromatization can occur sequentially within a single residence period. In the formation of a CS ring, direct ring closure to a Ca is the most important path (perhaps exclusively so), compared to the prior formation of rings of other sizes followed by ring contraction or expansion. Neopentane does not undergo isomerization (185) on chromia/alumina (non-acidic) a t 537"C, the only significant reaction been hydrogenolysis t o methane and iso-Cd. However, the reality of isomerization is made clear from, for instance, the formation of xylenes from 2,3,4-trimethylpentane. For 0- and p-xylene, the reactions are (24) and (25) (182,93). These processes are formally quite analogous to those we have described in previous c c c I) I I c-c-c-c-c

-[

v ]

c-c-c-c-c-c

(8

-&

(24)

sections as bond shift reactions over metal catalysts. It has frequently been proposed that the isomerization steps in reactions such as (24) and (25) occur via adsorbed cyclo-C, intermediates (89, 3, 185). We can do no more than offer the same comment here as was given in previous sections in relation to the high energy of a cyclopropane ring. This sort of reaction has been discussed in some detail by Pines and Goetschel (S), who compared it with vinyl insertion reactions studied by Raley and coworkers (i9Gi92)in the isomerization of hydrocarbons in the presence of iodine at about 500"C, and which are known to proceed by a type of free-radical mechanism. An important point of agreement is that both fail to give neopentane isomerization. There has long been evidence that the first step in alkane isomerization over chromium oxide is dehydrogenation to an olefin (not necessarily desorbed before further reaction). On this basis, Pines and Goetschel suggest that, starting with neohexane, dehydrogenation first occurs in the ethyl group, and skeletal rearrangement, ultimately to an iso-C5 skeleton, occurs via an adsorbed free radical of the type in (V) .


83

With the evidence available a t the moment, it is the author’s opinion that Pines and Goetschel’s free-radical mechanism is well founded, and thus this mechanism is quite distinct from the bond shift reaction occurring over metals. In order to preserve this distinction, we shall retain the term vinyl insertion for the type of isomerization exemplified in (24) and (25). Observed product distributions, however, make it clear that there must exist reaction pathways in addition to those of the sort in (24) and (25). Thus m-xylene is also a product from 2,3,4-trimethylpentane. This is illustrative of a reaction for which an adsorbed cyclo-C4 intermediate has been suggested (89, 96, 98,S, 185, 188, 189).

c c c I l l c-c-c-c-c

-I??

c-c+c-c-c

-[

1,s)

c-c-c-c-c-c-c

I

(S)

Although providing a satisfactory rationale for experimental facts, an adsorbed cyclo-C4 intermediate still suffers from the problem of high energy, although this should be of lesser importance on chromium oxide catalysts because the reactions are carried out at much higher temperatures than on platinum. The products for which the cyclo-C*isomerization intermediate has been suggested, can also be explained by a sequence of vinyl insertions. Thus, two vinyl insertions would be adequate to explain the formation of m-xylene from 2,3,4-trimethylpentane. Although we have seen in previous sections that extensive reaction sequences are possible on platinum, isomerization by a single vinyl insertion process on chromium oxide is relatively difficult, and the chance of two occurring in sequence would therefore be expected to be very low. In fact, the proportion of m-xylene is comparable to that of 0- and p-xylene. The occurrence of adsorbed cyclo-C7 and CS intermediates has been

84

J. R. ANDERSON

adduced from various studies using 14C-labeled molecules (87, 91-93, 95, 98, 3, 188). For instance, aromatization of 3-methylhe~ane-5-’~Cgave toluene of which some 5% had the 14Catom in the methyl group (188), and this has been interpreted as due to the sequence I

c-c-c-c-c-c

vinyl

14

insertion

* c-c-c-c-c-c-c

14

14

c while aromatization of various 14C-labeled octanes (91-93, 95, 98) has indicated some contribution from an adsorbed cyclo-C8 intermediate. Nevertheless, the occurrence of cyclic reaction intermediates of ring size larger than Cg, does not remain unquestioned. Feighan and Davis (88) examined the aromatization of n-heptane-4-14C and, by comparison with the results of Pines and Chen (87) for n-heptane-l-14C, concluded that an adsorbed cyclo-C, intermediate was not involved. The origin of this disagreement remain unresolved. Inasmuch as the importance of cyclic C7 and C8 intermediates decreases strongly with catalytic usage, it is likely that the disagreement reflects real differences in catalyst properties. The detailed mechanism of ring closure to a Cs ring has long been the subject of speculation (3, 184, 187). For the present purposes we merely note that all of the cyclization processes (11)-(15) already discussed for platinum, are potentially applicable. Although reactions over nonacidic metal oxide catalysts possess some superficial similarities to reactions over platinum catalysts, on the whole, the two systems are sufficiently distinct that, at a mechanistic level, they are worth treating independently.

REFERENCES 1. Weisz, P. B. Advan. Catal. 13, 137 (1962). 2. Gil’debrand, E. I., Znt. Chem. Eng. 6,449 (1966). 3. Pines, H., and Goetschel, C. T., J . Org. C h m . 30, 3530 (1965). 4. Germain, J. E., “Catalytic Conversion of Hydrocarbons.” Academic Press, New York, 1969. 6. Ciapetta, G. F., Dobres, R. M., and Baker, R. W., in “Catalysis” (P. H. Emmett, ed.), Vol. 6, p. 492. Reinhold, New York, 1958. 6. Steiner, A. H., in “Catalysis” (P. H. Emmett, ed.), Vol. 4, p. 529. Reinhold, New York, 1956.


85

7. Anderson, J. R., and Baker, B. G., in “Chemisorption and Reactions on Metallic Films” (J. R. Anderson, ed.), Vol. 2, p. 1. Academic Press, New York, 1971. 8. Brennan, D., Haywood, D. D., and Trapnell, B. M. W., Proc. Roy. SOC.Ser. A 256, 81 (1960). 9. Geus, J. W., in “Chemisorption and Reactions on Metallic Films” (J. R. Anderson, ed.), Vol. 1, p. 129. Academic Press, New York, 1971. 10. Suhrmann, R., and Wedler, G., 2. Angew. Phys., 14, 70 (1962). 11. Maire, G., Anderson, J. R., and Johnson, B. B., Proc. Roy. Soc. Ser. A 320, 227 (1970). 12. Baker, B. G., Johnson, B. B., and Maire, G., Surface Sci. 24,572 (1971). 13. Anderson, J. S., and Klemperer, D. F., Proc. Roy. SOC.Ser. A 256, 350 (1960). 14. Bouwman, R., van Keulen, H. P., and Sachtler, W. M. H., Ber. Bunsenges. 74, 32 (1970). 15. Bouwman, R., Ph.D. Thesis, University of Leiden, The Netherlands, 1970. 16. Baker, B. G., and Fox, P. G., Trans. Faraday SOC.61,2001 (1965). 17. Baker, B. G., and Bruce, L. A., Trans. Faraday SOC.64,2533, (1968). 18. Suhrmann, R., Gerdes, R., and Wedler, G., 2. Naturjorsch. A 18, 1208 (1963). 19. Drechsler, M., and Nicholas, J. F., J.Phys. Chcm. Solids 28,2609 (1967); Nicholas, J. F., Aust. J . Phys. 21, 21 (1968). 20. Evans, D. M., and Wilman, H., Acta Crystallogr. 5,731 (1952). 21. Adamsky, R. F., J.A p p l . Phys. 31,2895 (1960). 22. Bauer, E., in “Single Crystal Films” (M. H. Francombe and H. Sato, eds.), p. 43. Pergamon Press, London, 1964. 23. Jaeger, H., J . Catal. 9, 237 (1967). 24. Anderson, J. R., and Avery, N. R., J. Catal. 5,446 (1966). 25. Sella, C., and Trillat, J. J., in “Single Crystal Films” (M. H. Francombe and H. Sato, eds.), p. 201. Pergamon Press, London, 1964. 26. Macdonald, R. J., Ph.D. Thesis, Flinders University, Adelaide, Australia, 1970. 27. Anderson, J. R., and Macdonald, R. J., J . Catal. 19, 227 (1970). 28. Anderson, J. R., Macdonald, R. J., and Shimoyama, Y., J . Catal. 20, 147 (1971). 29. Anderson, J. R., and Shimoyama, Y., unpublished results, Flinders University, Adelaide, Australia (1970). 30. Shimoyama, Y., Ph.D. Thesis, Flinders University, Adelaide, Australia, 1971. 31. Maat, R. J., and MOSCOU, L., Proc. 3rd Znt. Congr. Catal., 1964 p. 1277 (1965). $2. Allpress, J. G., and Sanders, J. V., Surface Sci. 7, 1 (1967). 33. Jaeger, H., Mercer, P. D., and Sherwood, R. G., Surface Sci. 11,265 (1968). 34. Sanders, J. V., private communication (1970). 35. Moss, R . L., Platinum Metals Rev. 11 (4), 1 (1967). 36. Dorling, T. A,, Eastlake, M. J., and Moss, R. L., J. Catal. 14,23 (1969). 37. Cormack, D., and Moss, R. L., J . Catal. 13, 1 (1969). 38. Dorling, T. A., and Moss, R. L., J . Catal. 7,378 (1967). 39. Dorling, T. A., Lynch, B. W. J., and Moss, R. L., J. Catal. 20, 190 (1971). 40. Schuit, G. C. A., and van Reijen, L. L., Advan. Catal. 10,242 (1958). 41. Coenen, J. W. E., and Linsen, B. C., in “Physical and Chemical Aspects of Adsorbents and Catalysts” (B. G. Linsen, ed.), p. 471. Academic Press, New York, 1970. 42. Morikawa, K., Shirasaki, T., and Okada, M., Advan. Catal. 20,98 (1969). 43. Reinen, D., and Selwood, P. W., J . Catal. 2, 109 (1963). 44. Webb, A. N., Znd. Eng. Chem. 49,261 (1957). 45, Yates, D. J. C., and Sinfelt, J. H., J . Catal. 14, 182 (1969).

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NOTEADDEDIN PROOF Some further relevant material has been published since the body of this review was written. Somorjai and co-workers ( A l ,A 2 ) have examined the dehydrocyclization of n-heptane to toluene over some platinum single crystal surfaces which had been characterized by LEED. The reactions were carried out at 100°-4000C at total pressures in the region of Torr, and with hydrocarbon/hydrogen ratios in the range 1/1 to 1/4. Over a (111) surface the reaction was found to be detectable at 250"-350°C. However, the reaction becomes self-poisoned due to strongly adsorbed residues which may include surface carbon. These self-poisoning residues are also evident from an increased diffuse backscattering of electrons when the surface is examined by LEED. Over some high index surfaces of orientations nominally close to (997) and (911) which have a step and terrace structure identifiable by LEED, the dehydrocyclization reaction was not only a good deal faster than over a (111) surface, but there was little evidence for a diminution in reaction rate, although an appreciable concentration of strongly adsorbed hydrocarbon

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residues was undoubtedly present. It is also known (AS, Ah) that hydrogen chemisorption is more difficult at a (111) platinum surface than at a stepped, high index surface. It seems reasonable to conclude that two effects are operating. I n the first place, platinum step atoms offer more favorable dehydrocyclization sites than do platinum atoms in a (111) face. This is in agreement with the general concepts outlined earlier in this review. In the second place, self-poisoning is minimized by features such as surface steps which promote hydrogen chemisorption. It should be pointed out that dehydrocyclization was only a minor component of the total reaction which also presumably involved isorneriaation and hydrogenolysis. It will be most desirable to extend this work to the entire reaction, because we have seen in previous sections how the importance of hydrogenolysis is likely to be influenced by the concentration of surface hydrogen. Rooney and co-workers ( A 6 )have studied interconversions between protoadamantane (A’) and adamantane (B’), bicyclo[3.2.2]nonane (C’) and bicyclo[3.3.l]nonane (D’), and between 1,7,7-trimethylbicyclo[2.2.l]heptane (E’) and endo- or ezo-2-methyl-3,3dirnethylbicyclo[2.2. llheptane (F‘ and G’, respectively), over a 2% palladium/silica catalyst in excess hydrogen at temperatures in the range 15O0-35O0C. More recently still there has been a report of the platinum catalyzed conversion of A’ to B’ (-46). Because of alleged steric difficulties in effecting bond shift via 1-3 adsorbed intermediate from molecules A’-G’, it was suggested that the reactions involved intermediates (or perhaps transition states) containing a sort of cyclic C) configuration, with the operation of a sort of quasi-carbonium ion mechanism. Examination of molecular models suggests that one’s expectations about the feasibility of 1-3 bond shift processes with these molecules are not unequivocal. There is no apparent impediment with D’; with A’ and C‘ (both of which are somewhat strained) there is some steric impediment, but a 1-3 adsorbed intermediate does not appear impossible; with E’, F’, and G‘ (all of which are highly strained) the required 1-3 adsorbed intermediate is certainly impossible. However, we have already seen in previous sections that there is evidence for more than one sort of bond shift mechanism, and it is by no means inconceivable that in special circumstances, such as with reactant molecules which are strained or which have other special steric requirements, a particular mechanism may be forced into prominence, although it may be relatively unimportant in less unusual circumstances. Somorjai, G. A . CataZ. Rev. 7( l ) , 87 (1972). Joyner, R. W., Lang, B., and Somorjai, G. A., J . Catul. I n press. Lang, B., Joyner, R. W., and Somorjai, G. A., Surface Sci. 30, 454 (1972). Morgan, A. E., and Somorjai, G. A., Surface Sci. 12,405 (1968). Quinn, H. A., Graham, J. H., McKervey, M. A., and Rooney, J. J., J . CataZ. 26, 333 (1972). A6. Sarnman, N. G. Proc. 6th Int. Congr. CataZ.,197B Comment topaper 48.In press.

Al. A2. A3. A4. A5.

Specificity in Catalytic Hydrogenolysis by Metals J. H. SINFELT Corporate Research Laboratories

Esso Research and Engineering Co. Linden, X e w Jersey

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. General Discussion of Hydrogenolysis Reactions. . . ........ A. Mechanistic Aspects.. . . . . . . . . . . . . . . . . . . B. Kinetics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ............ 111. Comparison of Metals as Hydrogenolysis Cat A. Activity Patterns. . . . . . . . . . . . . . . . . . . . . . B. Product Distributions in Hydrogenolysis. . . . . . . . . . . . . . . . . . . . . . . . IV. Contrast between Ethane Hydrogenolysis and Other Reactions. . . . . . . A. Ethane Hydrogenolysis versus Cyclopropane Hydrogenation. . . . . . . B. Ethane Hydrogenolysis versus Cyclohexane Dehydrogenation.. . . . . V. Conclusion ................................... References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ..........

91 92 94

102 106 107 110 116 116

1. Introduction Hydrogenolysis reactions of hydrocarbons have been known for many years ( 1 4 ) . Such reactions involve the rupture of carbon-carbon bonds and the formation of carbon-hydrogen bonds. The reactions are exothermic and are catalyzed by various transition metals. There is a related type of reaction known as hydrocracking which combines the features of metal and acid catalysis (7-11). Bifunctional catalysts consisting of a metal component dispersed on an acidic carrier are commonly employed for this purpose. In general, the nature of the reaction on bifunctional catalysts is different from that on catalysts with purely metallic properties. Hydrocracking on bifunctional catalysts presumably involves a carbonium ion type of reaction mechanism generally associated with acid catalysis, whereas hydrogenolysis on metals is generally interpreted as involving adsorbed hydrocarbon radicals as reaction intermediates. The present article is limited to catalytic hydrogenolysis on metals, and does not consider the subject of hydrocracking on bifunctional metal-acidic oxide catalysts. *

I

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J. H. SINFELT

In this treatment of the metal-catalyzed hydrogenolysis of hydrocarbons, a general discussion is given first of mechanistic and kinetic aspects of the reactions. The remainder of the article is then concerned with the specificity of metal catalysts for hydrogenolysis reactions. This includes an extensive comparison of the catalytic activities of various metals and an attempt to relate the resulting “activity patterns” to properties of the metallic state. Variations in the distribution of products obtained with different metal catalysts in the hydrogenolysis of selected hydrocarbon reactants are also considered. Finally, the marked differences in the “activity patterns” of metals for hydrogenolysis and certain other hydrocarbon reactions are emphasized. When the results of these several types of comparisons are considered together, there emerges a striking example of specificity in heterogeneous catalysis.

II. General Discussion of Hydrogenolysis Reactions Hydrogenolysis reactions of hydrocarbons on metal catalysts have been investigated in some detail. Extensive studies have been conducted on both alkanes and cycloalkanes. While a number of questions still remain with regard to mechanistic and kinetic details of the reactions, the general features seem reasonably clear. A. MECHANISTIC ASPECTS In discussing the mechanism of hydrogenolysis of saturated hydrocarbons, it is logical to begin by considering the mode of chemisorption of the hydrocarbon reactant. Available evidence indicates that the chemisorption involves rupture of carbon-hydrogen bonds (1.2). At temperatures much lower than are required for hydrogenolysis, the chemisorption of hydrocarbons on metals is accompanied by evolution of hydrogen (13). Furthermore, exchange reactions between paraffins and deuterium to yield deuteroparaffins occur a t similar temperatures. These results indicate that carbon-hydrogen bonds are activated more readily than carbon-carbon bonds, and that dehydrogenative chemisorption of the hydrocarbon is the initial step in hydrogenolysis (14-16). The hydrogen deficient surface species formed from the hydrocarbon then undergoes carbon-carbon bond scission. This is followed by hydrogenation and desorption to complete the reaction. As an example, the hydrogenolysis of ethane to methane can be dissected into the following sequence of reaction steps: CzHa

CzHs (ads)

+ H (ads)

CIH, (ads) + adsorbed

CZH, (ads)

HZ C1fragments + CHI

+ aHz

SPECIFICITY IN CATALYTIC HYDROGENOLYSIS BY METALS

93

The symbol (ads) indicates an adsorbed species, and the quantity a is equal to ( 6 - x)/2. The initial step in the sequence involves rupture of a carbonhydrogen bond to form adsorbed CzHs, which then dehydrogenates further to yield the surface species CZH,. The latter then undergoes carbon-carbon bond scission to form monocarbon fragments which are subsequently hydrogenated to methane. A further consideration in the mechanism of hydrogenolysis of hydrocarbons is the structure of the chemisorbed species that undergoes scission of the carbon-carbon bond. In the case of ethane hydrogenolysis, one readily visualizes bonding of the two carbon atoms in the species C2H, to adjacent metal surface atoms. It is convenient to refer to the adsorbed C2H, as a 1,2-adsorbed intermediate (17). If the intermediate is adsorbed CzH4,then each carbon atom could form a single bond with a surface metal atom. However, if adsorbed CzHz is the intermediate in question, it is conceivable that each carbon atom could form a double bond with a surface metal atom. I n any case, ethane hydrogenolysis is visualized as proceeding via a 1,2-adsorbed intermediate. By contrast, Anderson and Avery conclude that the hydrogenolysis of higher molecular weight alkanes involves 1,3adsorbed intermediates (18). This is based on their observation that neopentane, which can form a 1,3- but not a 1 ,2-adsorbed species, exhibits reactivity about the same as that found for n-butane, isobutane, and isopentane, but much higher than that found for ethane. The isomerization reaction of alkanes on platinum films (18-20) is also considered to proceed via 1 ,&adsorbed intermediates, but for this reaction it has been proposed that adsorbed intermediates involving bonding of three carbon atoms to the surface also play an important role on certain surfaces (18). Such surfaces are characterized by the presence of sites consisting of triplets of equally spaced metal atoms, e.g., triplets of atoms in the (111) face of platinum. As a result of studies on the reactions of various hexanes and methylcyclopentane on platinum catalysts, Gault and associates (21, 22) have proposed that isomerization of the former and hydrogenolysis of the latter are interrelated, both involving a common cyclic intermediate. This was based on a comparison of initial product distributions in methylcyclopentane hydrogenolysis and in the isomerization of n-hexane, 2-methylpentane, and 3-methylpentane. In other work on n-heptane isomerization on a platinum powder ( 2 S ) , it has been concluded that a reaction sequence involving cyclic intermediates is not the primary one, although it may play some role. In any casc, there is now abundant evidence for alkane isomerization on platinum, involving a form of catalysis which is different from that observed on common bifunctional catalysts (d4-26), where acidic sites of the carrier are involved. It has recently been reported that iridium and gold also catalyze the isomerization reaction ( 2 7 ) .

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B. KINETICS The first reported work on the kinetics of hydrogenolysis reactions of simple hydrocarbons appears to be that of Taylor and associates a t Princeton (2-4, 14, 15), primarily on the hydrogenolysis of ethane to methane. The studies were conducted on nickel, cobalt, and iron catalysts. More recently, extensive studies on ethane hydrogenolysis kinetics have been conducted on all the group VIII metals and on certain other metals as well (16,28-33). Perhaps the most interesting result of kinetic studies of ethane hydrogenolysis is the strong inverse effect of hydrogen pressure on the reaction rate on most of the metals investigated (Table I). For example, the reaction rate is approximately inversely proportional to the second power of the hydrogen pressure on nickel, rhodium, palladium, and platinum. One explanation for this, originally suggested by the Princeton workers ( 1 5 ) , is that the concentration of the intermpdiate CZH, in the reaction sequence considered in the previous section decreases with inTABLE I Kinetic Parameters jor Ethane Hydrogenolysis on Silica-Supported Metals ( 16 )

Metal Fe co Ni Ru Rh Pd Re 0s

Ir Pt

Temperature range("C)

Ea

239-376 219-259 177-2 19 177-210 190-224 343-377 229-265 125-161 177-210 344-385

30 40.6 32 42 58 31 35 36 5-1

rolb

7LC

md

-

0.6 1.0 1.0 0.8 0.8 0.9 0.5 0.6 0.7 0.9

+ 0.5 -0.8 -2.4

3 . 0 x 1025 4 . 9 x 1031 1 . 3 X loz8 ~ i . 8x 1031 3.7 X 1 . 8 x 1028

7.0

x

1030

5 . 2 X lo2* 5 . 9 x 1031

-1.3 -2.2 -2.5 +0.3 -1.2 -1.6

-2.5

T ( n ,m ) ("C). 270 219 177 188 214 354 250 152 210 357

Apparent activation energy, kcallmole, determined from the temperature dependence of the rate T O a t ethane and hydrogen partial pressures of 0.030 and 0.20 atm, respectively. h Preexporiential factor, molecules/sec/cni2, in the equation, ro = To' exp (-E/RT). * Exponent on ethane pressure in experimental power rate law. d Exponent on hydrogen pressure in experimental power rate law. Temperature at which exponents on ethane and hydrogen pressures were determined.

