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ADVANCES IN CATALYSIS AND RELATED SUBJECTS

VOLUME 14

This Page Intentionally Left Blank

ADVANCES IN CATALYSIS AND RELATED SUBJECTS VOLUME 14 EDITED BY

D. D. ELEY

HERMAN PINES

Nottingham, England

Evanston, Illinois

PAULB. WEISZ Paulsboro, New Jersey

ADVISORY BOARD

A. A. BALANDIN Moscow, U.S.S.R.

P. H. EMMETT Baltimore, Maryland

J. H. DE BOER Delft, The Netherlands

J. HORIUTI Sapporo, Japan

G. NATTA

E. K. RIDEAL

M i l a w , Italy

London, England

P. J. DEBYE Ithaca, New York

W. JOST Gottingen, Germany

P. W. SELWOOD Evanston, Illinois

H. S. TAYLOR Princeton, New Jersey

1963 ACADEMIC PRESS, NEW YORK AND LONDON

COPYRIQHT

01963, BY ACADEMICP R E S S INC. ALL RIQHTS RESERVED

NO PART OF THIS BOOK MAY BE REPRODUCED IN ANY FORM BY PHOTOSTAT, MICROFILM, OR ANY OTHER MEANS, WITHOUT WRITTEN PERMISSION FROM THE PUBLISHERS

ACADEMIC PRESS INC. 111 FIFTHAVENUE NEW YORK3, N. Y.

United Kingdom Edition Published by

ACADEMIC PRESS INC. (LONDON) LTD.

BERKELEY SQUARE HOUSE, LONDON W. 1

Library of Congrese Catalog Card Number 49-7768

PRINTED I N THE UNITED STATES OF AMERICA

Contributors

MELVIN CALVIN, Department of Chemistry and Lawrence Radiation Laboratory, University of California, Berkeley, California

GERTEHRLICH, General Electric Research Laboratory, Schenectady, New York

MORRISFREIFELDER, Organic Chemistry Department, Research Division, Abbott Laboratories, North Chicago, Illinois

M. C. HOBSON, JR., Research and Development Department, M . W . Kellogg Company, Jersey City, New Jersey

H. P. LEFTIN,Research and Development Department, M . W . Kellogg Company, Jersey City, New Jersey

L. YA. MAROOLIS,Institute of Chemical Physizs, Academy of Sciences, Moscow, U.S.S.R.

P. MARS, Central Laboratory, Staatsmijnen in Limburg, Geleen, The Netherlands

J. J. F. SCHOLTEN, Central Laboratory, Staatsmijnen, Geleen, The Netherlands P. ZWIETERINO, Central Laboratory, Staatsmijnen, Geleen, The Netherlands

This Page Intentionally Left Blank

Vectors of Advance A Preface We have watched the acquisition of m recent books by the library, have fingered an influx of n copies of scientific journals, and have come away from p scientific meetings with an average of q simultaneous sessions. The number of bits of information N in the mass of knowledge is growing exponentially. Yet, N is just a number, a bare scalar quantity. As we search the mass for progress, for advances, we find ourselves examining its components for directional properties. A report, a publication becomes a vector, its length still proportional to the number of bits of information, but its direction oriented more or less in the direction of an ordinate x which we call progress. It is a simple picture, as long as we need not air any rigorous definition for this quality we have in mind as the x direction. We do not pretend to master such a definition, nor do we know if it is possible to design a universally acceptable one. We are, however, inclined to place on top of the list those developments which are recognizable advances in basic knowledge. Unfortunately, this utterance places us only a very small step closer to having explained ourselves. We could say, ((well,you readers know, of course, what we mean,” but feel obliged to do a IiOtle better. Perhaps one most important qualification for basic knowledge lies in its ability to allow derivation of larger number of bits of new and useful information than it contains. It then is located a t the (‘base” of a spreading tree of knowledge. We are inclined to use what may be more than a formal analog. A single point in coordinate space, which represents an observation, can be covered by a (‘basic” equation of one term, having one coefficient. If we wish to cover two points, we may have to devise an equationa “general theory”-with two terms, define two coefficients. This progression continues logically to embrace any number of observations, by the use of a sufficiently complicated general theory (single equation). Such curve fitting can always be done for a set of existing information. But it may disappoint us some day when we have moved away far enough from the presently conventional arena of data taking. Beyond the explanatory element, therefore, there must be a strong predictive ability. To achieve this end the path toward an explanatory theory must be treaded with the guidance from mechanistic models having vii

viii

PREFACE

(a) minimum complexity and (b) maximum contact with tested scientific principles. It is this essential process which introduces simplicity into the framework of knowledge and in this very manner differentiates it from the results of the mere “fitting” approach. We can see that the element of simplicity is related to the predictive ability of the information package or theory. While this describes one concept of the essence of basic advance, we cannot ignore advances in the development of information which increase our potential to develop basic knowledge. The development of basically new tools which allow deeper experimental probing is an example. One may be inclined to also consider as a possible candidate any development which creates an important new technological potential; but on closer examination a significant enough situation of this type will undoubtedly either be or follow one or the other of the previous two categories; otherwise i t is likely to constitute a purely engineering type advance. So much for our coordinate system, in general; we assume it to be a pretty universal one for any discipline of the physical sciences. It is ours since we span most of those “disciplines.” We hope that the contents of this volume will reflect this and some of the above discussed vector properties. M. Calvin’s discussion of the mechanism of the role of chloroplast in the photosynthetic processes seems to us to contain rare conceptional depth and broadness of implication: It demonstrates the novel electronic-and consequently chemical-properties attained by the ordering of many molecular entities into a macrostructure, somewhere between being the “molecule” of the chemist and the “solid” of the physicist. The tools which G. Ehrlich describes have opened up doors to new insights of gas-surface processes of adsorption and desorption. They approach “clean” experimentation and “clean” surfaces. H. P. Leftin and M. C. Hobson supply a critical view and review of spectroscopic methods and what information they have yielded. An old proving ground of mechanisms is that of formic acid decomposition. Activities a t the 1980 International Conference of Catalysis convinced us that a review of this subject, in calm print, would be interesting a t this time. For this difficult task the Editors are grateful to P. Mars, J. J. F. Scholten, and P. Zwietering. As in previous years, a portion of this volume is devoted to critical expert reviews of specific classes of catalyzed chemical reactions. Oxidation of hydrocarbons is reviewed by L. Ya. Margolis and the hydrogenation of pyridines and quinolines, by M. Freifelder ,

PREFACE

ix

I n reading the present contributions, there are many thoughts that come to mind. The inevitable ultimate involvement in matters of the electronic structure of the catalytic solids is becoming self-evident. M. Calvin’s work provides concrete and fascinating experiments in this arena in which chemistry, biology, and physics must learn to meet when each is forced to search its electronic heritage. I n seeing this arena form, sporadically so far, we cannot refrain from mentioning the daring suggestions of A. Szent-Gyorgyi, (“Introduction to a Submolecular Biology,” Academic Press, New York, 1960), or forget many past and current attempts of the catalytic chemist to search for such connections. The formic acid arena reflects both, the classical chemical attempts to equate surface species with simple bulk compounds of chemistry, as well as the viewing of the solid as a new electronic (and therefore chemical) species. Importantly, the authors of that review point to the danger of generalizing mechanistic observations made with one reaction to “heterogeneous catalysis.” To us such a step would indeed be analogous to the expectation that the reactivity of an acid with a base would supply the answer to chemical reactions, generally. We must realize that right now the specific skills of physics have grown up around a rather narrow choice of chemical materials, the semimetals (e.g., germanium, silicon), of interest in device technology, but almost completely afield from catalytic interests. The catalytic researcher, on the other hand, has tended to stick to his most familiar materials which are relatively too complex for the physicist t o choose, and/or too low in electrical conductivity to interest the device-influenced solid state investigator. There is a bothersome angle to having pointed to the “interdisciplinary” nature of catalysis, While we wish to thereby emphasize our strong relationships everywhere, such description appears superficially to place us nowhere in particular. What is the place of catalytic research when viewed in the context of the changes taking place in our scientific society, in the institutions of learning, and with respect to the methods by which we define areas of scientific endeavor? We are inclined to equate the disciplines of science with the corresponding teaching “departments” of our society’s universities (and other schools of higher learning). I n some countries the scholars are associated within analogous “institutes” of the university, this being a matter only of terminology. Indeed, we looked up discipline in the Encyclopedia Britannica Dictionary and found it defined by (‘a course of study.” Our disciplines were borne, grew up, reshaped themselves and subdivided like organisms, when the times were ripe and when i t became practical for them to do so. Our closest parent, chemistry, has long pro-

X

PREFACE

duced sons and daughters like inorganic chemistry, organic chemistry, physical chemistry, biochemistry, etc. Each has found departmental or institutional stature here and there, when the appropriate elements of research, teaching, and especially human inspiration reached a critical mass. Curiously enough, when a split-off occurs, a narrower segment seems to be removed from its parent, yet it now proceeds to freely and vigorously combine contributions from a broad spectrum of disciplines. Take radiation biology; what a beautiful rainbow of older disciplines are therein united, from theoretical physics to medicine! It becomes increasingly clear that Catalytic Chemistry is a t least as large a course of study as chemistry itself. In fact, last year, we discussed [Advan. Catalysis 13, Preface (1962)] what might well be called a rigorous proof that it must be larger in scope than noncatalytic reactivity alone, even before we add the many additional phenomena, which are peculiar to the special geometry of heterogeneous catalytic systems (e.g., diffusion, heat flow, etc.). With this in mind, and after 14 volumes, some 100 authors, and about 5000 pages of Advances in Catalysis, it is not difficult to see Catalytic Chemistry as a rising discipline in its own right.

September 1963

PAULB. WEISZ

Contents

.

.

CONTRIBUTORS

.

.

V

vii

PREFACE

1. Quantum Conversion in Chloroplasts BY MELVIN CALVIN

.

I. Introduction 11. Energy Sources for Photosynthesis 111. Photochemistry of Chlorophyll IV. Photophysical Effects in Model Systems V. Charge Separation Processes VI. The Relation t o the Photosynthetic Apparatus VII. Quantum Conversion in Biological Material VIII. Conclusion . References

. .

. . .

. I

.

&-

.

3 5 7 15 22 29 32 32

.

.

. . .

.

.

2. The Catalytic Decomposition of Formic Acid BY P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

.

.

.

.

I. Introduction 11. General Aspects of the Reaction 111. Dehydrogenation of Formic Acid IV. Dehydration of Formic Acid , V. Decomposition of Bulk Fonnates VI. General Discussion References

.

.

.

. . .

35 --, 38 39 90 103 106 110

.

3. Application of Spectrophotometry to the Study of Catalytic Systems BY H. P. LEFTINand M. C. HOBSON.JR.

.

I. Introduction . 11.Principles and Theory 111. Experimental IV, Surface Structure Studies V. Physisorption Studies VI. Chemisorption Studies References

.

.

.

.

. . . . . .

*

. .

.

.

.

.

_

115 116 122 126 130 150 198

4. Hydrogenation of Pyridines and Quinolines BY MORRIS FREIFELDER

I. Introduction . 11. Catalysts 111. Occurrences of Side Reactions

.

.

.

. . xi

. .

203 204 206

xii

CONTENTS

.

IV. Catalyst Poisoning V. Reduction of Pyridine Salts . VI. Effect of Substitutes on Hydrogenation VII. Pyridinecarboxylic Acids . VIII. Pyridylalkanoic Acids . IX. Esters and Amides in the Pyridine Series . X. Hydroxypyridines and Derivatives . XI. Pyridylalkanoils . XII. Aminopyridines . X I I I . Pyridylalkylamines . XIV. Dipyridyls and Related Compounds XV. Quaternary Compounds . XVI. Some Selective Reductions . XVII. Quinolines . XVIII. Isoquinoline XIX. Effect of Substituents . XX. Quinoline and Isoquinolinecarboxylic Acids XXI. Esters and Amides XXII. Hydroxyquinolines and Related Compounds XXIILAmino Compounds XXIV. Quinolyl Ketones . XXV. Concluding Remarks References .

.

.

.

.

.

.

. . .

.

.

.

.

.

.

. .

.

.

I

.

.

..

. .

. .

.

.

.

.

.

.

.

.

.

.

.

.

.

.

208 209 210 213 216 216 217 220 220 222 223 224 226 238 239 241 242 243 244 245 245 248 248

5. Modern Methods in Surface Kinetics: Flash Desorption, Field Emission Microscopy, and Ultrahigh Vacuum Techniques

BY GERT EHRLICH I. Macroscopic Rate Studies 11. Field Electron Microscopy 111. Field Ion Microscopy IV. Ultrahigh Vacuum Techniques References

.

. .

.

.

.

.

. . .

. . .

.

.

. . .

. .

. . .

. . . . .

. .

.

256 311 341 391 424

6. Catalytic Oxidation of Hydrocarbons BY L. YA. MARGOLIS

.

.

.

.

. . .

. . .

. . . .

I. Introduction I T . Catalysts . 111. Reaction Mechanisms IV. Reaction Kinetics V. Modified Catalysts VI. Mixed Catalysts References

.

AUTHORINDEX SUBJECTINDEX

.

. . . .

. . . .

*

.

.

.

.

.

.

.

.

. . . . . .

.

. .

429437 439 468 479 493 496 503 517

Quantum Conversion in Chloroplasts* MELVIN CALVIN Department of Chemhtry and Lawrence Radiation Laboratory University of California, Berkeley, Calqornia

I. Introduction ....................................................... 11. Energy Sources for Photosynthesis ................................... 111. Photochemistry of Chlorophyll. ...................................... A. Photochemistry of Chlorophyll in Solution.. ........................ IV. Photophysical Effects in Model Systems.. ............................. A. Energy Transfer in Model Systems. ................................ B. Phthalocyanine aa a Model for Chlorophyll Energy Transfer. C. Photoinduced Charge Separation and Quantum Conversion. .......... D. Dark Conductivity Measurements on Violanthrene Systems. V. Charge Separation Processes. ........................................ VI. The Relation to the Photosynthetic Apparatus.. ....................... A. Electron Spin Resonance in Chloroplast Materials. ................... B. Apparent SpectralEfficiency ...................................... C. Luminescence ................................................... D. Absorption Spectral Changes Induced by Illumination. ............... VII. Quantum Conversion in Biological Material. ........................... VIII. Conclusion ........................................................ References

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

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

1 3 6 6 7 7

9 10 16 16

22 23 26 27 28 29 32 32

I. Introduction I n this chaptert we may not use the terms catalysis or catalysts explicitly. However, we can state at the outset that the whole photosensitive system that we are discussing is a catalytic system for converting electromagnetic energy into chemical energy. The system itself, to a first approximation, remains unchanged and, therefore, fulfills the definition of a catalyst. Within the system there are a great many individual catalytic steps, some of which are clearly defined and recognized, and others as yet less known. Yet the primary quantum conversion process involves perhaps the most important case of radiation catalysis. It is followed by the numerous catalytic steps which constitute a large chapter of the chemistry of life processes.

*

The work described in this paper was sponsored by the U.S.Atomia Energy Commission. t The material of this chapter is based on lectures given at the Research Conference of the Department of Chemistry, University of California, Berkeley, California. 1

2

MELVIN CALVIN

One of our major interests for the past fifteen years has been the mechanism by which the green plants can convert electromagnetic energy into chemical energy. The over-all reaction by which plants do this, chemically a t least, is written in the following terms: COa

+ He0

hu

(CHaOh

+0 s

This is the over-a11 process in which light is converted into chemical potential and that light is the light which is first absorbed, in most cases, by chlorophyll. A good deal of work has gone into breaking the process of photosynthesis down into a whole series of physical and chemical steps, and I will briefly define those aspects which have led to the present subject matter. Very early in the work, the reduction of the carbon dioxide by a reducing agent to produce carbohydrate was shown t o be a separate and independent reaction, independent a t least of the primary light absorption and conversion process ( I ) . Using radiocarbon, we have drawn a rather detailed map of the route from CO, to carbohydrate and a variety of other plant materials (amino acids, proteins and the like) (see refs. 2-6; for a complete review of the path of carbon in photosynthesis, see 6a). The other half of the process, the primary quantum conversion, was and has been for many years associatedwiththeso-called photolysis of water. Here, we come to a matter of definition of terms. We can write the first step as the absorption of light by the chlorophyll t o produce an excited chlorophyll, and the excited chlorophyll then is presumed to react with some molecule t o produce two species of some sort-let us call them 0 and R. It is a t this point that the electronic excitation is converted, a t least hypothetically, into two chemical species which should have stored in them the energy corresponding to the original excitation. They should be able to liberate very nearly all of that excitation upon back reaction: Chl

+ hv

Chl*

> Chl*

+M

+0

+R+Chl

(OH)(H)

(+I Chl

* - - - - - --

r--O

)L

Excitation

E

Chl

I

M

(e)

+R

QUANTUM CONVERSION IN CHLOROPLASTS

3

I n other words, these two chemical species, 0 and R, whatever they are, upon recombination will liberate some quantity of energy approaching that which was originally used to create them. This can be said because the over-all efficiency of the process is relatively high as photoprocesses go. It is certainly not less than 20 to 30%, and there are those who believe it may be as high, over-all, as 60% or more. Even if we accept a number as low as 30% as the upper limit of the over-all efficiency of the quantum conversion, we require that the initial step be of very high efficiency, because in the succeeding reduction of the carbon dioxide we already know that we are not more than 80% efficient in the chemical steps. If we are going to end up with an over-all efficiency of not less than 30%, we must have a very efficient transformation in the very first step.

II. Energy Sources for Photosynthesis Commonly this transformation, as I said earlier, has been written in terms of the photolysis of water. The biochemists and biologists have tended to call this molecule (M) a water molecule, and call the product 0 some kind of a hydroxyl radical, in an unknown form, and the product R some kind of a hydrogen atom, again in some unknown form (7-10). The chemists who have been working on this area, on the other hand, have tended to go one step further in their generalization and call one of the products an oxidant (0) and the other a reductant (R). But that really doesn’t say much more than the other description, given by the biochemists and biologists. Still more recently, physicists have entered the picture and they have tended to write the reaction in a slightly different way, calling the reductant, the electron, and the oxidant, the “hole” (11,12, 13). These are completely parallel representations of the energy conversion process. This is all that was meant by the statement that the primary quantum conversion process involves the photolysis of water. The ultimate result must be, as we know, generation of molecular oxygen from the oxidant (whether we call it a hole, an oxidant or hydroxyl radical in some unknown form) and generation, on the reduction side, of electrons or chemical species which can be defined as a reduced species, or a hydrogen atom bound to some particular form, The reductant will eventually be used in the reduction of carbon dioxide to produce the carbohydrate. We already know from our earlier mapping of the path of carbon in photosynthesis that we need not only a reducing agent for the reduction of carbohydrate but also, in addition, a rather specific type of chemical. Since we know the exact steps for carbon reduction, we were able to

4

MELVIN CALVIN

define what this reducing agent is at the last stages a t which it is used. It is a reduced form of pyridine nucleotide (TPNH). I n addition to this molecule which carries the hydrogen, we know that we need some pyrophosphate linkage, generally i n the form of adenosine triphosphate (ATP): H

HC/c*C/c\ 11 I HCPN//CH I

fi0

y4

" 7 1

HC-OH

I

H I7 y o o 0I I HC- 0-P-0-P-0-CH H 11 8 0 HC

?Ha

N5c\C,N N% 1 II C 'H HCQ, N1

"/C' HC*

I

H :

/N

I

1-

N/

tC

'N

%li /

I

"HC-OH 7 1

HC-OPO,H I Hy-OH Hf:

"+-OH(;-

o-

HCI I I HC-0-P-0-P-0-P-OH H I1 II 0 0

H

Triphosphopyridine nucleotide (oxidized form) (TPN+) H,

HC/C' II

HC,

N I R

c+o

/H /

cII

oI

II 0

Adenosine triphoaphate (ATP). In adenosine diphosphate (ADP) terminal phosphate is replaced by -OH

\

NH,

,CH

Nicotinamide portion of TPNH (reduced T P N f )

The pyrophosphate linkage (in ATP) is required to assist in the final reduction of the carbohydrates. We know that this molecule must also be produced somewhere in the region of the initial photochemical act. There has been a good deal of work done which has demonstrated that both TPNH and ATP are indeed produced by the photochemical apparatus of the plant (14,15),and there has been a good deal of additional work trying to define exactly where along the chain of eventsthat is, on the route from the oxidant to oxygen, or from the electron (or the reductant) to carbohydrate-the polyphosphate (pyrophosphate) is produced. There is some suggestion that it may be produced as the electron (hydrogen atom, reductant) moves down from its initial high energy level. One could also conceivably produce the pyrophosphate as the "hole," or oxidant, drops down some sort of energy hill toward oxygen. Finally,' there could be some recombination of the separated products through an enzymic mechanism which would give rise to pyrophosphate here as well, All three types of transformations are

QUANTUM CONVERSION IN CHLOROPLASTS

6

potential sources for the generation of a pyrophosphate linkage. There have appeared in the course of breaking down the fragments of plants, a variety of experiments in which there has been some demonstration that the recombination process works (9, 10, 16). The most popular notion currently is that most of the pyrophosphate arises along the route of the hydrogen on the way down, and some by recombination. But all three are real possibilities.

111. Photochemistry of Chlorophyll What we are concerned with is not what happens after this primary quantum conversion act but only what the nature of the primary act (or acts) itself may be. There is a good deal of information developing in the region of the reaction following the primary process, particularly on the reducing side, and relatively little on the oxidized side yet, and some on the recombination. But it is the primary act itself that we will be concerned with here. I n order to see what our train of thought was over the course of the years, one should remember the structure of chlorophyll and a few things that have been done in an attempt to try and find out what the primary quantum conversion act of photosynthesis was. There are a number of peculiarities about the structure of chlorophyll which have led many chemists t o suggest a number of reactions which might possibly be involved in the primary quantum conversion act.

A. PHOTOCHEMISTRY OF CHLOROPHYLL IN SOLUTION Over the past fifty years, ever since Stoll first gave us the basic notion (17),and in the last twenty years since a good many attempts the structure has assumed its present form (18,19), have been made to study the photochemistry of chlorophyll in solution to try and find out if the nature of the primary quantum conversion act could not be worked out. For example, this author and his colleagues spent a good deal of time on the assumption that the “extra” pair of hydrogen atoms which is involved in the conversion of porphin to chlorin (ring IV) might be involved in the photochemistry, that is, chlorophyll might be functioning between the porphyrin, or protochlorophyll, and chlorophyll (20). I n t h e plant in the steady state it is, of course, mostly in the form of the dihydro compound, but upon illumination it might reduce something and then the protochlorophyll so formed might pick up the two extra hydrogens from the water molecule. This is one type of reaction which has been suggested. as to what chlorophyll was like

6

I XN U

MELVIN CALVIN

ZT V

QUANTUM CONVERSION I N CHLOROPLASTS

7

Another point of great interest which has attracted considerable attention is the enolizable hydrogen on carbon atom No. 10 in the isocyclic ring. There have been a number of schemes devised and a good deal of photochemistry done with a view toward possibly using this particular reactive point as the point at which the water molecule was spIit, or the GO, molecule was reduced (21-24). None of this work has been very fruitful, except to give more information about the photochemistry of chlorophyll in solution. It has not, however, led to any model reaction in which anything approaching the 35 kcal, which the excited chlorophyll bears, is stored in the reaction product. In fact, as far as I know, there is not any photochemical reaction in solution in which a very large fraction (more than 50%, say 80 to 90%) of the energy in the excited state is stored in the immediate reaction product. This is not surprising when you to stop to think about it in retrospect. The excited state of chlorophyll stores about 35 kcal. If the products also had to have in them an appreciable fraction of this, Bay 30 kcal, the products cannot have a barrier for their return in the back reaction of more than this difference, whatever it is-and it cannot be very large. For this reason, the back reaction generally does take place, and one is not successful in separating the energetic products in any photochemical model in solution, with chlorophyll or any other substance for that matter.

IV. Photophysical Effects in Model Systems A. ENERGY TRANSFER IN MODEL SYSTEMS After spending a good many years of our own lives, as well as having looked at the efforts of many others, there occurred a concatenation of circumstances which focused our attention on quite a different point of view. The first of these resulted from the examination of the photochemical apparatus of the plant itself, shown in Fig. 1, which is an electron micrograph of a tobacco chloroplast in which all of the photochemistry occurs. The layers are roughly 60 A thick and this will give some idea of the highly ordered structure in which the photochemistry is performed. The chloroplasts contain a highly ordered array of chlorophyll and many other molecules. Figure 2 shows a still higher magnification of one of the small lamina (lamella), and one can see that they are themselves composed of small spherical molecules which are packed in some as yet hypothetical way; it looks as if they are packed together to form the lamina, one sphere on top of the other (25). These photographs show that we are not dealing with molecules in solution in the ordinary sense at all when we consider the photochemical apparatus.

8

MELVIN CALVIN

FIG. 1. Electron micrograph of tobacco chloroplast.

FIG.2. Electron micrographof frozen dried spinach sonicate showing “Quantasomes”; 880 A diameter polystyrene latex markers (25).