SPECIFICITY I N CATALYTIC HMROGENOLYSIS BY METALS

95

creasing hydrogen pressure. In the original kinetic analysis of Cimino et al. (15), i t was assumed that equilibrium was maintained between adsorbed C2H, and ethane and hydrogen in the gas phase. It was further assumed that carbon-carbon scission resulted from the interaction of adsorbed CZH, with a molecule of hydrogen and that this step was rate determining. No allowance was made for a possible competition between hydrogen and hydrocarbon for the surface. The kinetic analysis led to a rate expression of the form

Here p~ and pH are the partial pressures of ethane and hydrogen, respectively, and the parameter a is equal to (6 - x)/2. This analysis was subscquently generalized to include cases in which equilibrium is not established betwccn adsorbed C2H, and gas phase ethane (16). Provided that surface coverage by adsorbed species is low, and that equilibrium is maintained between the surface species CZH5 and CZH,, and HZin the gas phase, a kinetic analysis leads to the rate expression

The parameter kl is the rate constant for the initial ethane chemisorption step leading to formation of adsorbed C2&, while the parameter b is equal to kl'lk3KZ. The rate constant kl' refers to the reaction step, CzHs (ads) H (ads) + CzH6, while the rate constant k3 applies to the step, CZH, (ads) Hz + adsorbed C1 fragments. The equilibrium constant KZ applies t o the reaction, CZH5 (ads) H (ads) CZH, (ads) aHz. The kinetic analysis accounts very well for the kinetic data on ethane hydrogenolysis on a cobalt catalyst over a wide range of temperatures and reactant partial pressures (16, 3 2 ) . I n this case the parameter b decreases with increasing temperature. At low temperatures the term bpH(a-l) is large compared to unity in the denominator, but a t sufficiently high temperatures it becomes negligible in comparison with unity. This means that a t low temperatures equilibrium is cff ectively maintained between adsorbed CZH, and gas phase ethane. However, at high temperatures the chemisorption of the ethane is effectively irrevcrsible, so that the overall reaction consists (wentially of a sequence of irreversible steps. This lattcr case is similar to the example of methylcyclohexane dehydrogenation on platinum discussed previously by the writer (26,34). In applying the foregoing kinetic analysis to data on ethane hydrogenolysis on the group VIII metals, one finds for most of the metals that the value of z in C2H, is equal to zero, i.e., the surface intermediate is a CZ species which is totally devoid of hydrogen ( 1 6 ) . This conclusion does not conflict with known facts. However, a value of zero for x for most of the

+

+

+

+

96

J. H. SINFELT

group VIII metals seems somewhat extreme, and may in part be a consequence of assumptions in the kinetic analysis. I n the analysis, it was assumed that a molecule of hydrogen participated in the rate-controlling step. If hydrogen were not involved in this step, ie., if the species CZH, decomposed into monocarbon fragments without interaction with a hydrogen molecule, rate equations (1) and ( 2 ) would become (34a, 34b)

With the revised kinetic analysis, a value of zero for x (i.e,, a = 3) gives the best fit for only platinum, palladium, and rhodium (Table 11). For iridium, osmium, ruthenium, and nickel, a value of 2 for x (a= 2 ) gives the best fit to the data, while for cobalt the best value of x is 4 (a = 1). I n considering the variation of x among the metals, it seems reasonable that x would decrease as the ratio of the dehydrogenation activity to hydrogenolysis activity of the metal increases. It is very likely that this ratio is higher for platinum, palladium, and rhodium than for the other metals in Table 11. In the case of platinum and palladium, the higher ratio is due primarily to the very low hydrogenolysis activities of these metals, TABLE I1 Analysis of Power Rate Law in Ethane HydrogenolysisComparison of Observed and Calculated Exponents on Hydrogen Pressure

Exponent on Hz Pressure

Catalyst Fe co Ni Ru Rh Pd Re 0s

Ir Pt 4

2"

ab

Observed (m)

4 2 2 0 0 2 2 0

1 2 2 3 3 2 2 3

+0.5 -0.8 -2.4 -1.3 -2.2 -2.5 +0.3 -1.2 -1.6 -2.5

Calculated (-nu)

-1.0 -2.0 -1.6 -2.4 -2.7 -1.2 -1.4 -2.7

Number of hydrogen atoms in the species C2H,. Defined by the expression, a = (6- 2)/2.


97

to be discussed in a subsequent section of this article. In the case of rhodium, which has high hydrogenolysis activity, the higher ratio must then be due to higher dehydrogenation activity. This is supported by the particularly high activity of rhodium compared to other metals for ethane-deuterium exchange (35) and ethylene hydrogenation (36-38).Thus, for platinum, palladium, and rhodium it would appear that any intermediate dicarbon surface species formed in the reaction path from CzHe to the final dicarbon species CzH, has a high probability of undergoing further dehydrogenation as opposed to carbon-carbon scission. The metals iron and rhenium in Table I1 present a special case, in that both exhibit positive dependency of the rate on hydrogen pressure. Such a dependency on hydrogen pressure is inconsistent with rate equations (3) and (4),indicating that the presently revised kinetic analysis is not applicable to these metals. Perhaps in these cases the rate of hydrogenolysis is determined by the hydrogenationdesorption of the monocarbon fragments resulting from the scission of the carbon-carbon bond (39). The kinetic analyses leading t o rate equations (1)-(4) are based on the simplifying assumption that competition between hydrogen and hydrocarbon for surface sites is insignificant. This assumption may be challenged on general grounds (40). If such an effect were significant, it might conceivably alter our conclusions on the composition of the surface species CzH, in the general reaction scheme described here. Although there is uncertainty concerning the detailed nature of a surface intermediate such as C2H,, the general nature of the steps involved in the reaction sequence would appear t o be well established. I n any case, the simplified kinetic analyses of the hydrogenolysis reaction have proved to be very useful in interpreting data on a variety of metal catalysts, and have effectively guided investigations in this area.

111. Comparison of Metals as Hydrogenolysis Catalysts A comparison of various metals as catalysts for the hydrogenolysis of hydrocarbons reveals a wide variation in catalytic activity, even among such closely related metals as the noble metals of group VIII of the periodic table. Striking differences in the distribution of hydrogenolysis products have also been revealed in studies on selected hydrocarbon reactants. These features are emphasized in the following discussion of activity patterns and product distributions in hydrogenolysis.

PATTERNS A. ACTIVITY Specific catalytic activities of a number of silica-supported metals have been determined for the hydrogenolysis of ethane to methane (16, 29-31,

98

J. H. SINFELT

33). Data for the metals of group VIII and for rhenium in group VIIA are given in Fig. 1, which is divided into three fields separating the metals of the first, second, and third transition series. The specific activity is defined as the activity per unit surface area of metal. Metal surface areas required for the determination of specific activities are derived from measurements

108 -

- 44

Ni

- 42

106-

lo4

-

Mn

- 40

d---------

A

A

102 -

4. ' \ -

1

I

I

I

-38

\

I

0s

- 52 - 50 - 48

-

46

- 44

1I

PERIODIC GROUP NUMBER

FIQ.1. Catalytic activities of metals for ethane hydrogenolysis in relation to the percentage d character of the metallic bond. The closed points represent activities compared at a temperature of 205°C and ethane and hydrogen pressures of 0.030 and 0.20 atm, respectively, and the open points represent percentage d character. Three separate fields are shown in the figure to distinguish the metals in the different long periods of the periodic table.


99

of the chemisorption of a simple gas, generally hydrogen or carbon monoxide (16, 29-31, 41), which adsorbs very selectively on the metal component of the catalyst. The chemisorption experiments are conveniently made a t room temperature. The method involves a determination of the number of molecules required to form a chemisorbed monolayer on a given amount of supported metal. This quantity is derived from an adsorption isotherm a t conditions such that saturation of the metal surface is achieved. If the stoichiometry of the adsorption process is known, i.e., the number of molecules adsorbed per surface metal atom, it is then a simple matter to determine thc number of metal surface atoms for a given amount of metal in the catalyst. By adopting an appropriate value for the area associated with a single atom in the metal surface, as derived from the lattice spacing of the metal, one can compute the surface area per unit weight of metal. The specific activities in Fig. 1 are relative reaction rates per unit surface area of metal at a temperature of 205OC and ethane and hydrogen partial pressures of 0.030 and 0.20 atm, respectively (16). Absolute values of the reaction rate ro at these conditions can be determined from the parameters E and ro' in Table I, using the experimentally determined relation

ro = rol exp( - E / R T )

(5)

Activities of the group IB metals (copper, silver, and gold) are not shown in Fig. 1, since they are too low to be measured satisfactorily in the same apparatus used for the other metals in Fig. 1. Measurable reaction rates could not be observed on these metals a t temperatures as high as 450°C, indicating that the activities arc pcrhaps several orders of magnitude lower than the activities of the least active group VIII metals. I n the case of copper, it had been reported previously that the hydrogenolysis activity, while small compared t o that of a metal such as nickel, was high enough to give measurable reaction rates at temperatures in the vicinity of 300°C (16, 29). Further work ( 4 2 ) ,however, has failed to confirm this result, and suggests that the copper catalyst used in the earlier work may have been contaminated with a small amount of an active impurity. Although the more recent work on copper indicates that the hydrogenolysis activity is much lower than was originally reported, this has had no important bearing on conclusions rcgarding the activity patterns of metals for hydrogenolysis. In comparing the catalytic activities of metals for ethane hydrogenolysis, it is instructive to consider the variation in activity as a function of the position of the metal within a given period of the periodic table. The data are most complete for the metals of the third transition series. Beginning with rhenium in group VIIA, and proceeding in the direction of increasing atomic number to osmium, iridium, and platinum in group VIII and on to gold in group IB, the hydrogenolysis activity passes through a maximum

100

J. H. SINFELT

value a t osmium. From osmium to platinum alone, the activity decreases by eight orders of magnitude. A similar variation is observed from ruthenium to palladium in the second transition series. It is also probable that the hydrogenolysis activity of the metals in the second transition series attains a maximum value a t ruthenium, much the same as it does a t osmium in the third transition series. This is supported by recent data on the group VIA metal, molybdenum, in the second transition series, which indicate that the hydrogenolysis activity of this metal is virtually negligible by comparison with that of ruthenium ( 4 3 ) . Indeed, the hydrogenolysis of ethane on molybdenum proceeds readily only a t temperatures high enough to cause carbiding of the molybdenum, i.e., about 375”-400°C. Thus, for the metals of the second and third transition series, there is a similar pattern of variation of catalytic activity for the hydrogenolysis of ethane. The pattern is observed also in n-heptane hydrogenolysis (23), as shown in Fig. 2, and in the hydrogenolysis of neopentane ( 2 7 ) . It is significant that the same pattern is observed with highly dispersed supported metals as with unsupported metals of much lower dispersion. The ethane and neopentane studies were made with supported metal catalysts, while the n-heptane studies were made with metal powders. While metal dispersion and support effects can have a significant influence in hydrogenolysis (16), they are still not of sufficient magnitude to have an important bearing on “activity patterns’’ observed in the comparison of various metals. I

I

I

Ru

Vllll

VlllZ

VlH3


FIQ.2. Catalytic activities of group VIII noble metals for n-heptane hydrogenolysis. The activities are compared at a temperature of 205” C at 1 atm pressure and a H*/nC? mole ratio of 5/1 (23).


101

I n considering the kinetic parameters for ethane hydrogenolysis compiled in Table I, it may be noted for the metals of the second and third transition series that the enormous decrease in catalytic activity from ruthenium to palladium and from osmium to platinum within group VIII is accompanied by a large increase in apparent activation energy. Thus, for these metals it appears that the variation in activity is due primarily to changes in activation energy, although there is a n indication of some variation in the preexponential factor and a corresponding ‘(compensation effect.” I n the first transition series, the group VIII metals (iron, cobalt, and nickel) are much more active for hydrogenolysis than copper in group IB. In this respect, the first transition series is very similar to the second and third transition series just discussed. However, maximum catalytic activity in the first transition series is observed for the metal in the third subgroup within group VIII, i.e., nickel, whereas in the second and third transition series the maximum activity is observed for the metal in the first subgroup within group VIII, namely, ruthenium or osmium. Thus, the pattern of variation of hydrogenolysis activity among the group VIII metals of the first transition series is somewhat different from that observed for the corresponding metals of the second and third transition series. This could be considered as a reflection of the known differences in chemical properties between elements of the first transition series on the one hand, and the corresponding elements of the second and third transition series on the other (44). Included with the plots of hydrogenolysis activities of the metals in Fig. 1 are plots of the percentage d character of the metallic bond, a quantity introduced in Pauling’s valence bond theory of metals to represent the extent of participation of d orbitals in the bonding between atoms in a metal lattice ( 4 5 ) . For the metals within a given transition series, the pattern of variation of hydrogenolysis activity from one metal to another is very similar to the pattern of variation of percentage d character. On further consideration of the relation between hydrogenolysis activity and percentage d character for all the metals of Fig. 1, it can be seen that the metals of the first transition series (which include iron, cobalt, and nickel) behave as a separate class from the metals of the second and third transition series (which include the group VIII noble metals). For the metals iron, cobalt, and nickel in the first transition series, the hydrogenolysis activities are comparable to those of metals with significantly higher d character in the second and third transition series. Thus, while there is a degree of correlation between hydrogenolysis activity and percentage d character, it is clear that this parameter alone is not adequate for characterizing the catalytic activity of transition metals for hydrogenolysis (16).

102

J. H. SINFELT

In studies of the activities of metals for the catalytic hydrogenolysis of alkanes, the available data indicate that activity comparisons are not strongly affected by the particular alkane used as a reactant. However, the rate of hydrogenolysis increases with the carbon number of the alkane, as indicated by the results of several investigations (19,23, 4 6 ) . Thus, the rates of hydrogenolysis of n-heptane on unsupported metals in Fig. 2 are several orders of magnitude higher than the rates of ethane hydrogenolysis a t 205°C on the same metals supported on silica. Rates for the latter reaction can be found in reference 30 or calculated from the Arrhenius parameters in Table I. The comparison of ethane and n-heptanc hydrogenolysis rates can only be considered as a rough indication of the effect of molecular size on the rate of hydrogenolysis, since the possibilities of metal dispersion and carrier effects have been ignored. Nevertheless, an independent study of the hydrogenolysis of a series of alkanes on a supported ruthenium catalyst showed an effect of similar magnitude for varying molecular size ( 4 6 ) .The increase in rate with increasing molecular size would seem to be attributable, a t least in part, to lower average dissociation energies of carbon-carbon bonds in larger alkane molecules ( 4 7 ) .

B. PRODUCT DISTRIBUTIONS IN HYDROGENOLYSIS I n the hydrogenolysis of saturated hydrocarbons possessing a number of carbon-carbon bonds, which are not all identical, there is the possibility of different rates of rupture of carbon-carbon bonds a t different locations in the molecule. Of particular interest in catalysis is the observation that the distribution of primary hydrogenolysis products (i.e., the "cracking pattern") depends on the metal catalyst used. A classical example of specificity in metal-catalyzed hydrogenolysis is the highly selective attack of nickel catalysts on the terminal carbon-carbon bonds in alkanes (5, 48, 4 9 ) . For example, in the hydrogenolysis of n-hexane on a nickel-silica catalyst a t lSO"C, the only products observed a t very low conversion are methane and n-pentane, formed in cquimolar amounts ( 4 9 ) . Similarly, the hydrogenolysis of 2-methylpentane and 3-methylpentanc yields methane, n-pentane, and isopentane as the only primary products a t low conversions, the mole fraction of methane in the product corresponding closely to the combined mole fraction of normal and isopentanes. The ratio of isopentane to n-pentane in the products of hydrogenolysis of 2-methylpentane and 3-mcthylpentane is very close to 0.5 and 2.0, respectively, which is expected on the basis of purely statistical considerations of the products resulting from scission of terminal carbon-carbon bonds ( 4 9 ) . The situation with platinum catalysts contrasts markedly with that on nickel catalysts, in that the rupture of different carbon-carbon bonds is


103

nonselective. Thus, in the hydrogenolysis of n-hexane over a platinumsilica catalyst, the rates of rupture of various carbon-carbon bonds are nearly the same, giving a spectrum of primary products comprising all of the normal alkanes from methane through n-pentane in comparable molar quantities ( 4 9 ) . The striking specificity possible in the hydrogenolysis of hydrocarbons on metals is clearly demonstrated by recent studies conducted in our laboratory on the hydrogenolysis of n-heptane on a number of group VIII noble metal catalysts ( 2 3 ) .To eliminate carrier effects of the kind which are operative in bifunctional metal-acidic oxide catalysts, unsupported metal powders were employed in the investigation. Hydrogenolysis was the predominant reaction on all the metals except platinum, on which extensive isomerization and dehydrocyclization were also observed. The extents of the various reactions at low total conversions are summarized for the different metals in Table 111. The products of the isomerization reaction are methylhexanes and dimethylpentanes, while the dehydrocyclization reaction yields dimethylcyclopentanes, methylcyclohexanc, and toluene. The data on the various metals in Table I11 were obtained a t widely different temperatures, as necessitated by the huge differences in the catalytic activities of the metals. At the conditions employed with the metals ruthenium, rhodium, and iridium, the thermodynamics are unfavorable for the occurrence of the dehydrocyclization reaction. However, in the case of platinum and palladium, for which the temperatures were high enough to obtain dehydrocyclization, the reaction occurred to a significant extent only on the platinum. In its ability t o catalyze isomerization and de-

TABLE I11 Reactions of n-Heptane on Metals in the Presence of Hydrogene ($3)

Metal Pd Rh Ru

a

Percent Percent Tempera- total con- hydroture ("C) version genolysis

rt

300 113 88 275

Ir

125

6.4 2.9 4.0 2.3 9.4 21.4 1.5

Percent isomerization

Percent dehydrocyclization

0.4 0.2 0.3 0.7 3.8 10.0 0.2

0.2 1.0 2.2 3.4 -

5.8 2.7 3.7 0.G 3.4 8.0 1.3

Conditions: 1 atm, H2/nC7mole ratio = 5.

104

J. H. SINFELT

hydrocyclization to a degree a t least roughly comparable to hydrogenolysis, platinum clearly distinguishes itself from the other metals. The distribution of products from the hydrogenolysis of n-heptane varies markedly among the group VIII noble metals, as shown clearly by the data in Table IV. On palladium and rhodium the terminal carboncarbon bond in n-heptane is attacked almost exclusively, yielding methane and n-hexane as the products a t low conversion levels. However, on the other metals, especially platinum and iridium, there is a much less selective type of cracking leading to a spectrum of primary products present in roughly comparable amounts. No simple relation exists between the level of hydrogenolysis activity and the type of “cracking pattern” observed. For example, platinum and palladium, despite their similar low hydrogenolysis activities, exhibit markedly different distributions of primary products. Furthermore, iridium and rhodium give widely different product distributions, although they are both highly active for hydrogenolysis. The highly selective rupture of the terminal carbon-carbon bond on palladium and rhodium has been observed by others, e.g., in n-hexane hydrogenolysis on a palladium-silica catalyst (50) and in n-butane hydrogenolysis on an evaporated rhodium film (19). The less selective type of “cracking pattern” on platinum has been observed repeatedly (23, 48-50). I n considering ways to rationalize the differences in product distribution in Table IV, it is speculated that the ability of the surface atoms of certain metals to exhibit different valence states may be important (2.3). It has previously been suggested by Boudart and Ptak (27) that the ability of platinum and iridium to catalyze the isomerization of neopentane could be a consequence of the variable surface valency of these metals proposed by TABLE IV Distribution of n-Heptane Hydrogenolysis Products Over Metalsa (23)

Metal Pd Rh Ru Pt Ir

Percent Tempera- total con- __ ture (“C) version C1 300 113 88 275 125

6.4 2.9 4.0 2.3 1.5

46 42 28 31 21

Distribution of hydrogenolysis products (Mole percent)* C2

CS

4 5 12 13 21

-

Conditions: 1 atm, H*/nC, mole ratio = 5. All products are n-paraffins.

4 13 17 15

nC4 nCs 3 12 16 14

nC6

4 46 5 4 1 10 25 9 14 14 15

SPECIFICITY IN CATALYTIC

/c C

I1

M

\ / C4H9 C

I

M

HYDROGENOLYSIS BY METALS

/ and

c\

/ C4H9

C

C

M

M

I

105

II

FIG.3. Structures of adsorbed intermediates formed from n-heptane on metals ($3).

Plummer and Rhodin ( 5 1 ) .The proposal of the latter workers was based on field desorption measurements of the energy of binding of various transition metal atoms to a tungsten single crystal tip. I n considering possible implications of this proposal with regard to “cracking patterns” on metals, it is useful to apply the ideas of Anderson and Avery (52) on the structure of the chemisorbed intermediates in the isomerization and hydrogenolysis of simple aliphatic hydrocarbons. According to these workers, a lf3-adsorbed intermediate is involved, with one of the carbon atoms being doubly bonded to a surface metal atom. In the hydrogenolysis of n-heptane, adsorbed species of the type shown in Fig. 3 may be visualized, where M refers to a surface metal atom. If the carbon-metal double bond is located primarily a t the terminal carbon atom, and if it is assumed that the carbon-carbon bond adjacent to the carbon-metal double bond cracks preferentially, it follows that methane and n-hexane will be the principal products of hydrogenolysis. If we propose that this is the case for palladium and rhodium, we are then left with the question of how these metals differ from metals such as platinum and iridium. If the surface atoms of these latter two metals can exist in different valence states, it seems possible that the carbon-metal double bond could shift readily from the terminal carbon atom to another carbon atom along the hydrocarbon chain, yielding a species such as the second one shown in Fig. 3. Scission of carbon-carbon bonds in the central part of the molecule, in addition to cracking a t the end of the molecule, could then occur readily. This leads to a nonselective “cracking pattern.” I n recent discussions of the mechanism of hydrogenolysis and isomerization of saturated hydrocarbons on platinum, it has been suggested that charge transfer from the adsorbate to the metal occurs (49, 52), with the result that a “carbonium ion-like’’ intermediate is formed. The extensive cracking of internal carbon-carbon bonds in alkanes, and the simultaneous skeletal isomerization reaction, are certainly consistent with such a suggestion. Reactions of this type are commonly observed on acid catalysts, and are generally associated with carbonium ion mechanisms ( 5 3 ) . I n the hydrogenolysis of the higher alkanes on the nonnoble group VIII metals (i.e., iron, cobalt, and nickel) , the mode of cracking is very different from that observed on the noble metals of group VIII ( 4 9 , 5 0 ) .On nickel,

106

J. H. SINFELT

for example, the products of hydrogenolysis are consistent with a reaction scheme involving successive demethylation of the hydrocarbon chain. According to this scheme, cracking occurs only a t terminal carbon-carbon bonds, so that one of the fragments is always a C1 species which is hydrogenated to form methane. The other fragment then has one of two fates. It can undergo further cracking at the terminal carbon-carbon bond to produce additional methane, or it can be hydrogenated and desorbed into the gas phase. I n the hydrogenolysis of various hexane isomers on nickel (@), the methane formation is approximated closely by the expression

Here C1,Cz, etc., represent the moles of alkanes of various carbon number in the product. According to this expression, one mole of methane is formed per mole of Cg, two moles are formed per mole of C,, etc. At the low to moderate conversions investigated, cracking of C2 fragments was apparently negligible. The expression also applies reasonably well to data obtained on cobalt, but not to data on iron (49).On moving from nickel t o cobalt to iron, one finds that the initial distribution of products shifts markedly in the direction of lower carbon number alkanes, suggesting that desorption of products becomes strongly limiting. In the case of iron, unlike that of nickel and cobalt, extensive cracking of Cz fragments must also occur, since the formation of methane is much higher than would correspond to the expression just considered. This is consistent with a low rate of desorption relative to carbon-carbon bond rupture. The successive demethylation scheme of hydrogenolysis just discussed for iron, cobalt, and nickel clearly does not apply to the noble metals of group VIII. This can be seen by examining the product distribution data in Table IV. The amounts of methane observed are much lower than would be expected if the hydrogenolysis occurred by successive demethylation steps. Thus, we have another indication that the noble and nonnoble metals of group VIII behave as two separate classes with regard to their catalytic properties in the hydrogenolysis of hydrocarbons.

IV. Contrast between Ethane Hydrogenolysis and Other Reactions Comparisons of the activities of various metal catalysts for reactions of hydrocarbons involving hydrogen as a reactant or product reveal activity patterns which are strongly dependent on the particular reaction considcrcd. Several examples are discussed in the following subsections of this article.

SPECIFICITY IN CATALYTIC

HYDROGENOLYSIS BY METALS

107

A. ETHANEHYDROGENOLYSIS VERSUS CYCLOPROPANE HYDROGENATION A reaction which provides an interesting contrast with ethane hydrogenolysis on the group VIII noble metals is the hydrogenation of cyclopropane to propane. The reaction has been investigated rather extensively by several groups of workers ( 5 4 4 5 ) . Cyclopropane hydrogenation occurs cleanly on platinum, palladium, rhodium, and iridium (63). However, on the rcmainder of the group VIII metals cyclopropane also undergoes a fragmentation reaction yielding methane and ethane as products (60, 61, 65-65). This fragmentation reaction appears to be a primary reaction occurring in parallel with the hydrogenation reaction. I n the case 106,

I

I

I

CYCLOPRDPANE HYDROGENATlOl -1O'C

-kI/-

1 -

I

I

I

I

I

a V

.'

n ul

Ru 0

w

0s

ETHANE HYDRDGENOLYSIS 205T

\

I-

K

\ lo4

I

\

\

\, I

1

Vllll

I

VII12

hp'

Pd

V11I3


FIa. 4. Comparison of activity patterns of the group VIII noble metals for cyclopropane hydrogenation and ethane hydrogenolysis. The activities were all determined at hydrogen and hydrocarbon partial pressures of 0.20 and 0.030 atm, respectively (63).

108

J. H. SINFELT

of the group VIII noble metals, it is interesting that the fragmentation reaction is limited to the two metals, ruthenium and osmium, which are most active for the hydrogenolysis of ethane. Comparisons of the group VIII noble metals as catalysts for the hydrogenation of cyclopropane to propane and for the hydrogenolysis of ethane to methane reveal striking differences in activity patterns ( 6 3 ) . Catalytic activities of the group VIII noble metals supported on silica are shown in Fig. 4.In the upper field of Fig. 4,the cyclopropane hydrogenation activities of the metals are compared, while in the lower field activities are compared for ethane hydrogenolysis. In the triad of metals comprising osmium, iridium, and platinum in the third transition series of the periodic table, the activity for cyclopropane hydrogenation increases in the direction of increasing atomic number from osmium to platinum. The pattern for ethane hydrogenolysis is exactly opposite, such that the activity decreases markedly from osmium to platinum. In the triad of metals comprising I

I

I

60 Pd a/

4 Pt 50

,/

ETHANE HYDROGENOLYSIS Rh

--

"/'

40

u

E

1 0

30 w

20

CYCLDPROPANE HYDROGENATION

\ 0s

-0

lr

\

Pd

o a

10

Ru

0 Rh

0

I Vllll

I

I

VII12

VlH3

Pt


FIG.5. Apparent activation energies of the ethane hydrogenolysis and cyclopropane hydrogenation reactions on the group VIII noble metals. The activation energies were determined a t hydrogen and hydrocarbon partial pressures of 0.20 and 0.030 atm, respectively (6'3).


109

ruthenium, rhodium, and palladium in the second transition series, the catalytic activity for cyclopropane hydrogenation passes through a maximum a t rhodium, while the activity for ethane hydrogenolysis decreases continuously from ruthenium to palladium. The cyclopropane hydrogenation reaction occurs much more readily than ethane hydrogenolysis on all of the group VIII noble metals, as is evident from the much lower temperatures required ( - 10 versus 205°C in Fig. 4). In addition, the range of activities of the metals in Fig. 4 is much smaller for cyclopropane hydrogenation than for ethane hydrogenolysis, spanning three orders of magnitude for the former as compared to eight orders of magnitude for the latter. Correspondingly, the apparent activation energies of the cyclopropane hydrogenation reaction on the various group VIII noble metals are virtually identical, whereas the activation energies of ethane hydrogenolysis vary markedly on these same metals (Fig. 5). Furthermore, for all the metals in Fig. 5 the apparent activation energy is much lower for cyclopropane hydrogenation than for ethane hydrogenolysis. Within each of the two triads of group VIII noble metals (Ru, Rh, Pd and Os, Ir, Pt) , the activity pattern of the metals for ethylene hydrogenation is the same as that shown for cyclopropane hydrogenation in Fig. 4, and the range of variation of activities is even smaller (66). Furthermore, the activation energy for ethylene hydrogenation is about the same on the various group VIII noble metals (679, and is similar in magnitude to that found for cyclopropane hydrogenation. Similar statements are applicable to the hydrogenation of benzene on these metals (68). Thus, the hydrogenation reactions of cyclopropane, ethylene, and benzene all provide the same general picture when the catalytic properties of the group VIII noble metals are compared. However, these reactions as a group present a marked contrast with the hydrogenolysis of ethane. The fact th a t cyclopropane hydrogenation is grouped with ethylene and benzene hydrogenation, despite the fact that cyclopropane is formally a saturated hydrocarbon, is not particularly surprising. It is well known that the bonding in cyclopropane is very different from that in alkanes in general (69).The view is commonly held that the electrons of the cyclopropane ring are partially delocalized, which is consistent with the classification of cyclopropane as a n unsaturated molecule. As a catalytic probe for investigating differences in the properties of the group VIII noble metals, ethane hydrogenolysis is much more sensitive to differences among the metals than are the hydrogenation reactions of cyclopropane, ethylene, and benzene. This conclusion is derived simply from a consideration of the magnitudes of differences in catalytic activities and apparent activation energies obtained in comparisons of the metals. It is interesting that various metal-catalyzed hydrogenation reactions of hydro-

110

J. H. SINFELT

carbons have been classificd as “facile” reactions (62, 70-72), since the state of metal dispersion has essentially no effect on the specific catalytic activity of the metal for these reactions. By contrast, the specific activity of a metal for ethane hydrogenolysis is a function of the state of metal dispersion (16). The nature of bonding of intermediates to the surface must be quite different for the two kinds of reactions. Ethane hydrogenolysis is believed to involve dissociatively chemisorbed hydrocarbon intermediates (16) which are presumably multiply bonded to the surface. Hydrogenation reactions, however, do not require the formation of dissociatively chemisorbed hydrocarbon intermediates, and the suggestion has been made that such reactions proceed via pi-bonded intermediates ( 7 3 ) .

B. ETHANE HYDROGENOLYSIS VERSUS CYCLOHEXANE DEHYDROGENATION A second reaction which contrasts markedly with ethane hydrogenolysis is the dehydrogenation of cyclohexane to benzene, as demonstrated in a recent detailed investigation on copper-nickel alloy catalysts ( 7 4 ) . The use of alloys to investigate the relationship between catalytic activity and the electronic structure of metals dates back to early ideas of Schwab (75) and Dowden ( 7 6 , 7 7 ) .Alloys of a group I B metal with a group VIII metal, such as copper-nickel, have received particular attention, especially for reactions such as the hydrogcnation of benzene (7’8-83) and ethylene (83-86). One might reasonably anticipate that the dehydrogenation of cyclohexane to benzene would exhibit behavior similar to benzene hydrogenation when copper is added to nickel. However, it does not follow that conclusions based on studies of these reactions would necessarily be applicable to ethane hydrogenolysis. Indeed, the recent data obtained on ethane hydrogenolysis indicate a very different effect of adding copper to nickel. Specific activities of a series of copper-nickel catalysts for ethane hydrogenolysis and cyclohexane dehydrogenation are shown in Fig. 6 as a function of the catalyst composition ( 7 4 ) . The alloys were prepared by a method involving coprecipitation of metal carbonates, conversion to mixed oxides, and reduction in flowing hydrogen at elevated temperature. The alloys were characterized by X-ray diffraction and magnetic measurements, and reaction rates werc determined at low conversion levels in a quasidifferential reactor. Kinetic parameters for the ethane hydrogenolysis reaction are summarized in Table V. Thc specific activities in Fig. 6 are reaction rates a t 3 1 6 ° C . From the figure it is clear that the catalytic activity of nickel for ethane hydrogenolysis decreases markedly and continuously as copper is alloyed with it. Addition of only 5 at.% copper decreases the hydrogenolysis activity by three orders of magnitude. With further addition


I

I

I

I

111

I

I

lot

lo5 I

0

T

CYCLOHEXANE DEHYDROGENATION

.-(

.-I 0

, " lo4 5 I

f

I

\

G

J

lo3

3

V w J

-z> 102 0

t > + V a

10

1

I

I

1

I

1

I

0

20

40

60

80

100

ATOM % COPPER

FIG.6. Activities of copper-nickel alloy catalysts for the hydrogenolysis of ethane to methane and the dehydrogenation of cyclohexane to benzene. The activities refer t o reaction rates a t 316" C. Ethane hydrogenolysis activities were obtained a t ethane and hydrogen pressures of 0.030 and 0.20 atm., respectively. Cyclohexane dehydrogenation activities were obtained at cyclohexane and hydrogen pressures of 0.17 and 0.83 atm, respectively ( 7 4 ) .

of copper, the activity continues to decline, such that the activity of a catalyst containing 63.3 at.% copper is five orders of magnitude lower than that of pure nickel. The activities of catalysts containing more than 95 at.% copper are too low to measure in the apparatus employed in the investigation. In marked contrast to the ethane hydrogenolysis results, the catalytic activity of nickel for cyclohexane dehydrogenation increases initially with addition of small amounts of copper. The activity is then fairly insensitive to alloy composition over a wide range, but finally decreases sharply as the composition approaches pure copper. Interestingly, the activity of a catalyst containing 95.6 at. % copper is about the same as that of the pure nickel catalyst, although the activity of the copper itself is about two orders of magnitude lower than that of the nickel. The initial promotional effect of

112

J. H. SINFELT

TABLE V Summary of Kinetic Parameters for Ethane Hydrogenolysis on Copper-Nickel Alloys ( 7 4 )

Composition ' Temperature (at. % CU) range ("C) 0 6.2 10.3 31.5 42.4 52.7 63.3 74.0

226-268 308-339 326-356 355-396 352-395 383408 383440 424455

Ea

rOf

n=

ma

T(n, m ) ("C).

43 51 51 50 51 50 48 47

2 . 1 x 1031 3 . 2 X 1031 1.1 x 1 0 3 1 6 . 0 X 1029 5 . 2 X 1029 2 . 0 X 1029 1 . 8 X lo** 6 . 5 x 1027

1.0 0.9 0.9 0.8 0.8 -

-2.1 -1.3 -1.3 -1.3 -1.2 -1.2

238 33 1 377 384 399 420

Apparent activation energy, kcal/mole. Preexponential factor, molecules/sec cm2, in the equation r,, = r: exp( -E / R T ) . c Order with respect to ethane. d Order with respect to hydrogen. Temperature at which the reaction orders were determined. a

b

copper on nickel has also been reported for the hydrogenation of benzene and ethylene (81, 85,8486). I n considering the effect of copper on the activity of nickel for ethane hydrogenolysis, we note first the previously mentioned correlation of hydrogenolysis activity with the percentage d character of the metallic bond for metals within a particular transition series (Fig. 1 ) . According to this correlation, the hydrogenolysis activity of nickel should decrease when copper is alloyed with it, since the percentage d character decreases (Fig. 7 ) . However, examination of Figs. 6 and 7 reveals that the hydrogenolysis activity decreases much more sharply than percentage d character with the initial incremental additions of copper to nickel. This could be reconciled if there was a surface region in the alloys with a much higher copper content than that which corresponds to the bulk composition. One might then consider the pcrcentage d character corresponding to the surface region, rather than the bulk. If copper concentrates strongly in the surface region of dilute copper-nickel alloys, the percentage d character in this region would also decline sharply on addition of small amounts of copper to nickel. The hydrogenolysis activities of copper-nickel alloys may also be considered from a different point of view. In studies of the kinetics of ethane hydrogenolysis on nickel and many other metals, it has been concluded that chemisorption of ethane occurs with extensive dissociation of carbon-


113

hydrogen bonds to give a highly unsaturated dicarbon surface residue as the reaction intermediate. It is probable that such an intermediate would be multiply bonded to metal atoms in the surface. If a site involving a number of adjacent nickel atoms is required for the chemisorption, the restriction in the number of such sites available on a surface in which copper concentrates so markedly could severely limit the formation of the intermediate. Evidence for a marked difference between the surface and bulk compositions of dilute copper-nickel alloys has been reported recently by a number of investigators (82, 87-90). Much of the experimental evidence comes from hydrogen adsorption data (7'4, 82, 87, 90). The conclusions of van der Plank and Sachtler were based on the premise that nickel chemisorbs hydrogen while copper does not (82, 87). The total adsorption of hydrogen at room temperature was taken as a measure of the amount of nickel in the surface. However, in hydrogen adsorption studies on the catalysts used to obtain the catalytic results in Fig. 6, the amount of adsorption on the copper catalyst, while small compared to the adsorption on nickel, is not negligible ( 7 4 ) . However, the amount of strongly adsorbed 40

39

E

+ W

38

V

a E

I

Y P

s

37

36

35

I

I

I

I

0

ATOM 70 COPPER

FIG.7. Percentage d character of the metallic bond in copper-nickel alloys as a function of composition (74, 84).

114

J. H. SINFELT

hydrogen on the copper catalyst is negligible, and may provide a better basis for estimating surface composition. Data comparing total adsorption with the amount of strongly adsorbed hydrogen are shown in Fig. 8. The strongly adsorbed hydrogen refers to the amount not removed by a 10-min evacuation a t room temperature following the completion of the initial adsorption isotherm a t room temperature. The amount is determined simply as the difference between the initial isotherm and a second isotherm obtained after the evacuation ( 7 4 ) . Figure 8 shows that on nickel the strongly adsorbed hydrogen constitutes a very high fraction of the total I

I

I

1

I

I

ADSORPTION

I

I

--

I

----

S T R O N G L Y ADSORBED H 2

\

1\

I

I

I

0

20

40

I 60

I 80

\A

100

A T O M % COPPER

FIG.8. The adsorption of hydrogen on copper-nickel catalysts as a function of the copper content. The circles represent the total amount of hydrogen adsorbed a t room temperature at 10-cm pressure. The triangles represent the amount of strongly adsorbed hydrogen, i.e., the amount not removed by a 10-min evacuation a t room temperature following the completion of the initial adsorption isotherm. The amount of strongly adsorbed hydrogen is determined as the difference between the initial isotherm and a subsequent isotherm obtained after a 10-min evacuation (74).

SPECIFICITY IN CATALYTIC HMROGENOLYSIS BY METALS

115

and declines much more sharply than the total adsorption when a small amount of copper is added to the nickel. This is consistent with a strong concentration of copper in the surface, since on copper the amount of strongly adsorbed hydrogen is clearly negligible. In the range of catalyst composition between about 10 and 70 at. % copper the variation in adsorption is relatively small, suggesting that the surface composition does not vary much in this region. As the overall composition of the catalyst approaches 100 at. % copper, the copper content of the surface, of course, increases correspondingly, and the amount of hydrogen adsorption again decreases sharply. I n summary, the hydrogen adsorption results appear to be consistent with the conclusion that copper concentrates markedly in the surface of copper-nickel catalysts of low overall copper content. In the analysis of the very different effects of copper on the catalytic activity of nickel for ethane hydrogenolysis and cyclohexane dehydrogenation, it appears that differences in the nature of the rate-determining steps may be involved. In proceeding with a discussion along these lines, it will be assumed that the strength of adsorption of hydrocarbons on nickel is affected directionally in the same way as the strength of adsorption of hydrogen when copper is added to nickel. If the surface coverage by the reaction product is very high, such that desorption controls the reaction rate, a decrease in the heat of adsorption would increase the rate. It appears that this may be the case in cyclohexane dehydrogenation, i.e., desorption of the benzene product controls the reaction rate. Thus, on addition of copper to nickel the heat of adsorption of benzene would be expected to decrease, leading to a corresponding decrease in the activation energy of the desorption step. The suggestion here is prompted by previous work on the dehydrogenation of methylcyclohexane to toluene on platinum (34), where it was concluded that the reaction rate was limited by desorption of the toluene product. This reasoning would account for the initial enhancement of the rate of cyclohexane dehydrogenation observed on addition of the first increments of copper to nickel. The range of composition (6-740/, copper) over which the rate was essentially constant is likely characterized by a somewhat smaller variation of the heat of adsorption of the hydrocarbon and by a gradual change in the rate determining step of the reaction, such that a t very high copper content the reaction rate is limited by a step prior to the final product desorption step. It is reasonable that the reaction rate could then be adversely affected by a decrease in the heat of adsorption of a hydrocarbon intermediate formed prior to benzene on the surface, thus leading to a decrease in rate as the catalyst composition approaches pure copper. For the case of ethane hydrogenolysis, in which the reaction intermediate is probably a highly unsaturated, dicarbon surface residue with both carbon atoms bonded to metal surface atoms, the results can also be

116

J. H. SINFELT

rationalized in a general way. It seems reasonable that the strength of bonding between the two carbon atoms in such a surface intermediate would vary in an inverse manner with the strength of bonding of the carbon atoms to the surface. Consequently, the rupture of the carbon-carbon bond would be facilitated by an increase in the heat of adsorption, and inhibited by a decrease. If such carbon-carbon bond rupture is the rate-limiting step in the reaction, the rate of reaction should decrease as the heat of adsorption decreases, corresponding to addition of copper to nickel. This would appear to be the case over the whole range of composition in the copper-nickel catalyst system. Presumably, the rate of rupture of the carbon-carbon bond would have to be even higher than that obtained on the base nickel catalyst to permit desorption of the methane product to be the ratecontrolling step.

V. Conclusion A strong clement of specificity is readily apparent in the hydrogenolysis reactions of hydrocarbons on metals. Extensive investigations of hydrogenolysis reactions on a variety of metals have revealed an enormous range of catalytic activities, amounting to a variation of seven to eight orders of magnitude among the group VIII metals alone. Characteristic activity patterns have been clearly identified, relating the activity of a metal to its position in the periodic table. For the metals of a given transition series, there is a close correspondence between the patterns of variation of hydrogenolysis activity and percentage d character of the metallic bond. Another aspect of specificity in hydrogenolysis is the striking variation among metals in the selectivity of rupture of specific carbon-carbon bonds in higher alkanes. The variation ranges from a completely selective rupture of the terminal carbon-carbon bond of an alkane to a completely random rupture of all carbon-carbon bonds a t similar rates. A final aspect of specificity is the marked difference in activity patterns among metal catalysts for alkane hydrogenolysis and various hydrogenation or dehydrogenation reactions of hydrocarbons. The marked difference in the two classes of reactions applies to activity patterns observed with various pure metals or with alloys of varying composition. REFERENCES 1. Zelinskii, N. D., Kazanskii, B. A,, and Plate, A. F., Chem. Ber. 66B, 1415 (1933). 2. Morikawa, K., Benedict, W. S., and Taylor, H. S., J . Amer. Chem. Soe. 58, 1795 (1936).


117

3. Morikawa, K., Trenner, N. R., and Taylor, H. S., J . Amer. Chem. SOC.59, 1103 (1937). 4. Taylor, E. H., and Taylor, H. S., J . Amer. Chem. SOC.61, 503 (1939). 6. Haensel, V., and Ipatieff, V . N., J . Amer. Chem. SOC.68,345 (1946). 6 . Haensel, V., and Ipatieff, V . N., Znd. Eng. Chem. 39, 853 (1947).

7. Archibald, R. C., Greensfelder, B. S., Holzman, G., and Rowe, D. H., Znd. Eng. Chem. 52,745 (1960). 8. Flinn, R. A., Larson, 0. A., and Beuther, H., Ind. Eng. Chem. 52, 153 (1960). 9 . Coonradt, H. L., Ciapetta, F. G., Garwood, W. E., Leaman, W. K., and Miale, J. N., Znd. Eng. Chem. 53, 727 (1961). 10. Larson, 0. A., MacIver, D. S., Tobin, H. H., and Flinn, R. A., Znd. Eng. Chem. Process Design Develop. 1(4), 300 (1962). 11. Coonradt, H . L., and Garwood, W. E., Znd. Eng. Chem. Process Design Develop. 3( l ) , 38 (1964). 12. Bond, G. C., “Catalysis by Metals,” p. 395. Academic Press, New York, 1962. 13. Wright, P. G., Ashmore, P. G., and Kemball, C., Trans. Faraday SOC.54, 1692 (1958). 14. Kemball, C., and Taylor, H. S., J . Amer. Chem. SOC.70, 345 (1948). 16. Cimino, A., Boudart, M., and Taylor, H. S., J . Phys. Chem. 58,796 (1954). 16. Sinfelt, J. H., Catal. Rev. 3(2), 175 (1969). 17. Burwell, R. L., Chem. Rev. 57, 895 (1957). 18. Anderson, J. R., and Avery, N. R., J. Catal. 5, 446 (1966). 19. Anderson, J. R., and Baker, B. G., Proc. Roy. SOC.(London), Ser. A271, 402 (1963). 20. Anderson, J. R., and Avery, N. R., J . Catal. 2, 542 (1963). 21. Barron, Y., Maire, G., Cornet, D., and Gault, F. G., J . Catal. 2, 152 (1963). 22. Barron, Y., Maire, G., Muller, J. M., and Gault, F. G., J . Catal. 5 , 428 (1966). 23. Carter, J. L., Cusumano, J. A., and Sinfelt, J. H., J . Catal. 20, 223 (1971). 84. Mills, G. A., Heinemann, H., Milliken, T. H., and Oblad, A. G., Znd. Eng. Chem. 45,134 (1953). 26. Sinfelt, J. H., Hurwitz, H., and Rohrer, J. C., J. Phys. Chem. 64, 892 (1960). 26. Sinfelt, J. H., Advan. Chem. Eng. 5, 37 (1964). 27. Boudart, M., and Ptak, L. D., J . Catal. 16, 90 (1970). 28. Sinfelt, J. H., J . Phys. Chem. 68, 344 (1964). 29. Sinfelt, J. H., Taylor, W. F., and Yates, D. J. C., J. Phys. Chem. 69, 95 (1965). 30. Sinfelt, J. H., and Yates, D. J. C., J . Catal. 8, 82 (1967). 31. Sinfelt, J. H., and Yates, D. J . C., J . Catal. 10, 362 (1968). 32. Sinfelt, J. H., and Taylor, W. F., Trans. Faraday SOC.64,3086 (1968). 33. Yates, D. J. C., and Sinfelt, J . H., J . Catal. 14, 182 (1969). 34. Sinfelt, J. H., Hurwits, H., and Shulman, R. A., J . Phys. Chem. 64, 1559 (1960). 34a. Sinfelt, J. H., Div. Petrol. Chem. Amer. Chem. SOC.,Preprints 17(3), A53 (1972). 34b. Sinfelt, J. H., J . Catal. 27, 468 (1972). 36. Anderson, J. R., and Kemball, C., Proc. Roy. Soc. (London) Ser. A 223,361 (1954). 36. Beeck, O., Discuss. Faraday SOC.8, 118 (1950). 37. Kemball, C., J . Chem. SOC.p. 735 (1956). 38. Schuit, G. C. A., and van Reijen, L. L., Advan. Catal. 10, 242 (1958). 39. Dowie, R. S., Gray, M. C., Whan, D. A., and Kemball, C., Chem. Commun. p . 883 (1971). 40. Kemball, C., Discuss. Faraday SOC.41, 190 (1966). 41. Sinfelt, J. H., Chem. Eng. Progr. Symp. Ser. No. 73 63, 16 (1967).