QUANTUM CONVERSION IN CHLOROPLASTS

9

We are dealing with a highly ordered array of substances. What that array is on the molecular level-i.e., how the plane of the chlorin ring is located with respect to these layers or with respect to the spherical globules of which the layers seem to be composed-is still a matter of controversy and remains one of the areas which this writer would like to see the physical chemists explore. It is necessary t o bridge the gap between ordinary statistical molecular behavior and individual structures that can be seen. We do not know, for example, what the arrangement of the chlorophyll molecules is in the lamellar layers. Order is, however, the major factor with which we are left. One of the other two factors which came together some years ago and which increased our interest in this kind of work was the first report by one of my former associates, D. D. Eley of the University of Manchester, England, of the fact that an organic molecule which had some of the structural features of a porphyrin, namely, phthalocyanine, had the properties of a semiconductor (26).He reported the nature of that conductivity and something about the energy gaps that were involved. The third piece of general knowledge which came our way and influenced us to change our point of view in an attempt to understand this primary quantum conversion act was the description of a quantum converter made up of a solid atomic lattice. This was the silicon photobattery in which the Si lattice had in it certain impurities (atoms which had either an excess electron, or had shortage of electrons), and a suitable structure was devised which was able to convert the quantum into electrical potential with a n efficiency at first of only 3%; this efficiency has since gone up to 10 to 12%. This fact, taken together with the other two (recognition of the ordered structure of the photochemical apparatus in the plant and the development of the concepts of organic semiconductors) on top of the theoretical arguments that we have no real way of chemically separating the high energy products that must be formed, and the failure of all experimental attempts to find any, led to the supposition that perhaps the quantum conversion act was not a photochemical process at all but something more resembling the photophysical process which occurs in the atomic lattices (14, 27, 28). Instead of using atomic arrays to achieve the separation of products we would use an organic molecular array, of which chlorophyll was one of the principal constituents.

B. PHTHALOCYANINE AS A MODEL FOR CHLOROPHYLL ENERGY TRANSFER I n order to develop this notion we undertook the model studies which will be discussed below, the determination of the electrical and

10

MELVIN CALVIN

magnetic properties of organic solids. The first type of solid which was studied was phthalocyanine itself, since this was the first one that was reported by Eley. The real problem was not the conductivity of the phthalocyanine but how it might interact with other molecules that might be present in a chloroplast system-electron donors and electron acceptors corresponding to the donor and acceptor systems in the atomic lattices. We here made use of the background of organic oxidationreduction which was available to us, both in terms of the knowledge of the nature of the oxidation-reduction reactions and also the knowledge of the nature of the interactions between donors and acceptors which has been growing steadily in the last decade, (The growth of our knowledge of the interaction of donor and acceptor molecules had one of its principal starting points here in Berkeley with the work of Hildebrand and Benesi (29) on iodine and benzene interactions).

C. PHOTOINDUCED CHARGESEPARATION AND QUANTUMCONVERSION The first thing we wanted to find out was whether or not it was possible to devise a system of donor and acceptor molecules, of which chlorophyll might be either, in which the primary quantum conversion act was not a production of molecules which themselves have to separate from each other, but rather purely a separation of charge rather than a separation of atoms (28). The charge could separate not by molecular diffusion but simply by charge diffusion. This kind of a separation process could occur much more easily and readily than the molecular diffusion process, and if there is a proper lattice arrangement of donors and acceptors, there will be a chance to achieve this kind of separation with concomitant quantum storage. Actually the work so far done has not demonstrated such a storage unequivocally, but it has been very suggestive and should be discussed further. The systems that were used by this writer and colleagues contained one of some four or five different donor molecules-phthalocyanine, violanthrene, perylene, etc.-we found that almost any highly aromatic system could serve as a conducting matrix and phthalocyanine was only one of these. The electron acceptor systems that we used were o-chloranil (the most commonly used), iodine and tetracyanoethylene. These donor and acceptor systems were used in various combinations. A pattern of results based on a conductivity cell is shown in Figs. 3a, b, and c. It is made of aquadag electrodes on which is coated the matrix (phthalocyanine or violanthrene, as the case may be); usually the samples are sublimed onto the electrode system. On top of the matrix is placed one of the “doping” agents, one of the donors or acceptors.

QUANTUM CONVERSION IN CHLOROPLASTS

11

In most of our work the matrices have been the donor molecules and the other part of the layer has been the acceptor molecules, which is either sublimed on top or sprayed on in a dilute benzene solution. These are the two principal ways in which the acceptor molecule is laid onto the donor.

Back face illumination

illumination

Fmnt fa-

(b) FIG. 3. Diagram of sample conductivity cells.

We have measured, in most cases, the conductivity between the electrodes, both in the dark and upon illumination, as a function of the added acceptor molecules (30, 31, 32). Figure 4 shows the results of ELECT

L

i

d

FIG. 30. Layered donor-acceptor system.

such an experiment using phthalocyanine. You will notice that the dark conductivity of the system rises by a factor of almost one million as small amounts of electron acceptor, o-chloranil, are added to, the

12

MELVIN CALVIN

phthalocyanine. Remember that the acceptor molecule is on the back face of the matrix and not on the electrodes-the matrix lies between the acceptor layer and the electrode. The photocurrent is the difference between the dark current and the current with the light on. The enormous increase in conductivity is found in almost every case for the combinations here described. I

I

I

-DARK

I

I

CURRENT

W---PHOTOCURRENT

2

I-

A W R E N T MOLE RAT0 OF ORTHO-CHLORANIL TO PHTHALOCYANINE

FIG.4. Variation of dark conductivity and photoconductivity of phthalocyanine with amount of electron acceptor (0-chloranil)added.

I n addition to increasing the conductivity in the matrix when a n acceptor layer is placed on top, another effect makes itself apparent. Unpaired electrons appear in the matrix when it is coated with the acceptor molecules. Figure 5 shows the appearance of a signal in the phthalocyanine (a similar signal appears within the other donor molecules as well). The g-value is very close to that of a free electron. When the light is turned on, the spin signal changes-i.e., the number of unpaired electrons changes. I n the case of the phthalocyanine-chloranil system, the number of unpaired electrons actually decreases when the light is turned on. I n other systems, it may increase or decrease, depending upon the combination, but the conductivity always increases. Figure 6 shows the growth and decay of the signal following illumina-

QUANTUM CONVERSION I N CHLOROPLASTS

13

tion and darkening. When the light is turned on, the spin signal is decreased and when the light is turned off, the spin signal comes back. The photoconductivity shows the same type of change, but increasing in the light and dropping back in the dark.

FIQ. 5. Electron spin resonance spectrum of o-chloranil “doped” phthalocyanine. Curve represents the first derivative of absorption.

1

1

Light on

z 0 +

5 E VI

I-1

Lighl off

mi”+

FIG.6. Effect of illumination on ESR signal of o-chloranil“doped” phthalocyanine.

We also performed measurements of photoinduced charge separation. This was done by placing a block of phthalocyanine, covered on one side with the o-chloranil, between the plates of a condenser and determining the direction of the induced polarization in this block when it is illuminated. It turns out that the polarization induced by illumination is

14

MELVIN CALVIN TRANSPARENT CONDUCTORS

r

1

PHTHALOCYANINE ORTHO-CHLORANIL I

I

VIBRATING REED ELECTROMETER

FIG. 7. Schematic diagram of polarization apparatus.

I

I

0 DECAY OF PHOTOCONDUCTIVITY AT 25OC AMCAY OF LIGHT INDUCED POLARIZATION25OC ORlSE OF ESR AT 25OC

-

XDECAY OF PHOTOCONDUCTIVITY -IOOOC oRlSE OF ESR AT -loO°C

20

60

100

TIME -SECONDS

FIG.8. Semilog plot of time dependence of photoconductivity, light-induced electron spin resonance and light-induced polarization in doped phthalocyanine.

QUANTUM CONVERSION IN CHLOROPLASTS

15

as shown in Fig. 7-the positive charges go to the phthalocyanine layer and the negative charge builds up in the o-chloranil layer. Here we have, in effect, induced a separation of charge by illuminating a sandwich layer of the type that we originally constructed. I n fact, if electrical contact is made between each side through a suitable load, a current can be taken off from this cell. This is, in effect, a photo-battery. One cannot take off much current because the internal cell resistance is too high and the cell was not made with a molecular thinness. Although in principle we have devised a quantum converter using the electron donors and acceptors, it is, as yet, not too efficient. Figure 8 shows the growth and decay curves (the kinetics) of these three phenomena-photoconductivity, light induced polarization and the electron spin resonance signal. All of them have the same unimolecular time constants a t 25°C. When the system is cooled to -loo", which has been done for the photoconductivity and the spin signal, they decay faster a t the lower temperature but again they are parallel; they have the same kinetic behavior. These results indicate that the separation of charge induced in the layer, the conductivity induced in the matrix, and the change in the number of unpaired spins introduced by the light, are all manifestations of the same process.

D. DARKCONDUCTIVITY MEASUREMENTS ON VIOLANTHRENE SYSTEMS It would be pertinent to describe a similar set of observations on a violanthrene system because there actually is more data on the violanthrene system. Figure 9 shows the dark conductivity in the violanthreneo-chloranil system under different types of illumination. The same general pattern appears as with the earlier phthalocyanine work. Figure 10 shows the light intensity dependence of the spin signal and the photoconductivity, and again they are exactly parallel. There is no question but what we are measuring the same thing in both cases. Figure 11 shows the kinetic behavior for violanthrene corresponding to that which we saw earlier for phthalocyanine and the results are identical. The parallel systems at the lower temperatures were not measured, but this author has no doubt that if they were measured they would give parallel results.

V. Charge Separation Processes The question now is: What is the actual process that is taking place in these laminar systems? We can without doubt say that the separation

16

MELVIN CALVIN

-

VIOLANTHRE~€ ORTHO-CHLORANIL

DARK CURRENT PHOTOCURRENT -WuX FACE ILLUMINATION

a

--o---

PHOTOCURRENT-FRONT FACE ILLUMINATION

-I

-I

w

-2

$

-3

a

/ - /

: -5 a I-

a

3 0

'

a

2

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

0

W W

0 - - - - 4 - -

-4

l0 W

a

UNPAIRED SPIN CONCENTRATION

I

In z 0

..........

5

........

IA

0

-6

K

-7

3

W

P z

a

3 -a -9

I

-10

FIG.9. Variation of violanthrene dark current, photocurrent, and unpaired spin concentration with the amount of o-chloraniladded.

/IOLA#THRE#E- ORTHO-CHLORA#IL

- --o---

WOTOCONDUCTIVITY PHOTO ESR

loo

10

I

RELATIVE

LIGHT

INTENSITY

FIG.10. Intensity dependence of the steady-state photocurrent and photoinduced electron spin resonance in violanthrene treated with o-chloranil.

17

QUANTUM CONVERSION I N CHLOROPLASTS

of charge has occurred, as suggested in the initial description, but what is the nature of this system which gives rise to the separation? At first we were thinking in terms of charge transfer complexes in which the absorption of light leads to an excited state which can be described in quantum terms as a state in which there has been effective charge transfer from the donor to the acceptor; in an excited state there is more

-

VIOLANTHRENE ORTHO-CHLORANU

BC

(-72. C )

20

40

80

60

TIME

IN

too

120

SECONDS

FIQ. 11. Semilog plots of the time dependence of photoconductivity and light induced in ESR in violanthrene treated with 0-chloranil.

of the electronic charge on the acceptors than there is in the ground state. I n solution these are simply excited states of a molecule which decay with the same kind of speed that any other excited state of singlet-singlet transitions would decay. I n this laminar system, however, a separation of charge is actually achieved as measured by the conductivity, the appearance of spin signals and the polarization. So

18

MELVIN CALVIN

this is quite another matter. I n the violanthrene we were able to make some definite deductions as to the nature of the condition of the interface between the violanthrene and the o-chloranil (the donor and the acceptor). I n order to get conductivity between these electrodes (see Fig. 3c) we would have to have a system which would allow charges to move very easily between them, and such mobile charges would presumably be induced by the addition of the acceptor molecules on the surface of the matrix. When the acceptor molecules are placed on the donor surface, there occurs a n actual transfer of an electron to form some of these acceptor anions, A', leaving behind in the matrix of donor molecules some of these donor positive ion radicals, DO. Presumably there will be a charge gradient developed, and the fact that there are a certain number of positive holes reaching clear back into the matrix improves the conductivity between the electrodes. There is also an equilibrium in the region where the anion gives the double negative ion as well as where the cation is produced to give the double positive ion. These are in equilibrium in the two respective phases, 2 A- < A= + AI " 2

Df

1 _

D'++ Do

This kind of a process is quite different from the change that results when a charge transfer complex absorbs in the charge transfer band. This is a separation of charge leading to two doublet molecules, whereas in the excited state of a charge transfer complex it is stiEb a singlet molecule and the charges are not actually separated. It seemed that the possible role of these radical ions could be determined by spectroscopic methods. We looked a t the absorption spectrum of a violanthrene film treated with the various acceptor molecules (0-chloranil, iodine and boron trifluoride), and Fig. 12 shows the spectra obtained with a variety of acceptor molecules. Figure 12b shows the spectrum of violanthrene dissolved in sulfuric acid and this has been shown to be the violanthrene positive ion radical in solution (33).Thus the band a t -7000 A (in Fig. 12a) would be the violanthrene positive ion radicals in the solid matrix. A charge transfer band of violanthrene and any of the added acceptors would not be the same point for each acceptor molecule, and with o-chloranil it would be expected out in the infrared, from all the other indications and it is not there. What there is, then, is a complete transfer, not a charge transfer complex a t all, but a complete transfer from the donor to the acceptor (34).However, this separation of charge appears to take place only in a

QUANTUM CONVERSION IN CHLOROPLASTS

19

heterogeneous system. It does not take place in the homogeneous crystal in which both the donor and the acceptor molecules are arranged in a good charge transfer type of complex.

t J

uca

h

0

FIQ. 12. (a) Curve A-solid state absorption spectrum of pure violanthrene; Curve B-difference between solid state absorption spectrum of violanthrene treated with BF, and that of pure violanthrene; Curve C-difference between solid state absorption spectrum of violanthrene treated with o-chloranil and that of pure violanthrene; Curve Ddifference between solid state absorption spectrum of violanthrene treated with iodine and that of pure violanthrene. Part B: (b) Absorption spectrum of violanthrene dissolved in concentrated sulfuric acid for one hour.

Eastman has examined the perylene-o-bromanil system and he was able to make a 1: 1 charge complex (35).When it was slowly crystallized out of a cool solvent, it had no spin signals and was a perfectly normal charge transfer complex. It shows this behavior in solution as well. However, if the quinone (0-bromanil) is sprayed onto the perylene crystals so that a two-phase system exists, the charge can actually separate when the transfer occurs and then the concomitant spin signals appear, conductivity is enhanced, the positive ion radical spectrum appears, etc. Figure 13 shows the hypothetical charge distribution in a lamellar donor-acceptor system which accounts for the observed changes in spin signal, Whether an increase or decrease in the spin signal is observed,

20

MELVIN CALVIN

depends upon the equilibrium between the mono- and dipositive ions on both sides of the interface. Always there is an increase in conductivity as the charge separation is increased by illumination. The charge equilibrium is shifted upon illumination, moving electrons from the

DONOR

W

0

LAYER

CONDUCTION R E G W

1.'I

'\\

\-'

DONOR LAYER BOUNDARY (ELECTRODES)

--O

4

I

M)NOR-ACCEPTOR INTERFACE

ACCEPTOR LAYER BOUNDARY

FIG.13. Hypothetical distribution of charges in a lamellar donor-acceptor system.

donor to the acceptor layer, thus changing these ionic equilibria and increasing the conductivity in both sides of the interface but not necessarily increasing the spin. The spin may increase or decrease, depending upon the nature of the equilibrium constants among the ions and the amount of charge separation. In quantum terms, we may describe the phenomena as follows (see Fig. 14). I n the acceptor layer, for example, there would be an anion radical with neighbors on either side with vacant orbitals equivalent to the one occupied by the extra electron. This electron in its migration, i.e., when it moves from one molecule to the next one, produces a state after the reaction [as in ( I ) , Fig. 141 shown in the top row of Fig. 14 which is identical in energy with the starting point. The same thing can be said of the donor layer with its hole, which can have electron migra-

QUANTUM CONVERSION IN CHLOROPLASTS

21

tion from a neighboring molecule to fill the hole, leaving a hole but in the neighbor. This final state after the migration of charge is exactly the same as the starting point. This situation does not obtain in the case of a free radical crystal. In a free radical crystal every one of the highest occupied orbitals is singly occupied, and in order to move an electron [as in process (3), Fig. 141 one would have to doubly occupy ACCEPTOR LAYER

A

A‘

A NARROW CONWIGTON BAND

DONOR LAYER

D

O+

D NARROW CONDUCTION BAND

NEUTRAL FREE RADICAL R R

R

FIG. 14. Charge migration in a molecular lattice. Schematic representation of donor and acceptor molecules and ions imbedded in a donor layer or an acceptor layer, respectively. From this diagram it is clear that process ( l ) ,the transfer of an electron from an acceptor negative ion to a neutral neighbor, produces a state of the system which is energetically identical with the initial state. Similarly, there is no net change in energy aa a result of process (2),which rearranges charge in the donor layer. In the case of a neutral free radical, however, the electron transfer process (3) does not result in a state energetically equivalent to the initial state. Since processes (1) and (2) simply change the location of negative and positive charges respectively with no net change in energy, we can consider the orbitals involved in the electronic rearrangements as forming conduction bands. If, however, the lattice were made up entirely of A- radical ions (no A’s) irrespective of the cations, or entirely of D+radical ions (no D’s) irrespective of the anions, there would be no identical vacant orbitals into which the charge carriers could move and hence no conduction bands (however narrow). This last situation would correspond to the completely filled free radical system aa in process (3) above.

22

MELVIN CALVIN

one orbital and vacate another, and there would have to be a different intermediate state than in the top two systems. This case, as represented by the upper two diagrams of Fig. 14, is what automatically appears in the donor-acceptor two phase systems. Quite recently this notion was dramatically illustrated by the report of the electrical properties of some complexes made from relatives of tetracyanoethylene, namely, tetracyanoquinodimethane (TCNQ) (3638). This material can be made into crystals of two different kinds. CN,

C

,CN

The first is one in which there exists an extra quinone molecule for every free radical ion, i.e., a crystal made up of two quinone molecules constituting one of these free radical ions and one cation. These materials have conductivities of the order of unity, varying with the cation, but they are all very good conductors. Thus, this material is a solid with more than one quinone for each odd electron. Presumably it has alternating neighbors with occupied and unoccupied orbitals so that one can get the electron to move from one into another. If, however, the crystals are made with only one quinone per odd electron and not the extra one, the conductivity drops by a factor of lo4. This is exactly the kind of thing that one would have predicted from the results on the two-phase system which was just described.

VI. The Relation to the Photosynthetic Apparatus How does the concept of organic semiconduction fit into our notions of what we know about the processes going on in the green plant? A number of similar physical measurements have been made on the chloroplasts themselves, and some of the measurements can be interpreted in language which is very similar to that used for the phthalocyanine model systems. The work on the biological system itself, designed with this type of quantum conversion in mind, has just begun. The remainder of the discussion is an attempt to see how many of the kinds of measurements which were performed on the model systems can be performed on the biological material and how they will tell us

QUANTUM CONVERSION I N CHLOROPLASTS

23

what goes on in the biological material in the same manner as they tell what is going on in the model systems. The first measurement which we made on the model system, namely, the conductivity, is the most difficult to make on this type of biological material because the lamella of the chloroplasts are very small (30 to 60 A thick). However, there have been conductivity measurements on dried chloroplast preparations which show that they are indeed photoconductive. Although these are subject to questionable interpretation in such a complex system (39, 40), the photoconductive behavior of a chromatophore film is shown in Fig. 15.

TIME (min)

FIU. 16. Photoconductivity of the chromatophore film is shown by the larger slope during the time that the exciting light is on. A sudden step in the charge can be seen at the beginning and at the end of the illumination. The wavelength of the exciting light was 850 mp.

A. ELECTRON SPIN RESONANCE IN CHLOROPLAST MATERIALS One of the principal types of experiments that we have done is to look for unpaired electrons which might be generated by the light in the biological material. The chloroplast suspensions are placed inside a resonance cavity to see whether or not unpaired electrons were generated when the light was turned on. Very early we found that there were unpaired electrons of this type. The first experiments were done with eucalyptus leaves in 1956, but we found the results were not reproducible due to the variability of the leaves themselves. Towards the end of that year the same kinds of observations were made a t St. Louis by Townsend, Heise, and Commoner (41, 42). This author and his colleagues made some chloroplast preparations and did a serious investigation of the same thing ( 4 3 , 4 4 ) .This particular type of experiment

24

MELVIN CALVIN

can be done with the whole organisms (whole algae or bacteria) or with pieces of chloroplasts or chromatophores (chloroplasts of bacteria). Figure 16 shows the light-produced signals from whole spinach chloroplasts. Light of 40 kcal/quantum is being shone on these materials and very few chemical bonds can be broken by as little as 40 kcal. The

1 0 GAUSS

FIQ. 16. Growth and decay of light signals from whole spinach chloroplmts.

signal indicates the appearance of unpaired electrons; any free radical will give this kind of signal, and most biological material which is undergoing rapid metabolism will show signals of this kind. The question, therefore, is: What kind of unpaired electrons are these? Are these ordinary free radicals, or are these electrons produced in photoprocesses such as have been described in the model systems Z If these were chemical free radicals produced by some secondary reactions, we might expect that if the system were cooled to a low enough temperature, the chemical reaction might stop and only the physical process of electron transport would remain. We attempted to do this by cooling the sample to - 150' and we still obtained light-induced signals.

QUANTUM CONVERSION I N CHLOROPLASTS

25

Figure 17 shows the kinetic behavior of such a signal for Rhodospirillum rubrum which use bacteriochlorophyll. This experiment was performed at a series of different temperatures, and the signals change in character with the variation in temperature. There is also a variation in the signal with time. At 25” after the light is turned on, the signal rises just as fast as the apparatus will follow it and reaches its equilibrium value immediately, and when the light is turned off, the

+2Y

P

I

w FIQ. 17. Electron spin resonance signals from Rhodoepirillurn rubrum; 5 min continuous illumination.

signal drops as rapidly as the equipment will follow it. In other words, the rise time and decay time that we have so far been able to see are not intrinsic to the electrons but rather they are limited by the apparatus. As the material is cooled from 25°C to - 15”C,a good deal larger signal appears, but there is a slow rising component in the signal; if the temperature is lowered still further, to - 55OC, some of the “extra” signal which

26

MELVIN CALVIN

is purely chemical (secondary, in other words) is frozen out, but not all of it. There is still a very fast rise and then there is a slow rise a t - 55", and the decay time shows the same characteristics-a fast decay and a slow decay. There are quite clearly several different kinds of unpaired electrons produced in this organism when the light is shone on it a t - 55°C. When the temperature reaches - 160"C, we have none of the slow signals left at all--only the fast signals; both the rise and the decay are fast. This phenomenon is most readily interpreted by the obvious notion that we are fimt making a conducting type of unpaired electron which is then undergoing chemistry inside the biological material via one electron reactions. We are seeing a t the various temperatures not only the physically-produced charge separation but chemical radicals as well, and as we cool the solution we freeze out the chemical reaction and have left only the physical process itself (45, 46). We really need something more to characterize the unpaired electrons. The rate of growth and decay, temperature dependence, etc., is not enough to identify these electrons as physically produced instead of chemically produced. We have tried to use one or two other ways of characterizing the electron, such as looking for hyperfine structure, that is, looking for the interaction of the unpaired electron with specific nuclei, but so far this has not been successful. Either there are so many nuclear hyperfine interactions as to overlap, or the quasi-solid matrix broadens the lines so that no very useful resolution has yet been possible (42).

B. APPARENTSPECTRAL EFFICIENCY The next characterization which we attempted was to measure the efficiency with which the light produces the spin signals, i.e., the quantum efficiency for the production of these electrons. I n absolute terms it can be stated that the quantum efficiency for the production of these electrons is in the same vicinity as the quantum efficiency of photosynthesis, i.e., of the order of one to one-tenth. However, how does the quantum efficiency vary with respect to the wavelength of light? The values obtained in these experiments are merely relative values for the production of unpaired electrons a t one wavelength to the production a t another, compared with the absorption of chlorophyll. Figure 18 shows the action spectrum for the production of free electrons and the absorption for the chloroplast. A thick layer of chloroplasts was used so that all the light was absorbed, and in those regions where the light is most strongly absorbed, the concentration

20

.1.0

?,

I5 .,/'

$

I?

-LEFT

--- RIGHT ORDINATE

I :

"t,

ORDINATE

,

.0.8

>

-.

'.

10-

I W

5-

2 v)

0

I,,-IO''QUANTA/SEC.

BAND WIDTH.

iood

FIQ.18. Absorption and action spectra of chloroplasts.

The idea that light of different wavelengths may collaborate in photosynthesis is now a current development of our knowledge of the behavior of plants with respect to light and it is known as the Emerson effect. I n simplest terms it may be defined by the following observations: Measure the number of oxygen molecules produced per quantum of red light; measure the number of molecules produced per quantum of green light; and then put both the red and the green light together on the same plant. This can be done in such a way that when the two wavelengths of light are together on the plant one gets more (or less) than the sum of the two separately. I n other words, there is a collaboration of the two wavelengths of light (48).The experiment can be done under conditions where there is a negative collaboration of the two wavelengths and they cancel each other, depending on the light intensities and other conditions of the experiment (49-51).

C. LUMINESCENCE If we have achieved a separation of charge in the molecular lattices and if the charge is allowed to recombine, light can be emitted a t low temperatures. Figure 19 shows the delayed light emission from Chlorella, spinach chloroplasts and Nostoc. The wavelength distribution is what might be expected, and the kinetics of decay of the light emission are exactly the kinetics of the decay of the spin signal (52-54).

28

MELVIN CALVIN

D. ABSORPTION SPECTRAL CHANGES INDUCED BY ILLUMINATION Another way of searching for the primary changes produced by illumination is to examine t.he changes, if any, produced in the absorption spectrum by illumination. This was done first for purposes of trying to identify the changes (55-59) and later to determine the kinetics 0

ACTION SPECTRUM

SPINACH CHLOROPLASTS

W M WIDTH -2Ornp

0

0

EMISSION SPECTRUM 25.C (THICK FILM) EMISSION SPECTRUM -45°C (THICK FILM) EMISSION SPECTRUM -70°C (THICK FILM)

ACTION SPECTRUM

1 EMISSION

SPECTRUM

0

ACTION SPECTRUM

x

EMISSION SPECTRUM 25% (THICK FILM)

X (mu)

FIG.19. Delayed light emission from a variety of biological materials.

of these changes (60-62). Still more recently, the effect of temperature (39, 63, 64) and of the color (65) of the exciting light have been examined. One of the principal results of this work seems t o be a consensus that illumination of most types of photosynthetic apparatus produces a,

QUANTUM CONVERSION IN CHLOROPLASTS

29

drop in absorption in at least three main regions of the spectrum: -420 mp, -530 mp and -680 mp. The two shorter wavelengths are believed to be associated with an oxidation of a cytochrome. Two pieces of work which have been done at the Oak Ridge National Laboratory by Arnold are important in this regard (39, 40). Arnold was measuring the change in the light absorption of the chromatophores from Rhodopseudomonas spheroides, induced by illumination with a second light, usually of longer wavelength. Figure 20 shows the change I-

30" -1

1-

30"

+I

1"

1 300°K

t J

t

ON OFF

LlGHX ON I

TI ME (X

5

6400A)

FIG. 20. Absorption changes at 4200A induced by illumination[Arnold and Clayton

(W1* a t absorption a t 4200 A, and you can see that the change occurs just as fast as the instrument can measure it a t 300°K. The decay is slow because part of it is chemical and part of it is physical. Specifically, one should .note the decay taken a t 1°K where there is very little light chemistry going on, and it can still be seen that spectral change is occurring just as fast as the instrument allows. This is clear evidence that the light is introducing a physical change, a change which can only be motion of electrons and not of atoms.

VII. Quantum Conversion in Biological Material This author will now explain what he believes, a t the moment, is the primary quantum conversion process that goes on in the layer of chlorophyll (and other pigments) in the lattice. We know something about the chemical composition of the chloroplast itself, i.e., it is a lipoprotein together with pigments. There are a number of specific molecules present in the chloroplast, for example, chlorophyll and

30

MELVIN CALVIN

carotenoid, and, in addition, there are two other rather important ones which are present in large amounts and which have a function in the energy conversion process. As you know, these laminar systems require not only the presence of an absorber but also the presence of an acceptor molecule for electron transfer to occur-to finish the process we need to have something present as a donor molecule. Two species that are well established in the chloroplasts as performing this donor function are a very important quinone called plastoquinone (66-69)and a variety of donor molecules, such as the iron heme (cytochrome) species which are always present in chloroplasts and chromatophores (70). Figure 21 shows the two boundaries of the pigment layer. Chl are the chlorophyll molecules in some array, perhaps including carotenoids. The first act of photosynthesis is the absorption of the quantum by the chlorophyll molecule to produce an excited chlorophyll molecule. If this were a perfect atomic lattice, this would be an absorption by the entire lattice, but this is not the case. It is a molecular lattice in which the interactions between molecules are relatively small compared with the interactions between atoms in an ionic lattice. The result is that the migration of this exciton occurs by resonance transfer between neighboring chlorophyll molecules until it arrives a t one which is bound, or adjacent, to an electron acceptor such as plastoquinone. The quinone, then, is a likely electron acceptor; there is one plastoquinone molecule present for about 400 chlorophyll molecules. The exciton reaches the chlorophyll which is bound by a charge transfer complex to the quinone, ionization occurs, the electron is transferred, leaving behind in this chlorophyll molecule an electronic vacancy, or “hole.” At this point, the idea of charge migration is introduced (see Fig. 21). After ionization occurs, it is necessary that there be a charge migration by an electron going from a neighboring chlorophyll molecule to the hole, so the hole moves down to the next chlorophyll molecule until it becomes adjacent to the iron heme (cytochrome). When the hole reaches this point, electron transfer occurs from the iron (or other donor) to neutralize it and the pigment layer is returned to its original condition (39, 63). The alternative step of electron transfer from donor to excited chlorophyll has also been suggested as a primary step (57). A discussion of these two alternatives originally led us to the present suggestion as the more likely of the two. The existence of two different primary quantum acts (48-51, 57, 65, 71-73) makes it not unlikely that both of these two alternatives may occur, although either one alone would be sufficient to support the whole process. Therefore, a separation of charge has been achieved and the oxidized

QUANTUM CONVERSION IN CHLOROPLASTS

31

donor becomes an oxidant and the electron in the quinone is the reductant. The reductant can go on to reduce carbon dioxide and the oxidant can go on to generate oxygen. ATP is required to help on the

Cyt -CYTOCHROME AND/OR OTHER ELECTRON DONOR SYSTEMS (AQUEOUS PHASE)

0 - PLASTOQUINONE AND/OR

OTHER ELECTRON ACCEPTOR SYSTEMS (TPN, LlPOlC ACID,ETC) LIPID PHASE

Chl- CHLOROPHYLL I.

2.

Chl

t

hu

-

C h l * t 0-

* +t F eP ----.

3. C h l

Chl*

b-+ &It m Fe t C h l

FIQ. 21. Schematic arrangement of chlorophyll and possible donor and acceptor molecules in the chloroplast. The system in the chloroplast might structurally bear some resemblance to the model in this figure, the chlorophyll having associated with it on the one side the electron acceptor, plastoquinone, in a lipid environment, and on the other side electron donor materials, such as the cytochromes, in an aqueous environment. Following the absorption of a quantum in chlorophyll [Eq. ( l ) ]it will migrate by resonance transfer to a suitable site near the quinone where electron transfer to the quinone will take place [Eq. (2)]. The resulting vacancy can migrate by hole diffusion, that is, electron transfer from normal chlorophyll, into the vacant orbital of the neighboring chlorophyll positive ion, This process is the one which most nearly resembles the properties of a semiconductor and it permits the oxidant (chlorophyll positive ion) to separate from the reductant (electrons in the quinone orbitals) by a very nearly temperatureindependent process. The oxidant then captures an electron from a suitable reducing agent, such as ferrocytochrome, thus producing a ferricytochrome and regenerating normal chlorophyll [Eq. (3)].

reduction of carbon dioxide and for many other energy-requiring operations. One possibility is that ATP may be generated during the passage of oxidant to oxygen, and ATP may also be generated on the reduction side and by recombination as well.

32

MELVIN CALVIN

VIII. Conclusion The primary quantum conversion act of photosynthesis in chemical terms is an ionization occurring in a charge transfer complex (28). But this cannot occur in isolated charge transfer molecules in solution or in a homogeneous single phase crystal because the products cannot escape from each other. The primary quantum conversion act as it occurs in modern photosynthesis can only take place in a laminated structure where the electrons and holes can escape from each other by electron migrations and not by atomic migrations (74). (For a discussion of the evolution of photosynthetic mechanisms, see Calvin, 74).

REFERENCES 1. Calvin, M., and Benson, A. A., Science 108, 304 (1948). 2. Bassham, J. A., and Calvin, M., “The Path of Carbon in Photosynthesis.” PrenticeHall, Englewood Cliffs, New Jersey, 1967. 3. Bassham, J. A,, and Calvin, M., in “Handbuch der Pflanzen-physiologie” (M. Ruhland, ed.), Vol. V, Part 1, p. 884, Springer, Berlin, 1960. 4. Bassham, J. A., and Calvin, M. “The Photosynthesis of Carbon Compounds.” W. A. Benjamin Co., New York, 1962. 5. Bassham, J. A., J. Chem. Educ. 86, 648 (1969); 38, 161 (1961); Soi. American June (1962). 6. Bassham, J. A., and Calvin, M., Comp. Biochem. Phyeiol. 4, 187 (1962). 6a Calvin, M., Science 135, 879 (1962). 7. van Niel, C. B., in “The Microbe’s Contribution to Biology” (A. J. Kluyver and C. B. van Niel, eds.), p. 155. Harvard Univ. Press, Cambridge, Massachusetts, 1966. 8. Amon, D. I., Nature 184, 10 (1959). 9. Arnon, D. I. i n “Light and Life” (W. D. McElroy and B. Glass, ads.), p. 489. Johns Hopkins Prkss, Baltimore, Maryland, 1961. 10. Stanier, R. Y., Bacteriol. Rev. 25, I (1961). 11. Calvin, M., J. Chem. SOC.1896 (1956). 12. Calvin, M., i n “Radiation Biology and Medicine” (W. D. Claus, ed.), Chapter 31, p. 826. Addison-Wesley, Reading, Massachusetts, 1958. 13. Calvin, M., Revs. Modern Phya. 31, 147 (1959). 14. Calvin, M., Brookhaven Symposia in B i d . 11, 160 (1968). 15. McCollum-Pratt Symposium on “Light and Life.” Johns Hopkins University Press, Baltimore, Md., 1961; see especially essays by A. Jagendorf, D. I. Amon, End A.

San Pietro. 16. Jagendorf, A. T., and Forti, G., in “Light and Life” (W. D. McElroy and B. Glass, ads.), p. 676. Johns Hopkins Press, Baltimore, Maryland, 1961. 17. Wilstattcr, H., and Stoll, A., “Untersuchungen iiber Chlorophyll.” Springer, Berlin 1913; Engl. ed., Science Press, Lancmter, Pennsylvania, 1928. 18. Fischer, H., and Stern, A., “Die Chemie des Pyrrols,” Vol. 11,Part 2: Pyrollfarbstoffe, Edwards Bros. Ann Arbor, Michigan, 1943. 19. Woodward, R. B. el al., J . Am. Chem. SOC.82,3800 (1960).

QUANTUM CONVERSION IN CHLOROPLASTS 20. 21. 22. 23. 24. 25. 26. 27.

33

Seely, G. R., and Calvin, M.,J. Chem. Phya. 23,1068 (1955). Gaffron, H., Biochena. 2.264, 251 (1933). Schenck, G. O., Angew. Chem. 69,579 (1957). Krmnovskii, A. A., Ann. Rev. Plant Phyaiol. 11, 363 (1960). Livingston, R. C., Radiation Reaearch Suppl. 2, 196 (1960). Park,R. B., andPon, N. G.,J. Mol. B i d . 3 , l (1961). Eley, D. D., Nature 162,819 (1948); Reaearch (London)12,293 (1959).

Calvin, M., i n “Light aad Life” (W. D. McElroy and B. Glass, eds.), p. 317. Johns Hopkins Press, Baltimore, Maryland, 1961. 28. Calvin, M., J. Theoret. B i d . 1, 258 (1961). Calvin, M. Science 138, 867 (1962). 29. Benesi, H. A,, and Hildebrand, J. H., J. Am. Chem.SOC.70,3978 (1948); Hildebrand, J. H., and Benesi, H. A., Nature 163, 963 (1949); Benesi, H. A., and Hildebrand, J. H., J . Am. Chem.Soc. 71,2703 (1949); 72,2273 (1950). 30. Kearns, D. R., and Calvin, M., J . Chem. Phys. 29, 950 (1958). 31. Kearns, D. R., Tollin, G., and Calvin, M.,J. Chem. Phys. 32,1013 (1960). 32. Tollin, G., Kearns, D. R., and Calvin, M., J . Chem. Phya. 32,1020 (1960). 33. Aalsberg, W. Ij., Hoijtink, G. J., Makor, E. L., and Weijland, W. P., J. Chem. SOC. p. 3055 (1959). 34. Kearns, D. R., and Calvin, M., J . Am. Chem.SOC.83,2110 (1961). 35. Eastman, J. W., Thesis, University of California (Berkeley), 1961; University of California Radiation Laboratory Report UCRL-9722, 1961. 35a. Eastman, J. W., Androes, G. M., and Calvin, M., J . Chem. Phya. 36,1197 (1962). 36. Chesnut, D. B., Foster, H., and Phillips, W. D., J . Chem. Phys. 34,684 (1961). 37. Acker, D. S., Harder, R. J., Hertler, W. R., Mahler, W., Melby, L. R., Benson, R. E., and Mochel, W. E., J . Am. Chem.Soc. 84, 3370, 3375, 3387 (1962). 38. Kepler, R. E., Bierstedt, P. E., and Merrifield, R. E., Phya. Rev. Letteva 5, 503 (1960). Kepler, R. E., Chem. Eng. News 39, 1961, 42; C. & E.N. Special Report, Organic semiconductor research. Ibid. 40, 86 (1962). 39. Arnold, W., and Clayton, R. K., Proc. Nutl. Acad. Sci. U.S. 46, 769 (1960). 40. Arnold, W., andMaclay, H. K., BrookhavenSympoeia i n B i d . 11, 1 (1958). 41. Commoner, B., Heise, J. J., and Townsend, J., Proc. Natl. Acad. Sci. U.S. 42, 710 (1956). 42. Commoner, B., Heise, J. J., Lippincott, B. B., Norbert, R. E., Passoneau, J. V., and Townsend, J., Science 126, 57 (1957). 43. Calvin, M., and Sogo, P. B., Science 125, 499 (1957). 44, Sogo, P. B., Pon, N. G., and Calvin, M., Proc. Natl. Acad. Sci. U.S. 43, 387 (1957). 45. Calvin, M., Reva. Modern Phya. 31, 157 (1959). 46. Sogo, P. B., Jost, M. R., and Calvin, M., Radiation Research Suppl. 1, 511 (1969). 47. Sogo, P. B., Carter, L. A., and Calvin, M., in “Free Radicals in Biological Systems” (M. S. Blois et al., eds.), p. 311. Academic Press, New York, 1961. 48. Emerson, R., Chalmem, R., Cederstrand, C . , Proc. Natl. Acad. Sci. U.S. 43, 135 (1957). 49. Govindjee, Rabinowitch, E., and Thomas, J. B., Biophys. J . 1 , Q l(1960). 50. Ichimura, S., and Rabinowitch, E., Biophya. J. 1, 99 (1960). 51. Govindjee and Rabinowitch, E., BiOphy8. J . 1, 377 (1960). 52. Tollin, G., and Calvin, M., Proc. Natl. Acad. Sci. U.S. 43, 897 (1957). 53. Tollin, G., Fujimori, E., and Calvin, M., Nature 181, 1266 (1958). 54. Tollin, G., Fujimori, E., and Calvin, M., Proc. Natl. Acad. Sci. U.S. 44, 1035 (1958). 55. Lundegardh, H., Biochim. et Biophya. Acta 35,340 (1959). 56. Duysens, L. N. M.,Science 121, 210 (1956).

34

MELVIN CALVIN

57. Kamen, M. D., in “Light and fife” (W. D. McElroy and B. Glass, eds.), p. 436. Johns Hopkins Press, Baltimore, Maryland, 1961. 58. Chance, B., and Smith, L., Nature 175,803 (1955). 59. Smith, L., in “Light and Life” (W. D. McElroy and B. Glass, eds.), p. 436. Johns Hopkins University Press, Baltimore, Maryland, 1961. 60. Duysens, L. N . M., PTOC. Third Intern. Congr. on Photobiol., 1960 in press. 61. Witt, H. T., Naturwissenachuften 42, 72 (1955);Witt, H. T.,Moraw, R., and Muller, A.,Z.physik.Chem. (Frankfurt)[N.F.]20,197(1959);Witt,H.T.,andMuller,A.,iMd. 21, 1 (1959);Witt, H. T.,Moraw, R., Muller, A,, and Zieger, G., 2. Elektrochem 64, 181 (1960);Witt, H.T.,Muller, A., and Rumberg, B., Nature 191, 194 (1961);192, 967 (1961);Zieger, G.,Muller, A. and Witt, H. T., 2.physik. Chem. (Frankfurt)[N.F.] 29, 13 (1961). 62. Chance, B., Nature 189,719 (1961). 63. Chance, B., andNishimura, M., Proc. Natl. Acad.Sci. U.S. 46,19 (1960). 6 4 . Muller, A., and Witt, H. T., Nature 189,944 (1961). 65. Duysens, L. N. M., Amesz, J., and Kamp, B. M., Nature 190,610(1961). 66. Bishop, N. I., Proc. Natl. Acad. Sci. U.S. 45, 1696 (1959). 67. Crane, F. L., Plant Physiol. 34, 128 (1959). 68. Crane, F.L., Ehrlich, B., and Kegel, L. P., Biochem. Biophys. Research Communa. 3, 37 (1960). 69. Lester, R. L., and Crane, F. L., J. Biol. Chem. 234,2169 (1959). 70. Karnen, M. D., in “Enzymes: Units of Biological Structure and Function” (0. H. Gaebler, ed.), p. 483. Academic Press, New York, 1956. 71. Haxo, F. T., in “Comparative Biochemistry of Photoreactive Systems” (M. B. Allen, ed.), p.339. Academic Press, New York, 1960. 72. French, C.S., in “Light and Life” (W. D. McElroy and B. Glass, eds.), p. 471.Johns Hopkins Press, Baltimore, Maryland, 1961. 73. Allen, M. B., Piette, L. H., and Murchio, J. C., Biochem. Biophys. Research Commum. 4, 271 (1961). 74. Calvin, M., Perspectives in Biol. and Med. 5, 147 (1962).

The Catalytic Decomposition of Formic Acid P. MARS, J. J. F. SCHOLTEN,

AND

P. ZWIETERING

Central Laboratory Staatsniijnen in Limburg, Qeleen, Netherlands I. ntroduction ........................................................ 5 11. General Aspects of the Reaction. ...................................... 38 111. Dehydrogenation of Formic Acid.. .................................... 39 39 A. Dehydrogenation on Metals.. ...................................... B. Dehydrogenation on Metal Alloys. .................................. 72 C. Dehydrogenation on Oxides. ....................................... 81 IV. Dehydration of Formic Acid.. ......................................... 90 A. Dehydration of Formic Acid in the Liquid Phase.. . . . . . . . . . . . . . . . . . . . 90 B. Dehydration of Formic Acid on Metals.. ............................ 92 C. Dehydration on Oxides. ........................................... 96 103 V. Decomposition of Bulk Formates.. .................................... VI. General Discussion. .................................................. 106 References ......................................................... 110

I. Introduction Research workers who try to arrive a t a better understanding of the fundamentals of catalysis favor the study of very simple reactions. The reason is of course that, in a first approach, they want as much as possible to minimize the problems, complicated enough as they are. Various reactions appeared to be very useful in this respect. Wellknown examples are the hydrogen-deuterium exchange and the hydrogenation of ethylene on metals and oxides. Such studies deepened the understanding of hydrogenation reactions in general, while these reactions could a t the same time serve as a means of comparing the activities of different catalysts. Other examples of reactions well suited for use in fundamental catalysis research, are the synthesis and decomposition of ammonia, and the decomposition of N,O. The study of the catalytic aspects of the decomposition of formic acid started already half a century ago with the work of Mailhe and Sabatier (I),who studied this reaction on various oxides and metals. They 35

36

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

observed that formic acid can decompose in three directions:

HCOOH

+ co,

4 H,

(1)

H,O

+ CO

(2)

H,O

+ CO, + CH,O

(3)

and investigated the activity and selectivity of their catalysts. Since that time this catalytic process has continuously been studied. One of the main reasons for the preference given to this reaction is obviously the C), while furthermore the simplicity of the reacting system (A --f B analytical determination of the reaction products is easily performed; a final important aspect is that on almost all metals and metal oxides the reaction proceeds a t a measurable rate in easily accessible regions of temperature and pressure. Although this research work, carried on for many years, resulted in a vast amount of valuable data, a closer study of the catalytic behavior of metallic compounds may be said to have been initiated by Schwab and Rienacker, who, from 1936 onwards, made extensive researches t o investigate the decomposition of formic acid on a large variety of metals and metal alloys. These workers primarily attempted t o find an answer t o the question whether or not the rate of catalytic reactions on metals in general is influenced by alloying these metals with others in such a way that the electron concentration increases. I n a later stage of their research, however, Schwab and Rienacker concentrated their attention 'more and more on the HCOOH decomposition as such, and finally arrived a t more pronounced ideas on the mechanism. Especially the conflict between the results of Rienacker and Hansen ( 2 , 3 ,3a) and those of Schwab (4)-which authors arrived a t a negatively and a positively charged intermediate, respectively-gave rise to a more general interest. I n the last ten years the exploration of the kinetics and the mechanism of these reactions has become very fruitful, owing t o the continuous effort of research workers in different countries all over the world, who made use of all the modern tools now available in catalytic research. Especially the simultaneous application of different measuring techniques, for instance the infrared study of adsorbed reaction intermediates combined with the study of kinetics and kinetic isotope effects, appeared to be very elucidating. As mentioned above, both dehydrogenation and dehydration of formic acid can occur. On metals the dehydrogenation reaction is observed almost exclusively, but the behavior of oxides is more complicated: on some oxides, such as A1,0, and SiO,, the dehydration

+

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

37

reaction prevails, whilst on others, e.g., MgO and ZnO, the dehydrogenation dominates. On catalysts like chromium or iron oxide both reactions occur simultaneously. Furthermore, the selectivity of some oxides may change in dependence on pretreatment conditions; preheating to high temperatures (say 1000°C) for instance, increases the tendency to dehydrogenation. An analogous situation exists in the decomposition of alcohols, where also a dehydrogenation (to aldehyde and hydrogen) and a dehydration (to olefin and H,O) is observed. I n an earlier publication by Mars ( 5 ) the close analogy between the selectivities of the different oxides for both decomposition reactions was already demonstrated. We shall make use of this analogy in our further discussions of the mechanism of the formic acid decomposition in those cases where sufficient direct information is lacking. As reported already in the original work by Sabatier, there is a third, though weak, reaction besides dehydrogenation and dehydration, which leads to the formation of formaldehyde. On specifically dehydrating catalysts only traces of formaldehyde can be found; more is formed on oxides which have also a dehydrogenating action, e.g., Cr,O, [Dalmai et al. (S)]; the formation of this compound on metals has not been reported, however. It is probable that the lowering of the activity of the catalyst during its action is caused by the formation of polymers of formaldehyde on the catalyst surface. If the formation of formaldehyde is hindered by the pressure of water not only on A1,0,, as shown by Menon, but on all oxide catalysts, it can be understood that the activity of Fe,O, is constant with time only if a high water vapour pressure is maintained [Mars (271. Sabatier and Mailhe (1)report that on Tho, also methyl alcohol is formed. As only very few data on the formation of these byproducts are available, and nothing is known about the mechanism involved, these reactions will be ignored in the following discussions. As it appeared that the main reactions on different catalytic systems had many features in common, the decomposition of formic acid will in the following chapters be discussed along the lines given by the two main reaction paths, Thus the dehydrogenation both on metals and oxides will be treated in the first part and the second part will be devoted to the dehydration reaction. As the latter reaction is also largely catalyzed in the liquid phase by strong acid, a short discussion of this phenomenon has been included. I n the final chapter we shall give a discussion on the kinetics and mechanisms occurring in the catalytic decomposition of HCOOH, together with our view on the significance of the research carried out in this field for the theory of heterogeneous catalysis in general.

38

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINU

II. General Aspects of t h e Reaction I n many cases encountered in literature the catalytic decomposition was studied in the gas phase:

This can easily be done in view of the relatively high volatility of the acid; the normal boiling point is 100.75"C, and a t 25OC the vapor pressure is 43 mm Hg. The heat of association needed to form the dimer (HCOOH),, is rather low, viz., 14.11 kcal/mole of dimer*; calculation shows that a t reaction

FIG. 1. Reactor for studying the dynamical decomposition of formic acid, according to Schwab ( 1 0 ) . The explanation of the symbols is given in the text.

temperatures above 150 "C and a relative formic acid pressure of ca. 0.1 atm the dimerization in the gas phase is virtually negligible. Furthermore reaction ( I ) is slightly exothermic, AH," being -7.38 kcal/mole.

* The thermodynamic data in this paragraph are taken from W. Waring ( 8 ) .

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

39

The standard free energy of reaction, AF,", amounts to - 14.02 kcal/ mole. The dehydration, however, is slightly endothermic; AH," = +2.45 kcal/mole. The standard free energy for dehydration amounts to - 7.21 kcal/mole. The experimental arrangements used for kinetic studies are generally quite simple. Rienacker (S),for instance, studied the decomposition in a static way by bringing a known amount of HCOOH vapor over the catalyst and measuring the change in pressure as a function of time. Analysis was performed by observing the pressure change on cooling the product gas in a trap, first a t - 79"C, when HCOOH and H,O were frozen out, and next down to - 183"C, which resulted in the condensation of CO,. A beautiful device for a dynamical study is the so called "Schwab reactor" (20) (see Fig. 1). Liquid formic acid is contained in vessel A, which communicates with the evaporator B filled with glass particles. At a certain point, the liquid in B starts to boil, owing to the heating by furnace C, and the formic acid vapor passes a t atmospheric pressure over the catalyst in E. Undecomposed acid is recirculated via the condenser G. The gaseous reaction products are taken away a t J for analysis; they pass through a flowmeter K. M-M is the outer tube of a thermocouple, removably connected by a ground joint a t the top of the apparatus. The reaction velocity can be directly noted a t different catalyst temperatures. Another method well suited for dynamic studies is the use of an inert carrier gas which is loaded, in a vaporizer and a t constant temperature, with formic acid of a certain partial pressure, and passed over the catalyst [Riengcker and Dietz ( I I ) ] . The analytical determination of the reaction products does not present great difficulties. Besides classical gas analysis, mass spectrometric analysis with the spectrometer directly connected to the reactor is very convenient; if this method is used, all the products of the reaction can be analyzed simultaneously [Fahrenfort et al. (12)].