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42. Sinfelt, J. H., and Cusumano, J. A., unpublished data. 43. Sinfelt, J. H., and Yates, D. J. C., Nature Phys. Sci. 229, 27 (1971). 44. Cotton, F. A., and Wilkinson, G., “Advanced Inorganic Chemistry,” pp. 661, 760. Interscience, New York, 1962. 45. Pauling, L., Proc. Roy. Soc. (London),Ser. A 196, 343 (1949). 46. Kempling, J. C., “Hydrogenolysis of Small Paraffins Over Ruthenium,” Ph.D. Thesis, McMaster University, Hamilton, Ontario, 1971. 47. Semenov, N. N., “Some Problems in Chemical Kinetics and Reactivity” (M. Boudart, transl.), Vol. I, p. 19. Princeton Univ. Press, Princeton, New Jersey, 1958. 48 Myers, C. G., and Munns, G. W., Ind. Eng. Chem. 50, 1727 (1958). 49. Matsumoto, H., Saito Y., and Yoneda, Y., J. Catal. 19, 101 (1970). 60. Matsumoto, H., Saito Y., and Yoneda, Y., J. Catal. 22, 182 (1971). 61. Plummer, E. W., and Rhodin, T. N., J. Chem. Phys. 49, 3479 (1968). 59. Anderson, J. R., and Avery, N. R., J. Catal. 7, 315 (1967). 63. Greensfelder, B. S., in “The Chemistry of Petroleum Hydrocarbons” (B. T. Brooks et al., eds.), Vol. 2, Chapter 27, pp. 137-164. Reinhold, New York, 1955. 64. Bond, G. C., and Sheridan, J., Trans. Faraday SOC.48, 713 (1952). 65. Addy, J., and Bond, G. C., Trans. Faraday SOC.53, 368 (1957). 66. Addy, J., and Bond, G. C., Trans. Faraday SOC.53, 383 (1957). 57. Addy, J., and Bond, G. C., Trans. Faraday SOC.53,388 (1957). 58. Bond, G. C., and Newham, J., Trans. Faraday Sac. 56, 1501 (1960). 69. Benson, J. E., and Kwan, T., J. Phys. Chem. 60, 1601 (1956). 60. Taylor, W. F., Yates, D. J. C., and Sinfelt, J. H., J . Catal. 4, 374 (1965). 61. Sinfelt, J. H., Yates, D. J. C., and Taylor, W. F., J. Phys. Chem. 69, 1877 (1965). 62. Boudart, M., Aldag, A., Benson, J. E., Dmgharty, N. A., and Harkins, C. G., J. Catal. 6, 92 (1966). 65. Dalla Betta, R. A., Cusumano, J. A., and Sinfelt, J. H., J. Catal. 19, 343 (1970). 64. Knor, Z., Ponec, V., Herman, Z., Dolejsek, Z., and Cerny, S., J . Catalysis 2, 299 (1963). 66. Anderson, J. R., and Avery, N. R., J. Catal. 8, 48 (1967). 66. Bond, G. C., “Catalysis by Metals,” p. 246. Academic Press, New York, 1962. 67. Bond, G. C., “Catalysis by Metals,” p. 242. Academic Press, New York, 1962. 68. Bond, G. C., “Catalysis by Metals,” pp. 315, 320. Academic Press, New York, 1962. 69. Lukina, M. Y . , Russ. Chem. Rev. 31, 419 (1962). 7’0. Boudart, M., Aldag, A. W., Ptak, L. D., and Benson, J. E., J. Catal. 11, 35 (1968). 71. Boudart, M., Advan. Catal. 20, 153 (1969). 72. Boudart, M., Amer. Sci. 57 ( I ) , 97 (1969). 73. Rooney, J. J., and Webb, G., J. Catal. 3, 488 (1964). 7’4, Sinfelt, J. H., Carter, J. L., and Yates, D. J. C., J. Catal. 24,283 (1972). 7’5. Schwab, G. M., Discuss. Faraday Soc. 8, 166 (1950). 7’6. Dowden, D. A., J. Chem. Soc. p. 242 (1950). 7’7. Dowden, D. A., and Reynolds, P., Discuss. Faraday SOC.8, 184 (1950). 7’8. Long, J. H., Fraser, J. C. W., and Ott, E. J., J . Amer. Chem. SOC. 56, 1101 (1934). 7’9. Emmett, P. H., and Skau, N. J., J. Amer. Chem. Soc. 65, 1029 (1943). 80. Reynolds. P. W., J . Chem. Soc. p. 265 (1950). 81. Hall, W. K., and Emmett, P. H., J . Phys. Ghem. 62, 816 (1958). 82. van der Plank, P., and Sachtler, W. M. H., J . Catal. 12, 35 (1968). 83. Best, R. J., and Russell, W. W., J . Amer. Chein. SOC.76, 838 (1954). 84. Hall, W. K., and Emmett, P. H., J. Phys. Chem. 63, 1102 (1959).

SPECIFICITY I N CATALYTIC HYDROGENOLYSIS BY METALS

86. Gharpurey, M. K., and Emmett, P. H., J . Phys. Chem. 65, 1182 (1961). 86. Campbell, J. S., and Emmett, P. H., J . Catal. 7, 252 (1967). 87. van der Plank, P., and Sachtler, W. M. H., J . Catal. 7, 300 (1967). 88. Sachtler, W. M. H., and Jongepier, R., J. Catal. 4, 665 (1965). 89. Sachtler, W. M. H., and Dorgelo, G. J. H., J. Catal. 4, 654 (1965). 90. Cadenhead, D. A., and Wagner, N. J., J . Phys. Chem. 72, 2775 (1968).

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The Chemisorption of Benzene R. B. MOYES AND P. B. WELLS Department of Chemistry, The University, Hull, England

I. Introduction.. . . . . . . . . . . . . . . . . . . . . . . . . . 11. Chemisorption. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. Volumetric and Gravimetric Methods B. Flow and Radiotracer Methods.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . C. Spectroscopic, Magnetic, and Other Instrumental Methods. . . . . . . . D. Evidence from Field Electron Emission Microscopy and from LowEnergy Electron Diffraction. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111. Exchange Reactions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ... A. Reactions of Benzene with Molecular Deuterium.. . . . . . . . . . . . . . . B. Reactions of Benzene with Deuterium Oxide and with DeuteriumLabeled Benzene. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . IV. Some Aspects of Benzene Hydrogenation V. Conclusions. . . . . . . . . . . . . . . . . . . .......... References. . . . . . . . . . . . . . . . . . . .

121 122 122 124 128 131 133 134 141 148 152 154

1. Introduction In any branch of catalysis, the chemisorption and reactivity of the simplest member of a family of substances is an intriguing field of study. For example, examinations of the interactions of simple alkanes and of ethylene with metal surfaces have provided much material for debate. The accumulation of large bodies of information has enabled many fundamental questions to be answered, although others remain unresolved. In this review we survey the literature that describes the chemisorption of the simplest of the aromatic substances, benzene. A cursory consideration will reveal that its behavior is likely t o be enigmatic. It is an unsaturated hydrocarbon, and as such it may undergo associative chemisorption in a manner not dissimilar to that of olefins and diolefins. However, disruption of the six r-electron system, leading to loss or partial loss of the resonance energy of stabilization, might be so unfavorable as to reduce the likelihood of, say, hydrogen atom addition, and so the typical reactions of aliphatic unsaturated hydrocarbons may not be observed. Dissociative chemisorption with no interaction of the r-electron system with the surface is the extreme alternative, in which case the basic reactivity is formally similar to 121

122

R. B. MOYES AND P. B. WELLS

that exhibited by the alkanes. Thus, from the standpoint of their chemisorption, aromatic compounds are best considered in a class on their own. The literature to be reviewed is restricted to reports concerning the chemisorption of benzene itself and its reactivity in the adsorbed state. The chemisorption of substituted benzenes is considered only insofar as it throws light on the mode of reaction of the benzene ring. For example, the reactivity of hydrogen atoms of the ring ortho to a substituent will not be discussed (because this is primarily an effect of substitution) even though this topic has been given much attention, and interesting disputes concerning the experimental evidence and its interpretation exist in the literature. There are few reviews on this subject. That by Bond (1) deals only with material to 1961, and that by Garnett and Sollich-Baumgartner ( 2 ) sets out to present only a proportion of the information covered in the present account. Nevertheless, each should be read along side this review, because the detail in the latter will not be repeated here, and the former remains a useful account of benzene hydrogenation. Section I1 will summarize experimental work designed to determine the extent and manner of benzene chemisorption by the use of physical methods and the radioisotope carbon-14. In Section 111, evidence obtained by use of deuterium as a tracer is examined. The relatively small amount of information concerning the mode of benzene chemisorption that is afforded by studies of its hydrogenation is presented in Section IV.

II. Chemisorption A. VOLUMETRIC AND GRAVIMETRIC PI/IETHODS Few studies have been made of benzene chemisorption by the volumetric method. Zettlemoyer et al. (3) have examined the adsorption of benzene vapor a t 0°C on powders of nickel and of copper. First, the monolayer coverage of argon ( v , ) ~ was ~ measured. The argon was then removed by pumping and the amount of benzene required to form a monolayer, (vml) B ~ , was measured. Weakly adsorbed benzene was then removed by pumping, after which further benzene adsorption provided the value ( v , ~ ) B ~ Some . results are reproduced in Table I. On the assumption that the same extent of surface is accessible both for argon and for benzene adsorption, i t is clear that complete monolayers of benzene were not achieved, that some (Ni) or all (Cu) of the benzene was adsorbed reversibly. It was considered that only the irreversibly adsorbed benzene was chemisorbed, the remainder being physically adsorbed. Thus chemisorption of benzene on copper appeared not to occur. The heat of adsorption of benzene on nickel a t zero

123

CHEMISORPTION OF BENZENE

TABLE I Adsorption of Benzene at 0°C on Nickel and Copper Powders

Metal

Ni CU

(Vrn~)~#/(Orn)~r

0.68 0.34

(Um2)Bz/(Vrn)Ar

0.51 0.34

Percent of monolayer reversibly adsorbed 76 100

coverage was low, being about 25 kcal mole-', and the value diminished with increasing coverage. The authors subscribed to the then current view that chemisorption occurred at the cost of the resonance energy of benzene, and on this basis the primary interaction may have involved a heat of about 65 kcal mole-l. Vacua in the region of Torr were employed in this work, and thus the cleanliness of the surfaces, as defined by present-day standards, is questionable. The heat of chemisorption of benzene on nickel fell with increasing coverage such that, a t about 0.5(u,) B ~ it , had fallen to the value for benzene adsorption on oxidized metal. It would be wise, therefore, to consider the values in Table I to be an indication rather than a quantitative measure of the extent of benzene chemisorption on these metals. The chemisorption a t 0°C of benzene on evaporated films of Ti, Cr, Mn, Fe, Co, Nil and Mo has been examined by Moyes, Baron, and Squire who also used a volumetric method ( 4 ) . The films were prepared a t pressures in the range 10-6-10-7 Torr. Surface areas of Fe, Co, and Ni films, as measured by deuterium chemisorption, were similar to those reported by other workers. The amount of benzene chemisorbed per milligram of metal film decreased on passing from titanium to nickel although, as will be shown in Section 111,the reactivity of adsorbed benzene (as displayed by the rate of hydrogen exchange a t 0°C between CaH6and CeD6)increased on passing from titanium to nickel. Thus, the extent of chemisorption and the reactivity of the chemisorbed material are not simply related. Gravimetric determinations of benzene adsorption using McBain balances have been made by Volter et al. (6) and by Shopov et al. (6). The former authors examined benzene adsorption on cobalt-magnesia. Isotherms (obtained a t 89°C over the narrow relative pressure range p / p " = 0.01 to 0.08) were discussed in terms of the Langmuir equation but, in common with other investigations, it was found that adsorption was not completely reversible. The latter authors (6) investigated benzene adsorption a t 25°C on nickel-silica, the sample being prepared in normal high-vacuum before

124

R. B. MOYES AND I?. B. WELLS

benzene admission. Weight changes showed that up to 24% of the metal surface was occupied by benzene, and that about one-third of this benzene could not be removed by hydrogen at 25°C. Slow deposition of carbonaceous residues occurred at room temperature, but this process could be partially reversed by the action of hydrogen. B. FLOW AND RADIOTRACER METHODS Benzene chemisorption on platinum-alumina in the range 26"-470°C has been measured in a flow system by Pitkethly and Goble ('7). A small dose of benzene was injected into a stream of inert carrier gas and transported to the reactor; the efRuent was then sampled repeatedly and analyzed by gas-liquid chromatography. Information concerning the adsorption and desorption of benzene was obtained from the shape of the subsequent benzene concentration versus time curves. Evidence was obtained for four types of adsorption of benzene: (i) a small extent of reversible adsorption (less than 5% of the total) on platinum; (ii) extensive irreversible adsorption on the platinum the extent of which was independent of benzene pressure and of temperature over a wide range (26"430"C for some catalysts) ; (iii) fast reversible adsorption on the support below 200°C; (iv) additional slow adsorption on, or reaction with, the support. Chemisorption was accompanied by some cracking to C1-C4 hydrocarbons a t high temperatures and by a diminution in the surface area available for benzene chemisorption. Again, partial regeneration of the surface was achievable by purging with hydrogen at elevated temperatures, typically 475°C. 0.22 molecules of benzene were chemisorbed per surface platinum atom on certain of the catalysts studied. The authors calculated that chemisorption of benzene on a (100) face of platinum surface should accommodate 0.185 molecules per surface atom. For adsorption on the (111) face, the number would be 0.143 (carbon atoms over interstices, no interference between hydrogen atoms) or 0.333 (dehydrogenation or distortion to accommodate the carbon skeleton over the interstices). For comparison, physical adsorption of benzene in a monolayer would accommodate 0.166 molecules of benzene per surface platinum atom. Thus, chemisorption with the geometry described above is consistent with the observed extent of chemisorption. It should be noted that these calculations assume that full surface coverage by benzene is achieved, the experimental evidence being as described in (ii) above. However, the investigation next

125

CHEMISORPTION O F BENZENE

described suggests that benzene achieves only low surface coverage on platinum over a wide range of temperature. The use of 14C-labeledbenzene provides a sensitive and powerful method for estimating the extent of benzene chemisorption on metal surfaces. It has, however, been little used until recently. I n principle, the method provides a means of estimating all fates of adsorbed molecules. I n favorable cases, measurements can be made of the total amount adsorbed by counting the surface layer, of the number of molecules displaced from the surface by various treatments, and of the concentration of material that is irreversibly chemisorbed. Tetenyi and Babernics have examined the adsorption of 14C-labeledbenzene on powdered nickel, platinum, and copper (8) and on cobalt (9). Nickel was examined in the greatest detail, and this system will be considered first. A flow system was used in which 14C-labeledbenzene was injected into an argon stream and carried over the catalyst. The system was operated in the range 100"-350°C and at one atmosphere pressure. As temperature was varied, the extent of adsorption was observed to pass through a maximum at 140"-180°C for all but one of the four nickel powders studied (Fig. 1). The exception was Raney nickel, for which the authors suggested that the maximum might have been observed below 100°C had experimentation been extended to that region of temperature. The maxi-

I

1

I

0

g

0.5-

E

6

0.4-

k 1

0.3-

P

3

n

3 ag

.--6

0.2

n

0.1-

o 0.0 100

150

m

200 Temperature

=

o U

-

n

U

I

I

250

300

("C)

FIG.1. The variation of the amount of benzene adsorbed on nickel and on platinum as a function of temperature. From Tetenyi and Barbernics (8),with permission.

126


mum coverage thus achieved was calculated to be 23-63% of the total surface area depending upon the method of preparation of the adsorbent. However, in the range 250"-3OO"C, the fraction of the surface covered by chemisorbed benzene appeared to be only 2-10%. Displacement of this chemisorbed benzene was achieved using unlabeled benzene or hydrogen. I n one experiment, 20% of the radioactive benzene chemisorbed a t 150°C could be replaced by unlabeled benzene at the same temperature, and a further 54% was removed by hydrogen at 150°C; the remainder (6%) was removable by hydrogen at 380°C. Alternatively, the entire quantity of adsorbed material could be removed by hydrogen a t 380°C. Thus, benzene appeared not to achieve full surface coverage on nickel. Indeed, its surface coverage was less than that achieved by hydrogen as measured in a separate series of experiments. Values of the ratio &/0Benzene were generally in the range 2-5. The temperature dependence of the extent of adsorption was not interpreted, except that the results were considered to be consistent with the magnetic measurements of Selwood (see Section I1,C) which indicate that the number of carbon-metal bonds between adsorbed species and the surface increases threefold between 120"and 200°C due to extensive dissociative chemisorption. The authors proposed that two forms of chemisorbed benzene exist a t the nickel surface, (i) an associatively adsorbed form which can be displaced by further benzene, and which may be T- or hexa-aadsorbed, and (ii) a dissociatively adsorbed form that requires the presence of hydrogen t o bring about its removal from the surface. Analogous studies of the chemisorption of I4C-labeledbenzene on copper powder and on platinum powder showed no chemisorption on the former, and the establishment of a low surface coverage over a range of temperature on the latter (Fig. 1).Again the surface coverage of hydrogen on platinum appeared to be about three times that achieved by benzene. Babernics and Tetenyi (9) used the flow method described above (7) together with the radioactive labeling technique, for their examination of benzene adsorption on cobalt powder in the range 35"-200"C. At the lower temperatures, 35"-85"C, the surface was almost fully covered with reversibly adsorbed benzene when the equilibrium pressure was 0.05 atm. The adsorption obeyed the Freundlich equation, and the enthalpy change on adsorption was 8.6 kcal mole-' a t a fractional coverage of 0.125 falling only slowly with increasing extent of adsorption to 7.5 kcal mole-' a t a fractional coverage of 0.500. This adsorption was considered to be entirely physical in nature. I n the range 160"-200°C the extent of adsorption, which was reduced by a factor of ten, again obeyed the Freundlich equation and the enthalpy change for adsorption decreased rapidly with increasing coverage from 10.5 kcal mole-' at a fractional coverage of 0.025 to 5.5 kcal mole-'


127

at a value of 0.225. The use of 14C-labeledbenzene showed that about half of the benzene adsorbed at 200°C was irreversibly chemisorbed at the surface and was only removable by hydrogen treatment at 380°C. According to these authors the extent of benzene chemisorption falls in the series Ni > Pt > Co > Cu = zero. This sequence differs from that for activity in benzene hydrogenation, which is Pt > Ni > Co. A careful study of the kinetics of the exchange reaction between I4Clabeled benzene and unlabeled benzene catalyzed by platinum powders at temperatures up to 150°C has been made by Brundege and Parravano (10).The heterogeneity of the surface was examined in the range4O0-100"C and the activation energies relevant to a wide range of surface sites were estimated, the values being about 9 kcal mole-' at the lower limit and greater than 50 keal mole-' at the upper limit. Pre-exponential factors and activation energies showed a compensation effect. Carbon-14 exchange between labeled benzene and a variety of Cehydrocarbons (e.g., cyclohexane, cyclohexene, n-hexane, 1 , Zdimethylbutane, and methylcyclopentane) was studied using alumina- and silica-supported noble group VIII metals as catalysts (11). For the benzene-cycIohexane combination, for example, the relative activity fell in the sequence Pt > Pd > Ir > Ru > Rh and the rate at 117°C was not dependent upon crystallite size over the range 12-2000 A. This study serves to demonstrate the wild range of interrelated reactions that benzene can undergo, and highlights some of the difficulties involved in attempting to identify the various adsorbed species formed when benzene chemisorbs at a metal surface. The chemisorption of 14C-labeledbenzene at 0°C on evaporated films of Ti, Cr, Mn, Fe, Co, Ni, and Mo has been studied by Moyes et al. ( 4 ) .The extent of adsorption was determined by recovering the material not adsorbed and measuring its activity. The quantity so determined agreed with that obtained by the standard volumetric procedure. Only S-lO% of the chemisorbed material was recoverable by heating the films in hydrogen to 200" or to 350°C. A larger proportion of the chemisorbed material was displaced from the surface, even at O"C, by a further dose of (unlabeled) benzene. For example, the fraction of the chemisorbed material so displaced in 15 min was found to be 20-30% depending upon the metal used. However, for several of these metals, the rate of this displacement is much slower than the rate a t 0°C of hydrogen-deuterium exchange in C&b-C&, mixtures (see Section 111). This constitutes further evidence that the strength of chemisorption of benzene varies widely from one region to another for a given metal surface. Furthermore, there are clear disparities between this work and that of Tetenyi concerning (i) the fraction of the chemisorbed 14C-labeled

128


benzene which was removable from a nickel surface by hydrogen a t about 300°C and (ii) the relative proportions of chemisorbed benzene removable by hydrogen and displaceable by further benzene. This shows that the surfaces of hydrogen-free evaporated nickel films and of nickel powder formed by hydrogen reduction of oxides differ markedly in respect of the sites available for chemisorption.