Ill. Dehydrogenation of Formic Acid A, DEHYDROGENATION ON METALS 1. Qeneral Aspects of the Kinetics

In their early researches on the decomposition of HCOOH on metals and oxides Hinshelwood and Topley (13) as early as in 1923 found re-

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINU

40

action rates of zero order in formic acid pressure, a t low temperature and pressures of the order of one atmosphere. A gradual change to first order kinetics could be effected by increasing the reaction temperature and/or lowering the pressure, They related this phenomenon to the adsorption of a reaction intermediate obeying the simple Langmuir kinetics. This idea seems to be a good starting point for the description of the kinetics of the dehydrogenation on most metals. Assuming an irreversible adsorption of formic acid on the metal surface and an irreversible decomposition of the adsorbed intermediate, we arrive a t the following reaction scheme: kl+

HCOOHgasphase

-

adsorbed intermediate

’

+ -

cos -!- Hs

For the steady state we thus get or ei

=k

klP

+ k,p

where

ei = coverage by intermediate,

k, k

= rate constant of formic acid adsorption,

= rate constant of intermediate decomposition, p = formic acid pressure. The rate of reaction r is:

r = ke,

so that 1

1

-==rc+-

r

1

klP

At high pressure the whole surface is covered by intermediates

(8,= l ) , so that

r=k

(8)

In this case the reaction rate is pressure-independent and the activation energy measured is the activation energy of the decomposition of the intermediate. At Iow pressure (1 - Oi) i5 about equal to unity so that i t follows from Eqs. (4)and (6) that r

=

k,p

(9)

Under these conditions the reaction is first-order in formic acid pressure and the activation energy measured equals that of the adsorption step.

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

41

A quantitative elaboration of these simple kinetic equations will be given here for the reaction on platinum [Block and Vogl(14)]; however, the kinetics on other metals can, in the first instance, be handled in the same way. Figure 2 shows that a plot of l/r as a function of l/p at

-1r (rnin rnrn

t

20

10 --c

30

l ( r n r n Hg’‘) P

FIG. 2. The kinetics of the decompositionof formic acid on platinum at 200 and 260°C. The points are derived from Table I in Block and Vogl’s psper (14).The reciprocal value of the rate of reaction is plotted aa a function of the reciprocal pressure.

reaction temperatures of 200 and 25OoC gives straight lines, which is in accordance with Eq. (7). From Fig. 2 the temperature dependence of k and k, can be calculated:

I n this way the energy of activation of the formic acid adsorption,

Endsis found to be 6.2 kcal/mole, whereas the activation energy for the is 15.8 kcal/mole. decomposition of the intermediate, Edeoomp,

42

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

From separate experiments performed exclusively in the first-order pressure region, Eadswas found to be 5.5 kcal/mole, a value which is almost equal to that found from the temperature dependence of Ilk, in Eq. (7). If zero order behavior is found it is still possible that only part of the surface is involved in the reaction; in this case Oi in Eq. ( 5 ) represents the coverage on an exclusive part of the surface which is active only. The experiments of Tamaru (15)are interesting in this respect. This author measured the rate of the decomposition of formic acid on a nickel sample in the zero-order region. By combining pressure readings and gas analysis as a function of time, and working with a closed reaction system in which the gas circulates, he was able to determine the formic acid surface coverage under reaction conditions. It appeared that the whole surface was covered during the reaction, every surface Ni-atom adsorbing two formic acid molecules. On heating, the surface complex formed decomposed into two CO, molecules and one H, molecule; this observation is of importance in view of determination of the structure of the adsorbed reaction intermediate, which will be discussed in Section 111,A, 7. Tamaru’s results with nickel were beautifully confirmed by Fahrenfort et al. ( l a ) ,who also arrived a t an adsorption of two HCOOH molecules per Ni-site. The question whether all surface sites, or only special centers, such as angular points or crystal edges, are active, was also raised by Rienacker and Volter (16, 17, 18). They measured the rate of the decomposition of formic acid on two polycrystalline copper samples widely differing in crystallite size; and hence in the length of their crystallite boundary lines. The sample having the longest boundary lines showed a somewhat lower activity, which argues against the view that special active centers are involved. Using the same measuring technique as in his experiments with Ni, Tamaru (19)made measurements on copper in the zero-order region and found a coverage of about one formic acid molecule on two surface copper sites. Consequently, for Ni and Cu apparently all surface sites are involved in the reaction. When working with silver, however, Tamaru (15) arrived at a different result; in this case he found a saturation ratio of one formic acid molecule to seven silver surface sites in the zero-order pression region a t 134OC, as against a ratio of one to ten a t 160°C. This would mean that only part of the Ag surface were active. With regard to this last result the present authors would express some doubt. Perhaps it might be accounted for by part of the surface being

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

43

composed of oxidic contaminations, or by a slight poisoning of the Ag surface. The high activation (30.7 kcal/mole) observed by Tamaru seems to indicate such poisoning: Schwab ( 2 4 ,Rienacker and Dietz (U), and Block and Kral (21)arrived a t 19, 18.8, and 19.8 kcal/mole, respectively, whereas Sosnovsky (22)for a polycrystalline Ag specimen, found a value of 25.5 kcal/mole. I n an analysis of the kinetics of the dehydrogenation of formic acid on gold Sachtler and Fahrenfort (23) also met with the difficulty presented by an exact determination of the amount of HCOOH adsorbed during the reaction. They assumed that the coverage in the zero-order pressure region would be equal to the saturation value found from an adsorption isotherm. The precision of their adsorption measurements was poor, due to formic acid decomposition accompanying adsorption even at 25"C, but as a rough indication it may be said that about 1 HCOOH was adsorbed on four surface gold atoms at 25°C and 20 mm Hg. I n conclusion it may be stated that the study of surface coverage during synthesis deserves more research, especially the coverage on the noble and seminoble metals. Special precautions have to be taken to prevent poisoning or contamination by non-metals invalidating the observations. Finally, our ignorance about the precise distribution of exposed crystallographic planes-which determines the number of metal surface sites per cm2-and our lack of knowledge about the number of sites necessary to the adsorption of one formic acid molecule, have to be taken into account, Only if the number of adsorbed molecules appears to be much (say, ten times) lower than the estimated number of exposed metal sites, would the conclusion be justified that only part of the metal surface atoms are involved in adsorption and catalysis. Studies of the decomposition of formic acid on both base and noble metals showed that generally the change in reaction order, attending it change in experimental conditions, may be explained as has been done above for platinum. For the particular case of nickel this was shown by Fahrenfort et al. (12). This suggests that the first reaction step on all metals is the adsorption of formic acid, and that the adsorbed intermediate decomposes in the reaction products which desorb from the surface. The coverage (lifetime) of the intermediate on the surface will determine the reaction order observed under certain conditions of temperature and pressure. As regards this last point, the study of the decomposition of formic acid on gold has led to a better understanding of the difference between noble and base metals. I n the temperature range 110-190°C Eley and Lueti6 (24)observed an

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINQ

44

exclusive decomposition into CO, and H, on gold wires reduced and ,cleaned in a hydrogen discharge. At a pressure of 6 x lo-* mm the reaction rate r was found to answer to the following expression:

r

=

2 x lo6 exp (-- 14,30O/RT) molecules site-' sec-l

(12)

In the calculation of this expression from the observed reaction rate it was assumed each formic acid molecule was held on two lattice points, and that the number of double sites per cm2 was 0.6 x 1015 for gold. The pressure-dependencewas not determined and no information about the coverage under reaction conditions was obtained. In view of the nearly first-order behavior found by Sachtler and Fahrenfort at much higher pressure (see hereafter) we think it most plausible that the reaction at 6 x 10-4 mm was first-order in pressure. In that case, their experimental result can be recalculated to be :

r = 2 x loz3 x p x exp ( - 14,30O/RT) molecules ( p in mm Hg).

sec-l

(13)

Sachtler and Fahrenfort (25) studied the decomposition of HCOOH on a silica-supported gold catalyst. Working at a pressure of 15 mm Hg and at temperatures between 100 and 26OoC, they arrived at the following rate expression :

r

= 0.9

x lozoexp ( - 14,30O/RT) molecules

sec-1

(14)

The pressure-dependence was studied only at 2OO0C,between 0.12 and 16 mm Hg; the rate appeared to be proportional to PO.'^. If this deviation from first-order kinetics is neglected, their result can be rewritten as: T

= 6 x l0ls

x p x exp ( - 14,30O/RT) molecules cm-, sec-'

(15)

( p in mm Hg).

It is seen that whereas excellent agreement is found for the value of the activation energy, the pre-exponential factors appear to differ very considerably [compare Eqs. (13) and (15)]. A certain degree of formic acid dimerization [Sachtler (26)] in Sachtler and Fahrenfort's experiment at 15 mm Hg can hardly be the cause of this difference in behavior, since at the lowest reaction temperature (100°C) the degree of dimerization is only 9%. In the case of a rate-determining adsorption step, the theory of absolute reaction rates [Glasstone et aE. (27)l gives at 180°C r = 0.72 x l O l 4 x p x exp (-E,,,/RT) molecules cm-, sec-l (16) ( p in mm Hg)

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

45

for an immobile activated complex,

r = 2.47 x loz1 x p x exp (-E,,,,/RT) molecules cm-2 sec-l ( p in mm Hg)

(17)

for a mobile activated complex with completely preserved rotation, and r = 5.3 x lOle x p x exp (-E,,,/RT) molecules cm-2 sec-1 ( p in mm Hg)

(18)

for a mobile activated complex with only one degree of rotational freedom. Theoretical value ( 18') is in rather good accordance with the rate found experimentally by Sachtler and Pahrenfort* ; however Eley and Lueti6's result is far too high, even for the case of fullest mobility. An explanation of this last discrepancy is not available. The activation energy of the adsorption step is rather high, as it is, for instance, in the case of adsorption on the seminoble metal copper (see Section 111,A, 3). This is in strong contrast with the behavior of the less noble metals. On Ni, for instance, E,, is found to be much smaller than 5 kcal/mole (see Section 111, A, 7 ) . This difference in behavior between noble and base metals will be further dealt with in the final discussion. I n the zero-order pressure region only some measurements by Schwab (28)are on hand. The rate of decomposition in this range is given by:

r

=

1021.5x p o

x exp ( - 12,500/RT)molecules cm-2 sec-l

(19)

Hence, the activation energy for decomposition of the adsorbed intermediate is nearly as high as the activation energy for the adsorption of HCOOH on gold. In view of the discrepancies mentioned before, a thorough re-investigation of the kinetics on gold seems desirable. 2. The Influence of the Reaction Products on the Kinetics I n a static decomposition of formic acid on nickel, in the zero-order pressure region, the reaction rate may be expected to remain constant as long as the formic acid pressure remains high enough and full coverage is maintained. Hence, in such a case the pressure in the reaction vessel will be observed to rise in linear relation to time. However, Volter (291, working with nickel, noticed that under these conditions a deviation from this linear relationship occurred a t 60-90 yo

* Sachtler and Fahrenfort concluded upon immobile adsorption. However, in their calculation these authors neglected the disappearanoe of the rotational entropy (24 e.u.).

46

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINO

conversion (dependent on pressure) which implies a deviation from zeroorder kinetics. This phenomenon was further studied by Rienacker and Hansen (3), and by Rienicker and Volter (30). On copper sheets, silver powder and iron powder, the decomposition was always perfectly zero-order, but on nickel “inhibited zero-order kinetics” were found (see Fig. 3), and the A D 61 (mm Hg)

t

=-

4-

P ~ = 3 . 6rnrn 0 Hg

32-

P ~ = l 6 0 m r nHg

1I 10

20

30

--.

40

rnin

FIG.3. Pressure increase aa a function of time in the static decomposition of formic acid on Ni-films at 190°C with different initial pressures P, [according to Rienacker and Hansen (3)].

authors tried t o correlate this with the adsorption of reaction products on the surface. It appeared, however, that preadsorption of H,O, CO,, and HCHO had no influence on the kinetics, whereas CO preadsorption gave only a weak retardation, certainly much too weak to account for the effect observed. Hydrogen preadsorption produced an effect of the same order of magnitude as CO preadsorption, but in the form of an acceleration. Thus, an explanation of the inhibited zero-order kinetics could not be found. Walton and Verhoek (31) re-examined this problem in the staticdecomposition on nickel-filmsa t 189°Cby adding largeamounts of product gas to the formic acid and measuring the influence on the initial reaction rate. Some of their results are given in Table I. This table shows that in fact none of the product gases has a pronounced influence. From the figures it can be estimated, however, that in a mixture of, for instance, 1 CO,, 1 H,, and 1 HCOOH, the retardation would be in the order of a factor two. This retardation effect is comparable with the retardation observed by the same authors in a static

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

47

decomposition experiment in the absence of added gases, a t 50% conversion [Walton and Verhoek (32)]. I n our opinion, therefore, chemisorption of H, and/or CO, on the nickel surface seems the cause of the inhibited zero-order kinetics. TABLE I

The Injluence of Reaction Products o n the Kinetics, According to Walton and Verhoek (31) Pressure of added gas (mm Hg)

Initial pressure of HCooH

Initial rate molec cm-a sec-110-9

0.05 CO 8.20 co 17.70 CO 0.09 CH, 6.80 CO, 17.30 CO, 14.7 H, 11.6 H,O

23.2 19.8 8.4 19.2 12.7 6.0 17.1 24.2

2.80 3.02 2.02 3.04 2.67 1.24 1.76 1.65

(mm Hg)

3. The Ejj’ect on the Kinetics of Dijj’erent Modes of Pretreatment and of the Crystallographic Surface Orientation It is well known to research workers that the behavior of catalysts is sometimes dependent on the way in which the samples are pretreated. Reduction time and temperature, ion bombardment, rolling, electropolishing, the predominance of certain crystallographic planes of the surface, all these factors may have a great influence on activity and kinetics. I n this field, again, the simple formic acid decomposition appears t o be a widely-used test reaction. I n this conjunction it is in particular the interesting studies by Rienacker and his school that deserve to be mentioned. Initially Rienacker (33) found that heavy rolling of a nickel sheet raised the activation energy of the formic acid decomposition from 20 to 22.5 kcal/mole, whereas heating of the sheet t o 1000°C resulted in a decrease to 13 kcal/mole. This work was afterwards duplicated by Schwab (20). I n a static decomposition on hammered Ni he found 21.1 kcal/mole; a dynamic experiment gave 25.5 kcal/mole. Heating up to 500” resulted in a value of 18 kcal/mole in a static experiment. I n this way Rienacker’s original results were qualitatively confirmed.

48

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

At a later stage the question was reinvestigated by Rienacker and his school. First they compared the behavior of silver and copper powders with that of sheets of the same metals [Rienacker et al. (34, 35)]; they found that after reduction in hydrogen above 400°C and heating a t 900°C, each metal gave the same activation energy in powder and sheet form, From this result they concluded that obviously the activation energy is a “material constant’’ when a rigorous pretreatment has beer1 applied. Next, a study was made of the effect of heating and rolling on nickel, iron, and nickel-iron alloys in the form of powders and sheets [Rienacker and Volter (18, 30)]. Here, again, it was found that the rolled samples, both of the pure metals and of the alloys, gave activation energies which were 3-5 kcal/mole higher than those for the heated samples. According to Rieniicker and Volter different explanations of the rolling effect can be given. It is known that rolling effects a crushing of the crystallites a t the surface, which results in a lengthening of the crystallite boundaries; in the case of copper, however, this had no influence on

1.9

2: 0

-g 2.1

2: 2

T

FIG.4. Decomposition of formic acid on different planes of copper single crystals, according to Rienlcker (17).

the activity (see Section 111,A, 1) and this result seems to rule out this explanation. Another effect of rolling, well known to metallographers, is the occurrence of a preferential orientation of certain crystallographic planes parallel to the surface. Hence, if different crystallographic planes exhibit different catalytic activities this may be understood as a rolling effect. I n connection with what has been said above it was interesting to compare the catalytic activity of different single-crystal planes. For this

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

49

purpose several discs were cut from a copper single crystal in the (111) and (100)directions [Rienacker (17)l. After electropolishing an electronmicroscopic inspection showed the surfaces to be smooth and undisturbed. Figure 4 gives an Arrhenius plot of the reaction rate for three HCOOH decomposition experiments with (111) and'three with (100) discs. I n all runs the decomposition was zero-order and the activation energy amounted to (23.6 f 0.5) kcal/mole. The interesting feature was that on the octahedral planes the reaction rate was three to four times higher; yet, the number of sites per cm2 on the two types of planes does not differ very much: (100) 1.44 x 1015sites, (111) 1.78 x l O l 5 sites. I n the decomposition of hydrazine, however, the opposite was observed (Fig. 5) ; log,, r

(molecules cm-* stc.')

I

7.0-

6.5

-

b.04

I

1.9

. 2.0

- lo'

~

2.1

T

FIG. 6. Decomposition of N,H, on different planes of copper single crystals, according to Rieniicker (17).

for different planes the activation energy is constant (ca. 23 kcal/mole), but now the cubic plane is the more active one. Thus, for both reactions a catalytic specificity of the crystal planes was found. Cunningham and Gwathmey (36)studying ethylene hydrogenation on different planes of a nickel single crystal, also arrived at the conclusion that planes cut in different crystallographic directions exhibit different activities. Formic acid decomposition on copper single crystals with different surface orientation was also studied by Crocker and Robertson (37). Their results are not directly comparable with those of Rienacker and Volter ; the experiments were performed in the low-pressure range mm Hg, and at higher temperatures, viz. 250-450°C (see Fig. 6). The rate of reaction was measured by a method similar to that

50

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

used in mass-spectrometric research work. Formic acid flows continuously into the reaction vessel containing the catalytic sample, and the formic acid and the reaction products are pumped off simultaneously. From the change in pressure observed in the vessel with and without the catalyst the reaction rate can be calculated. Under all experimental conditions first-order behavior was found, so that it may be concluded that the activation energy is an apparent one, equal t o the activation energy of the adsorption step. The reaction rate r is expressed as the reaction probability P per collision: T = P X

P

___-

2/277mkT

where: p/z/2xmkT = the Hertz-Knudsen expression, giving thenumber of gas molecules striking 1 cm2 of a surface per sec; this number is equal t o the adsorption rate for a mobile activated complex if Eads= 0 [Glasstone et al. (27)]. The probability of reaction P is given by

P

= B exp

(-E,/RT)

(21)

where

B E,

= the pre-exponential factor, and = the apparent activation energy.

For log,& and E, the following values were found: Surface orientation

log,,B

E, (kcal/mole)

_____

.~

-

19 16 8.5

Apparently, under these conditions the specificity of the different orientations manifests itself in a variation of the En-values, which is perfectly compensated by the change in the pre-exponential factor B, SO that the actual rates are not very different. The values of B support the idea that in the adsorption step of the reaction a mobile transition state is involved. On polycrystalline copper the rate was of the same order as on the single-crystal faces, which is in accordance with the results of Rienacker and Volter mentioned in Section 111,B, 1. I n a later stage of his research, Robertson, in cooperation with Duell, found results deviating widely from those reported here (see the dotted line in Fig. 6 ) . I n Section 111, A, 4 we shall return to these measurements.

61

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

Further evidence of the specificity of different crystal planes emerges from the work of Sosnovsky (22, 38, 39). This author studied the decomposition of HCOOH on single crystals of silver which had been grown from the melt in the form of thin sheets in such a way that the surface had a predetermined orientation. The crystal surfaces, cut in

‘0910

‘t

I

16

18

20

--

i

22

10’ T

FIU. 6. The decomposition of formic acid on different crystal faces of copper single crystal discs, after electropolishing and prolonged reduction of the discs in hydrogen at 400°C. The logarithm of reactionprobability P is plotted aa a function of 10*/T.0:copper (111) face; A : copper (110) face; 0 : copper (100) face. The dotted line refers to a result mentioned in Section 111, A, 4. [According to A. J. Crocker and A. J. B. Robertson (37).]

the ( l l l ) , (110), and (100) directions, were polished chemically and electrolytically, and then bombarded with positive argon ions in a low pressure discharge a t voltages varying between 14 and 4000 volts. Immediately after this pretreatment the specimens were transferred to a reaction vessel, annealed in vacuo a t 25OOC for 18 hr, and then the decomposition rate of formic acid was measured dynamically between 150 and 250°C. I n the pressure range 50-90 mm the reaction was zeroorder on all specimens investigated and the sole decomposition products were CO, and H,. Sosnovsky’s data are plotted by us in Fig. 7. I n the calculation of the pre-exponential factor we assumed that the roughness factor of the surfaces is of the order of unity; this is in accordance with Sosnovsky’s observation that under all conditions of pretreatment the surfaces remained perfectly smooth. From Fig. 7 it is seen that, dependent on the surface orientation and the energy of the bombarding ions, the activation energies are found to vary from 12 to 35 kcal/mole. However, as was also found by Crocker and Robertson (37) in the case of copper, the change

62

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINQ

of E is strongly compensated by a linear increase of log A with increasing E , resulting in a difference in reaction rate by a factor of only 10 a t 150"C, and 100 at 25OOC. Sosnovsky explains his results in the following

I

A

: surfac* p u . l l e l

t 0 (110)

0 : aurfaco p u d ~ e LO l (1oC)

15

20

25

-

30

35

E (kcallmole)

FIG.7. The logarithms of the pre-exponentialfactor veraus theactivetion energy, for the decomposition of formic acid on single crystals of silver with surfaaes oriented parallel to ( l l l ) , (loo), and (110). Pointeare takenfromdataof Soenovsky (22,38,39). Thenumbera in the graph refer to the voltage applied during ion bombardment. X -samples without special pretreatment: A-( 100)orientation;B-polycrystalline specimen; C( 11 1) orientation. See further text in figure.

way. It is known that owing to ion bombardment disoriented regions are introduced on the surface which are bounded by stable arrays of dislocations. If it is assumed that reaction occurs only at sites where dislocation lines intersect the surface, it may be imagined that owing to ion bombardment a second type of intersection is created, which intersections vary in number according as different voltages are applied. For the (1 11) face, for instance, we then have: reaction rate = k (nle-EIIRT n,e-E*'RT) where k = constant, n1 = number of intersections initially present, E = activation energy on the n 1sites, n2 = number of intersections introduced by ion bombardment, E , = activation energy on the n , sites. Analogous equations hold for the (100) and (1 10) directions.

+

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

53

It has been shown by Constable (40)and Cremer ( 4 1 )that an equation like (22) may, indeed, lead to the compensation effect, if samples with different nl/nzratios are studied, and provided that the measurements are performed in a limited temperature region (dT l0OOC). A detailed picture of the single-crystal Ag surfaces used by Sosnovsky was given by Ogilvie (42). From electron-diffraction patterns he concluded that after ion bombardment the smooth oriented surfaces, though still having a low roughness factor, are composed of small crystallites which expose crystallographic planes in different ratios, depending on conditions of pretreatment and ion-bombardments. Furthermore, it seems likely (see the preceding sections) that all surface sites are active and that each crystallographic plane has its own activation energy. Hence, contrary to the explanation offered by Sosnovsky, we may just as well assume that for a certain distribution of planes the reaction rate is given by:

-

where k = pre-exponential factor, a(,,,, a(111),. . . . . etc. = concentration of sites in (100) plane, (111) plane, etc., Eflll),E(,,,,, . . . . . etc. = activation energy for the plane indicated. For the other combinations of planes, occurring after pretreatment under other conditions, analogous equations will hold. Cremer (41)has shown that for different distributions of a and E values, equations of the shape of (23) lead to the conclusion that the compensation effect is involved, as did equation (22). The interpretation of the compensation effect given above has the disadvantage that it does not make clear that there is also a compensation effect between the E - and log A-values found with perfectly smooth single-crystalsurfaces.Therefore,it is likely that other factors may play a role (see Section VI). Finally it should be remarked that the highest value of log,, A observed is 1 0 2 9 , which is in rather good accordance with the “expected value” 102s [Glasstone et at?. (27)l. N

4. “Catalytic Superactivity” on Nickel and Copper

In the foregoing section we reported already on the research of Crocker and Robertson, who found a catalytic specificity for different surface orientations of copper.

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

54

I n a later stage of his research Robertson, now working on nonoriented samples in cooperation with Duell (43), arrived at results deviating from his original findings (see Fig. 8, where also the mode of

f

-I-

-2-

- 3- 4I

.

.

9

11

0

13

'

15

17

19

-21

lo4 1

FIG. 8. Arrhenius plot of the reaction probabiIity per coIIision, for the decomposition of HCOOH on Cu. According to Duell and Robertson ( 4 3 ) . 0-copper wire outgassed and reduced at 700°C; .-copper disc, electropolished, outgassed, and reduced at 400OC; +--copper wire, flashed for two minutes at 1000°C; - - - - -copper single crystal with (111) surface orientation (taken from Fig. 6). Measurements performed between and 1O-$mm H g . All rates were first order in formic acid pressure.

pretreatment is indicated). The figure gives a comparison of the results with a previous one (dotted line), which shows that especially the copper wire flashed at 1000°C has a somewhat higher activity. The deviation occurs from 425OC upwards to 1000°C; such temperatures are rather high for the study of an organic decomposition reaction. Figure 9 gives the results obtained with nickel wires and discs. Nonflashed samples give a high value for the activation energy (see open circles and points; E = 12.5 kcal/mole up to 850°C). For the flashed samples (crosses in Fig. 9) P was found t o be equal to unity between 330 and 1200°C, which means that every molecule striking the surface is adsorbed without activation energy, However, Duell and Robertson report that especially at the lower temperatures this high activity persisted only for a short time. Duell and Robertson account for the superactivity of the samples above the temperatures where a steep rise in the Arrhenius plots occurs by supposing that at these high temperatures the equilibrium concencentration of different defects (surface vacancies, adsorbed metal atoms etc.) increases considerably, and that on such defects there is a

55

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

much higher probability of decomposition than on the nondefective part of the surface. Flashing produces a large amount of defects, which remain frozen in for a time which is longer than the time of reaction at the lower temperatures; this would be the reason for the constant P = 1 value on nickel.

1

- 1O

+-+-+ +

loglo p

I

+

+-+

-2-

-3-4.

E = 12 k c a l / r n o l c

-5.

7

9

11

13

15 4

lp4 T

FIG.9. Arrhenius plot of the reaction probability per collision, for the decomposition of HCOOH on Ni. According to Duel1 and Robertson (43). 0-nickel wire (nonflashed); .-nickel disc (nonflashed); +-nickel wire after flashing for 2 min at 1330°C. Measurements performed between and mm Hg. Only first-order behavior was found.

I n the preceding sections we have already expressed our objections to the acceptance of special active points. Therefore we would suggest another interpretation. Owing to the high reaction temperatures the samples become covered with carbon and/or other cracking products; this explains the high activation energy found with nickel ( E = 12.5 kcal; see open circles and points in Fig. 9); Fahrenfort and Sachtler (25) report a value smaller than 5 kcal/mole. At temperatures above 850°C the activity rises considerably, owing to the fact that now rearrangement of surface atoms and diffusion of carbon can occur. Finally, the question arises if the adsorbed formic acid molecule decomposes on the surface of the flashed sample a t a higher rate than may normally be expected for the decomposition on a Ni surface, in other words if the flashed sample is actually “catalytically super-active”. Unfortunately not many measurements have been performed at the high temperatures used in these experiments, but we can extrapolate the results of Sachtler and Fahrenfort (25), who studied the reaction between 80 and 200°C. We then arrive a t Table 11.The figures in the second column have been

56

P. MARS, J. J. 5'. SCHOLTEN, AND P. ZWIETERINO

calculated from the Knudsen-Hertz equation [and hence are the rates observed by Robertson and Duell ( 4 3 ) ] ;those in the third column have been derived from the rate of decomposition between 80 and 200°C found by Sachtler and Fahrenfort (25): r = lO'O*l e-~z3~ooo*z000~RT) molecules per site and per second (24) The table shows that only at the lowest temperature the number of molecules striking the surface is higher than the number decomposing, TABLE I1 Rate of the Decompoaitwn of HCOOH on Ni,According to Duell and Robertson ( 4 3 ) ,Compared with the Ratea Extrapolated from Sachtler and Fahrenfort'a Reeulte

.T ("C) 1200 700 330

Number of molecules striking the surface per em*and per second p = 10-'mm 7.3 x 10'8 6.9 x lo1*

4.7 x

1018

Number of adsorbed complexes decomposing per cmr and per second 1.3 X loaa 2 x 1020 1.8 x 1017

but, as a consequence of the uncertainties in the last extrapolated value, we may say that catalytic superactivity of flashed samples is, to say the least, still a dubious point. 5 . Isotope Effects

It is general experience that for establishing the mechanism of a catalytic reaction an extensive analysis of the kinetics is needed, although it is never sufficient by itself. Additional information is always neoessary, and it has often been found that a determination of kinetic and thermodynamic isotope effects, in addition to an analysis of isotope distributions in the decomposition products, can give rather conclusive information. Only a few publications have been devoted to the kinetic isotope effects in the decomposition of formic acid, and only the metals Ag and Au were investigated. The most extensive work was published by Block and Kral (Zl),who measured the rate of the decomposition of formic acid over silver wires between 190 and 25OoC, using the isotopic compositions HCOOH, HCOOD, DCOOH, and DCOOD. The kinetic isotope effect a, which is defined as the number of times the reaction

67

THE CATALYTIU DECOMPOSITION OF FORMIC ACID

-

rate is slower when H in the formic acid molecule is displaced by D, was found to be pressure-dependent at 200°C below p = 10 mm Hg. For instance, aDCOOU was 6.05 at 210°C and 10 mm Hg, but at 0.17 mm and 210°C it was only 2.76. This could be explained from the fact that upon a lowering of the pressure the kinetics gradually change over to first-order kinetics (HCOOH adsorption is rate-determining ; see also Section 111, A, 1); in view of the Hertz-Knudsen equation, the ultimate value of aDCOOD at very low pressures and high temperature-when every collision leads to decomposition-has to be 448/46 1.02. Block and Kral studied the isotope effect in the zero-order pressure region as a function of temperature, using partially deuterated formic acid molecules. The results are summarized in Table 111. TABLE I11 Iaotope Effectafor Different Iaotopic Compo&ona of HCOOH,According Block and Kral (21) %Din Substanoe

C-H-group

%Din 0H.group

HCOOD

0

66.8

DCOOH

66.6

0.8

DCOOD

66.8

42.6

to

GL

190°C

210°C

230°C

250°C

1.35 (1.93) 1.69 (1.71) 4.92 (-2.6)

1.36 (1.91) 1.69 (1.69) 4.77

1.36 (1.86) 1.61 (1.66) 3.99

1.32 (1.82) 1.48 (1.62) 3.60

The values between brackets are calculated from Eyring’s reaction rate theory; in this calculation it was supposed: (a) that the adsorbed reaction intermediate and the transition state are so strongly bound on the surface that partition functions for translations and rotations may be neglected; only vibrational partition functions play a part; (b) that the transition state has the structure of a formic acid molecule in which two H-atoms are “activated”: H . . . . C//

0

‘0..

..H

The vibrational partition functions were calculated from vibrationfrequencies known from the infrared spectra of the different HCOOHisotopic compositions.

58

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

The values thus calculated fitted in best with the measured values, whereas transition states with only one “activated” H-atom: 0

// H . . ..C

0

or H-C

// \O . . . . H

‘OH

showed large deviations from the experimental values. However, an activated complex

was not in full accordance with the experiments as shown in the table. Experimentally the following sequence in reaction rate was found: HCOOH > HCOOD > DCOOH > DCOOD,

whereas according to the theory it should be: HCOOH > DCOOH > HCOOD > DCOOD

The cause of the deviation of the experimental sequence from the theoretical one has not been sought in experimental errors or inaccuracies; it was beautifully confirmed by Otaki’s experiments (44), who reported the following ratio of decomposition rates r (90% deuterated complexes on reduced Ag a t 100°C): rHCOO1l

: rHCOOU : rDCOOH = lo :

:

It would be worth while t o try other structures with two “activated” H-atoms for the transition state, for instance:

... . . H c=o c=o I I0 0 H..

a particle composed of two formate ions adsorbed on the surface. For such a bimolecular .mechanism we would expect: rHCOOH

:r

HCOOD: TUCOOH

: rDCOOD = 1 : 1 : x : 5,

with x < 1. Neither is this sequence in accordance with observation, However, the experimental kinetic isotope effects are perhaps obscured by a secondary kinetic isotope effect, the operation of which is due t o the adsorption of the decomposition product H, or D,. Indeed, Block and

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

59

Kral report that preadsorption of hydrogen accelerates the decomposition of formic acid on silver. If the accelerating effect of D, were weaker than that of H the experimental result would be understandable.

,

The most extensive observations on the isotope distribution in the decomposition products of deuterated formic acid have been made by Hirota and co-workers ( 4 5 ) . On Ag-films the following reactions were studied between 160 and 245°C: ( a ) the static decomposition of HCOOD (OD group 100% deuterated); (/?) the exchange reaction H, D, 2 HD; ( y ) the exchange reaction

+

HCOOH

+ D,

-+

HCOOD

+ HD

( 5 ) : (1)

All reactions were performed under as nearly as possible the same conditions of pressure, temperature and reaction time. Under these conditions exchange according to reaction (/?)could be completely neglected. Under the conditions of reaction ( y ) only H, enrichment in the gas phase was observed, whereas the formation of HD was negligible. From these first results Hirota et al. conclude that i f HD formation were observed in reaction ( a ) this would certainly not take place via exchange reaction (p) or via the reaction

...........

(8)

+ HiH +. HCOOH ..............

HCOOtD

+ HD,

which is an analog of (y)*. After H D formation by exchange reactions had been excluded in this way, reaction ( a ) was studied. It was found that after completion of the reaction the number of D atoms in the gas phase was greater than the number of H atoms. Furthermore, the amount of HD was much larger than would be expected from the equilibration of reaction (/?). From these two observations it follows: (a) that there remains an intermediate on the surface in which more H atoms than D atoms are present, for instance a HCOO- ion; (b) that in reaction ( a )HD (which is formed besides H, and D,) is a primary product, However, this last conclusion is not fully established; it is also possible [Hirota et al. (4611 that HD is formed by exchange between HCOOD

* Note: reaction ( y ) was studied instead of (S), because in (6)HD can be split off from the formic acid molecule, but can also be formed by exchange.

60

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

in the gas phase (or physisorbed HCOOD) and, for instance, adsorbed formate ions in this way: HCOOD,,,

+ HCOO,,

HCOO-

3

+ DCOO-

HCOOH,,, --f

2 CO,

+ DCOO,,

+ HD

Working with gold, Sachtler and Fahrenfort (23) studied the distributions of the hydrogen isotopes in the decomposition products of HCOOD in mixtures with HCOOH. At 150°C and at different HCOOH/ HCOOD ratios, the three molecular hydrogen isotopes are formed in their equilibrium proportions. As under the same conditions HCOOH did not exchange with D,, and no H, - D, exchange occurred, the authors concluded that the equilibrium mixture is formed as a primary product. This means that the equilibrium must be reached while the hydrogen is still in the atomic form on the metal surface, which points to an independent fission of the H atoms from the molecule. Mechanistic conclusions from the work on gold and silver are facilitated by the low exchange activity of these metals. Therefore, Hirota’s (45) research with nickel powders and films-showing a high exchange activity-do not provide very important data in this respect; consequently only brief mention of these results will be made here. DCOOH was decomposed at 100°Con Ni powder with about the same number of surface Ni sites as there were D-formic acid molecules in the gas phase. Less deuterium than hydrogen was formed in the gas phase, which points to the formation of an intermediate with excess D on the surface. On fresh powders the CO,/H, D, HD ratio was greater than unity, but this was not so after one decomposition run had been carried out. Obviously only fresh nickel absorbed hydrogen. On Ni-films, as on powders, DCOOH decomposition resulted in an excess of hydrogen in the gas phase and HCOOD decomposition in an excess of D,, which again points to the formation of an intermediate in which the acid hydrogen atom is split off. The exchange reaction HCOOD + H, -+ HCOOH + HD did not initially occur at room temperature on a fresh nickel-film, and in a second run, in which the intermediates formed in the first run were still on the surface, only 8% conversion occurred after 90 hr, while there was no formic acid decomposition at all. It has to be remarked that in the exchange work described above the formation of exchanged formic acid molecules (e.g., HCOOH from HCOOD and H,) was not shown unambiguously. Thus, although this is assumed by a number of investigators, it is not certain if the formic

+

+

T H E CATALYTIC DECOMPOSITION O F FORMIC ACID

61

acid molecule has a real chance of following a “way back” when the chemisorbed intermediate is formed from it. In our opinion almost all the experimental evidence can be interpreted without taking this “way back” into account. 6. The Change in the Electricat Resistance of Nickel Films during Formic

Acid Decomposition Rienacker and Hansen (3a),and also Suhrmann and Wedler ( 4 7 , 4 8 ) , have tried to gather information on the mechanism of the decomposition of HCOOH on evaporated nickel films by measuring the change in electrical resistance during reaction, and also the changes occurring upon adsorption of formic acid and of possible reaction products. Hydrogen adsorption resulted in a resistance drop of about 3% at 20°C and at 200°C. This points to film contamination by oxygen, as was shown by Sachtler (49);when the surface is reduced owing to hydrogen addition Ni atoms are added at the contact points between the Ni crystallites of the film, and hence the resistance drops. The initial resistance increase of about 3%, found upon addition of formic acid at room temperature and low pressures, may be interpreted in the same way. Adsorption of the reaction products CO,, H,O, and CO at 200°C had no appreciable influence on film resistance. At 2OoC there is some influence in a positive or negative direction, but, as was the case with H, adsorption, the fact that the films were contaminated makes it difficult to draw conclusions. During static formic acid decomposition at 200°C there is an initial increase of the resistance of only 1.5%, obviously owing to a lower coverage of the reaction intermediates at this temperature, but after a few minutes the resistance decreases slowly again, owing to the decomposition of the intermediates. In our opinion it is impossible to draw conclusions from such measurements as to the direction in which electrical charge is transported upon adsorption of the intermediate. According to the present insight into the physics of resistance changes in metal films, chemisorption on clean films will always result in an increase of resistance, irrespective of the direction in which electrical charge is transported [Sachtler and Dorgelo (SO)]. 7 . Infrared Studies of the Chemisorption and Decomposition of Formic

Acid on Metals In the foregoing sections we learned that from classical kinetic studies and from the determination of isotope effects some indications can be obtained about the mechanism of the reaction and the structure of the

62

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

adsorbed reaction intermediate. It was made obvious, for instance, that on Ni the adsorbed complex preponderates at the surface, whereas under most experimental conditions the adsorption of decomposition products is low and has hardly any influence on kinetics and mechanism (see Sections 111,A, 1 and 111,A, 2). Furthermore, kinetic isotope effects in the decomposition on silver [Block and Kral ( Z I ) ] seem to make it likely that the activated complex in the rate determining reaction step contains two H atoms. Concerning the structure of the adsorbed reaction intermediate on Ni, Tamaru (15) made the important observation that two formic acid molecules are taken up per surface Ni site, and that the intermediate decomposes into 2 CO, and 1 H, molecule. However, about the exact structure of the intermediate only guesses could be made; Schwab, from his work with alloys (see Section 111,B), and also Dowden and Reynolds (51),concluded that it was a positively loaded particle, whereas RienLcker (3) came to suppose that it was a (negative) formate ion, on account of his film resistance work published earlier [RienLcker and Hansen (3a)l. I n this respect it is interesting that from their study of the decomposition of bulk formates Hofmann and Schibsted (52, 53) already in 1918 arrived a t the view that also in the catalytic decomposition of HCOOH, the formation of formates may play a role. The application of infrared spectroscopy to the elucidation of this disputed question has clarified the situation considerably. I n the study of the adsorbed state this technique was first used by Yaroslavskii and Terenin ( 5 4 ) , while Eischens improved the method and demonstrated its wide applicability to the study of chemisorption and catalysis [Eischens and Pliskin ( 5 5 ) ] . The first infrared evidence of the possible occurrence of formate ions as reaction intermediates on the surface of metals resulted from the research of Hirota and his colleagues (46, 56), who showed formate ions to be present on powders of silver, copper, nickel, palladium, rhodium, platinum, and zinc, after adsorption of formic acid at room temperature. Moreover, in the far-infrared region they observed bands at 410 and 130 cm-1, which is an indication of bonding between metal atoms and formate ions via oxygen atoms [Hirota and Nakai (57)l. A more detailed study was made by the same Japanese workers of the static decomposition reaction with Ni as a catalyst [Hirota et al. ( 4 5 ) ] . An infrared cell, constructed according to the principles indicated by Eischens (55),was horizontally mounted in a Hilger H 800 spectrometer. To avoid interference by adsorbed groups occurring on the carrier material, finely divided Ni particles were deposited on a thin rock-salt plate by decomposition of Ni formate suspended in ethanol.

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

63

The spectra scanned between 1250 and 2000 cm-1 are represented in Fig. 10. Adsorption of a slight amount of HCOOH resulted in the appearance of an intensive absorption band a t 1580 cm-1 and a less intensive one a t 1360 cm-1. From a comparison of these frequencies

40. I

- 1

201

.

i .j ;

I

I

t

I

I

1

& I b.

1250

1500

1750

2000

ern-'

FIG.10. Infrared spectre, scanned during the static decomposition of formic acid on Ni at lOO"C, according to Hirota (45)i (1)Spectrum of adsorbed-state HCOOH a t 100°C end 10-* mm Hg. (Z), (3), end (4)are the spectra taken successively at two-minute intervals after gaseous formic acid of about 10 mm Hg had been added to the adsorbedstate HCOOH at 100°C. After 30 min a spectrum similar to (1) could be obtained.

with those found in the spectrum of powdered crystalline nickel formate, the investigators concluded that these bands had to be ascribed t o the symmetric and antisymmetric vibrations of the resonance hybrid of the OCO group in a formate ion:

After pumping t o p = 10-2 mm Hg a t lOO"C, spectrum 1in Pig. 10 was measured. The band intensities were slightly lower than in the spectrum taken a t room temperature (not given in the figure), but there was no shift in the frequencies. When formic acid vapor is introduced into the cell a t 100°C and a t pressures up to 10 mm Hg, the following changes in the spectrum occur: I n the first place, a strong band at 1740 cm-l with a small shoulder

64

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINO

at -1800 cm-1 appears, which has to be ascribed to the formic acid (dimer) vapor. Secondly, the formate-ion bands already present remain in the spectrum, but the 1580 cm-1 band becomes broader. Thirdly, a new band a t 1690 cm-1 appears, which does not occur in the gas phase spectrum of formic acid, and hence points to the existence of a second adsorbed phase by the side of the formate ions. The origin of this band is not completely clear. I n the study of infrared spectra of the physisorbed molecules [Eischens and Pliskin (55)] the frequencies in the adsorbed state are always found to shift to a lower value. Hence, the 1690 cm-l band, being closely situated to the 1740 cm-l gas phase band, may be ascribable to physically adsorbed HCOOH, or to another undissociated adsorbed state. The change of the spectrum with time is represented by Nos. 2 , 3 , and 4 of Fig. 10. Owing to decomposition of the formic acid the gas phase bands disappear slowly, whereas the bands of the adsorbed states remain practically constant, which indicates a constant coverage of the reaction intermediates. A mass-spectrometric analysis of the gas phase in the cell gave a production of 1.0 part of hydrogen to 1.8 part of CO,. This ratio is in fair accordance with the value found by Tamaru (14,and is compatible with the picture of formate ions as decomposing intermediates (if H, formed in the adsorption of HCOOH does not remain on the surface). From the work of Hirota and his colleagues we may conclude that the presence of formate ions as intermediates in the decomposition on Ni is a t least highly probable. The broadening of the 1580 om-l band during decomposition and the origin of the small 1690 cm-1 band, however, are not fully understood. Even stronger arguments for the decisive role of a two-dimensional layer of formate ions on Ni during the decomposition of HCOOH are furnished by the research work of Fahrenfort, van Reijen, Sachtler, and Hazebroek (12, 25, 58, 59), which was published shortly after Hirota et al. arrived at their first results. The infrared spectrum obtained by the Dutch workers for formic acid adsorbed on a Ni-on-silica catalyst generally confirms the results of Hirota et al. (see Fig. 11, spectrum A). After the disappearance of bands owing t o evacuation a t 55OC of gaseous and physically adsorbed formic acid, two pronounced bands remained in the spectrum, a t 1575 cm-l and 1360 cm-l (spectrum B), which indicated the presence of formate ions. Also some bands of lower intensity occurred at 2080,2200, and 3320 cm-l, presumably owing to the adsorption of small quantities of CO, CO,, and H,O on the surface of the carrier or of the metal. Analogous results were arrived a t by Clarke and Pullin (60).

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

65

That the formate ions observed on the surface indeed form the reaction intermediate-the decomposition of which is rate-determining -and not an adsorbed by-product, was beautifully demonstrated by the "spectrokinetic method," as it was called by Fahrenfort et al. The spectrometer was set a t the frequency of the strong formate band at 1575 cm-l and the change in optical density was measured as a function of time. The time of formation of the band upon adsorption wavrlcngth, p 3

25

100-

*O-@'

60-

40-

' .

4

1''-

' . . I

i

0 1 ' '

..'I

Cn

*.-.

8

6

-*-" \

.V' .ebp'"'

\

23.

5

7

' I

"

.I'

I

..do

P

' '

I

' . 1B

.

o$o~-~eo.o *.o

;P

\'

0.

.,*'

9 ;,

re**

t

r

ti$

HCOOH '

;

t co

t HCOOH

f n20

t HCOOH .

,

of HCOOH was equal to or shorter than the response time of the spec< 5 kcal/mole. The trometer (10 sec), from which it followed that EallR kinetics of the disappearance of the band gave an energy of activation of 23 kcal/mole a t the higher coverages; this was in beautiful accordance with the energy of activation of the decomposition of HCOOH on the same sample a t full coverage; which was also 23 kcal/mole. It was thus made likely that the slow step in the over-all reaction is obviously the decomposition of the formate ions. The frequency factor of the reaction was somewhat lower than the value t o be expected from Eyring's theory, viz., 101O (instead of 1013;for this problem see Section 111,A, 3, and the general discussion).

66

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINO

The superficial formate decomposed a t lower temperatures than does bulk nickel formate. The authors believe that the reason is probably to be found in the much more complicated and less favorable situation for the growth of the metal phase in the case of pyrolysis. We shall have to return to this problem in the final discussion. It seemed of interest to apply the same powerful research tool to the study of the decomposition on a noble metal, and to find out if here, too, the same intermediate and the same type of kinetics play a role. For this purpose a gold catalyst with a mean crystallite size of ca. 100 A was prepared [Fahrenfort et al. ( I Z ) ] from a solution of AuC1,. The infrared spectrum of formic acid adsorbed on this sample a t 25°C) showed, in addition to the formic acid bands, only one weak absorption band in the 1530-1550 cm-1 region. Hence, a t first sight, it seems questionable if also on gold formate ions are present. It must be borne in mind, however, that the surface area of the sample was small and most probably the degree of coverage low; in this case it would be understandable that only the strongest antisymmetric 0-C-0 stretching band became visible. Another argument in favour of the view that also on gold a formate mechanism may be operative can be found in the isotope effect, mentioned already in Section 111, A, 5. Unfortunately, a spectro-kinetic study of the decomposition on gold was impossible, owing to the weakness of the formate band at 15301550 cm-l. However, the kinetics of the reaction are certainly not in conflict with the acceptance of a formate mechanism (see Section 111, A, 1). It is a pity that the beautifully corresponding results of the Japanese and Dutch research groups are not fully confirmed by the work of Eischens and Pliskin (61).Like Eischens (62))we believe, however, that the deviations are largely to be ascribed to differences in experimental conditions. Eischens and Pliskin’s results can be summarized as follows: Formic acid adsorbed at -60°C on a 9% Ni-on-silica catalyst produced spectrum A of Fig. 12. The formic acid was completely adsorbed up to a coverage of ,- 0.1-0.2. A strong band a t 1359 cm-l appeared (together with a smaller band a t 2941 cm-1 in the C-H-stretching region), which was attributed to “covalently chemisorbed formic acid”. Hence, the exclusive appearance of formate bands (reported at high temperatures by the other research workers) was not found a t this low temperature; only a weak band a t 1587 cm-l was observed, which may point to the presence of some formate on the surface. After heating up to 25°C) spectrum B was found; the band due to “covalently chemisorbed acid” disappears and, beaides the small band

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

67

a t 1587 cm-1, two new bands are found a t 2020 and 1820 cm-l, which are due to “linear” and “bridged” chemisorbed CO on Ni [Eischens et al. (63)].Finally, after 3 hr a t 25”C, a further increase of the amount of chemisorbed CO and a further decrease of the amount of the covalently chemisorbed acid occurs. frcqucnce

3000

2500

4 0

2000 1800

crn-1

1600

1400

70

60

50

w a v o l e n g t h in m i c r o n s

FIG. 12. Infrared spectra of formic acid adsorbed on a 9% Ni-on-silica catalyst. according to Eischens and Pliskin (61). Spectrum (A): at - 60°C; Spectrum (B): after heating up to 26°C; Spectrum (C): after 3 hr at 25°C.

Y

QO: --TP

;I 8 0

._ 7 0 E Y

YI

c

0

L

CI

n r

I

I I

60-

I I

I

1

I

F I ~13. . -6O”C, according to Eischens and Pliskin (61).

According to Eischens and Pliskin, “these facts show that formic acid chemisorbs at - 60°C as a species which decomposes a t a moderate rate a t room temperature, with formation of chemisorbed CO.” We

68

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINQ

should like to suggest that perhaps this decomposition proceeds via the formation of formate ions as a low-coverage-intermediate, since in the three spectra of Fig. 12 the 1587 cm-l band is always present. What happens upon HCOOH chemisorption at -6O"C, after the Ni sample has been oxidized in 6 mm 0, at 35OoC, is shown in Fig. 13. frequency ( c m - 1 I

-

r

80-

I

I

A

70-

6 0-

c

50-

C

0

c

.c- 40E ul

C U L

'3d + 30-

1

I

1

6.4

6.6

6.8

70

1

7.2

w a v e le ng th

FIQ. 14. The infrared spectra of butyric acid on (A) Ni and (B) NiO 8t room temperature, according to Eischens (private communication).