C. SPECTROSCOPIC, MAGNETIC, AND OTHERINSTRUMENTAL METHODS In principle, direct evidence concerning the formation of adsorbed benzene should be obtainable by infrared spectroscopy. Changes in the infrared spectrum of benzene physically adsorbed on silica surfaces have been observed ( I d ) and interpreted in terms of the interaction of the a-electron system of the aromatic molecule and surface siloxyl groups. An ultraviolet spectrum of benzene adsorbed on Vycor glass has also been observed ( I S ) . However, attempts to obtain an infrared spectrum of benzene chemisorbed at a metal surface have been mostly unsuccessful, e.g., Shopov et al. recorded failure when using nickel-silica (14). The magnitude of the problem is appreciated when one considers the evidence presented by Erkelens and Eggink-du Burck (16a) who claim that a broad band of weak intensity and showing no fine structure extending from 2700 to 3100 cm-l represents absorption by benzene chemisorbed on nickel and on copper. The band was not observed when benzene was chemisorbed on iron, palladium, or platinum. The catalysts were prepared in the form of pressed disks and reduction of salt to metal was carried out in the infrared cell for 16 hr in a stream of hydrogen. Chemisorption at a clean palladium surface resulted in carboncarbon bond fission. At platinum, either clean or hydrogen covered, chemisorption of benzene gave rise to CH2 bands which increased in intensity with time. The authors claim that the characteristic band observed with nickel and copper is due to benzene chemisorbed with loss of aromatic character. Comparison with the spectra of metal-benzene complexes does not support the formulation of the chemisorbed species as a r-complex. A variety of chemisorbed species involving the rupture of some carbonhydrogen bonds, varying in the number of such ruptures up to a maximum of six, may explain the width of the band. A tentative explanation of the chemisorption on iron, platinum, and palladium involved proposing the formation of a metal-carbon skeletal complex which would not be expected t o give an infrared spectrum in the region studied. This spectroscopic detection of chemisorbed benzene is remarkable in that it remains virtually the only evidence for the chemisorption of benzene on copper. Sheppard (15b) has, however, reported the spectrum of chemisorbed benzene on silica-supported platinum. A weak, broad band around 3040 em-' was

129


observed, attributable to aromatic C-H bonds. It was not, however, possible to use this evidence to differentiate between ?r- and a-adsorbed species. The change in magnetization that occurs when benzene is chemisorbed on silica-supported nickel has been examined by Selwood by use of the,ac permeameter method (16). The isotherm obtained by plotting magnetization against volume adsorbed a t 15OoC (Fig. 2) showed that the effect for benzene was three to five times that for an equal number of hydrogen molecules. (Figure 2a also shows, incidentally, that the amount of benzene chemisorbed was one-fifth or one-sixth of the amount of hydrogen chemisorbed.) To interpret these results, it was assumed that each hydrogen atom forms one bond to nickel, and that NCH and Ni-C bonds exhibit identical magnetic behavior. On this basis the number of metal-adsorbate bonds formed on the chemisorption of each molecule of benzene at 150°C is in the range six to ten. Six-point attachment suggests that the benzene molecule chemisorbs with its plane parallel with the surface. A subsequent examination of this system (17) confirmed that approximately six metal-

0.101

I

S

15

10 cc H,(or

C,H,) /o

Ni

50

100 150 Adsorption temperature

200

(%)

FIG.2. (a) Magnetization-volume isotherms for the chemisorption of hydrogen and of benzene on kieselguhr-supported nickel at 150" C (16).(b) Average number of bonds formed by benzene adsorbed on nickel-silica as a function of temperature (17). From J. Amer. Chem. SOC.79, 4637 (1957); 83, 1033 (1961). Copyright by the American Chemical Society. Reprinted by permission of copyright owner.

130


adsorbate bonds are formed on chemisorption of benzene at 120°C and that the number of such bonds formed increases to about eighteen a t 200°C (Fig. 2b), indicating substantial dissociation and carbon-carbon bond rupture. The latter was demonstrated by the hydrogenation of these carbonaceous residues and the characterization of the products so formed. This method is both direct and convincing; it is necessary, however, to reconsider these results in the light of recent advances in knowledge. First, the only type of nickel-carbon bonding considered was the covalent a-bond. We need now to consider also the possibility of benzene chemisorption as a a-bonded species involving both donation of electrons from benzene to nickel, and back-donation of electrons from nickel to benzene. The extent of back-donation (which cannot be determined experimentally by the magnetic method) will modify the interpretation of the results. We merely note that the model proposed approximates perhaps to zero back-donation. Secondly, it is now well known that hydrogen migration between metal crystallites and support (commonly termed “hydrogen spillover”) occurs when many supported metal catalysts are treated with hydrogen. Care must therefore be exercised when interpreting the magnetic effects accompanying benzene chemisorption (which may release hydrogen by dissociation) and when these results are compared with those obtained for hydrogen chemisorption. Thirdly, there is now evidence that the coordination number of metal atoms in a surface influences the reactivity of benzene, and presumably, therefore, the types of chemisorbed states formed. The nickel-silica catalysts used by Selwood contained very high weightings of metal (37.5 and 52.8%) and hence the average particle size is likely to have been several hundreds of angstroms. Thus, the behavior reported should be considered to apply to metal atoms having a coordination number approaching nine. Shopov et al. (6) have used infrared and EPR spectroscopy and gravimetry together to examine benzene chemisorption on nickel-silica (9-12% nickel by weight) a t 25°C. It was demonstrated that benzene chemisorption results in the formation of residues tightly bound to the surface; such residues required treatment in hydrogen a t temperatures in excess of 300°C to bring about their removal from the surface. These authors considered that the sites occupied by the residues are not those which participate in the hydrogenation-dehydrogenation reactions observed when benzenehydrogen mixtures were admitted to the catalyst. It should be noted that the average metal particle size in these catalysts was probably below 40A, i.e., in the range where coordination numbers of substantially less than nine are likely to be met. Suhrmann and co-workers have examined the effects of benzene chemisorption upon a number of physical properties of evaporated metal films

131


(18). Work function and resistance changes have been related to surface coverage on films of Fe, Ni, Cu, Zn, Pd, and Ag (19).The changes were insignificant with the metals of groups IB and IIB, indicating that chemical bonding had not occurred. Benzene adsorption at 90°K on transparent nickel films caused changes in electrical resistivity and photoelectric sensitivity from which it was calculated that 6.2 electrons per molecule of benzene were donated to the metal for covalent bonding. Resistance changes upon benzene chemisorption have also been used by Gryaznov et al. (20) in an attempt to chaoracterize adsorption sites at the surfaces of thin ( < 10, 10-20, and 20-30 A) platinum films. FROM FIELD ELECTRON EMISSION MICROSCOPY AND D. EVIDENCE LOW-ENERGY ELECTRON DIFFRACTION

FROM

Field electron emission coupled with flash-filament s t d i e s have been employed by Condon and Hansen to study benzene chemisorption on tungsten ( 2 1 ) .Evidence was obtained for the chemisorption of benzene by a single bond (probably of ?r-character) to the surface. This form of associatively adsorbed benzene [(I), Scheme 11 appeared to exist in equilibrium with o-adsorbed-CsHs (11) and adsorbed atomic hydrogen.

i

*

(1)

(11)

Scheme 1

Evidence of dissociative chemisorption resulting in the final formation of atomic carbon and its incorporation in the metal lattice at lOOO"I< to form the carbide was reported. These methods when used in combination are informative, and the surfaces studied are clean and well defined. It is to be hoped that other metals will be so studied in the future. Low-energy electron diffraction has been used by Pitkethly and others to investigate the chemisorption of benzene, at pressures up to lo-' Torr and at temperatures ranging from ambient to about 500"C, on the (100) ( 2 2 ) , (110) (23,24) and (111) (22,24) faces of nickel single crystals. Disoriented adsorption occurred at room temperature. Disorientation was complete at the (110) face, and partial at the (111) face. Adsorption was weak and reversible, the benzene being removable by pumping to below 10-lo Torr. On Ni(100), a diffuse ( 2 4 2 X 2 4 2 ) pattern was observed corresponding to some ordering of the adsorbed layer, and desorption

132


at very low pressures occurred more slowly than from the (111) face. T h e diffraction patterns have been interpreted in terms of the associative adsorption of benzene with the plane of the ring parallel with the (111) face, but the molecules seem to be sufficiently compressed a t the (100) face to suggest that the adsorbed species may be a Dewar form ( 2 2 ) .When a crystal with benzene adsorbed a t the surface was heated, some desorption occurred and a proportion of the benzene underwent carbon-carbon bond rearrangement. Direct chemisorption of benzene in the range 175"-375°C on the (100) face of nickel gave the Ni(100) (2 X 2)-carbon pattern which is produced also by the chemisorption of other unsaturated hydrocarbons ( 2 2 ) . For the (111) face, the diffraction pattern observed was interpreted as arising by a complex superposition of coincidence patterns from six domains given by a layer of slightly distorted Ni( 100) (2 X 2)-carbon lying over the hexagonal nickel ( 2 6 ) .This structure was stable to about 400°C. The structures formed on the (11 1) and (100) faces of nickel were clearly related, and it was concluded the adsorption of benzene and its conversion to carbonaceous residues a t the (111) face may cause reorientation of the surface layer of metal atoms to a square structure. Similar reconstruction accompanies 25% coverage of Ni(ll1) by sulfur ( 2 6 ) . The carbonaceous residues are extensively dehydrogenated with respect to CaHa,but do not represent the formation of surface carbide, Ni3C. Linear and cross-linked unsaturated polymeric structures (which are extensions of the unsaturated Cd-species formed on palladium after self-hydrogenation of ethylene (27)) can account for the observed diffraction patterns. Polymers which satisfy the Ni(100) (2 X 2)-carbon pattern (28) are shown in Fig. 3. Chemisorption of benzene a t 297°C on Ni( 110) occurred in a rather different manner. Several patterns, some streaked, were observed, and they followed the same sequence and showed the same behavior as those obtained when acetylene was chemisorbed on this surface (29). These structures have not been fully elucidated, but the streaked patterns suggest (i) that the mobility of adsorbed species along the "furrows" of the (110) face is easier than their mobility across them, and (ii) that dissociation of the carbon skeleton of benzene and the formation of other structures occurs. Partly disoriented layers of graphite were formed when each of these faces was dosed with sufficient benzene and annealed a t temperatures between 375" and 425°C. At higher temperatures, the graphitic and other structures broke down and carbon diffused into the bulk of the metal cr yst a1. Adsorption of benzene on sulfur- and oxygen-contaminated (110) faces of nickel revealed that the ordered layers formed differed from those obtained a t the clean Ni (110) surface (23,29).


133

FIQ.3. Linear and cross-linked polymers suggested as an interpretation of the LEED patterns obtained by admitting benzene t o Ni( 100) in the range 175'375" C (99).

Ill. Exchange Reactions In this section we consider the information that may be obtained about the chemisorbed state of benzene from reactions in which hydrogen atoms of C6H6 are exchanged for deuterium atoms, the source of deuterium atoms being either deuterium gas, or deuterium oxide, or deuterium-labeled benzene. A knowledge of the mechanism of exchange provides information concerning the reactivity of chemisorbed benzene. Published mechanisms of exchange fall naturally into various groups, depending upon whether chemisorption is formulated as an associative or a dissociative process, and whether exchange occurs by the loss of a hydrogen atom followed by the acquisition of a deuterium atom (abstraction-addition mechanism) or by the aquisition of D and the subsequent loss of H (addition-abstraction mechanism). The experimental information is summarized and discussed. In each subsection, the investigations are described in historical sequence, since this allows the development of ideas to be followed.

134

R.

A. REACTIONS

B. MOYES AND P. B. WELLS

OF

BENZENE WITH

n/lOLECULAR

DEUTERIUM

Horiuti, Ogden, and Polanyi (SO) examined the hydrogen exchange reaction between CeHs and gaseous deuterium-enriched hydrogen (enrichment -2% D) , and between CsHs and heavy water using platinum and nickel as catalysts. Combustion and micropyknometry were used to determine the extent of exchange. They recorded, as have many workers since them, that exchange of hydrogen for deuterium in benzene was much faster than the addition of deuterium-enriched hydrogen to benzene, and that exchange with deuterium gas was very much faster than with heavy water. They considered three mechanisms, one involving the dissociative chemisorption of benzene and hence an abstraction-addition mechanism, the second involving a concerted hydrogen switch between chemisorbed deuterium and physically adsorbed benzene, and the third involving associatively adsorbed benzene and an addition-abstraction mechanism. The abstraction-addition mechanism was later rejected (31), on the basis that exchange with deuterium gas and with deuterium oxide should proceed a t the same rate if dissociative chemisorption of benzene was the rate-determining step, it being assumed that both deuterium sources were able to provide chemisorbed deuterium atoms in a fast step. The process of exchange in benzene was considered to proceed as shown in Scheme 2. The

Scheme 2. Hydrogen exchange in benzene by an addition-abstraction mechanism involving associative di-u-adsorption of the reactant [Polanyi ( S l ) ] .

geometrical requirements for the removal of the hydrogen atom from the half-hydrogenated state (IV), which will be referred to in detail below, were not considered. The addition of a further hydrogen or deuterium atom to the half-hydrogenated state would give cyclohexadiene and hence exchange and hydrogenation proceed via a common intermediate. Of course, by definition, no hydrogen atoms of the benzene are released onto the surface in the primary adsorption step when this is associative in nature. Farkas and Farkas (32) examined the kinetics of the exchange and hydrogenation of benzene catalyzed by platinized platinum foil a t room temperature. The occurrence of isotope exchange was detected by the thermal conductivity technique. They reported (i) that the exchange reaction was only a little faster than hydrogenation and (ii) that exchange


135

was zero order in deuterium and of order 0.4in benzene, whereas hydrogenation was zero order in benzene and of the first order in deuterium. This difference of kinetic form for the two competing reactions suggested that they might not proceed via a common intermediate. This was supported by the observed variations in the rates of hydrogenation and of exchange as the temperature was varied. The mechanism proposed for the exchange reaction involved the dissociative chemisorption of benzene to give a-adsorbed CsH6, this abstraction step being followed by deuterium atom addition to give exchanged benzene (Scheme 3).

(11)

Scheme 3. Hydrogen exchange in benzene by an abstraction-addition mechanism involving dissociative chemisorption of the reactant [Farkas and Farkas (3.291.

Thus, both the associative and dissociative chemisorption of benzene gained their devotees. The early work, which has been presented here in the briefest outline, has been reviewed in greater detail by Taylor (33). Many of the conclusions were based on measurements of reaction rates, and on comparisons of the rate of reaction of benzene with, for example, that of the ortho-para hydrogen conversion. These comparisons and conclusions may require some qualification in view of more recent knowledge, such as the mechanistic complexity of hydrogen-deuterium exchange and of the ortho-para hydrogen conversion [see pp. 149-181 in Bond ( I ) ] . Nevertheless, the very fact that argued cases both for associative and for dissociative chemisorption of benzene appeared before the era of the mass spectrometer represents notable experimental achievement. Following the Second World War, hydrogen very highly enriched in the isotope of mass 2 became available, and the mass spectrometer appeared as an analytical tool for the chemist; the time was ripe for very detailed studies of catalyzed isotope exchange in hydrocarbons. The technique of continuously monitoring the reaction by means of a mass spectrometer linked directly to the reaction vessel has been used for many of the studies now to be described. The method by which the experimental data are treated is well known (34);it is reproduced briefly in the footnote (p. 136). Anderson and Kemball (36) examined the reaction between gaseous deuterium and benzene catalyzed by evaporated films of iron, nickel, palladium, silver, tungsten, and platinum. The order of reactivity (estimated from the temperature a t which the addition reaction achieved an initial rate of 1% per minute for a 10 mg film a t certain specified reactant

136


pressures) was W > Pt > Ni > Fe > Pd. Detailed information was obtained for reaction over palladium (0-58°C) platinum (-43.5' to -22.5"C) and silver (293"-373°C). For the first time, it was appreciated that all possible deuterium-labeled benzenes were formed as initial products. For example, for palladium a t 29.5"C the distribution was CaH6D, 61.8%; C ~ H ~ D17.7%; Z, CeH3D3, 7.1%; CtjHzD4, 3.8%; CsHD6, 3.5%; CP,D~, 6.1%. The multiplicity factor' M was always greater than unity (Table 11). The observation that multiple exchange occurred was used as an argument in support of dissociative chemisorption. A process involving "repeated second-point adsorption" was proposed; the formation of di-a-adsorbed C6H4 (V) from phenyl meets this requirement (Scheme 4 ) . The rapid interconversion of (11) and (V) provides multiply exchanged phenyl

Scheme 4 . The formation of di-cr-adsorbed-C6H, by the dissociation of two hydrogen atoms from benzene [Anderson and Kemball (SS)].

species which may then react with chemisorbed hydrogen or deuterium to give multiply exchanged benzene. When the relative chances of phenyl being converted to phenylene and to benzene were specified by a parameter P, a calculated distribution of deuterium in benzene was obtained. For the product distribution quoted above, it was claimed that some 20% of the product was formed at sites where P = 14.8 and about 80% a t sites where P = 0.3 (no single value of P provided a satisfactory calculated distribution). What distinguishes these two types of site was not specified. Let u, 0, w, . . ., z denote the percentage of total benzene present as the species C E H ~ DC6HIDZ,. , . ., CEDE.A function +, which is a measure of the deuterium content of a sample, is given by $I = u 2v 3w 42 5y 62. Provided isotope effects on the reaction rate are ignored, d+/dt = k+(1 - +/+*) where +* is the value of + when isotopic equilibrium is achieved in the system. A second process, namely the disappearance of C6HE,is described by the equation

+ + + + +

-d(CEHe)/dt = ~ ~ [ ( C E H -(c~He.)m]/[lOO E) - (CsHe),] where (CEHe), is the percentage of total benzene present as CEHE when isotopic equilibrium is achieved. Values of k+ and ka can be obtained graphically by use of the integrated forms of the above two rate expressions, and their ratio k$/ko = M represents the mean number of deuterium atoms entering each benzene molecule at the beginning of the reaction. When M = 1 a reaction is said to undergo stepwise exchange when M > 1, a t least a proportion of the reaction occurs by multiple exchange.


137

TABLE I1 M-Values Observed for the Hydrogen Isotope Exchange i n Benaene

Metal (film) Pt

Pd Ag

Temperature ("C) -43.5 -22.5 0.0 50.4 293 373

M 1.5 1.4 1.8 2.9 1.2

1.4

The mechanism described in Scheme 2 was rejected on the grounds that the steric requirement for the abstraction of a hydrogen atom from -CHDof species (IV) could not be met. Assuming an atomically flat surface, and sp3 hybridization of the carbon atom bonded to the surface, the plane of the C6-ring in (IV) is in such a configuration that the hydrogen atom of -CHD- is directed away from the surface, and the deuterium atom toward the surface. Thus, unless the species is adsorbed near a step in the metal lattice, the loss of this hydrogen and the formation of a second carbonmetal bond would require a very considerable distortion of adsorbed species. Lastly, hydrogenation and exchange appeared to be independent processes at palladium and platinum surfaces. Cyclohexane formation involved the addition of six deuterium atoms to benzene, and thus the two processes appeared not to share a common intermediate. Furthermore, although each reaction was of approximately zero order in benzene, exchange was of negative order (-0.5 f 0.2) in deuterium whereas hydrogenation was of positive order (+0.8 f 0.2). Thus, the independence of exchange and of hydrogenation a t the surfaces of these metals appeared to be firmly established. Experimental work published in the years following Anderson and Kemball's report (1957) , have revealed the complexity of the situation. A study of the exchange and hydrogenation of liquid benzene catalyzed b y Raney nickel (36) suggested that the two processes might, in fact, proceed by a common mechanism. However, entry of deuterium into the aromatic hydrocarbon was not measured; instead, the argument was based on kinetic measurements, Langmuir expressions being used to relate surface coverages of their reactants to their pressure or concentration. This work is not a

138

R . B. MOYES AND P. B. WELLS

convincing demonstration that exchange and hydrogenation occur via the common intermediate adsorbed CBH7. The exchange of alkylbenzenes with deuterium catalyzed by nickel (37) provided information that is not easily understood in terms of the mechanism shown in Scheme 4. Taking n-propylbenzene as an example, the hydrogen atoms

A

B

C

D

may be divided into the four groups A, B, C,D.Exchange of this hydrocarbon with deuterium at the surface of unsintered nickel films a t 0°C revealed that hydrogen atoms in groups A and C underwent exchange more rapidly than those of group B which, in turn, were exchanged more rapidly than those of group D ( k A , C : k B : k D = 137:7:2). However, exchange a t 30-50°C a t the surface of sintered nickel films showed that the groups of hydrogen atoms underwent exchange a t decreasing rates in the sequence C > D > A > B. The high reactivity of hydrogen atoms of group A is associated with the low bond dissociation energies of these carbon-hydrogen bonds, and is of no particular significance here. The important feature, for present purposes, is the marked lowering of the exchangeability of group A hydrogen atoms that occurred as the film was sintered. Sintering reduced the exchange rate of hydrogen atoms of alkyl groups by a factor of about 600, but for hydrogen atoms of group A, the factor was about 30,000. [A similar selective deactivation of hydrogen exchange in the ring has been observed for the reaction of tolaene with deuterium catalyzed by unsintered and sintered cobalt films (58).] Thus, it was considered that mechanism of exchange in the benzene ring must differ from that in an alkyl side-chain. A problem was thereby posed. The alkyl side-chain can undergo exchange only by a mechanism involving its dissociative chemisorption. Is it tenable, therefore, to suppose that the hydrogen atoms of the benzene ring also become exchanged by a mechanism involving dissociative chemisorption (according to Scheme 3 or 4)if sintering so disproportionately reduces that rate of exchange in the ring? Crawford and Kemball thought not, and accordingly proposed that exchange of hydrogen atoms of the benzene ring occurred by the preliminary addition of a deuterium atom and the subsequent abstraction of a hydrogen atom. The intermediate was


139

conceived to be either (IV) of Scheme 2, or a Ir-bonded intermediate of identical composition which features as species (VI) of Scheme 5 below. Now, supposing that a single metal atom can act as a site for species (VI) , it must necessarily be of low coordination number since two or three “vacant ligand sites” will be required for the establishment of a ligand with such extensive electron delocalization. Such metal atoms must, therefore, occupy metastable situations in the environment of the surface, and a drastic reduction in their concentration is to be expected when the surface is sintered. Further evidence that the sites for the exchange of hydrogen atoms of group A are present in very low concentrations a t the sintered surface is seen in the fact that not only are the observed rates low, but also the activation energy for their exchange is some 5 kcal mole-’ lower than for those of group C. The mechanism of hydrogen exchange in the benzene ring was developed further by Harper and Kemball in their account of the exchange and

Scheme 6. Hydrogen exchange in benzene by an addition-abstraction mechanism involving associative r-adsorption of the reactant [Harper and Kemball (SS)].