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

69

Now a strong band is seen at 1587 em-' and a weaker one at 1351 em-*, which can be attributed to the antisymmetric and symmetric carbonoxygen stretching vibrations. Hence, the adsorption of HCOOH on the nickel oxide must lead immediately and exclusively to the production of formate ions on the surface at - 60°C. The same difference in behavior between Ni and NiO was observed in the adsorption of butyric acid and acetic acid. The result for butyric acid at room temperature is given in Fig. 14.* Here, again, a covalently bound acid is observed on the metal, whereas on the oxide obviously butyrate ions are present. 8 . The Adsorption of Formic Acid on Nickel at Higher Temperatures

When excess formic acid was brought over a Ni-sample a t -5°C the bands of the covalently bound acid were not observed (see Fig. 15, spectrum A), but strong bands were found a t 1587 and 1351 em-', which indicated the presence of formate ions on the surface. The strong band at 1724 em-' was due to physically adsorbed acid and is also found on silica samples without Ni. Furthermore, narrow bands were present at 1639 cm-l, 2041 cni-' and 1887 em-'; these were due to adsorbed wat.er and chemisorbed carbon monoxide. Spectra B and C were taken after waiting times of 22 and 68 hr, respectively. I n these periods the physically adsorbed formic acid was used up. From what has been said above it appears that Eischens and Plisltin's results a t -5°C are in accordance with those of Fahrenfort et al. and Hirota et al. I n Fig. 15, spectrum B, including even the small band a t 1690 cml, is present. This band is identical with that present in the Hirota's spectrum (see Fig. l o ) , the significance of which was still obscure. However, the difficulty lies in the interpretation ofthe result. Eischens and Pliskin think that in the -5°C experiment formate ions were formed, instead of covalently chemisorbed acid, because of the fact that in this experiment a n excess of formic acid was supplied to the catalyst. The consequence is, they say, that first the Ni-surface is oxidized by HCOOH, and that thereafter formic acid is adsorbed dissociatively on the NiO (as formate ions and protons), as was also the case in the - 60°C experiment on NiO. The acceptance of oxidation of the Ni-surface by forniic acid is criticized by Sachtler (64) and by Hirota (G5) on thermodynamic grounds; they argue that formic acid is a suitable reducing agent for NiO. If their view is correct, the results of Eischens and Pliskin can

* This figure was kindly placed at our disposal by Dr. Eischens. The result was communicated at the 138th meeting of the American Chemical Society in 1960.

70

P. MARS, J. J . F. SCHOLTEN, AND P. ZWIETERING

perhaps be understood if i t is assumed t h a t the temperature a t which formic acid is adsorbed determines the mode of binding; a t - 60°C the acid is covalently bound on Ni and at - 5°C it is bound dissociatively. It is also possible t h a t the formic acid adsorption was studied on a Ni-H layer instead of on free Ni sites. I n their preparation technique, Eischens et al. (63, 66) cool down their sample t o room temperature

5.0

60

70

w a v e l e n g t h in m i c r o n s

FIG.15. Infrared spectra of formic acid adsorbed at -5°C on silica-supported-Ni, according to Eischens and Pliskin (GI). Spect)rum (A): after 1 hr at -5°C; Spectrum (B) after 22 hr at - 5°C. Spectrum (C) aftcr a n addition 68 hr at 25°C.

in a stream of hydrogen after reduction, and next evaporate the H, from the cell. I n this case a monolayer of hydrogen is left on the surface. It is quite possible t h a t at -60°C this hydrogen is not removed from the surface b y the competing adsorption of formic acid, and that a t higher temperatures the formic acid can drive away the hydrogen from the surface. If this is so, the covalent chemisorption of formic acid is a form of adsorption on a Ni-H layer. It will be clear t h a t further experimental corroboration of these views is necessary and highly desirable. I n a static decomposition experiment a t a much higher temperature -viz. 200°C-and at a constant formic acid pressure of 4 mm. Hg, Eischens and Pliskin observed strong bands at 2020 cm-I and 1820 crn-l, which were due t o linear and bridged CO chemisorption. There were only

71

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

barely detectable bands at 1587 cm-l an d 1351 cm-l, so that, if formate ions were present on the surface, the concentration had to be very low. After evacuation at 200"C* and cooling to - 6 0 " C , the covalently chemisorbed formic acid bands were again observed upon adsorption of HCOOH. Hirota (65) rightly remarked th at in this experiment a low steadystate coverage of formate ions as intermediates is understandable in view of the relatively high reaction temperature. From the fact th at a t -60°C the covalently chemisorbed species is detected again, Eischens concludes th at a t higher temperatures formic acid acts as a reducing agent with respect t o Ni. 9. Infrared Xpectra of Deuterated Formic

Acid

Infrared spectra of deuterated formic acid chemisorbed on Ni were studied at -6O"C, in order to elucidate the structure of covalently crn-1

frequency

2500

3000

50

1

"

40

2000

1

53

1600

1800

1

1

,

I

I

I

,

1

8

1

70

GO wavelength

1400

in

microns

FIG.16. Infrared spectra of DCOOD chcmisorbed on nickel at - 60°C, according t o Eischens and Pliskin (GI).A-spectrum a t - 60°C on Ni; B-after addition of 0.02 em3 of 0, at -60°C; C-after addition of 0.5 em3 of 0, at - 60°C.

chemisorbed acid a t this temperature. The spectrum of DCOOD on Ni is given in Fig. 1 6 ~ This . spectrum differs from th a t of HCOOH (compare Fig. 1 2 A ) in two respects:

* After evacuation at 200"C, hydrogen and CO may still be chemisorbed on the surface.

72

P. MARS, J. J. F. SCHOLTEN, A N D P. Z W I E T E R I N G

-

(a) The C-H stretching band a t 2941 em-1 is now found a t 2188

cm-l as a C-D stretching band (shift: 42); (b) The bands occurring a t 1359 cm-l and 1587 cm-l in Fig. 1 2 are ~ now found a t 1325 cm-l and 1546 cm-l. Moreover i t appeared t h a t HCOOD gave the same spectrum as HCOOH, and DCOOH the same as DCOOD. From this last fact, and from the frequency and intensity shifts occurring in the 1359 and 1587 cni-l hands as a result of deuteration, Eischens and Pliskin derive the following structure for the covalently chemisorbed acid: D

-Ni.Ni--Ni.Ni-

Upon adsorption of 0, a t - 60°C (spectra B and C of Fig. 10) adsorption bands of deuterated formate ions are found. For the formation of a n amount of formate ions equivalent t o 0.05 ml of formic acid, about 0.02-0.05 ml of oxygen has t o be added.

B. DEHYDROGENATION O N METAL ALLOYS A very large amount of work has been done on the decomposition of formic acid on metal alloys, in particular by Rienhcker and his school, a n d b y Schwab and his collaborators. We shall not here give a full account of all this work, but sclcct some typical results which demonstrate the general trend of the influence of alloying. To obtain a good understanding of this matter, i t is of rather great importance t o get a knowledge of the studies made of the influence of the mode of alloy preparation, for which reason this will be dealt with first. The system Aq-Cu was studied hy Schwab and Schw ah-Agallidis ( 2 0 ) , and by Rienarker et aZ.* The activation energies observed in the zeroorder dehytlrogenation reaction arc reproduced in Fig. 17 ; the observations of both laboratories are in perfect accordance, except for the alloy with the eutectic composition. According t o Schwab, the reason may be sought in the fact t h a t the surface of his alloy (highest E') had a n appearance like copper, whereas Rienacker's sample had a surface which was more like silver a t the eutectic composition, and this difference could be reduced t o a difference in pretreatment, Schuab's sample had

* The publication b y S r h w ab and Schxab A g a h l e i (LO) nicntions obscrx ations of Rienacker, pcrsonally corriinuiiicatcd to these authors

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

73

been oxidized and reduced and then preserved under formic acid, whereas Rienacker’s sample had been reduced only. The same phenomenon was found by Rieniicker and Bade ( 6 7 ) in the case of Cu-Ni

I

E u t e c t ic composition F I G . 17. Dehydrogenation of formic acid on Cu-Ag alloys, according to Schwab and Schwab-Agallidis(ZO), and according to Ricnucker. 0-Ricnacker, static decomposition; 0-Schuab, static decomposition; A-Schuab, dynamic decomposition,

alloys. Up t o 80% Cu, the alloy has the white nickel color and E is only slightly higher than for Ni. Above 80% Ni there is a sudden jump in the activation energy and a t the same time the surface assumes the red copper color. Another striking example of the decisive role of the mode of preparation may be found in the behavior of Cu-Yd alloys [Rienacker et al. (Gbi)] (see Fig. 18). When copper is alloyed with Imlladium the actibation energy remains practically constant up to liSO/, P d ; on further addition of Pd, E falls rapidly down to the low value of Pd. This is only true, however, for normally prepared, disordered alloys. If, by tempering, the so-called ordered alloys are prepared, in which the Pd-atoms have definite crystallographic positions, the result is essentially different (see Fig. 18);now the activation energy of the alloys is nearly additively composed of the contributions of the components in the alloy. Analogous results were obtained for the system Cu-Pt [Rienacker (6571. From the foregoing w c may c.onclutle th at the activation energy on a n alloy catalyst is a h a) s mnieu here beta een th a t of the components (this is a general experience) [Eieiiaclier et ul. ( 6 8 ) ]I. n general, E deviates

74

P. MARS, J . J . F. SCHOLTEN, AND P. ZWIETERING

from thevalne found by linear interpolation and the most simple explanation a t hand is t h a t there are possibly strong deviations of the surface composition from thc bulk composition. Much research has been performed, especially by Schuab and his school, on the HCOOH decomposition on €iume-Kothery ulloys [Schwab (4, 70)l. As from this work conclusions have heen drawn also with regard t o the mechanism of the formic acid decomposition, i t will here be given special attention.

(Kcallmole) E 301

cu (100%1

5 0 % Cu

5 0 % Pd

Pd (100%)

FIG. 18. Dehydrogcnation of formic acid on Cu-l’d alloys, according t o Rienacker, IVcssing, and Trautrnarni (68)(Cu and I’d forin a continuous scrics of solid solutions). 0-disordcred alloys; x -ordered alloys.

I n the Hume-Rothery alloys [see Mott and Jones ( 7 1 ) ;for instance Cu with Zn, Al, or Sn as solute element] the electron concentration may be increased by raising the concentration of the higher valence atoms, u ithout bringing a radical changc in tlic crystallographic character of the alloy, up to a certain limit only. This liniit is given by that electron concentration a t 15 hich the f a s t a t electrons begin to be reflected hy the most intensely reflecting lattice 1jlnncs, according t o l3ragg’s la\% (the first Brillonin zone is then filled). If the limiting elcctron concentration is exceeded, the phase becomes unstable in favor of a new one, in which the electronic concentration can agnin increase u p to a certain limit. The increase in electron concentration (e.c.)is given by:

d (e.c.) = x (n - 1)

(25

75

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

where

x

=

molal fraction of solute element,

n = valency of the solute atom.

Schwab ( 4 , 7 0 ) studied the decomposition of formic acid on a large number of a-phase Hume-Rothery alloys consisting of Ag and elements of periods VB and VIB, and having different electron concentrations. He tried t o establish a relationship between E , the activation energy, assumed t o be the characteristic catalytic parameter, and the electron concentration, e.c. The experimental results are given in Table IV. TABLE IV

The Injuence of Alloying on the Activation Energy, within the a-Phase of Different Hume-Rothery Alloys; According to Schwab ( 4 , 7 0 ) ( T h e Letters in BTackets Correspond to the Points in Fig. 19) Alloys of Ag with elements of period VB

Alloy

Ag Ag-Cd Ag-Cd Ag-In Ag-In Ag-In Ag-Sn Ag-Sn Ag-Sb Ag-Sb

At. yo solute element 0 15.8 38.4 5.0 7.9 14.6 3.6 7.5 5.5 9.7

Alloys of Ag with elements of period VIB

E

(E)

Alloy

At. Yo solute

E

mole 17.6 20.0 19.0 17.8 18.5 20.8 20.0 22.6 31.0 33.0

Ag Ag-Hg Ag-T1 Ag-TI Ag-TI

0 2.6 1.12 3.60 4.50

17.6 21.0 20.0 20.0 20.0

Ag-Pb Ag-Pb Ag-Pb Ag-Pb Ag-Bi Ag-Bi Ag-Ri Ag-Ui

0.69 1.07 1.12 1.43 0.86 1.62 2.00 2.12

35.0 30.0 35.0 37.0 31.0 29.0 35.0 34.0

An about linear relation appears t o cxist between AE, the increase in activation energy with respect t o pure silver, arid (Ae.c.)Z/x, the square of the increase of electronic concentration per solute atom. Instead of plotting AE as a function of (Ae.c.)2/x,4Asj.x was plotted against ( n - 1) (see Fig. 19). To bring the points for the VR alloys on a common line with those for the VIB alloys, d A E / n was arbitrarily multiplied b y respectively a = 10 and 01 = 10/3.

76

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

Schwab ( 4 , 70) concludes th at an increase of electron concentration has a retarding effect on the reaction (increase of E ) , and this means th a t in the rate-determining step electrons from the formic acid molecule must intrude into the metal. This entry requires more energy the more the metal is already saturated with electrons. The experiments performed by Schwab and his school ( 7 2 ) with liquid alloys of the bivalent mercury, with, successively, Pb("'), Cd'") and Ag") as additional elements, are in contradiction with this theory. Addition of Pb, Cd, or Ag always resulted in a decrease of the activation energy, irrespective

1

2

3

4

-

5

6

(n-1)

FIG. 19. The relation between a function of the activation cnrrgy, and a function of the electronic conccritration, for different Hurne-Rothcry alloys, according t o Scliwab ( 4 , 70). The explanation of the symbols is given in the tcwt. The lotters in brackets correspond t o those used in Table IV.

of whether the addition effected an increase or a lowering of the electron concentration. Schwab also studied the catalytic 1)ehavior of intermetallic Hume-Kothery phases and mixtures of such phases. I n the

77

T H E CATALYTIC DECOMPOSITION O F FORMIC ACID

system Ag-Sb [Schwab (28);Fig. 201, the activation energy rises abruptly with increasing Sb content in the a-phase. I n the €-phase,however, there is first a decrease and afterwards a smaller increase of E . I n the heterogeneous mixtures of the <-phase an d Sb crystals it reaches a maximum at the composition of the eutectic mixture. It is seen th a t in this system activation energies are found th at are higher than th a t of the component with the highest E. I n this conjuction we should draw attention t o the fact t ha t “the activation energy” for a certain element or alloy 401 E

I I I I

I

I I \ /

I

I

(Kcallmole)

30-

20 -

Ag

(loo./.)

20

40

60

-

a0

, 100

%Sb atoms

FIG.20. The activation cnergy for the decomposition of HCOOH on intermetallic phases of Ag and Sb, according to Schwab ( 2 8 ) .

is not a well-defined quantity; in Section 111,A, 3 different examples were reported of the great influence on E of the pretreatment and of the exposition of crystallographic planes. Of course this last factor in particular may have had a n influence on the measured activation energy of intermetallic phases, where changes in lattice constant and lattice type (Ag : face-centered cubic, but Sb rhombohedral!) can occur. Also the system Cu-Sn [Schwab (28);Fig. 211 shows a rise in activation energy within the a-phase, although it is here much less pronounced. The other intermetallic phases of this system have such narrow fields of existence, t ha t they could be prepared in a crystallographically pure state only with difficulty; changes within these regions could not therefore be studied. The deformed body-centered cubic lattice of the a-phase corresponds t o a high maximum of E , the hexagonal close-

78

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINU

packed E lattice to a much lower value, and the hexagonal 7 lattice to a still lower one. I n this system the changes in hardness, activation energy and electrical resistance show the same trend (see Fig. 21). A

B

C

300

1

I

I

I

€1

\

71

I

\

100

25 1

15

20

-

25

e.c.

FIQ. 21. The activation energy for the decomposition of HCOOH on intcrmetallic phases of the system Cu-Sn, according t o Schwab ( 2 8 ) . A -hardness-kg mm-2; B- '0'- activation energy-kcal mole-'; C - - - - - resistance-Q cm lo4.

Alloys of sp and d metals were studied by different authors. The system Ni-Cu is interesting, as it was studied by Schwab and SchwabAgallidis (ZO),Rienacker and Bade (67) by Dowden and Reynolds (51) and by Fuderer-Lueti6 and Brihta ( 7 3 ) .The results are plotted in Fig. 22, and form a clear demonstration of the inconstancy of the results and hence of the relative significance of the conclusions drawn from them. According to Dowden and Reynolds (51),alloying of nickel with copper results in a filling-up of the 3 d band holes in nickel, and as a consequence electron transfer from the substratum to the metal is obstructed and the decomposition rate decreases (see Fig. 22). Many other systems composed of an s-, a p - , and a d-metal have been

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

79

studied, e.g., Pd-Ag [Rieniicker ( 7 4 ) ] ,Pd-Au [Eley and Luetid ( 7 5 ) ]and Pt-Cu [Rienacker and Hildebrandt (7671. I n all these cases the activation energy increases when Pd or Pt is alloyed with the s-metals, which would appear to support the d-band theory of Dowdcn and Reynolds. To sum up, it may be said that in a number of cases different authors have found that filling-up of the electronic levels of the metal catalyst results in an increase of the activation energy of the decomposition reaction; from this they concluded that: (a) The adsorbed intermediate could be a positively charged species; this is contrary to the negative formate ion accepted by many others on the basis of a vast amount of evidence described in the earlier chapters;

N, (100 * I . )

cu ( 100 * I .

1

+

FIG. 22. Formic acid decomposition of formic acid on Ni-Cu-alloys. Dowdon and Reynolds ( 5 1 ) ,on foils; x 0Rieniicker and Bade ( G 7 ) , on foils; Schwab and SchwabAgallidis ( Z O ) , on foils; @I Fuderer-Luetid and Brihta ( 7 3 ) , alumina-supportccl alloys.

(b) The “electronic factor,” that is the electronic structure of the

whole metal, is of deciding importance for the catalytic properties of the metal; this is contrary to the views of others who claim that only the chemical properties of the single metal atoms taking part in the formation of a surface formate determine the catalytic properties.

80

P. MARS, J. J. F. SCIIOLTEN, AND P. ZWIETERINO

A short discussion of this last point, although somewhat outside the scope of this review, is given in the general discussion, but about the first point, which is of direct importance in connection with the mechanism of the decomposition reaction, the following remarks may be given here : The correlation between the change of the activation energy and the increase in electron concentration, as shown in Fig. 19, is not very convincing. It is to be regretted that, as far as we know, no information is available about the change in the electronic work function, 4, of the alloys involved. If values of 4 had been determined-preferentially of the same specimens as used for the catalytic investigation-a direct correlation between A + and AE might have been expected, which could have enforced the argument considerably. However, if it is accepted for the time being that a rise in electron concentration decreases the work function, and increases the energy needed to put an electron into the metal, the fact that this is accompanied by a rise of the activation energy of the decomposition reaction is certainly not in contradiction with the concept of negatively charged formate ions forming the intermediate at the surface. From a consideration of the experimental procedures used it is most plausible that all investigations described in this section were performed in the zeroorder region, in which the decomposition of the adsorbed complex is rate-determining. I n this step obviously the metal phase has to take back the electron originally given to the adsorbate. Finally, we should like to repeat again that it is most dangerous to accept a priori that the bulk composition of the alloy is equal to that of the surface layer. I n many cases large differences may be expecteciaccording to Gibb’s law-especially when small percentages of additional foreign atoms are present. I n all cases investigated the activation energy of the reaction was lower on the original metal than on a pure specimen of the metal added; for bismuth, for instance, values of 18-25.5 kcal/ mole [Rienacker (33)]and for Sb 27 kcal/mole were reported [Schwab ( 2 8 ) ] ,which values are both higher than the value of 17.6 kcal/niole found for pure Ag in Table 111. Also the activation energies for Ag, Au, and Cu, which metals were added to Pd, are considerably higher than those for pure palladium. In view of the possible difference in surface composition, a simple additive effect cannot be excluded with certainty; such an effect may account for the important influence of the preparation method found in the experiments. All in all, we do not believe that the experiments described in this section invalidate the concept of a formate ion being the adsorbed intermediate in the dehydrogenation of formic acid on metals.

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

81

C. DEHYDROGENATION ON OXIDES 1. The Kinetics of the Reaction The kinetics of the dehydrogenation of formic acid on oxides are not extensively investigated; only about the reactions on MgO and on Fe,O, are some data available. For the reaction on MgO Cremer and Kullich (77) found that the reaction products do not retard the reaction.* I n our laboratory it was found that on MgO formed by decomposition of the oxalate, both at 372OC and 425OC, and between 0.2 and 0.8 atm pressure, the reaction is almost first-order. On a different MgO-sample (containing about 70% of oxalate) at 327°C first-order kinetics were observed. On Fe,O, the reaction is zero-order at 3 1 5 T between 0.12 and 0.33 atm [Mars (7)l. For the dehydrogenation of alcohols it was found by Balandin et al. (78)that on many oxides the rate can be formulated as T"

+

balc *

*

Palo

Palc

+

bald

*

paid

'

[where b and p are adsorption coefficients and partial pressures, respectively, of alcohol (alc) and aldehyde (ald)] which shows that in this case the adsorption of one of the reaction products influences the rate of the reaction. 2. Activities for the Dehydrogenation of the Various Oxides

As mentioned already in the introduction, the dehydrogenation of formic acid proceeds particularly well on MgO and ZnO. It is, however, always accompanied by a dehydration reaction. The following scheme gives the order of selectivity: MgO

HCOOH<

-

ZnO /Ha Fed4 Cr,O, TiO, SiO, 40,

+ CO,

H,O

+ CO

Eucken (79, 80), studying alcohol decomposition, was the first to make an attempt at relating the properties of the oxide to the selectivity. He succeeded in establishing a correlation between the selectivity (i.e., the ratio of the rate of dehydrogenation to that of total decomposition) and the quantity r ,"

"=v.z,'

* Erroneously they stated that the reaction is zero-order.

82

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

where re = radius of cation, V = molar volume per cation, 2, = valency of cation. With minor approximations this relation can be replaced by

with ra = radius of anion. Thus, the dehydrogenation appears to occur in particular if the radius of the cation is large and the valency small. Simple electrostatic considerations reveal that this is essentially the situation where the basicity of the hydroxides is large. Indeed, the more basic oxides show a high selectivity for the dehydrogenation reaction. However, the selectivity of an oxide is related not only to the nature of the cation, but also to the pretreatment temperature. Heating to

molecules H2/sec cm2

\

t

\

s!yasl

1.5

1.6

I.?

1.8

1.9

2.0

2.1

2.2

FIa. 23. Specific dehydrogenation rates on various oxides.

high temperatures enhances the dehydrogenation activity of the oxide for a shorter or longer time. Even a specifically dehydrating oxide like A1,O can be given an appreciable dehydrogenation activity lasting

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

83

TABLE V Results of Investigations on the Catalytic Decomposition of Formic Acid on Oxides (-Ha : dehydrogenation; -CO : dehydration) :enation mposition

370" 670' 900"

-

380" 420' 1200" 400' 1500'

60 100 100 100 60-70 90 85-93 100 80 80

20"

19-27

600'

72-82

x 100

Activation energy (kcal/mole)

-

Literature

Sabatier Cremer Cremer Cremer Komarov Schwab Schwab

62 28 30 18 24-26 11

Szabb 24 Stone 23 Mars 40 Mars 47 -Ha: 34 Schwab I-CO: 271

1000"

?

26

Szabb

-

Sabatier

loooo

low low

28

Szabb

-

Adkins Sabatier

(85)

Adkins Schwab Freidlin Fischer Frachon de Pradel

(85) (81)

Adkins Fischer Menon Graeber

(85) (87) (90)

-

(decomp. ~ulfate)

-

(decamp. nitrate)

0 13

-

50 0-6 0-2 4 3 60-70 1-30

llooo

90-94

-

360' 400' 900"

350" 400"

-

970'

1200°

-

23 15 14 9-10 27 21

5

30-50 0-10

16-16

21 21 4 -Ha: 34 -CO: 241 '-CO: 27'

[

(1)

(86) (87) (88) Blake (89) Frachon de Pradel (88) I

0-1 1 4 10

(82,931 (1) (82)

schwab Schwab

(81)

30-73 high

view of the method of preparation used, the authors take this catalyst to be SiO.

84

P. MARS, J. J. I?. SCHOLTEN, AND P. ZWIETERINQ

for some time. I n Table V this effect is shown on several oxides. It is very probable that in most cases not only the selectivity changes, but also the absolute activity (i.e. the activity per unit surface area). The activities for the dehydrogenation of formic acid given in literature are compiled in Fig. 23; here the reaction rate per sq cm of the accessible catalyst surface area is given as a function of temperature. If not given by the authors, the value of the accessible surface area is estimated from experience, This figure shows that the specific activity as well as the activation energy of the reaction depend largely on the nature of the oxide (see also Table V). The bundle of lines for (mineral) MgO was observed by Cremer and Kullich (77) and refers to various temperatures of preheating (indicated in the figure).In our laboratory we obtained confirmatory evidence that selectivity and activation energy can vary appreciably, depending on the history of the MgO sample. Also the presence of foreign metallic ions can enhance the dehydrogenation activity. Freidlin and Levit (86) observed that addition of 5 to 30% K,CO, to silica promotes this reaction. Dalmai et al. ( 6 ) showed that not only potassium, but also lithium and magnesium ions act as promoters. Also on ZnO the addition of alkali (Li or Na) increases the rate of dehydrogenation [Shabrowa (94)l. In this connection it should be mentioned that Na,CO, itself has no measurable activity for the reaction at reasonable temperatures, owing to the formation of the stable sodium formate. The presence of Fe or carbon black, however, renders the formate unstable [Royen and Erhard (95)]. Possibly, this effect is the key to the promotion mentioned above. Some authors made investigations in the field of semiconductors, with the aim of seeing the influence of the properties of these compounds on those of the reaction. Szab6 and Solymosi (82,93) studied the change in activation energy of the decomposition reaction on various doped semiconducting oxides. On MgO-Cr,O, mixtures the influence of reaction temperature and preheating temperature on conductivity, catalytic activity and selectivity of the oxide was studied. After sintering at 600-8OO0C (resulting in spinel formation) high activities of the catalyst were measured. The authors interpret their results in terms of the Hedvall-Huttig hypothesis of active intermediates. 3. Infrared Measurements Infrared measurements on dehydrogenating oxides were performed on ZnO by Hirota et al. (46)and, in our laboratory, on MgO by Scholten.

86

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

Hirota et al., using the infrared technique described above (Section 111, A, 7), measured the infrared spectrum of formic acid chemisorbed at room temperature on zinc oxide, after pumping off the excess formic acid in the gas phase and on the surface. The spectrum observed was practically identical with the spectra of zinc formate. This suggests that the formation of zinc formate on the surface plays a role in the mechanism. Scholten, using the same technique, adsorbed formic acid on MgO a t 100°C (after pumping for 20 hr a t 400OC; surface area 45 meter2/gm). After pumping off the excess formic acid he found that one formic acid molecule was present on 2-3 surface sites. I n the infrared spectrum the stretching-vibration bands of OH groups (at 3500-3600 cm-l, with a tail up to 3100 cm-l) and of the formate ion (CO stretching vibration at -1600 and -1350 cm-1, a C-H stretching-vibration band a t 28902870 cm-1) were easily detected (see Fig. 24).* The sample was heated

\--

40-

20-

u

b

ing to Scholten. The relative transmission is given as a function of the wave number.

for some time at a distinct temperature where the formate ions on the surface decompose slowly. From the pressure increase and mass-spectrometric analysis of the gaseous products the data given in Table VI were calculated. The evolution of gaseous products is obviously accompanied by a proportional decrease in intensity of the formate ion bands. It was remarkable that the intensity of the OH-bands decreased more rapidly, which is in agreement with the initial rapid formation of water. The observed total decomposition rates of formate ions are indicated in Fig. 23 and are near the Iine marked MgO (infrared). It is clear that

* At 780 cm-l a band appeared the meaning of which is not yet clear. The spectrum of bulk Mg formate is unknown, but bulk Ba formate shows a GO-bending band at 770 cm-l.

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

86

these rates are fairly well within the region of the rates of the formic acid decomposition measured in our laboratory, but lower than those measured by Komarov and by Cremer. It should be added, however, that the latter authors did not measure the surface area of their MgOcatalysts, for which reason we estimated the value of that quantity. TABLE VI

Decomposition of Pormate Ions on the Surface of Magnesium Oxide

Heating temperature ("C) 250' 305" 322" 341" a The

Decomposition products (yo)

'HCOO-

before heating 1 0.85 0.64 -0.03

HSO

CO

59.1 9.2 14.7 <0.1

26.6 61.4 58.0 90.7

Ha

H, Reaction rate a _.~ CO molec./cma sec COa

9.2 5.1 14.2 15.2 10.4 16.8 9.2 <0.1

0.35 0.23 0.18 0.10

1.2 1.5 1.1 7.8

x 109

x 1O'O x 10" x 1O'O

reaction rates are corrected for the decrease in coverage

From the HJCO ratio observed (Table VI) it follows that the selectivity of the decomposition reaction amounts to 0.2-0.3 (at high coverage). This value is appreciably lower than the selectivities observed during continuous formic acid decomposition experiments : Schwrtb and Petroutsos (92a) found 0.9; K.omarov et at. (92)reported values between 0.6and 0.7, whereasour own experimentsgavevalues between 0.3 and 0.8. 4. Considerations on the Mechanism Considering the position of the equilibrium CO

+ H,O f CO, + Ha

at 3OOOC (i,e., K,, = 40) one might think that the formation of CO, and H a as the main reaction products could take place through the formation of CO and H,O, followed by the carbon monoxide-steam reaction. On account of the striking analogy with the selectivity of the alcohol decomposition, however, this does not seem so probable, but it is worth while to consider this possibility in the case of the decomposition of formic acid on those catalysts which are known to accelerate the carbon monoxide shift reaction; i.e., Fe,O, and MgO. Directly connected with this is the question whether the three reactions (i.e., the carbon monoxide-steam reaction and the formic acid

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

87

decomposition in the two directions) have a common intermediate step on the catalyst surface: COs

+ H,

These problems have been studied on Fe,O, by Mars (7). It was found that at 300°C the reaction proceeds as to 70435% in the direction of dehydrogenation, giving a H,/CO ratio which is lower than would have to be expected if the watergas shift reaction had reached equilibrium. Even a long residence time of the gas in the reactor has no effect upon the selectivity, which shows that a secondary reaction does not take place. Moreover it was shown that formic acid completely poisons the conversion of additional CO + H,O. Evidently, formic acid is adsorbed very strongly on the surface: as a result, the decomposition of formic acid is zero-order. The situation on MgO is not so clear: the hypothesis of a secondary reaction cannot be rejected. It might be regarded as an argument in favor of this view that, as described above, the H,/CO ratio is smaller during decomposition experiments in vacuo than during continuous experiments, in which, naturally a secondary reaction is more likely to occur. That the H&O ratio always differ from the equilibrium ratio of the watergas reaction [Komarov et al. (92)] also shows that further investigations about this subject are necessary.* As, however, certainly in the case of Fe304, the dehydrogenation reaction cannot be considered as the effect of a secondary reaction following upon the dehydration, a search will have to be made for a special mechanism according to which this dehydrogenation on oxides proceeds. I n this respect not much information can be obtained from investigations made to relate the semiconduction properties of the catalyst to its selectivity. Both the dehydrogenating ZnO and the dehydrating Ti0 are n-type semiconductors, whereas the dehydrogenating MgO and the dehydrating SiO a are well-known insulators. Eucken and Heuer (79), Euken (80), and Wicke (96), not yet aware of the spectroscopic evidence described above, based their ideas about the mechanism of the analogous dehydrogenation of alcohols on their observation that dehydrogenation takes place in particular on oxides with a large cation radius and a small cation valency. As in such oxides

*

Recent results of a detailed study of the watergas reaction and the HCOOH decomposition over various MgO samples will be given in a paper by Scholten, Mars, Menon and van Hardeveld, to be presented at the Third International Congress on Catalysis, Amsterdam, July, 1984.

88

P. U S , J. J. F. SCHOLTEN, AND P. ZWIETERINO

the cation can only be incompletely surrounded, these authors supposed that the reaction proceeds via hydrogen addition to metal ions present on the surface; in analogy with this view, the dehydrogenation of formic acid could be visualized as follows: 0

HCOOH

+ M e 0 -+

0

I

i

Me-----0

: -+ i

’

! ~

1 --3 i

H,

+ CO,+MeO

Me----O

As discussed before there are, however, indications of the importance of formate ions as reaction intermediate. The work of Scholten revealed that the rate of decomposition of formate ions situated on the surface is well within the range of the rate of the decomposition of formic acid on magnesium oxide. The importance of formate ions is also stressed by Komarov et al. (92). These workers analyzed the composition of the catalyst after this had been used for a certain period in the decomposition of formic acid. The analyses showed Fe, Mg, and Mn formate to be present on the separate catalysts. From the fact that the selectivity for the formate decomposition resembles that for formic acid decomposition they concluded that in the last reaction formate ions occur on the surface as an intermediate. Naturally, one is inclined then to look for a mechanism comparable with that for the same reaction on metals, where most probably a formate ion deoomposesinto gaseous carbon dioxide and a chemisorbed hydrogen atom. It is not very difficult to use this picture in explaining the dehydrogenation on semiconductors like ZnO, the hydrogenation activity of which is commonly attributed to the presence of interstitials (Zn or Zn+)with a “metallic” character, and the occurrence of Zn-H bonds on which has been definitely shown by Eischena et al. (97). Also on oxide surfaces freed of OH groups by a drastic heat treatment H atoms can apparently be bound, as is shown by the appreciable activity for H,-D, exchange found on highly heated MgO, ZnO, A1,0,, V,O,, ZrO,, CeO,, Tho,, and WO, by Holm and Blue (98). In some cases even ethylene hydrogenation is shown to occur on A1,0,, Cr203, and MOO, [Hindin and Weller (99)] after heat treatment above 450°C. The explanation of this activity towards hydrogen is attributed either t o the presence of minute amounts of metal ions of lower valency as a result of the high temperature treatment, or to the formation of “oxygen

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

89

bridges," giving rise to highly stressed chemical bonds, as suggested by Frachon de Pradel and Imelik (88). It is certain, however, that this activity is strongly poisoned by minute amounts of water vapor, so that in those cases where the dehydrogenation on oxides like MgO or ZnO seems to occur even in the presence of appreciable amounts of water vapor, the reaction must proceed either by the mechanism presented by Eucken and Heuer (79), Eucken (80), and Wicke (96) or via the formation of carbonate groups: H 0

Me CO 0

/

+ HOOCH---t

/ /"

+ CO,

Me-OCH 0

0

Me-OCH

0 CO

Me

A

0

+ H,

0

or oxalate groups: 0 /O-CH

It

Me \O-CH

__.-

"r'

0

4-

-+ Me

+H,

\o-c

0 I1

NO 0

/o-C \O-

Yo

+ HCOOH-4

/O-CH Me \O-CH

II

+ GO2

Whether the reaction will proceed via the carbonate or via the oxalate cycle will depend on, among other things, the stabilities of the carbonate and the oxalate. These are not known exactly, but the stabilities of the salts derived from basic (dehydrogenating) oxides are of the same order as those of the corresponding formates. It is clear that there is a great lack of experimental data that can be used to decide between the different possibilities.

90

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINO

IV. Dehydration of Formic Acid A. DEHYDRATION OF FORMIC ACID IN

THE

LIQUIDPHASE

I n contrast with the generally accepted stability of formic acid in the gas phase-spontaneous decomposition occurs onIy a t temperatures well above 450" [Blake and Hinshelwood (89)J-it was reported by Barham and Clark (100) that the decomposition of pure anhydrous formic acid in the liquid phase could be observed already a t 40". Between 51 and 99" the decomposition is rather fast, producing CO and H,O. Other investigators state that the dehydration of formic acid in the same temperature range occurs only in the presence of strong acids such as sulfuric and phosphoric acid [Hammett (101) and Gel'bshtein et al. (102)l.The observation of Barham and Clark can be explained from their experimental method; these authors prepared their anhydrous formic acid by distillation of 96% HCOOH after addition of magnesium perchlorate a t 30". Since the vapor pressure of perchloric acid a t that temperature is near that of formic acid, i t is probable that the distilled 99.9% HCOOH (determined by titration) contained perchloric acid, which, being a strong acid, catalyzes the decomposition. The decomposition of formic acid in the liquid phase influenced by acids was intensively studied by Hammett. He found a beautiful correlation between the rate of decomposition of formic acid in concentrated sulfuric acid and the acid function [Hammett (101), Long and Paul (103)]: H , = - log aH+ . f B

(fuH'-)

where aH+ f B fBH+

proton activity , activity coefficient of the base, = activity coefficient of the corresponding acid, = =

The investigators showed that the logarithm of the reaction rate is proportional t o H,, with a slope equal t o unity. They concluded that protonation of formic acid is an essential step in the decomposition, and that the rate of the over-all process will be determined by the monomolecular decomposition of a protonized complex: HCOOH

' ] : ; C H [

+ H,SO,+

[

HC

-+ HCO +

I;]+

+ H,O

+ H*SO,

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

91

and HCO --f H+

+

or

HCO

+

+ H,SO,

+ CO --f

H,SO,+

+ CO

Gel’bshtein, Shcheglova, and Temkin (102)confirmed the results of Hammett, also for the decomposition in phosphoric acid. The nature of the solute has an appreciable influence, however; a t the same value of H,, the reaction is much faster in phosphoric acid than in sulfuric acid. They concluded that reaction (a) is in equilibrium (as already assumed by Hammett), and that one of the decomposition reactions (b), (c), or (d)is rate-determining. To decide between these alternative possibilities, Ropp et al. (104,105)and Biegeleisen et al. (106)studied the decomposition of labeled formic acid compounds in concentrated sulfuric acid. From their observations that the 13C- and 1%-isotope effects are rather high (viz., 7 and 9% respectively), and that of H-D is low, they concluded that a carbon-oxygen bond, and not a carbon-hydrogen bond, is broken in the rate-determining step. This means that reaction (b) is the slowest. The fact that the activation energy is always in the same range (Table VII)is in accordance with the mechanism described: the reaction TABLE VII

Activation Energiea of Acid-Catalyzed Formic Acid Decomposition I n phosphoric acid with I n phosphoric acid with I n 96% H,SO, In 8 6 4 7 % H,SO, In < 0.1% HCIO,

7 2 4 3 % P,O, 81-98~, P,O,

28-25 27-25 20 20-22 26

kcal/mole kcal/mole kcal/mole kcal/mole kcal/mole

[Gel’bshtein (ZOZ)] [Gel’bshtein (ZOZ)] [Ropp (104)] [Schierz (107)] [Barham and Clark (10011

rate is influenced by the value of the equilibrium constant of reaction (a), but the activation energy is principally determined by that of the monomolecular decomposition of the protonated formic acid. The mechanism accepted for this decomposition is in accordance with those of esters and ethers, in which likewise is formed a protonized compound which undergoes decomposition. The fact that the decomposition of formic acid in the gas phase on basic oxides proceeds in the direction of dehydrogenation prompted an

92

P. MARS, J. J. F.

SCHOLTEN, AND

P. ZWIETERINa

interest in the decomposition of formic acid in the liquid phase in the presence of basic compounds. Buiten (108), studying distillation problems, observed that a mixture of pyridine and formic acid evolved carbon monoxide when heated to about 150°C. We repeated this experiment, using a 1 : l (weight) mixture of y-pycoline and formic acid and a mixture of triethylene diamine and formic acid. I n both cases the gas evolved principally contained CO. I n 4% CO, was present in the CO evolved. the latter case also 4% H, Although these basic compounds can form salts with formic acid (as is evident from the high boiling points), the decomposition products are practically identical with those formed in acid media.

+

B. DEHYDRATION OF FORMIC ACID ON METALS I n the chapter dealing with the dehydrogenation on metals we remarked that in practically all publications a predominant decomposition of formic acid into hydrogen and carbon dioxide was reported. I n this respect, however, two problems remain: (a) Are CO, and H, primary products? (b) Does the water-gas shift reaction play a part in the formation of CO, and H, (or CO and H,O) as secondary products? A special study of these questions was made with nickel as a catalyst, while in one case also platinum was used. Already Platonov and Tomilov ( l o g ) , using Ni as a catalyst in a flow system, reported a CO,: CO ratio of about 2.5 in the exit gas, increasing with rising water content above 250°C. They related this observation to the water gas reaction CO,

+ Ha 2 Ha0 + CO Ni

which is catalyzed by nickel, simultaneously with the formic acid decomposition. Walton and Verhoek (31) studied the static decomposition on Nifilms in the zero-order pressure region. The reaction products obtained with different degrees of decomposition were analyzed. A typical result is given in Table VIII. The table shows that in the very first stage of the reaction more CO, than CO is formed. Afterwards the ratio shifts to an about constant value, which depends on the water gas equilibration. This result strongly suggests that CO, was the primary product and that, already after a short reaction time, the water gas reaction starts to play a part. However, it is not impossible that CO formation is a weak primary side reaction.

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

93

This question was studied in detail by Fahrenfort, Sachtler, and van Reijen (12). First, the product distribution in the pyrolysis of bulk nickel formate was studied at different temperatures, both with the reaction gas remaining in contact with the decomposing product, and with the reaction gas rapidly drawn off. TABLE VIII The C0,:CO Ratio in the Decomposition of HCOOH on Ni-Films, According to Walton and Verhaek ( 3 1 )

Temperature ("C)

% HCOOH decomposed

Ratio CO, : CO

189 187.9 188 188.2 188.5

17 37 45 55 68

4 2.6 2.5 2.2 2.3

The results are given in Table IX, together with some results obtained by Kornienko (110)and by Bircumshaw and Edwards (111). TABLE IX The CO:CO, Ratio in the Decompoaition of Nickel-Formate According to Fahrenjort, van Reijen, and Sachtler ( 1 2 )

::;;:

637 572 543

0.156 0.120 0.086

0.083

509 453 438

0.144 0.316 0.316

0.062 0.035 0.029

I I

reaction products remain in contact with Ni

reaction gas drawn off rapidly

I n the second and third columns the experimental CO : CO, ratio is compared with the theoretical value (CO/CO,),, following from tabulated values for the water gas equilibrium. For the static decomposition the values in two columns appear to be practically equal, but when the gas is rapidly drawn off much more CO is formed than would be compatible with the equilibrium concentration,

94

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINU

This result points to a decomposition of bulk formate with both CO and CO, as primary products, and to a further equilibration of these products via the water gas reaction, if the contact time is long enough. It was of course interesting to see if the same conclusions could be drawn [Fahrenfort et al. (12)]from the static decomposition of formic acid on nickel. Formic acid vapor was to this end brought into contact with a 13% Ni-on-silica catalyst, and the reaction products were analyzed after a long contact time and also after they had been rapidly removed from the catalyst by flushing with formic acid. The results are given in Table X. TABLE X The CO:CO, Ratio in the Decomposition of HCOOH on N i , According to Fahrenfort, van Reijen, and Sachtler ( l a )

463 431 381 368

0.146 0.138 0.062 0.058

0.067

473 419 376 351 292

0.244 0.228 0.313 0.489 0.260

0.066 0.034 0.018 0.012 0.003

0.020 0.015

reaction gas remains in contact with Ni

reaction gas rapidly removed

It follows from the first four results tabulated that, unlike in the experiment with bulk formate, the product gases did not react to establish the water gas equilibrium in a measurably short time. This may have been due to the fact that the decomposition temperatures were lower than in the bulk experiment. I n two respects the results differ from those obtained by Walton and Verhoek with Ni-films: (a) more CO was formed than would be expected in view of the water gas reaction; in Walton and Verhoek’s experiment the amount was less; (b) the water gas equilibrium was not reached in a measurable time, whereas in Walton and Verhoek’s experiment it was. I n the experiments in which the reaction gas was rapidly removed (see Table X) the water gas equilibrium was not reached a t all; here the CO concentration was much higher.

95

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

I n conclusion we may say about Fahrenfort, Sachtler and van Reijen's experiments that (a) the results of the decomposition reaction show the same trend as those of the decomposition of bulk formate; the decomposition temperatures are lower in the catalytic reaction than in the bulk-formate decomposition; (b) CO and CO, are both primary products; more CO is formed than would be expected in view of the water gas equilibrium; this is in contrast with Walton and Verhoek's results with films; (c) the water gas reaction modifies the product composition. We think that in the experiments with the Ni-on-silica catalyst (and perhaps also with the bulk-formate) the situation is somewhat obscured by the fact that the CO : CO, ratio may be influenced by adsorption of H,O on the silica carrier, and of H, and CO on the nickel. A reinvestigation with a pure nickel surface is desirable. The expression ([CO] [H,O]/[CO,] [H,]} will be a better guide in this last case, than the CO : CO, ratio as the former is insensitive to adsorption of the reaction products. Eischens and Pliskin (112) studied the product distribution in the decomposition a t - 60°C on a platinum-on-silica catalyst. Infrared analysis of the adsorbed phase was combined with gas-phase analysis. The decomposition of 0.06 cm3 of HCOOH added to the catalyst was so rapid that bands caused by the adsorbed acid could not be observed. At - 60"C, 0.021 cm3of CO was found on the surface and 0.032 om3 of CO, in the gas phase. The CO is not formed via dissociative CO, adsorption, for, in a separate experiment carried out under the same conditions, addition of a mixture of 0.027 om3 of CO, and 0.027 em3 of H, gave only 0.003 cm3 of CO on the surface. It was further observed that the water gas reaction is not operative in this case, as the equilibrium leans heavily to the CO, side a t this low temperature. It was concluded that the reaction over Pt a t -60°C proceeds for one third part as a dehydration and for two thirds as a dehydrogenation. We wonder if the highly evacuated silica carrier did not cooperate in the decomposition to H,O and CO; Mars found that highly activated silica is an extremely active dehydrator, even a t low temperatures. I n addition, working on platinum a t higher temperatures (200-250°C), Block and Vogl(14) found an exclusive decomposition into CO and H ,. Finally we may say that more research needs to be done on the possible dehydration on metals. The most conclusive information would be obtained from studies on the dehydration on metals with a low activity for the water-gas reaction, in the absence of carrier materials.

,

96

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINO

C. DEHYDRATION ON OXIDES 1 . The Kinetics of the Reaction

Although various authors report activities of oxidic catalysts for the dehydration of HCOOH, only some of them studied the kinetics of this reaction. The results of some relevant investigations summarized in Table XI are in some cases conflicting. TABLE XI Kinetic Data on the Dehydration of Formic Acid on OxirEee

Catalyst

Temperature ("')

TiO, A1 and Ti silicate Cr,O, and Cr-Mg spinels Fed4 A1 8 0 s SiO, and glass SiO, SiO,

HCOOH pressure (mm mercury)

140-180

Reaction order

1

400-1600? 760?

200 380-460 316 200

760 80-260 1-16 760 300-600 10-8-10-~

300-400 376

180-300

>O

>o, < 1 0 0 0 1 1

References

Hinshelwood Rastrenenko

SzabB Mars Menon Dalmai Mars Crocker

Rastrenenko and Szab6, who observed an order higher than zero (with respect to formic acid), experimented under conditions from which it can be calculated that very probably pore diffusion influenced the reaction rate. Hinshelwood and Mars observed first-order reactions on TiO, and SiO,, respectively, whereas Dalmai et al. state that on SiO, and glass the reaction is zero-order. On Alto,, Menon (90) found a retarding effect of water a t 200°C; this could be described by the Langmuir expression rate = k

'

bf * P f

1 +bf.Pf +bW.P,

where p f = formic acid pressure, and pw = water vapor pressure, with bf.pf > > 1 and b,/b, = 0.05. This result shows a competition between the adsorption of the reaction intermediate and water on the same surface sites. On SiO, the retarding effect of water was observed at 300°C by Dalmai et aE. (6). The activation energies of the dehydration reaction are compiled in Table V, which shows that in most cases the value is somewhere between 15 and 25 kcal/mole. They are always lower than the values of the dehydrogenation reaction. N

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

97

By studying the formic acid dehydration on AI,O,-SiO, samples with Al/Si ratio Fischer and Sebba (87) observed that the activation energy decreases with increasing acidity of the surface (measured by means of titration with n-butylamine). More extensive information is available about the kinetics of the dehydration of alcohol on oxides. Owing to the work of e.g., Balandin et al. (114)and of Balaceanu and Jungers (115)i t has become clear that the rate of this reaction is proportional to

which shows that, here too, the reaction rate is determined by the concentration of an adsorbed intermediate competing with the adsorption of water on the same sites. 2. The Activities of the Various Catalysts Figure 25 shows the specific activities of various oxides plotted versus 103/T. The accessible surface area, if not given by the authors, has been

estimated from experience gained with catalysts of the same composition and mode of preparation. It is clear that the specific activities of the various oxides do not differ by more than a factor fifty, substantially less than was found above (Fig, 3) for the dehydrogenation reaction. On A1,0, not only water vapor but also added ammonia and piperidine retard the reaction, The activity of Peso, is retarded by the presence of piperidine in the gas stream [Mars (7)l. Addition of metallic ions to the catalysts can also influence their activity. In the case of alcohol decomposition a retardation of the dehydration reaction is always observed, but in the decomposition of formic acid the result depends on the type of ion added. On silica gel addition of potassium promotes the reaction [Dalmai et al. (S)]. The large decrease in activation energy (12 to 2 kcal/mole) points to a change of the mechanism and to the possibility of the potassium itself acting as a catalyst. Addition of Ca or of 1% Li has no influence on the activity, addition of 1-2% Mg retards the reaction [Dalmai et al. (S)]. Addition of 3-574 Li,O to AI,O, did not change the reaction rate and the activation energy, which amounted to 19 to 20 kcal/mole [Noller et al. (116a)l. 3. Infrared Investigations

Investigations of the mechanism of the dehydration reaction by means of the infrared technique are reported for alumina and for silica gel. On alumina Hirota, et al. (45,117, 118) observed absorption bands

98

P. MARS, J. J. F. SCIIOLTEN, AND P. ZWIETERING

both of formic acid and of the formate ion a t room temperature. On heating to 220°C the physically adsorbed acid disappeared, as did the gaseous acid; the intensity of the formate ion absorption, however, remained unchanged during a long heating time. 10”

lot4

1013

lo=

molecules

10”

1010

t

14

16

18

20

22

FIG.25. Specific dehydration rates on various oxides.

I n our laboratory also Scholten applied the infrared technique to Also,, preheated to 450”C, using a cell which could be heated to induce reaction, but remained a t room temperature when the absorption spectrum was being measured. After addition of formic acid a t 200°C up to a coverage of two per cent, he observed the symmetric and antisymmetric c-0 stretching-vibration band and the C-H band of the formate ion (Fig, 26). Furthermore it was observed that no shift in the position of the stretching-vibration band of the OH groups occurred upon adsorption of formic acid.

T H E CATALYTIC DECOMPOSITION O F FORMIC ACID

99

After heating a t distinct temperatures between 25OOC and 330°C for 25 min he observed a decrease in intensity of the formate bands. The decrease of the three bands was strictly proportional, showing that they belonged to the same adsorbed species. The amount of reaction proH,O) was measured manometriducts (practically exclusively CO cally. From this the amount of adsorbed formic acid can be calculated;

+

*/. transmission

O!