140


hydrogenation of para-xylene catalyzed by palladium, tungsten, and platinum (39). Intermediate (VI) was envisaged as being formed and removed in two types of process, one involving molecular deuterium, either gaseous or physically adsorbed, and the other atomic hydrogen (Scheme 5). Both processes are required in order to achieve exchange of hydrogen for deuterium in benzene. Identical intermediates were proposed at the same time by Hartog, Tebben, and Weterings (40) to account for hydrogen exchange in benzene catalyzed by ruthenium, palladium, and platinum. Unfortunately, exchange was slow in comparison with hydrogenation, and so the kinetic behavior of the former could not be measured. These authors rejected the dissociative mechanism [Eq. (3)] on the ground that it was inconsistent with their observed orders for the nickel-catalyzed reaction (36)-a somewhat slender argument, since the catalytic behavior afforded by one metal is not necessarily mirrored in that of its neighbor in the periodic table. Nevertheless, the authors were able t o account for their observed distributions of deuterium in benzene on the basis of Scheme 5 provided that they assumed that two different types of site are active in the exchange reaction. Catalysis is well known as a field in which apparent contradictions abound in the literature. One interesting example is obtained by comparing the work just described with that reported by van Hardeveld and Hartog concerning the relative rates of hydrogenation and of exchange of benzene catalyzed by various nickel-silicas (41). The weightings of nickel on the support and the reduction conditions were varied so that the mean nickel particle size ranged from about 20 to 200 A on passing from one catalyst to another. Over this size range, the mean coordination number of surface metal atoms in perfect microcrystals is expected to rise from about 7 to a limiting value of 9. Table I11 shows the rates of hydrogenation and exchange referred to unit area of surface. That for hydrogenation is independent of crystallite size, but the exchange rate increased as the nickel TABLE I11 Specijc Activities of Various Nickel-Silicas for Benzene Hydrogenation ( A H )and Exchange ( A E ) Range of crystallite size ( A) -200 mostly < 70 all < 50 all < 50

105 A~ lo5 A E (mole hr-1 m-*) (mole hr-1 m-*)

4.7-5.3 11.0 9.0-12.5 9.5

77-90 20 0.7-3.8 0.28

141

CITEMISORPTION OF BENZENE

particle size increased. This is surely inconsistent with the spectacular loss of activity for exchange observed on sintering nickel films. van Hardeveld and Hartog tentatively attributed their observed increase in exchange rate to the pressure of stacking faults in the larger crystals, a view which was supported by the observation in electron micrographs of twinning in the larger crystals. However, it is not clear how sites created a t stacking faults should differ fundamentally from those at the surfaces of the smallest crystallites. Further investigations of these systems would be valuable, to see whether this apparent paradox can be resolved. Thus, evidence has accumulated in support of hydrogen exchange in benzene by a mechanism involving associatively chemisorbed benzene, and without the necessity to postulate the participation of chemisorbed CeH6. One attractive test of these ideas which, so far as we know, has not been made, would be to repeat, for example, the reaction of para-xylene with deuterium using as catalyst a palladium thimble. This system would allow the exchange reaction to proceed either in the presence of molecular deuterium (both reactants on same side of the thimble) or in the presence of atomic deuterium only (xylene and molecular deuterium on opposite sides of the thimble, so that the hydrocarbon reacts only with chemisorbed atomic deuterium that arrives a t the surface after diffusion through the metal). Careful reading of references (56-40), and of recorded discussion where this exists, indicates that authors who favor exchange by an additionabstraction mechanism seldom reject the alternative entirely. Indeed, since evidence from subsection B supports the abstraction-addition mechanism, it may well be that both mechanisms operate simultaneously when molecular deuterium is present, and that only when one predominates can telling experimental evidence be obtained. Exchange in benzene catalyzed by alloys has been little studied. Reaction a t 41°C over a nickel-copper alloy containing 23 f 401, Ni has been examined by van der Plank and Sachtler ( 4 2 ) . Values of the multiplicity factor M in the range 1.4-1.7 agree with that of 1.6 reported by Moyes and co-workers for nickel films ( 4 ) . The rate of exchange exceeded that of hydrogenation by several orders of magnitude. The poisoning of the surface by dissociatively adsorbed species was noted. The mechanism of exchange was not discussed.

B. REACTIONS OF BENZENE WITH DEUTERIUM OXIDE AND DEUTERIUM-LABELED BENZENE

WITH

We now turn to examine several reports of hydrogen exchange in the benzene ring in which the deuterium source is either heavy water or a

142


deuterium-containing hydrocarbon. In these systems, no hydrogenation can occur, and the concensus of opinion is that exchange occurs via an abstraction-addi tion mechanism. I n the early 1960s, while the above-mentioned exchange reactions employing molecular deuterium were being examined, Garnett and his school were making an extensive study of exchange reactions between aromatic hydrocarbons and deuterium oxide or deuteriated benzene. I n this work the effects upon the exchange rate of substituents in the benzene ring ( @ ) , catalyst preparation (44)) and poisons (46) were studied. Evidence from changes in reactivity within a series of alkylbenzenes, and an observed inverse relation between the effectiveness of various poisons and their ionization potentials strongly supported the proposition th a t associative adsorption as species (I) (ie., ?r-complex formation) occurred. The strengths of adsorption of the wide variety of aromatic molecules studied varied by a factor of fifty; this was difficult to understand in terms of the classical mechanisms [Eqs. (2) and (3)] but is interpretable in terms of the likely strengths of the resultant metal-ring ?r-bonds. This work has been summarized by Garnett and Sollich-Baumgartner ( 2 ) and hence will not be reviewed in detail here. Attention will now be confined to one paper (46) in which the rate of isotope exchange between DtO and COHO was compared with the rate of exchange between c6& and C6Ha.These reactions were catalyzed by platinum at 32°C. That the latter reaction occurred a t all was construed as compelling evidence for the dissociation of benzene during the course of, or after, its chemisorption. The process was envisaged to occur as shown in Scheme 6; the primary act of chemisorption is associative, exchange occurs as a result of the dissociation of benzene, and in the proposed transition state the plane of the benzene ring is inclined a t an angle of 45" to the catalyst surface. This mechanism was proposed by analogy with published mechanisms of homogeneous electrophilic aromatic hydrogen exchange ( 4 7 ) . For a given catalyst, a randomization rate constant for the C6HE-CaDsexchange reaction of 5.9 X lop2hr-' was observed, and this compared with a value of 8.60 x lov2hr-' for hydrogen exchange

(n)

(1)

[X = H o r D] Scheme 6. Hydrogen exchange in benzene by an abstraction-addition mechanism involving associative *-adsorption of the reactant [Garnett and Sollich-Baumgartner

(-@)I.


143

between C ~ H and G DzO. From the similarity between these values it was concluded that, since the former reaction must proceed by an abstractionaddition mechanism, so must the exchange between benzene and water. The factor of 1.5 between these two constants was attributed to modification of the physical character of the catalyst by the benzene-water reaction, since the catalyst was “transformed from a coarse and coagulated powder into a finely divided filmlike state” during this reaction. By measuring the relative rates of exchange of deuterium and of tritium the rate-determining step was identified as the recombination of a chemisorbed hydrogen atom and a-adsorbed CeH6. The hydrogen exchange reaction a t 100°C between para-xylene and deuterium oxide catalyzed by cobalt, nickel, ruthenium, rhodium, palladium, iridium, and platinum has been studied by Hirota and Ueda (48). Nickel and cobalt catalyzed exchange only in the methyl groups (although exchange of hydrogen for deuterium in the ring occurred if benzene was used in place of para-xylene) . For the remainder of the metals, exchange in the ring occurred although it proceeded a little more slowly than exchange in the methyl groups. The authors postulate a mechanism involving dissociative chemisorption for the exchange of hydrogen atoms of the methyl groups, but a process in all essential respects identical to Scheme 2 was considered responsible for exchange of hydrogen in the ring. This mechanism is thus open to the criticism already applied to Scheme 2 (above), namely that the geometrical requirements of the step in which a hydrogen atom is removed from species (IV) are very stringent. These two studies of the exchange of aromatic compounds with deuterium oxide have been the subject of much discussion. The claim (46) that the platinum-catalyzed hydrogen exchange between CsH6 and C6D6 must, without doubt, proceed by an abstraction-addition mechanism has been questioned (49) on the ground that the platinum used was formed by a procedure involving reduction in hydrogen, and that the evacuation techniques used were not stringent enough to remove chemisorbed or occluded hydrogen before the admission of benzene. Such hydrogen atoms might then propagate exchange by an addition-abstraction process. According to Pliskin and Eischens (50) some hydrogen may remain chemisorbed to platinum supported on silica even after evacuation for short periods to Torr a t 35°C) whereas the platinum used by Garnett was evacuated, a t best, t o pressures no less than lo-* Torr (44a). Moreover, the comparison of rate constants described above appears to pay little regard to the effect on the rate of the occupation of a fraction of the surface by chemisorbed water. Fraser and Renaud (49), having made these criticisms, examined the platinum-catalyzed hydrogen exchange reaction between deuterium oxide and several monosubstituted benzenes (fluorobenzene, chlorobenzene,

144


anisole, aniline, phenol, and others). The relative rates of exchange a t ortho, metal and para positions were determined, and interpreted in terms of steric effects which were found to be all-important; evidence for electronic effects was not obtained. The authors concluded that an abstractionaddition mechanism was the only one that would interpret adequately the effect of substituents on the relative rates of exchange in the three distinguishable positions of monosubstituted benzenes. Thus, their mechanism is adequately described by Scheme 6 except that they considered that no decision concerning the nature of associatively chemisorbed benzene (whether it was T - or di-a-adsorbed) could be made. To summarize, the use of heavy water as a deuterium source has provided a wealth of experimental information. Evidence for the associative r-adsorption of benzene [species (I)J is secure ( 2 ) .Evidence for hydrogen exchange in the benzene ring by an abstraction-addition mechanism is less well established, partly because of uncertainties that surround the mode of chemisorption and reaction of water at metal surfaces. Nevertheless, it would be wrong to deny that Scheme 6 is consistent with a large body of experimental work. The complexities of using heavy water as a deuterium source having thus been appreciated, attention has been directed once again to the incorporation of the deuterium label within the aromatic compound itself. Hirota and co-workers (61) have examined the isotopic redistribution that occurs at 100°C when monodeuteriotoluene containing deuterium in the ortho, or meta, orpara position is admitted to powdered nickel or platinum catalysts. The catalysts were pumped for unspecified periods before use; platinum was examined a t 150”C, but no temperature was quoted for nickel. Deuterium became distributed throughout the toluene molecules, both in the methyl group and in the ring. The mechanism proposed for hydrogen exchange in the benzene ring was essentially that shown in Scheme 6. It is important, however, to note the grounds on which an addition-abstraction mechanism was ruled out. These authors observed that, since there was no net transfer of hydrogen from the catalyst to toluene, and no net loss of deuterium from toluene to the catalyst, “the role of occluded hydrogen in the catalyst, if present, can be ruled out from discussion.” Unfortunately, this is not so. The steady state concentration of chemisorbed hydrogen atoms required to propagate exchange by the associative mechanism might be extremely low; certainly it cannot be supposed that its concentration would have been detectable if it had appearcd in the 0.5 gm of hydrocarbon used. Thus, a n assessment of this work turns on a value judgement as to the likelihood that the catalysts were hydrogen-free. In the reviewers’


145

opinion, the results are likely to be valid because (i) the methods of catalyst production did not involve reduction of salts in molecular hydrogen (platinum was prepared by Willstatter’s method, and nickel b y decomposition of the formate), and (ii) of the Group VIII metals nickel and platinum occlude the least hydrogen (38). Moyes and co-workers ( 4 ) have examined the hydrogen exchange reaction that occurs a t 0°C when equimolar mixtures of C6H6and CsD6 are admitted to a wide range of evaporated metal films. The rates of entry of deuterium into CeH6,to give CsH6D, C6H4D2,and C6H3D3,and of hydrogen into C6D6 to give C6HDs, C6HD4,and C6H3D3 were measured, it being assumed that equal proportions of C6H3D3were formed by each process. Suitably modified equations of the type presented in the footnote (p. 136) allow the calculation of values for k F and k ~the , velocity constants for the initial rate of entry of deuterium or of hydrogen respectively into 100 molecules of benzene per minute per milligram of catalyst a t the beginning of the reaction. The sequence of activity, as presented by the values of k F was Rh

> I r > Mo > Re > W = Co > Ni = Fe > Pt > Mn > Cr > Pd > T a > V > Ti > Ag

No exchange was observed in the range 0”-200°C a t the surfaces of copper, hafnium, or gold. The films were formed under carefully controlled conditions. Wires were rigorously degassed before evaporation, and films were thrown in pumped vessels at pressures in the range 10-6-10-7 Torr, (for Rh, Ag, and Re, 10-9 Torr). The apparatus contained no greased taps so that contamination by adventituous hydrocarbon was avoided. I n this way the authors endeavored to ensure that the surfaces so obtained were free of chemisorbed hydrogen atoms. Confirmation that this was so was obtained from the mass balances. The quantity of benzene used in each reaction (30 micromoles) was that required to form about ten monolayers; thus, the presence of a very small fraction of a monolayer of chemisorbed hydrogen atoms a t the surface of a newly formed film would have been detectable in terms of a change in the hydrogen:deuterium balance of the gas phase benzene early in the reaction. No such displacement of the massbalance was observed. The values of the multiplicity factor M for the “forward” reaction (the exchange of hydrogen in C6H6for deuterium) were Ti, 3.0; Ta, 2.0; Ir, 2.0; Co, 1.9; Mo, 1.9; Ag, 1.9; Re, 1.9; Rh, 1.8; W, 1.8; Mn, 1.7;V, 1.7; Cr, 1.6; Fe, 1.6; Ni, 1.6;Pt, 1.2; and Pd, 1.0. Thus, multiple exchange occurred at each metal surface with the exception of

146


25

30

35

Percentage d- character

45

40

50

3

FIG.4.Hydrogen isotope exchange between CGHGand CGD,. Correlation of randomisa, percentage d-character of the metallic bonds ( 4 ) . tion rate constant k ~ with

palladium and platinum. These last-mentioned metals were exceptional in a further respect, namely that their surfaces became poisoned, probably by highly dissociated forms of benzene, as reaction proceeded. A linear corrclation of the logarithm of k~ with the percentage d-character of the metalmetal bonds was observed (Fig. 4).It must be remembered that values of ICF refer to unit weight of film and not to unit surface area. However, for Ti, Cr, Mn, Fe, Co, and Ni, the linear correlation with percentage d-character also holds when kF is referred to unit surface area as measured by the

147


chemisorption of benzene. Surface areas of films of the transition elements of the second and third series were not measured. The conclusions from this work were (i) that the mechanism that operates is of wide applicability, (ii) that exchange proceeds by either the dissociative chemisorption of benzene OT by the dissociation of benzene which has previously been associatively chemisorbed, and (iii) that M values of about 2 indicate that further dissociation of a-adsorbed-C6H6 to give di-a-adsorbed-C6H4 occurs. The process shown in Scheme 7 is that presented in Scheme 4 with the inclusion of species (I). Evidence for the formation of (I) was obtained from surface-area measurements. Surface areas of metal films determined by the chemisorption of hydrogen, oxygen, carbon monoxide, or by physical adsorption of krypton or of xenon concur

(1)

* (n)

*

‘* (V)

[X = H or D]

Scheme 7. Hydrogen exchange in benzene by double abstraction-addition, benzene being initially associatively chemisorbed [Moyes et al. ( 4 1 .

with those obtained by benzene chemisorption provided it is assumed that the area of surface occupied by a chemisorbed benzene molecule is 42 ( 6 2 ) .This value is usually that associated with a benzene molecule chemisorbed with its plane parallel with the surface, and hence it is concluded that ?r-bonded benzene may well achieve high surface coverage and that the intermediates in the exchange process are present in low concentration. Alternatively, the surface area occupied by the a-adsorbed species will approach the value of 42 Az, because of the “thickness” of the Ir-electron system. Thus, this work should be considered to demonstrate only the formation of entities by the dissociative adsorption of benzene. When the logarithm of k F is plotted against metallic radius (Fig. 5) a correloation is observed for those elements with radii in the range 1.351.45 A. The correlation does not extend to those Zlements of the first transition series for which the radii are less than 1.30 A (with the exception of titanium, which obeys the correlation). This correlation lends some support t o the view that there may be a critical intermediate in the exchange process the facile formation of which requires the matching of

A2

148


0 co 0 Fe 0 Ni

0 Mn 0 Cr

.V

I

I

I

0.13

0.14

0.1!

Metallic Radius (nm) -b

FIG.5. Hydrogen isotope exchange between C6H6and CGDG. Correlation of randomization rate constant k ~with , metallic radius ( 4 ) .

the geometry of the intermediate to the interatomic distances available in the metal.

IV. Some Aspects of Benzene Hydrogenation I n the two previous sections, evidence has been presrnted conccrning the chemisorbed states formed when benzene interacts with metal surfaces. It is not the intention in this Section to discuss benzene hydrogenation in detail, but rather to enquire whether studies of this hydrogen-addition reaction provide information about the chemisorbcd state of benzene.


149

Benzene hydrogenation is generally found to be of about the first order in hydrogen and of approximately zero order or of slightly negative order in hydrocarbon. In this respect it is a typical example of a metal-catalyzed hydrogenation of an unsaturated hydrocarbon. Moreover, in the region of room temperature, the unsaturated products of its hydrogenation are themselves usually hydrogenated very rapidly indeed, and hence they may not be formed in measurable quantities. This c6 system is unusual, however, in that the reverse process, namely the conversion of cyclohexane to benzene, proceeds virtually to completion a t about 300°C and atmospheric pressure; such is not the case for the Cg- or C7-cyclic systems or for the straight-chain hydrocarbons. The question to be asked is this: do the processes of benzene hydrogenation and of hydrogen exchange in benzene involve common intermediates, and in particular do these processes share a common form of chemisorbed benzene? If the answer is in the affirmative, then the relevant surface species are described in Section 11. The work of Anderson and Kemball (35) reported in Section I11 concerning the reaction of benzene with molecular deuterium catalyzed by evaporated films of platinum and of palladium included an examination of the kinetics of cyclohexane formation. The kinetic form of the hydrogenation reaction differed from that for the exchange reaction. Moreover, a t the palladium surface, cyclohexane formed inhibited the rate of hydrogen exchange in the benzene ring without influencing the rate of hydrogenation. It was thus concluded that the processes of hydrogenation and exchange occurred by separate mechanisms the former involving CeX7(ads) and the latter CsXs(adS) [X = H or D]. The deuterium distribution in the products was consistent with this view, although the distribution of deuterium in benzene has since been shown to be consistent also with a mechanism involving C6H7(ads) as the intermediate (40). Just as Garnett ( 2 ) argued for the participation of ?r-bonded aromatic hydrocarbons in the exchange reaction from considerations of ionization potential data, etc. so somewhat analogous arguments have been advanced with respect to benzene hydrogenation. Volter (53) examined the hydrogenation of several aromatic hydrocarbons using nickel supported on magnesium oxide as catalyst. The temperature range was 90"-200°C. It was observed that the activation energy for hydrogenation E , diminished with the first ionization potential of the hydrocarbon and changed with the stabilities of the corresponding complexes of the aromatic hydrocarbons with hydrochloric acid, picric acid, or iodine. According to this argument, those aromatics which can establish the strongest r-bond at the catalyst surface (mesitylene in the present context) should be hydrogenated with the lowest activation energy, as observed (see Table IV) . Thus, the mecha-

150


TABLE IV Activation Energies for Hydrogenation, and First Ionization Potentials, for some Aromatic Hydrocarbons

Hydrocarbon Benzene Toluene Ethylbenzene p-Xylene Mesitylene

Ionization potential (eV)

Activation energy (kcal mole-') Ni-MgO

Rh-MgO

Co-MgO

9.24 8.82 8.77 8.44 8.39

14.2 13.5 10.4 11.1 8.0

9.2 7.4 8.7 8.0

11.3 13.1

-

14.2 14.9

nism of hydrogenation was thought to be as proposed by Rooney and Webb (54) and shown in Scheme 8. This was supported by Shopov and Andreev (55) who demonstrated that the change in activation energy reported by Volter correlated well with the energy of bonding of the hydrocarbon to the surface as calculated by molecular orbital theory. However, this whole matter must be viewed with caution because the activation energies subsequently reported by Volter and co-workers ( 5 ) for reaction over rhodium and over cobalt do not fall smoothly on passing from benzene to mesitylene (see Table IV) .

*

*

*

tl Scheme 8. Mechanism for the hydrogenation of benzene [Rooney and Webb (541.


151

Evidence was presented in Section 11, from experiments in which the activities of unsintered and sintered films were compared, that hydrogen exchange in benzene requires special sites. There is complementary evidence that the hydrogenation of benzene is not demanding as to site requirement. The report of van Hardeveld and Hartog ( 4 l ) ,that the specific activity of nickel for benzene hydrogenation does not depend upon crystallite size within the range 20-50 A where mean coordination number varies markedly has been mentioned in Section 11. Aben et al. (56) have confirmed and extended this finding. For nickel, palladium, and platinum supported on a variety of refractory oxides, the activity per exposed metal atom was found to be independent of the metal crystallite size over the range lO-200& and independent of the support used. Thus, the exchange reaction is clearly more demanding as regards the constitution of necessary sites than is hydrogenation. Thus it is certainly established that exchange and hydrogenation are not different aspects of a common process under all conditions. Some reports of benzene hydrogenation record that poisoning of the catalyst surface by hydrocarbon residues occurs, resulting in a diminution of hydrogenation activity. For example, for nickel and tungsten films the rate and extent of poisoning each increased with increasing temperature ( 6 7 ) .Kubicka (68) has observed that hydrogenolysis accompanies benzene hydrogenation over alumina- and silica-supported ruthenium, technetium, and rhenium. The products of the ruthenium-catalyzed reaction, for example, were mostly hexane, pentane, and butane in the range 180”-195°C and propane, ethane, and methane above 200°C. No such products were observed when palladium and platinum were used. Wells and Bates (59) have reported that iridium wire at 245°C loses its activity for hydrogenation and acquires activity for the hydrogenolysis of several unsaturated hydrocarbons of low molecular weight, including benzene. Clearly, a t these elevated temperatures, the dissociative chemisorption of benzene extends not only to the rupture of carbon-hydrogen bonds but also to that of the carbon-carbon bonds. It is tempting to suppose that, on these less well studied metals of groups VII, VIII1, and VII12,the “carbonaceous residues’’ referred to by Anderson (5 7 ),or the highly dissociated species reported by Silvent and Selwood (17’) can react with hydrogen and leave the surface as identifiable products. Kubicka further reported that the specific activities of the metals for benzene hydrogenation fell in the sequence Ru > Pt > Tc = P d > Re. We note that, for the elements of the second transition series, the maximum activity was observed for the element of group VIIIl (group VIIIz was not studied). This should be compared uith the results in Fig. 4 which show that the activities for the exchange reaction pass through a maximum a t

152


about group VIIIz for each transition series (group VIIIl only being represented by iron). There is thus, apparently, a common pattern of activity for hydrogenation and for exchange. Whether or not this constitutes evidence for a common mechanism cannot be so simply determined, but the common pattern should not be ignored especially if future work confirms Kubicka’s findings. Thus, interpretations of benzene hydrogenation do not require the formation of additional adsorbed states of benzene.

V. Conclusions The following conclusions may be drawn from the foregoing sections (only leading references quoted). (i) Associative chemisorption of benzene as a ?r-complex occurs ( 2 ) . (ii) Chemisorption of benzene on clean metals in the absence of molecular hydrogen leads t o the fission of at least two carbon-hydrogen bonds a t the surfaces of the majority of transition elements ( 4 ), Scheme 7. (iii) Further dissociation of benzene has been detected by magnetic measurements (17 ), field-electron emission microscopy (21) and LEED (22-29) and may be inferred from the behavior of adsorbed I4C-labeled benzene (4, 8-11). Such further dissociation increases in extent as the temperature is raised, but varies in extent from one metal to another at a given temperature. This is realized by observations of poisoning and of hydrogen01ysis. (iv) Chemisorption of benzene in competition with molecular hydrogen leads to hydrogenation, Scheme 8. When molecular deuterium is employed, the resulting hydrogen exchange in benzene can be interpreted in terms of the reversible formation of CaX7(ads)from benzene (39, 40) provided a modified Rideal-Eley mechanism operates, Scheme 5 . (v) Hydrogen exchange in benzene that accompanies hydrogenation depends on the crystallite size of the metal (41) or the degree of sintering of the catalyst (37). Thus, this process may be “structure-sensitive” according to the terminology of Boudart (60). (vi) The aromatic character of the benzene ring is retained during exchange via the dissociative chemisorption of benzene. On the other hand, the number of dclocalized *-electrons is reduced from six to five and then restored to six during exchange in the mechanism described in Scheme 5 . (vii) The hydrogenation of benzene does not require the formation of a special chemisorbed state of benzene. However, the possibility must not be overlooked that dissociatively adsorbed species derived from benzene


HYDROGENATION

EXCHANGE e.g. in the C,H,-C,D,

C, H,

system

SPECIES RESPONSIBLE FOR CATALYST POISONING

153

HYDROGENOLYSIS

in the - D, system

FIG.6. Schematic representation of the range of chemisorbed species formed from benzene and the reactions that benzene undergoes at a transition metal surface.

may be hydrogenated to cyclohexane via routes and involving species not discussed here. All of these processes are displayed schematically in Fig. 6. The adsorbed species responsible for poisoning have not been determined experimentally and hence are represented by a query. Similarly the precursors of the hydrogenolysis products are not known; the methine groups shown should be regarded merely as an example. (viii) These conclusions, although apparently of wide validity, are inevitably influenced b y the fact that the majority of studies of benzene chemisorption or exchange have employed nickel or platinum as adsorbent or catalyst. Further studies utilizing other metals, particularly those of cph or bcc structure, would reveal whether or not these conclusions are a n oversimplication. Finally, a comment must be made concerning the nature of the sites for benzene chemisorption. The description of chemisorbed benzene as a *-bonded species carries with it certain implications as to the nature of the site, implications which the symbolism of the asterisk too easily obscures. In the language of the organometallic chemist, benzene is a six-electron ligand, and would occupy three ligand positions in an octahedral metal complex. Thus, if a single metal atom is to constitute a site for the associative wadsorption of benzene, then that metal atom must have a rather low coordination number. But need the site be a single metal atom? Might not the asterisk signify a site comprising two or more metal atoms of higher coordination number? Certainly, compounds of the type shown below have been reported ( G I ) , and such structures might well serve as models for the chemisorption process.

154


L-Pd----Pd-L

Furthermore, ?r-arene complexes of transition metals are seldom formed by the direct reaction of benzene with metal complexes. More usually, the syntheses require the formation of (often unstable) metal-u-aryl complexes and these are then converted to 17-arene complexes. The analogous formation of ?r-adsorbed benzene a t a metal surface via the initial formation of a-adsorbed phenyl, merits more consideration than it has yet been given. It is to be hoped that the recognition and study of structure-sensitive reactions will allow more exact definition of the sites responsible for catalytic activity a t metal surfaces. The reactions of benzene, using suitably labeled materials, may prove to be useful probes for such studies. ACKNOWLEDGMENT We thank Dr. K. Baron for writing a preliminary draft of Section 111, and Professor R. C. Pitkethly for communicating some unpublished work. REFERENCES 1. Bond, G. C., “Catalysis by Metals,” pp. 311-334. Academic Press, New York, 1962. 2. Garnett, J. L., and Sollich-Baumgartner, W. A., Advan. Catal. 16, 95 (1966).* S. Yu, Y.-F., Chessick, J. J., and Zettlemoyer, A. C., J . Phys. Chem. 63, 1626 (1959). 4 . Moyes, R. B., Baron, K., and Squire, R. C., 6th Int. Congr. Catal. Palm Beach, 1972 Paper No. 50; J . CataZ. 22, 333 (1971). 6. Volter, J., Hermann, M., and Heise, K., J . Catal. 12, 307 (1968). 6 . Shopov, D., Palazov, A., and Andreev, A., 4th Int. Congr. CataZ., Moscow, 1969, Paper No. 30. 7 . Pitkethly, R. C., and Goble, A. G., Proc. 2nd Int. Congr. Catal. Paris, 1960, Vol. 11, p. 1851 (1961). 8. Tetenyi, P., and Babernics, L., J . Catal. 8, 215 (1967). 9. Babernics, L., and Tetenyi, P., J . Catal. 17, 35 (1970). 10. Brundege, J. A., and Parravano, G., J . Catal. 2, 380 (1963). 11. Parravano, G., J . Catal. 16, I (1970). 12. Galkin, G. A., Kiselev, A. V., and Lygin, V. I., Trans. Faraday SOC.60, 431 (1964). 13. Ron, A., Folman, M., and Schepp, O., J . Chem. Phys. 36,2449 (1962). 14. Palazov, A., Andreev, A., and Shopov, D., C . R. Akad. Bulgare Sci. 18, 1145 (1965). 16a. Erkelens, J., and Eggink-du Burck, S. H., J . Catal. 15, 62 (1969).

* A further review [J. L. Garnett, Catal. Rev. 5 , 229 (1972)l has appeared since the completion of this article.


155

16b. Sheppard, N., Avery, N. R., Clark, M., Morrow, B. A., Smart, R. St. C., Takenaka, T., and Ward, J. W., Proc. Conf. Mol. Spectrosc., 4th, 1968. p. 97. Institute Petroleum, London, 1969. 16. Selwood, P. W., J . Amer. Chem. Soc. 79, 4637 (1957). i 7 . Silvent, J. A., and Selwood, P. W., J . Amer. Chem. SOC.83, 1033 (1961). 18. Suhrmann, R., Advan. Catal. 7, 303 (1955). 19. Suhrmann, R., Kruger, G., and Wedler, G., 2.Phys. Chem. 30, l(1961). 20. Gryaenov, V. M., Shimulis, V. I., and Yagodovskii, V. D., Dokl. Akad. Nauk SSSR 132,1132 (1960). 21. Condon, J. B., Diss. Abstr. B 29(4), 1317 (1968); Ph.D. thesis, Iowa State University, Ames Iowa, 1968. 22. Edmonds, T., McCarroll, J. J., and Pitkethly, R. C., paper presented at the Discussion on Carbon Deposition on Metals, Glasgow, March 1972. 23. Hopkins, K. N., Duckworth, R., and Pitkethly, R. C., in press. 24. Dalmai-Imelik, G., and Bertolini, J. C., paper presented at the International Conference on Solid Surfaces, Boston, 1971. 26. McCarroll, J. J., Edmonds, T., and Pitkethly, R. C., Nature (London) 223, 1260 (1969). 26. Edmonds, T., McCarroll, J. J., and Pitkethly, R. C., Ned. Tijdschr. Vacuumtech. 8, 162 (1970); J . Vuc. Sci. Technol. 8, 68, (1971). 27. McCarroll, J. J., and Thomson, S. J., J . Cata2. 19, 144 (1970). 28. Pitkethly, R. C., private communication, 1972. 29. Pitkethly, R. C., i n “Chemisorption and Catalysis” (P. Hepple, ed.), p. 98. Inst. Petroleum, London, 1971. 30. Horiuti, J., Ogden, G., and Polanyi, M., Trans. Faraday SOC.30, 663 (1934). 31. Horiuti, J., and Polanyi, M., Trans. Faraday SOC.30, 1164 (1934). 32. Farkas, A., and Farkas, L., Trans. Faraday SOC.33,678 (1937); 33,827 (1937). 33. Taylor, T. I., i n “Catalysis” (P. H. Emmett, ed.), Vol. V, pp. 257-403. Van Nostrand-Reinhold, New York, 1957. 34. Kemball, C., Advan. Catal. 11, 223 (1959). 36. Anderson, J. R., and Kemball, C., Advan. Catal. 9, 51 (1957). 36. Hartog, F., Tebben, J. H., and Zweitering, P., Proc. 2nd Int. Congr. Catal. Paris, 1960, Vol. I, p. 1229 (1961). 37. Crawford, E., and Kemball, C., Trans. Faraday SOC.58, 2452 (1963). 38. Moyes, R. B., and Wells, P. B., unpublished work. 39. Harper, R. J., and Kemball, C., Proc. 3rd Int. Congr. Catal. Amsterdam, 1964, Vol. 11, p. 1145 (1965). 40. Hartog, F., Tebben, J. H., and Weterings, C. A. M., Proc. 3rd Int. Congr. Catal. Amsterdam, 1964, Vol. 11, p. 1210 (1965). 41. van Hardeveld, R., and Hartog, F., 4th Int. Congr. Catal., Moscow, 1968, Paper No. 70. 42. van der Plank, P., and Sachtler, W. M. H., J . Catal. 12, 35 (1968). 43. Garnett, J. L., Henderson, D. J., Sollich, W. A., and Tiers, G. V. D., Tetrahedron Lett. 15,516 (1961); Garnett, J. L., and Sollich, W. A., Aust. J . Chem. 14,441 (1961). 44. (a) Garnett, J. L., and Sollich, W. A., J . Catal. 2,339 (1963); (b) J . Phys. Chem. 68, 436 (1964); (c) Calf, G. E., and Garnett, J. L., ibid. 68, 3887 (1964). 46. Garnett, J. L., and Sollich, W. A,, Aust. J . Chem. 15, 56 (1962); Ashby, R. A,, and Garnett, J. L., ibid. 16, 549 (1963); Calf, G. E., and Garnett, J. L., J . Catal. 3,461 (1964); Garnett, J. L., and Sollich-Baumgartner, W. A., Aust. J . Chem. 18, 993 (1965).

156


46. Garnett, J. L., and Sollich-Baumgartner, W. A., J . Phys. Chem. 68, 3177 (1964). 47. Melander, L., Spec. Publ. Chem. SOC.(London) 16, 77 (1962). 48. Hirota, K., and Ueda, T., Bull. Chem. SOC.Jap. 35, 228 (1962); Proc. Srd Int. Congr. Catal., Amsterdam, 1964, Vol. 11, p. 1238 (1965), and references contained therein. 49. Fraser, R. R., and Renaud, R. N., J . Amer. Chem. Soe. 88, 4365 (1966). 60. Pliskin, W. A., and Eischens, R. P., 2. Phys. Chenz. (Frankfurt am Main) 24, 11 (1960). 61. Hirota, K., and Ueda, T., Tetrahedron Lett. 2351 (1965); Hirota, K., Ueda, T., Kitayama, T., and Itoh, M., J. Phys. Chem. 72, 1976 (1968). 6%’. Baron, K., Ph.D. Thesis, University of Hull, Hull, England, 1971. 65. Volter, J., J . Catal. 3, 297 (1964). 64. Rooney, J. J., and Webb, G., J . Catal. 3, 488 (1964). 66. Shopov, D., and Andreev, A., J . Catal. 6 , 316 (1966). 66. Aben, P. C., Platteeuw, J. C., and Stouthamer, B., 4th Int. Congr. Catalysis, Moscow, 1968, Paper No. 31; published as RecueiZ89, 449 (1970). 67. Anderson, J. R., Aust. J. Chem. 10, 409 (1957). 68. Kubicka, H., J. Catal. 12, 223 (1968). 69. Wells, P. B., and Bates, A. J., J . Chem. SOC.A 3064 (1968). 60. Boudart, M., Advan. Catal. 20, 153 (1969); Wells, P. B., in “Specialist Periodical Reports: Surface and Defect Properties of Solids” (M. W. Roberts and J. M. Thomas, eds.), Vol. I, p. 236. Chemical Society, London, 1972. 61. Allegra, G., Immirai, A,, and Porri, L., J. Amer. Chern. SOC.87, 1394 (1965).

The Electronic Theory of Photocatalytic Reactions on Semiconductors TH. WOLKENSTEIN Institute of Physical Chemistry Academy of Sciences Moscow. USSR

Introduction. . . . . . . . , . , . . . . . . . . . . . . . . , . . . , . . . . . . . , , . . . , . , . . , . , , I. The Mechanism of the Influence of Illumination on the Adsorption and Catalytic Properties of a Surface. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. Various Forms of Chemisorption.. . . . . . . . . . . . . . . . . . . . . . . . . , , . , B. Relative Content of Various Forms of Chemisorption in the Ab............................ s of Chemisorption on IlluminaC. Relative Content ..................................... tion. A General C D. Relative Content ious Forms of Chemisorption on Illumination. The Case of Strong Excitation. . . . . . . . . . . . . . . . . . . . . . . , . , . 11. The Photoadsorptive Effect. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. Summary of Experimental Data. . . . . . . . . ............... B. Theory of the Photoadsorption Effect. . . . . . . . . . . . . . . . . . . . C. Comparison of Theory with Experiment. . . . . . . . . . . . . . . . . . 111. The Reaction of Hydrogen-Deuterium Exchange. . . . . . . . . . . . . . . . , . . A. Summary of Experimental Data. . . , . . , . , . . . . . . . . . . , . , . . . . . . . . B. The Reaction Mechanism. . . . . . . . . . . . . C. Comparison of Theory with Experiment. . . . . . . . . . . . . . . . , . . . . . . IV. The Reaction of Oxidation of Carbon Monoxide. . . . . . . . . . . . . . . . . . . A. Summary of Experimental Data. . . . . . . . . . B. The Reaction Mechanism. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . C. Comparison of Theory with Experiment. . . . . . . . . . . . . . . . . . . . . . . V. The Reaction of Synthesis of Hydrogen Peroxide. . . . . . . . . . . . . . . . . . . A. Summary of Experiment B. The Reaction Mechanism. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . C. Comparison of Theory with Experiment. . . VI. Conclusions. . . . . . . . . , , . . . . . . . . . . . . . . . . . . , , . . . . . . . . . . . . . . . . . . . . References ..... ... . ., ...

157 158 159 161 164 167 170 171 173 176 179 180 182 185 189 190 191 194 197 197 198 20 1 203 206

Introduction

It is now well known that a catalytic reaction taking place on the surface of a semiconductor can be considerably accelerated (and sometimes re157

158

TH. WOLKENSTEIN

tarded) under the influence of illumination, i.e., when light is being absorbed by the semiconductor. This phenomenon (the change in catalytic activity of a semiconductor under the influence of illumination) is termed here the photocatalytic e$ect. This is a new phenomenon, discovered and experimentally studied relatively recently. It is this effect that we are concerned with in the present article. It should be noted immediately that not all the frequencies absorbed b y a semiconductor are photocatalytically active, but only those that are also photoelectrically active, i.e., that cause an internal photoelectric effect in the semiconductor. Note further that the sign and magnitude of the photocatalytic effect depend on the past history of the specimen exposed to illumination; i.e., they depend on the external influences to which the specimen in question was subjected in the course of the whole of its life, and also on the conditions of the experiment (temperature, intensity of illumination, etc.) . For example, by introducing into the semiconductor an impurity of any concentration or by adsorbing foreign gases on its surface it is possible t o render its catalytic activity more or less sensitive to illumination. Our aim is to disclose the mechanism of the photocatalytic effect. It is necessary first to understand why and how illumination, in general, influences the course of a heterogeneous catalytic reaction by stimulating or, on the contrary, retarding it. One has to understand why the effect is positive in some cases (acceleration of the reaction) and negative in others (retardation of the reaction), and how the sign of the effect is determined. Furthermore, i t is necessary to find out upon what factors, and in what manner, the magnitude of the effect depends. We shall try to answer all these questions. 1. The Mechanism of the Influence of Illumination on the Adsorption and Catalytic Properties of a Surface

A clue to the understanding of the photocatalytic effect is the electronic theory of catalysis on semiconductors ( I ) . As will be seen later, the existence and the basic regularities of the photocatalytic effect follow dircctly from the electronic theory of catalysis. Whereas the theory of the photoadsorptive effect (the influence of illumination on the adsorption capacity of a semiconductor) has received much attention in the literature, the theory of the photocatalytic effect based on the electronic theory of catalysis has almost escaped the attention of investigators. The purpose of the present work is to fill in the gap to a certain extent. We shall naturally start by recalling certain principal concepts of the electronic theory which will be needed later.

THE PHOTOCATALYTIC EFFECT

159

A. VARIOUS FORMS OF CHEMISORPTION According to the electronic theory, a particle chemisorbed on the surface of a semiconductor has a definite affinity for a free electron or, depending on its nature, for a free hole in the lattice. In the first case the chemisorbed particle is presented in the energy spectrum of the lattice as an acceptor and in the second as a donor surface local level situated in the forbidden zone between the valency band and the conduction band. In the general case, one and the same particle may possess an affinity both for an electron and a hole. In this case two alternative local levels, a n acceptor and a donor, will correspond to it. By capturing an electron or a hole the chemisorbed particle passes from the electrically neutral to the charged state. It is very important that the trapped electron or hole is forced to take part in the chemisorption bonding. Thus, three forms of chemisorption should be distinguished: (a) The neutral form realized without the participation of a free electron or hole. This form is usually called the “weak” form. (b) The negatively charged form involving a free electron from the crystal lattice, localized on the chemisorbed particle. This is the so-called “strong” acceptor form. (c) The positively charged form involving a free hole localized on the chemisorbed particle. This form of chemisorption is termed the “strong” donor form.

It is important that these forms differ in the strength of the chemisorption bonding, i.e., in the heat of adsorption. The charged form is always stronger than the neutral form. Indeed, in the first case, unlike the second, desorption must be accompanied by the delocalization of an electron or hole; this is always an endothermic process. It is also essential that in certain cases the charged form is practically an irreversible form [see, e.g., reference ( a ) ] .By subjecting the specimen to evacuation, we remove the neutral form from the surface, while the charged form remains practically on the surface (it leaves the surface very slowly). The desorption of a particle being in the charged state is an act, in which an electron (or hole) localized on the chemisorbed particle is delocalized and the particle itself becomes neutral and leaves the surface. It is this hindered delocalization of the electron (hole), i.e., the discharging of the charged particle, that is responsible for the fact that the charged form of chemisorption often assumes the role of a practically irreversible form. It is also of importance that among the various forms of chemisorption there are, besides valency-saturated, radical forms in which the chemisorbed

160

TH. WOLKENSTEIN

FIG.1. “Weak” (a) and “strong” (b, c) chemisorption forms of the H atom.

particle retains unsaturated valency. This is particularly important to catalysis because, unlike the valency-saturated forms, the radical ones are much more reactive. Note that the free electrons and holes of a crystal lattice involved in chemisorption and catalytic processes play the role of free valencies (positive and negative, respectively). In a number of cases it is these electrons and holes that provide the appearance of radical forms on the surface. Figures 1, 2, and 3 show, as examples, different forms of chemisorption (presented as valency lines) for H and 0 atoms and the CO, molecule, respectively. Figure l a shows the “weak” (electrically neutral) form of chemisorption of a H atom; the chemisorption bond, as can be illustrated, is provided in this case by an electron of the H atom which is drawn, to a greater or lesser extent, from the atom into the lattice; this is the radical form of chemisorption. The “strong” acceptor and donor forms are presented in Fig. 1 (b and c , respectively) ; these are electrically charged and valency-saturated forms. The “weak” and “strong” acceptor forms of chemisorption of an 0 atom are shown in Fig. 2 (a and b, respectively). I n the first case the chemisorbed particle, is a dipole with a negative pole directed outward (Fig. 2a) ;this is an electrically neutral formation as a whole, it being valency-saturated. I n the second case (Fig. 2b) the chemisorbed particle is a negative ion radical.

FIa. 2. “Weak” (a) and “strong” (b) acceptor forms of chemisorption of the 0 atom.


161

FIG.3. “Weak” (a) and “strong” (b, c) chemisorption of the CO, molecule.

Figure 3 shows different forms of chemisorption for a COZ molecule. In the “weak” form of chemisorption the COZmolecule is bound to the surface by two valency bonds, as shown in Fig. 3a. This is a n example of adsorption on a Mott exciton which is a pair of free valencies of opposite sign (i.e., an electron-hole pair). This may be either a free exciton wandering about the crystal or a virtual exciton generated in the very act of adsorption. As seen from Fig. 3a, in the case of the COZ molecule the “weak” form of chemisorption is a valency-saturated and electrically neutral form. As a result of electron capture, this form is transformed into a “strong” acceptor form shown in Fig. 3b, while as a result of hole capture it becomes a “strong” donor form shown in Fig. 3c. Both these forms are ion-radical ones. It should, however, be noted that the ion-radicals formed in these two cases are quite different and, having entered into a reaction, may cause it to proceed in different directions. OF CHEMISORPTION IN B. RELATIVECONTENTOF VARIOUSFORMS ABSENCEOF ILLUMINATION

THE

Suppose that N particles of a definite species are chemisorbed on unit surface. Of these NO, N - and N + particles are, respectively, in the electrically neutral, and the negatively and positively charged states. Obviously,

No

+ N- + N+ = N .

The quantities

NO/ N , 7- = N - / N , 7+ = N + / N (1) characterize the relative contents of the various forms of chemisorption on the surface. These quantities play an important role in the electronic theory of chemisorption and catalysis. It is obvious that 70

=

$+7]-+q+=

1.

(2)

162

TH. WOLKENSTEIN

If an electronic equilibrium is set up on the surface, the parameters

qo,

v-, and q+ are strictly fixed. Their values are determined by the position of the Fermi level at the crystal surface, which will be characterized here by the quantity e,- or ca+. These latter quantities are the distances from the Fermi level to the bottom of the conduction band or, accordingly, to the top of the valency band in the plane of the surface. Evidently, c,

+

= u,

(3) where u is the width of the forbidden region between the bands. Let us find the dependences of qo, q-, and q+ on cg- or e,+. According to Fermi statistics, we have

- v-

N-

€+,

-1

)'

N+ No+N+

=

(1

+ exp

kt

(4) where k is Boltzmann's constant; T is the absolute temperature; v- is the distance from the acceptor level A, representing the particle in question, to the bottom of the conduction band; W+ is the distance from the donor level D, corresponding to the same particle, to the ceiling of the valency band. (See Fig. 4 which shows the energy spectrum of the crystal: the xf axis is directed into the bulk of the crystal at right angles to the surface, which is assumed to coincide with the x = 0 plane; FF is the Fermi level; the bands are shown bent near the surface since the crystal surface bears, as a rule, a charge of adsorption or "biographical" origin.)

F

F" F

FP

5

F"

x=o FIQ.4. Energy spectrum of a crystal with acceptor level A and donor level D representing a chemisorbed particle.

163


FIG.5. Dependence of 7 0 , ?-, and ?+ of various forms of chemisorption on the position of the Fermi level.

From Eqs. (4),on the basis of (1) , one has the required relationships: qo =

+ exp[-

{1

q- = qo expcq+ =

qoexp[-

(6,-

- v-)/kT]

+ exp[-

(e,+

- v-)/kT], (s+ - w + ) / k T ] ,

(5)

(6,-

where

v-

+ v+

=

w-

- w+)/kT]]-l,

+ w+ = u.

(6)

The parameters qo, 7-, and q+ as functions of 6,- or c,+ are schematically presented in Fig. 5 in accordance with ( 5 ) . We see that when the Fermi level is displaced from bottom to top in Fig. 5 (ie., as it moves away from the valency band and approaches the conduction band), the quantity qincreases monotonically and q+ decreases monotonically, i.e., the relative number of particles in the negatively charged state increases, and the relative number of particles in the positively charged state decreases. As to the quantity qo characterizing the relative content of the neutral form of chemisorption, it passes through a maximum when the Fermi level is monotonically displaced. Formulas (5) refer to the general case when the chemisorbed particles are both acceptors and donors. In the particular case of acceptor particles, putting (E,+ - w+)/kT = 00

164

TH. WOLKENSTEIN

in (5) one obtains

(1

9' =

q- = ( 1 q+ =

+ exp[-

(E,-

+ exp[(e,-

- v-)/k~])-l,

- v-)/k~]]-l,

(7)

0.

In t.he particular case of donor particles, putting (eB-

- v-)/kT

= co

in ( 5 ) we have

+ exp[q+ = ( 1 + exp[(e,+ qo =

q- =

(1

(E,+

- w+)/kT])-l,

- w+)/kT])-l,

(8)

0.

As we have already noted, the parameters qo, q-, and t+ are of special significance in the electronic theory. They enter into all the basic formulas of the theory. For one thing, they are the quantities on which the adsorption capacity and the catalytic activity of a surface depend. The adsorption capacity of a surface with respect to molecules of a given species is characterized by the total number N of molecules of the particular species retained by unit surface area under the conditions of equilibrium with the gas phase under the given external conditions (i.e., a t a given pressure P and temperature T). An expression for N as a function of qo, q-, and q+ will be derived in Section 11. The catalytic activity in relation to a given reaction occurring on the surface is characterized by the rate g of this reaction, i.e., by the amount of reaction products formed under the given external conditions per unit time on unit surface area. An expression for g has different forms for different reactions. For the reactions of hydrogen-deuterium exchange, oxidation of CO, and synthesis of HzOz, this expression will be derived in Sections 111, IV, and V, respectively.

C. RELATIVE CONTENT OF VARIOUS FORMSOF CHEMISORPTION ON ILLUMINATION. A GENERAL CASE For a crystal illuminated by a photoelectrically active light, the quantities to,q-, and q+ have values different from those for a crystal in the dark. Thus, the effect of illumination is to change the relative content of different forms of chemisorption on a surface for particles of cach particular spccics; in other words, it changes the population of electron and holes on the surface local levels corresponding to chemisorbed particles. A change in the quanti-


165

ties q0, q-, and 7+ under the influence of illumination results in a change in the adsorption capacity and catalytic activity of the surface. Let us determine the quantities qol q-, and 7+ for an illuminated specimen (5, 4 ) . The same quantities for a specimen in the dark are denoted by qoo, 70- and qo+ (hereafter in the text the subscript 0 signifies the absence of illumination). From the condition of electronic equilibrium for the levels A and D, representing a particle of the species under discussion, we have, respectively,

al-No - az-p,N-

=

a3-N-

- a4-n,N0,

a1+No - az+n,N+ = w+N+ - a4+psN0,

(9a) (9b)

where n, and p , are the concentrations of free electrons and holes in the plane of the surface in the presence of illumination. The first term on the left-hand side of Eq. (9a) represents the number of electron transitions from the valency band to the level A referred to unit time and unit surface area (see Fig. 4); the second term corresponds to reverse transitions. The first term on the right-hand side of Eq. (9a) expresses transitions from the level A to the conduction band, while the second term corresponds to transitions in the opposite direction. Equation (9b) describes, in an analogous manner, electronic transitions between the level D and the conduction band (the left-hand side of the equation) and from the level D to the valency band (the right-hand side). From Eqs. (9a) and (9b), one has, respectively,

+ a4-ns)/(a3- + m-p,), (al++ a4+ps)/(a3+ + m+n,).

N-/No =

7-/vo = (a1-

(10a)

N+/No=

7+/vo =

(lob)

A relation between the coefficients q-,az-, a3-, and a4- as well as between and ad+ can be obtained from the conditions of equilibrium prior to illumination, which have the following form (the principle of detailed equilibrium) : a1+, a2+, as+,

al-Noo - a2-psoN0-= ff3-NO- - a4-nsON~0 = 0,

Wa)

c~l+NoO- az+n,oNo+ = ff3+Nof- ff4+psONO0= 0,

(1lb)

where Noo,No-, and No+ are the surface concentrations of neutral, and negatively and positively charged chemisorbed particles; n,o and pSoare the concentrations of free carriers before illumination. From ( l l a ) , on the basis of (1), one has a- =

~ 1 - ( 7 o 0 / 7 0 - )(pso)-'

a4- = a3-

-

(7o-/7o0)

(n,o)-'

where al-

=

where

=

a3-

pl- exp ( -u+/kT), exp( - v - / k T ) .

(12a)

166

TH. WOLKENSTEIN

I n a similar manner, from (1lb ) , we obtain az+ =

a1+( qoo/go+)

where al+ = P1+ exp ( -w-/kT)

(n,o)-l

where a3+ = P3+ exp( -w+/kT).

a4+ = a3+(g0+/qO0) ( p s O ) - l

, (12b)

I n Eqs. (12a) and (12b) it may be assumed, in order of magnitude, that

01-

=

and

p3-

PI+

=

(13)

P3+.

Substituting Eqs. (12a) and (12b) into Eqs. (10a) and (lo b ), respectively, and adopting, according to (12a), (12b) and (5), the notation

p1-

a1- 700

a-=--=(113-

a+ =

qo-

€a~

P2-

a1f qo0

-- =

a3+

exp

qo+

p1+

- exp

E3+

- v+ kT -

)

w-

kT

P3+

'

we obtain, respectively, 4-/v0

= (90-/to0)

t+/q0 =

/.-,

(11O+/17O0)/.+l

where

I n these equations the following notation is used: An,

=

n, - n,o,

Aps

=

p, - p , ~ .

(17)

Evidently, An, and Ap, represent the excesses due to light in the corresponding concentrations. Note that the quantities An,/nSoand Ap,/p,o characterize the degree of excitation and increase with intensity of illumination I . As reported in the literature ( 4 ) ,we have An,/n,o

= y1I

and

AP,/P,o =

YJ

( 18)

(the proportionality coefficients y1 and yz may be ignored here). From Eqs. (15a) and (15b), we have, on the basis of Eq. (2), the follow-


167

ing final result: 90/900

= [l

+ 90-

(p-

9-/90-

= (9°/900)c1’-,

?+/?lo+

= (O0/9O0)c(+.

- 1)

+

90+ (p+

- 1) 1-1,

D. RELATIVECONTENTOF VARIOUSFORMS OF CHEMISORPTION ON ILLUMINATION. THECASEOF STRONG EXCITATION Let us now calculate the coefficients p- and p+ contained in Eqs. (19). We Will confine ourselves to the case of fairly strong excitation, where An,/n,o

>> 1, a-, l/a+,

Ap,/p,o

>> 1, a+,l/a-.

(20)

In this case Eqs. (16a) and (16b) assume the form

The excesses due to light An, and A p , contained in these equations require estimation ( 4 , 5 ).l Assuming that the electron and hole gases are nondegenerate, one has

n,o

=

n,

=

C, exp ( - e,-/kT),

p,o

=

C , exp ( -e.+/kT).

(22a)

p,

=

C, exp ( - e+.JkT).

(22b)

and C, exp ( - e-,,/kT)

,

Here the coefficients C, and C, are of no interest; the meanings of the remaining symbols are clear from Fig. 4, where FF is the Fermi level a t a thermodynamic equilibrium (in the dark); F,F, and F,F, are Fermi quasi levels (in the presence of illumination) for electrons and holes, respectively; V , in Fig. 4 denotes the bending of the bands near the surface ( V , is taken to be greater than zero if the bands are bent upward). Suppose ( 5 ) that the Fermi quasi levels for electrons and holes remain constant throughout the bulk of the crystal (for all 2), as shown in Fig. 4 (the straight lines F,F, and F,F, are horizontal). This occurs with a crystal of fairly small size and with a sufficiently low coefficient of light The quantities AnBand A p 8 have been calculated in references ( 4 , 5 )using different approximations: in reference (4) the excitation is supposed to be weak [the condition (20) is not observed], and in reference ( 5 )any level of excitation is possible.

168

TH. WOLKENSTEIN

absorption. It may be assumed here (see Fig. 4) that = e-vn

e-sn

+ v,,

-

= ,+,€

e+,,

v,.

(23a)

Besides, note that (see Fig. 4)

+ vso,

= c,

=

-

vso,

(23b) where T', and V,o denote the bending of the bands near the surface in the presence and in the absence of illumination. Assume that €8-

€+,

A V 8 = V8 - V,o

EV+

$(v+ w+) the reaction belongs to the class of the so-called donor reactions, i.e., reactions which are accelerated as the Fermi level is lowered. When the region of es+ < +(u+ w+)is reached, the reaction becomes one of the so-called acceptor reactions which are decelerated as the Fermi level is lowered. The rate g of the photoreaction is given by expression ( 6 2 ) , in which, according to (19))

+

+

+

+

+

go = c1 7lo-(cl- - 1) 17o+(cl+ l)-J-'qo0* (64) We shall limit our attention to the case where the Fermi level in an unilluminated specimen is situated fairly deeply below the D level, which can be brought about, for example, when the bands are sufficiently bent upward, as shown in Fig. 8b. This corresponds to the acceptor branch of the curve go = go (e,+), i.e., the hydrogen and deuterium atoms on the surface fulfill, in this case, the role of donors. Here we may suppose that (see Fig. Sa) go- = 0 and qo+ = 1, and expression (64) takes the form

qo = (1//.I+)qo0.

(65)


185

For the photocatalytic effect K we have, according to (52), (62), (63), and (65), K = (TO/TOO) - 1 = (l/p+) - 1. (66) Substitution of (16b) into (66) gives

where a+ has the form (14b). In the case of strong excitation, substituting (31b) into (66), we shall have, instead of (67),

K

=

exp[(2ev- - es-

- w-)/kT] - 1.

(68)

This equation yields the dependence of the magnitude of the photocatalytic effect K on the position of the Fermi level a t the surface (e.-) and in the bulk ( ev-) of the unilluminated specimen. C. COMPARISON OF THEORY WITH EXPERIMENT We shall first consider the influence of various factors on the rate of a dark reaction go, which is implicitly present in formula (63). 1. Pressure

The pressure P is contained in formula (62) not only in an explicit form but also in terms of the parameters es- and es+, as seen from (5), because eB- and E ~ + are, generally speaking, functions of pressure. I n our model, however, es- and es+ may be regarded as independent of P since the surface is supposed to be saturated with hydrogen and deuterium atoms (all the adsorption centers are assumed to be occupied). Thus, the hydrogendeuterium exchange proves, in accordance with (63), to be a reaction of the first order with respect to hydrogen (deuterium), which is consistent with numerous experimental data (see Section 1II.A). 2. Impurities The introduction of an impurity into a crystal causes a displacement of the Fermi level both inside the crystal and, generally speaking, a t its surface [in this case the Fermi level is displaced in the same direction both a t the surface and in the bulk of the crystal, see reference ( I ) ] . This results, according to (63) and ( 5 ) ,in a change of go. A donor impurity displaces the Fermi level upward, while an acceptor impurity shifts it in the opposite direction. The same impurity exerts diametrically opposite influences on the catalytic activity in acceptor and donor reactions.

186

TH. WOLKENSTEIN

The great majority of experimental data (see Section 1II.A) indicate that the hydrogen-deuterium exchange reaction belongs to the class of acceptor reactions ( i t . , reactions that are accelerated by electrons and decelerated by holes). This means that the experimenter, as a rule, remains on the acceptor branch of the thick curve in Fig. 8a, on which the chemisorbed hydrogen and deuterium atoms act as donors. Here a donor impurity must enhance the catalytic activity, while an acceptor impurity must decrease it. This is what actually occurs, as we have already seen (see Section 1II.A). Emphasis should here be placed on the observations of Holm and Clark ( 3 4 ) )according to whom the reaction rate go passes through a maximum when the concentration of a donor impurity is monotonically increased. This maximum may be due, as shown in Fig. 8a, to the transition from the acceptor to the donor branch of the go = go ( E , - ) curve as e8- monotonically decreases.

3. The State of the Surface Any treatment of the surface, in particular, the adsorption of foreign gases on it, causing a change in E,- (i.e., a change in the bending of the bands V , near the surface), must, according to ( 5 ) and (63), lead to a change in go. As a result of the adsorption of a donor gas, we are transferred up the curve go = go(c8-) in Fig. 8a. The adsorption of an acceptor gas, on the contrary, transfers one down this curve. If one remains on the acceptor branch of the curve, this will mean that the catalytic activity must increase when a donor gas is adsorbed and fall upon adsorption of an acceptor gas. This is in accord with much experimental data (see Section 1II.A). Special emphasis must be made on the experiments carried out by Voltz and Weller (35) who observed a fall in activity caused by the adsorption of water which usually acts as a donor. To understand this result, one must suppose that the authors were dealing with the donor branch of the curve in Fig. 8a. Or else that they remained on the acceptor branch but the water molecules acted as acceptors. It should be noted in this connection that the acceptor functions of watcr (the negative charging of the surface upon adsorption of water) had also been observed (before Voltz and Weller) in certain cases by Yelovich and Rlargolis (46). 4. Correlation with Electrical Conductivity

The displacement of the Fermi level downward (increase of ev- and es-) always diminishes the electronic component and increases the hole component of conductivity. The upward displacement of the Fermi level

T H E PHOTOCATALYTIC E F F E C T

187

(decrease of EV and cs-) has an opposite effect. From this follows, as seen from Fig. 8a, a characteristic parallelism between the changes of electrical conductivity and catalytic activity. The changes in catalytic activity and conductivity on the acceptor branch of the curve (Fig. 8a) are directly related in the case of a n n-semiconductor and inversely related in the case of a p-semiconductor. It is this correlation that has been found in many experimental works, as noted in Section 1II.A. We see that the correlation between the electrical conductivity of a specimen and its catalytic activity established by the electronic theory ( I ) must show up distinctly and in fact reveals itself in the case of the hydrogendeuterium exchange reaction. We now turn our attention to a photoreaction. Let us consider the influence of various factors on the photocatalytic effect K , which is contained in formulas (67) or (68) in an implicit form. 5. Impurities

The influence of the treatment of a specimen on the photocatalytic effect can be investigated with the aid of Fig. 9. This figure, which is similar to

i

FIG.9. Sign and magnitude of the photocatalytic effect of the hydrogen-deuterium exchange.

188

TH. WOLKENSTEIN

Fig. 7, shows, according to (68), the isophotocatalytic curves tv- = f (es-) corresponding to different values of K . The curves are numbered in order of increasing K :

K1

< 0 < Kz < Ka < Kq.

The region for which formula (68) is valid is enclosed by a heavy line. The straight line K = 0 divides this region into the areas of the positive and negative photocatalytic effects. The introduction of an impurity into a specimen (accompanied by a change in ev- and ts-) will transfer us from one point to another in Fig. 9. Suppose that when a donor impurity is introduced into the specimen (decrease in e-, and ts-) , we are transferred from the point A to the point B. This involves a decrease in K , as can be seen from Fig. 9. Such a decrease in the photocatalytic effect caused by the addition of donor impurities has been observed by Kohn and Taylor (40) who studied the photoreaction of hydrogen-deuterium exchange on zinc oxide exposed to y radiation. Suppose now that the introduction of an acceptor impurity (increase of e-, and t8-) brings us from the point A to the point C (Fig. 9). This involves an increase in K , as seen from Fig. 9. This is in agreement with the results obtained by the same authors ( d l ) , who observed an increase in the photocatalytic effect on silica gel when acceptor impurities were added to the catalyst, and also with the data of Lunsford and Leland (42) who found that the effect was enhanced on MgO with increasing concentration of V-centers (acceptors). 6. The State of the Surface

A change in the state of the surface accompanied by a change of c,must also exert an influence on the photocatalytic effect. Thus, the preliminary chemisorption of a foreign donor gas causing a fall in cs- (at e y = const) must increase K (transfer from the point A to the point D in Fig. 9). The chemisorption of an acceptor gas accompanied by an increase in es- (at e y = const) must weaken the effect (transfer from the point A to the point E). If the positive effect is observed on a specimen deposited in the hydrogen atmosphere, then after the specimen is calcined in vacuo, this being accompanied by an increase of es-, it is replaced by the negative effect (transfer from the point A to the point F in Fig. 9). Such an inversion (change of sign) of the photocatalytic effect due to the calcination of the specimen in vacuo (after it is annealed in hydrogen) was observed by Kohn and Taylor (40) who worked with hydrides of various metals.

189


7. Pressure Generally speaking, the quantity e,- in (68) depends on pressure P. However, as we have already noted, in our model we may assume that e,- = const. Thus, according to (68) ,K is independent of P. As can be seen from (62) and (63) the order of the reaction upon irradiation must remain the same as in the dark. This agrees with the experimental data (42, &I), according to which the irradiation does not alter the reaction order.

8. Temperature The quantities ev- and e,- may be regarded as constant over fairly wide temperature ranges. Thus, as is evident from (68), the positive photocatalytic effect (the case where 2ev- - e,- - w- > 0 ) must decrease, and the negative effect (the case where 2ev- - e,- - w- < 0 ) must increase (in absolute value) with rising temperature. Indeed, Freund (44) who dealt with the region of the positive effect, observed a decrease in the effect with increase of temperature (the hydrogen-deuterium exchange on zinc oxide in the presence of illumination by ultraviolet light). 9. Intensity oj Illumination

Substituting (18) into (67) yields K = AI/(B

+ CI),

(69) where I is the intensity of illumination and the following notation is adopted [see (14b) and (IS)]:

A = (Y+71 - 7 2 ,

B

= 1

+

C

= yz.

Formula (69) is in agreement with experimental data according to which the velocity of a photocatalytic reaction g a t low radiation doses (CI > B ) reaches saturation, i.e., ceases to be dependent on the intensity of illumination ( 4 5 ) . (Note that it is in this region of saturation that the high levels of excitation, which we discussed above and a t which formula (67) is transformed into (68), are attained.)

IV. The Reaction of Oxidation of Carbon Monoxide The heterogeneous reaction

2co + 0

2

+ 2c02

has received much attention in the literature. This reaction may proceed by different mechanisms depending on the conditions. As has been shown,

190

TH. WOLKENSTEIN

illumination in a number of cases speeds up and sometimes slows down the reaction. Reaction (70) in thc dark has been discussed in the literature ( 1 ) from the viewpoint of the electronic theory of catalysis. The photoreaction (70) has also been considered in the literature (3)) though briefly and purely qualitatively. I n the present article we shall proceed from the mechanism which has been discussed in the literature ( 1 ) as one of the possible mechanisms. Let us examine the influence of illumination on the rate of the reaction [see reference ( @ ) I .

A. SUMMARY OF EXPERIMENTAL DATA The experimental papers devoted to the exidation of CO in the dark will not be considered here. This has been done in a paper by Takaishi (48) and in Germain’s book ( 4 9 ) .We shall limit our consideration to the basic experimental results pertaining to the photocatalytic reaction. (1) A large number of works ( 1 1 , 50-59) have been devoted to the investigation of the dependence of the rate of the photocatalytic reaction (70) on the partial pressures of the reagents. Most investigators (11,46-48,50-52,54-57) came to the conclusion that the reaction of photooxidation of CO is first order with respect to CO and zero order with respect to 0 2 . This result has been obtained, in particular, by Doerfler and Hauffe (57) for a reaction mixture enriched in oxygen; for a reaction mixture enriched in carbon monoxide, however, the same authors observed the zero order for CO and the first order for 0 2 . Steinbach (54) has found that in the case of ZnO and NiO specimens the reaction is first order for CO and zero order for 0 2 , and in the case of Co304 specimens it is first order for CO and of the order of 0.5 with respect to 0 2 . As noted by this author, the order of the reaction for both reagents was thc same as in the dark (as in the case of the hydrogen-deuterium exchange, illumination did not change the order of the reaction). Fujita (59) working with ZnO obtained the zero order with respect to CO and order 0.6 for 0 2 . (2) It has been shown that the irradiation by light in the main absorption band may either accelerate the oxidation of CO [the positive photocatalytic effect ( 1 1 , 50-56)] or decelerate it [the negative photocatalytic effect ( 1 1 , 5 3 ) ] . The magnitude and sign of the effect are determined by experimental conditions. For example, Romero-Rossi and Stone ( l l ) , who worked with ZnO, point out that the magnitude and sign of the effect depend on the ratio of ). the partial pressures of 0, and CO in the reaction mixture ( P o ~ / P c oThe magnitude of the positive effect decreases with increase of this ratio, and at a certain value of PO2/Pco the reaction is retarded by light.

191


(3) It has been shown in a number of papers that the magnitude of the effect can be changed by alloying the sample. Thus, Romero-Rossi and Stone (11) have found that the effect is enhanced on ZnO when an acceptor impurity (Li) is introduced into the specimen. The increase of the effect on CUZOupon the introduction of acceptor impurities (S and Sb) has also been observed by Ritchey and Calvert (58). The addition of a donor (Cr) to ZnO, as reported ( l l ) , lowers the magnitude of the effect. (4) The positive photocatalytic effect has been observed in the works of Doerfler and Hauffe (57) and of Lyashenko and Gorokhovatsky (53)who studied the influence of visible and ultraviolet light on the oxidation of CO on zinc oxide. It has been shown that the magnitude of the effect falls with increasing temperature (at a temperature of about 250°C the absorption of light becomes practically inactive). It should be noted that in some papers (53, 57) the specimens of zinc oxide were preliminarily calcined in an atmosphere of oxygen, i.e., the surface of the catalyst was enriched in the adsorbed oxygen. B. THEREACTION MECHANISM We shall now consider one of the possible mechanisms of the reaction (70). It should be emphasized here that this is one of the possible mechanisms, but not the only possible one. We shall assume that the surface of the catalyst contains chemisorbed atomic oxygen and that it is these chemisorbed oxygen atoms that act, when in the ion-radical state, as adsorption centers for CO molecules. I n this case, during the adsorption of CO molecules, surface ion radicals COzare formed as intermediate compounds, which, after being preliminarily neutralized, are desorbed in the form of COz molecules. The course of the reaction is depicted in Fig. 10 by means of valency lines. Figure 10a shows a chemisorbed oxygen atom in the ion-radical co

I 0

L/ 0=

FIG.10. Mechanism of oxidation of carbon monoxide.

co

192

TH. WOLKENSTEIN

state; Figs. 10b and 1Oc illustrate the negatively charged (radical) and electrically neutral (valency-saturated) forms of chemisorption of a COZ molecule (cf. Fig. 3 ) . Neglecting the adsorption of COz molecules and assuming the surface coverage by C02 molecules t o be insignificant, we have

dNo/dt

=

alPo,(N*o - N o ) 2 - bl(Noo)2- a2PcoN-o

+ bZN-CO,, (71a)

dNco,/dt

=

azPcoN-0 - bzN-co, - cNk02,

(71b)

where N*o is the surface concentration of adsorption centers for oxygen atoms. The first terms on the right-hand sides of Eqs. (71a) and (71b) represent the number of 0 2 and CO molecules, respectively, adsorbed per unit time on unit surface area; the second terms are the number of 0 2 and CO molecules, respectively, desorbed during the same time from the same surface area. It is assumed here that both atoms of oxygen which recombine with each other to give an 0 2 molecule must be in the electrically neutral state [see ( I ) ] . The last term in the right-hand side of Eq. (71b) is the number of C02 molecules that are the product of the reaction and are transferred to the gaseous phase from unit surface area per unit time; obviously, g = cN&o,. (72) Under steady-state equilibrium we have, from (71a) and (71b),

alPo,(N*o - NO)^

=

b1(Noo)2- cN&,

azPcoN-o = b2N-co

+ cN&.

(73)

Assuming here that

bzN-co

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