4000

3500

3000

-

2500

2000

1500

1000

wave number (cm-'1

FI5. 26. Infrared spectrum of formic acid, chemisorbed on aluminium oxide, according to Scholten. The relative transmission is given as a function of the wave number.

the latter proved to be strictly proportional to the intensity of the absorption bands. Table XI1 gives the results of the measurements of the decomposition of the formate ions on the surface, together with the rates of the continuous decomposition of formic acid on A1,0, measured by Menon (90). The fact that the two rates are in the same range indicates not only that the dehydration of formic acid on alumina proceeds via a surface formate, as already shown by Hirota, but also that the decomposition of these formate ions determines the rate of the process. This is compatible with zero-order kinetics. On silica, Hirota et al. (45, 118) detected only physically adsorbed formic acid in the form of a monomer at room temperature. The shift in the absorption frequency of the OH-groups indicates an interaction with a proton of the surface, but the C=O bands are not involved and disappear on heating to 100 and 25OOC. No new bands appear and formate bands could not be detected. Imelik et al. (119) also concluded that adsorption via OH-groups takes place but they believe that also the C-H-group interacts with the

100

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINU

surface. Clarke and Pullin (60)likewise found that there must be a bond involving the OH group of the acid, but could not detect an important change in the absorption of the C-H bond. In accordance with these investigations, Fahrenfort and Hazebroek (58)found that no specific absorption band appeared after adsorption of formic acid on silica. TABLE XI1

Comparison of the Rate of the Decomposition of Formate Iona on the Surface of Al,O with the Rate of the Continuous Decomposition of Formic Acid on AllO,

T

e

("C)

(HCOO-) ads decomposition rate (niolec./cm*sec)

Formic mida decomposition rate (molec./cm*sec)

~

260 260 276

0.0116 0.0080 0.0066

x 108 6 X lo8 26 X lo8 7

140 X lo8 100 x 10' 160 X lo8

Calculated for the value of 0 indicated, by means of the rate formula found by Menon:

r = 1.26 x 1081 x 0 x e-21700/RTmolec./cm* sec.

4. Considerations on the Mechanism

The main lines of the mechanism of the dehydration of formic acid in the liquid phase under the influence of acids are clear. As pointed out in the paragraph in question the decomposition of formic acid is completely analogous to that of alcohol and proceeds via proton addition, water elimination, and the decomposition of an (unstable) carbonium ion. I n considering the mechanism occurring on a solid surface the analogy with the last reaction is an obvious starting point. If the surface contains OH-groups and if these groups have proton donor properties, the reaction might proceed along the same route. For the dehydration of alcohol on oxide surfaces this mechanism has been suggested by Whitmore (120),Day (121)and de Boer (122).The importance of OH-groups for the dehydration of formic acid on the surface of silica has become evident from the work of Dalmai, Frachon de Pradel, and Imelik (6). The possible steps in this reaction are visualized in route I of Fig. 27. The bonding between the formic acid and the catalyst surface, however, is apparently so weak that on the acidic surfaces (SiO,, TiO,) mostly first-order kinetics are observed. It is possible that this proton addition mechanism plays a secondary role on A1,0,, where it may involve acid sites on the surface.

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

101

The analogy between the reaction in the gas phase and that in the liquid phase is beautifully demonstrated by Andrianova and Bruns (123).Their catalyst was an ion exchange resin containing S0,H-groups. A t 150°C the reaction is first-order and slowed down by water; the

iHC [.-“-B [6AJMeOH

+

+

H

[ H - 4

I

-H,O

Ib

1

[i]+ 0

IC

co

J.-

(slowest step)

1t[q Hi

-Hzo

[JJ IIc

1

(slowest step)

Me+

FXO.27. Proposed reaction paths of dehydration on oxides (Me = metal ion).

activation energy amounts to 25 kcal/mole. These data. are comparable with those found in the liquid phase. This mechanism provides a harmonious explanation of the poisoning effect observed upon addition of gaseous bases (NH,, piperidine). I n

102

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERINO

analogy with the reaction in liquid acids, reaction step I c (see Fig. 27) is expected to be faster than step Ib, because of the instability of the protonized carbon monoxide group. In the liquid phase the nature of the anion is found to exert a strong influence on the rate of the dehydration reaction. On oxidic surfaces the role of the anion is taken over by the metallic ion surrounded by oxygen ions. Because of the proton-donating property which must be ascribed to the hydrogen of the OH group this mechanism can only be valid for oxides the OH-groups of which have acidic properties, but not in cases where the proton is bonded too firmly, or where no OH groups are present. I n these instances the formation of formate ions may proceed as shown by infrared measurements (see above). Komarov et al. (92) confirmed that this occurs not only on the more basic (dehydrogenating) oxides of Mg, Fe, Mn (see Section 111,C), but also on Al,O,. It may thus be taken for granted that in the dehydration of formic acid on A1,0, formates play an essential role. This mechanism is visualized in Fig. 27, branch 11. The infrared measurements made by Scholten showed that the decomposition of the formate ion determines the overall reaction rate, step IIc being the slowest step. This mechanism is analogous to that of the dehydration of alcohol on A1,0,; Rubinshtein (124)showed alcoholate groups to be present on the surface. Additional evidence that also in this case the proton of the OH group is involved was given by Leugering (125).Passing alcohol over y-AlzO, the surface of which was covered with OD groups instead of with OH groups, this worker found that much more deuterium was eliminated from the surface than when water had been led over the catalyst under identical conditions. To sum up, there seem to exist two mechanisms for the dehydration: one proceeding via proton transfer from the surface to the formic acid on surfaces with acidic OH groups, the other via formation of a formate ion on surfaces of a neutral or basic character. I n both dehydrogenation and dehydration formate ions on the surface thus seem to play a central role. It is therefore of interest to study the decomposition of bulk formates, although it is obviously to be expected that there will always be differences between two-dimensional surface compounds and bulk compounds. The decomposition reactions of various formates will be discussed in the following section. Knowledge of the properties of these compounds will enable us to deal more extensively with the reaction mechanisms of the dehydration and the dehydrogenation of formic acid.

THE CATALYTIC DECOMPOSITION O F FORMIC ACID

103

V. Decomposition of Bulk Formates The decomposition of bulk formates has drawn special attention, in connection with the function of promoters in iron and cobalt catalysts for the Fischer-Tropsch synthesis. I n the investigations made on this point the interest was, naturally, focused on the organic products formed during the decomposition of the formate [Hofmann and Schibsted (53), Marec and Hahn (126)l. Hofmann (53) found an appreciable amount of formaldehyde (about 25%) and small amounts of methyl formate during the decomposition of zinc formate. Lithium formate produced acetone (about 20%); from lead formate, formaldehyde and methyl alcohol were formed. Pichler (127) found that during the decomposition of calcium formate, oxalate was formed. I n general it appeared that the nature and the amount of the organic by-products depended largely on the reaction conditions [Hofmann (53)]. Our interest is of course directed more towards the formation of inorganic products. Four main reaction routes are possible : M(OOCH),

-+

M

+ (Ha + 2 CO,) n2

(14

3

M

+ (H,O + CO + C O , ) 9

(Ib)

+ (Ha + CO + CO,)

(IW

+ K O + 2 CO)9

(IIb)

+ MO,,

--f

MO,

Kornienko (128) in some cases reports the formation of carbon, whereas in other cases the formation of carbonates seems highly probable, in view of the stability of these compounds under the experimental conditions. The affinity for oxygen of the metal involved determines whether reactions of type I or type I1 occur. Nickel formate produces nickel, magnesium formate produces magnesium oxide. Of special interest to us, however, is the extent to which reactions of type a or b occur, and whether the Ia/Ib or the IIa/IIb ratio is in any way related to the selectivity of the catalytic decomposition of formic acid on the metals and oxides in question. Furthermore it is’worth while to investigate whether the stability of the bulk formates (e.g., the decomposition temperature) bears a relation to the catalytic activity of the corresponding metals or oxides.

104

P. M A R S , J. J. F. SCHOLTEN,AND P. ZWIETERING

I n Table XI11 a survey is given of data on the decomposition of bulk formates and of the direction of the formic acid decomposition on the corresponding metals or oxides. TABLE XI11 Data on the Decompoaition of Formates

Metal Temperature of ion decomposition

Main Reaction

("C)

Literature

Selectivity of metal or oxide for formic acid decomposition; dehydrogenation rate x 100 rate of total decomposition

Ag+ cup+

90-106 170-180

Sna+ Ni2+

180 166 170-220 210 226 240 270 280 300 240 260 176 240 260 260 200 290 296

Ala+ Zna+

coa+

Fea+

Mna+

CrJ+

Mg2+

246 280 338 296 300 338 310 290 340

330 380 390 446

Keller (129) Kornienko (128) Komarov (92) Beilstein (130) Ia/Ib = 1 Bircumshaw (111) Ia/Ib = 1 Krogmann (131) Hofmann (53) I& Rienacker (132) Zapletal (133) Ia/Ib = 6 Fahrenfort (59) Ponec (134) Ia/Ib = 4 Kornienko (128) IIb Mars (84) IIb Komarov (92) Beilstein (130) IIa Hofmann (53) IIa/IIb 2 Mars (84) Stone (83) I and I1 Caujoile (135) Zapletal (133) I and 11; Kornienko (128) Is/IIa 1 +Fe,O, Hofmann (53) Ib Komarov (92) IIa Kornienko (128) Hofmann (53) IIa Komarov (92) IIa Kornienko (128) IIb Mars (84) IIa Hofmann (53) Zapletal (133) Hofmann (53) IIa Rienacker (132) IIa Mars (84) Ponec (134) IIa Komarov (92)

Ia

-

-

about 96 large about 90

0-1 1 86-93

large

80

19-27

19-27

30-80

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

105

As regards selectivity it is clear that there exists a general correlation, as was already concluded by Komarov et al. (92). Exceptions were observed in the cases of the metals nickel and cobalt. For the catalytic reaction on both metals practically complete dehydrogenation is found, whereas in the decomposition of the bulk formates CO is also produced in appreciable amounts. I n interpreting this result caution should be exercised, however. Dependent on the experimental conditions, the water-gas shift reaction can occur on the metal phase already formed. Careful experiments by Fahrenfort et nl. (12) have shown that under conditions which exclude a secondary reaction the product distribution of the decomposition of Ni-formate and of the catalytic reaction on Ni are equal within the experimental error (see also Section IV, B). BOO-

700.

600.

500.

400.

300i 300

400

500

600

FIG. 28. Relative activity versus stability of the bulk formates. (For the definition of parameters see text).

No direct correlation exists between decomposition temperature and selectivity; thus, the dehydrating A1,0, is located between mainly dehydrogenating specimens. As to the second point introduced above, that of a relation between the stability of the bulk formates (expressed in the value of the decomposition temperature) and the catalytic activity of the relevant metal or metal oxide, it is clear that no more than a qualitative agreement may be expected. I n the decomposition of formates sometimes appreciable amounts of oxalates and carbonates are formed, which, a t least on the dehydrogenating (basic) oxides, are more stable than their formates. The presence of these anions during the catalytic reaction has, however, never been clearly demonstrated. Furthermore, during the decomposition of formic acid the surface intermediates may

106

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

have properties more or less like those of a basic formate, and different from those of the normal formates. Nevertheless it is encouraging that in an appreciable number of cases a fair correlation can be demonstrated, as is shown in Fig. 28. Here the temperature at which the total reaction rate amounts to 1.6 x l O I 4 molecules/cmZ sec, ( T r ) ,is plotted versus the temperature of decomIt is clear that the equal-rate temperaposition of the bulk formates ( Td). tures of the various catalysts are related to the decomposition temperatures of the corresponding bulk formates. The decomposition temperatures of Au, Pt, and Pd formates are not known, these compounds being unstable a t room temperature. I n accordance with the general picture, the T, values on these metals are far below those observed on Ag, Ni, and Cu. The correlation found suggests that the activation energies of the surface reactions will be symbatic with those of the decomposition of the bulk compounds. Summarizing we may say that in a great number of cases a strong analogy can be demonstrated between the chemical properties of bulk formates and the catalytic properties of the corresponding metals or oxides in the decomposition of formic acid.

VI. General Discussion From the foregoing discussion of the available data it has become clear that although the general pattern of the mechanism of the catalytic decomposition of formic acid seems to be firmly established, there still remains a lack of reliable data from which to derive a full understanding of the details. It seems worthwhile to summarize here the most important features of kinetics and mechanism, and t o indicate those problems which remain to be solved. Making an exception for those surfaces which, as a result of their acidic nature, seem to catalyze this reaction via proton addition, in analogy with the mechanism established for the catalytic effect of liquid acids, we can conclude that in most cases the reaction proceeds via a formate ion as surface intermediate: HCOOH

T H E CATALYTIC DECOMPOSITION O F FORMIC ACID

107

The kinetic behavior observed corresponds with this picture; dependent on conditions of temperature and pressure, a transition from first-order kinetics (step A rate-determining) t o zero-order (step B or C rate-determining) is found in a number of cases. As demonstrated already by Fahrenfort, van Reijen, and Sachtler (12) by means of a I< volcano curve,” it is in particular the noble metals like Au and Ag which show a high activation energy for step A, whereas on the other metals and also on oxides the stability of the formate ion is so high that, generally speaking, step B or C governs the over-all reaction rate. Figure 29 illustrates the energy pattern for a number of typical cases. En t rgy ( Kcallmolt)

i

207 15. 10.

5

0.-5-

-10-15Adsorbed intt rmtdiatt

R t a c t i o n - coordi n a t t

-

FIG. 29. Possible potential energy curves for the decomposition of formic acid on Ni (full line), Cu (dotted line) and Au (dashed line). I n view of the diversity of values given for the activation energies, and the uncertainty as to the heats of adsorption, the figure has only qualitative significance.

I n kinetic experiments in the first-order region (e.g., on Cu and Au) pre-exponential factors were found from which, on the basis of the absolute reaction rate theory, the conclusion was reached that a mobile transition state is involved. I n the zero-order region the pre-exponential factors show a large compensation effect, but are generally much lower than the theoretical

108

P. MARS, J. J. F. SCHOLTEN, AND P. ZWIETERING

value of 1013 molecules/site sec, which holds for a monomolecular decomposition [G.C. Bond (136)l. It is certain, however, that both for metals and metal oxides in the zero-order region there is a close correlation between the catalytic activity and the stability of the bulk formates. All substances investigated show a reaction rate of the order of l O I 4 molecules/cm2 sec, at a temperature just above the decomposition temperature of the corresponding bulk formate. However, a closer look a t the details of the decomposition of a formate ion at the surface reveals a number of problems. It is not certain whether the decomposition is monomolecular or bimolecular. From their work on nickel, Schuit et at?.(137) came to support the assumption of a monomolecular process, on the strength of energy considerations, and most other authors make this assumption more or less intuitively. The possibility of a bimolecular process can, in our opinion, not be definitively excluded, however. I n a monomolecular decomposition the direction of the process (the selectivity) must be determined from the structure of the transition state. It is acceptable that on those surfaces which show an affinity for hydrogen atoms dehydrogenation will be favored, and proceed via an intermediate in which the hydrogen of the formate ion has already formed a bond with a metal atom. I n those cases where the oxygenmetal bond is much stronger than the Me-H-bond a n OH group will be formed, which leads to dehydration. If such a monomolecular mechanism is accepted a possible explanation for the low pre-exponential factors observed is, that in the initial state (the formate ion) there are internal degrees of freedom+.g., rotation with the axis perpendicular to the surface, and bending vibrationswhich are lost in the transition to the activated state. This can be schematically represented as follows:

initial state (rotation and bending vibrations possible)

activated state (restricted possibility of motion)

THE CATALYTIC DECOMPOSITION OF FORMIC ACID

109

I n a bimolecular process it seems possible that interaction takes place on the surface either between two formate ions, resulting in the formation of oxalate groups, or between a formate ion and a hydroxyl group, giving carbonate ions, as was indicated in the discussion of the dehydrogenation on basic oxides (see Section 111, C). On a metal surface a bimolecular reaction might take place between a formate ion and a chemisorbed H-atom, producing molecular hydrogen and CO in one reaction step. An argument in favor of this mechanism may be found in the kinetic isotope effects on Ag observed by Block and Kral (21) (see Section 111,A, 5). The basic observation that it might be possible to identify a welldefined chemical entity as the intermediate in this catalytic reaction has highly stimulated the search for an explanation in terms of a twodimensional chemical reaction. I n this connection it will be necessary at some time to answer the question as to how far the metal atom taking part in the surface formate has lost its links with the metal structure, in other words to what extent the intermediate has to be considered as a “metal formate” physisorbed on the metal surface; this last extreme view is favored by some investigators [e.g., Fahrenfort ( 1 2 ) ] .No doubt there are a number of data which are in agreement with this picture, e.g., the close correspondence between reaction conditions of the catalytic reaction and the stability region of the corresponding bulk formates, and the analogy in product distribution. Especially the large influence of surface orientation on catalytic activity, which was settled unambiguously in a number of cases, can, in our opinion, only with difficulty be reconciled with the view that the metal atom of the formate “molecule” should have completely “forgotten” the metal structure to which it first belonged. I n the case of formic acid decomposition on a surface with a particular distribution of crystallographic planes one would expect a redistribution of planes, if the decomposition occurred via a physically adsorbed formate. Hence, a slow change in activation energy with reaction time would be expected to take place, simultaneously with the change in the plane distribution. As far as we know this has never been observed; perhaps a closer study could reveal whether or not such a phenomenon actually occurs. It is clear that the considerahle advances made towards a complete understanding of this catalytic reaction has its important consequences for our knowledge of catalysis in general, and obviously this has been the main reason for investigating this reaction a t all. The question may be raised, therefore, whether this model reaction is typical of a large number or group of catalytic reactions. Also in other catalytic reactions intermediates are formed whose properties may be close to those of

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bulk compounds. Probably this is the case in reactions leading to the formation of CO, or SO,, where superficial carbonates [Garner et al. (138),Courtois et al. (139)]or sulfates [Pewny (140)]seem to play a role. On the other hand, we must realize that in a large number of catalytic reactions surface complexes must be accepted without a straightforward chemical analogy being a t hand. Well-known examples are found in double-bond hydrogenation, hydrocarbon synthesis and ammonia synthesis. Therefore, it is as yet impossible to give a correct estimate of the significance the research in the field of catalytic formic acid decomposition has for our understanding of heterogeneous catalysis in general, but in our opinion a warning against over-estimation is justified.

REFERENCES 1. Mailhe, A., and Sabatier, P.,Compt. rend. mad. sci. 152, 1212 (1911). 2 . Rienacker, G., and Hansen, N., 2. Elektrochem. 60, 887 (1956). 3. Rieniicker, G., and Hansen, N., 2. anorg. u. allgem. Chem. 285,283 (1956). 3a. Rienacker, G., and Hansen, N., 2. anorg. u. allgem. Chem. 284, 162 (1956). 4. Schwab, G.-M., Discussions Faraday Soc. 8, 166 (1950). 5 . Mars, P., i n “The Mechanism of Heterogeneous Catalysis” (J. H. de Boer, ed.), p. 49. Elsevier, Amsterdam, 1960. 6 . Dalmai, G., Frachon de Pradel, A., and Imelik, B., i n “Actes du DeuxiQmeCongrQs Internationalc de Catalyse, Paris, 1960,” p. 865. Edition Technip, Paris, 1961. 7. Mars, P., 2. physik. Chem. (Prankjurt) [N.S.] 2 2 , 309 (1959). 8. Waring, W., Chem. Reas. 51, 171 (1952). 9. Rienacker, G., 2. anorg. u. aEEgem. Chem. 227, 353 (1936). 10. Schwab, G.-M., J . Phys. Chem. 50, 427 (1946). 11. Rienacker, G., and Dietz, W., 2. anorg. u. allgem. Chem. 228, 65 (1936). 12. Fahrenfort, J., van Reijen, L. L., and Sachtler, W. M. H., i n “The Mechanism of Heterogeneous Catalysis” (J.H. de Boer, ed.), p. 23. Elsevier, Amstcrdam, 1960. 13. Hinshelwood, C. N., and Topley, B., J . Chenz. Soc. 123, 1014 (1923). 14. Block, J., and Vogl, J., 2. Ebektrochem. 63, 3 (1959). 15. Tamaru, K., Trans. Paraday SOC.55, 824 (1959). 16. Rieniickcr, G., and Voltcr, J., Angew. Chem. 69, 545 (1957). 17. Rieniickcr, G., i n “Bericht von der Hauptjahrestagung der Chemischen Gesellschaft in der D.D.R.,” p. 97. Chem. Ges. D.D.R., Berlin, 1959. 18. Voltcr, J., Dissertation, Humboldt Universitat, Berlin, 1958. 19. Tamaru, K., Trans. Faraday Soc. 55,1191 (1959). 20. Schwab, G.-M., and Schwab.Agallidis, E., Ber. 76, 1228 (1943). 21. Block, J., and Kral, H., 2. Elektrochem. 63, 182 (1959). 22. Sosnovsky, H. M. C., J . Chem. Phys. 23, 1486 (1955). 23. Sachtler, W. M. H., and Fahrenfort, J., i n “Actes du Deuxibme Congrbs Internationale de Catalyse, Paris, 1960,” p. 838. Edition Technip, Paris, 1961. 24. Eley, D. D., and Luetid, P., Trans. Faraday Soc. 53, 1483 (1957). 25. Sachtler, W. M. H., and Fahrenfort, J., i n “Actes du Deuxikme CongrOs Internationale de Cetalyse, Paris, 1960,” p. 831. Edition Technip, Paris, 1961. 26. Sachtler, W. M. H., i n “Actes du Deuxikme Congrks Internationale do Catalyse, Paris, 1960,” p. 863. Edition Technip, Paris, 1961.

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27. Glasstone, S., Laidler, K. J., and Eyring, H., “The Theory of Rate Processes,” p. 351. McGraw-Hill, New York, 1941. 28. Schwab, G.-M., Trans. Faraday Soc. 42, 689 (1946). 29. Volter, J., Diplomarbeit, Rostock, 1953. 30. Rienbcker, G., and Volter, J., 2. anorg. u. allgem. Chem. 296, 210 (1958). 31. Walton, D. K., and Verhoek, F. H., Advances in Catalysis 9, 682 (1957). 32. Walton, D. K., and Verhoek, F. H., Advances in Catalysis 9 , 6 8 8 (1957). 33. Rienlicker, G., 2. Elektrochem. 46, 369 (1940). 34. Rienlicker, G., Bremer, H., and Unger, S., Naturwissenschaften 39, 259 (1952). 35. Rienlicker, G., and Bremer, H., 2. anorg. u. allgem. Chem. 272, 126 (1953). 36. Cunningham, R. E., and Gwathmey, A. T., Advances in Catalysis 10, 57 (1958). 37. Crocker, A. J., and Robertson, A. J . B., Trans. FaradaySoc. 54,931 (1958). 38. Sosnovsky, H. M. C., Ogilvie, G. J., and Gillam, E., Nature 182, 523 (1958). 39. Sosnovsky, H. M. C., J . Phys. Chem. Solid8 10, 304 (1959). 40. Constable, F. H., Proc. Roy. SOC.A108, 355 (1925). 41. Cremer, E., Advances in Catalysis 7, 75 (1955). 42. Ogilvie, G. J., J . Phys. Chem. Solids 10, 222 (1959). 43. Duell, M. J., andRobertson, A. J . B., Trans.FaradaySoc. 57, 1416 (1961). 44. Otaki, T., J . Chem. SOC. Japan 80,255 (1959). 45. Hirota, K., Kuwata, K., Otaki, T., and Asai, S., in “Actes du Deuxibme Congrbs Internationale de Catalyse, Paris, 1960,” p. 809. Edition Technip, Paris, 1961. 46. Hirota, K., Kuwata, K., and Nakai, Y., Bull. Chem. SOC.Japan 31,861 (1958). 47. Suhrmann, R., and Wedler, G., Angew. Chem. 67, 655 (1955). 48. Suhrrnann, R., and Wedler, G., Advances in Catalysis 9, 223 (1957). 49. Sachtler, W. M. H . , J . Chem. Phys. 25,751 (1962). 50. Sachtler, W. M. H., and Dorgelo, G. J . H., 2. physik. Chem. (Frankfurt)[N.S.] 25, 69 (1960). 51. Dowden, D. A., and Reynolds, P. W., Discussions Faraday SOC.8, 184 (1950). 52. Hofmann, K., and Schibsted, H., Ber. 51, 1389 (1918). 53. Hofmann, K., and Schibsted, H., Ber. 51, 1398 (1918). 54. Yaroslavskii, N. G., and Terenin, A. N., Doklady Akad. Nauk S.S.S.R. 66, 885 ( 1949). 55. Eischens, R. P., and Pliskin, W. A,, Advances in Catalysis 10, 2 (1958). 56. Hirota, K., Otaki, T., and Asai, S., 2. physik. Chem. (Frankjurt) [N.S.] 21, 438 (1959). 57. Hirota, K., and Nakai, Y., J . Chem. SOC. Japan 80,700 (1959). 58. Fahrenfort, J., and Hazebroek, H. F., 2.physik. Chem. (Frankfurt) [N.S.] 20, 105 (1959). 59. Fahrenfort, J., van Reijen, L. L., and Sachtler, W. M . H., 2. Elektrochem. 64, 216 (1960). 60. Clarke, J. K. A., and Pullin, A. D. E., Trans. Faraday Soc. 56,534 (1960). 61. Eischens, R. P., and Pliskin, W. A., in “Actes du Deuxibme Congrbs Internationale de Catalyse, Paris, 1960,” p. 789. Edition Technip, Paris, 1961. 62. Eischens, R. P., in “Actes du Deuxibme Congrbs Internationale de Catalyse, Paris, 1960,” p. 854. Edition Technip, Paris, 1961. 63. Eischens, R. P., Francis, S. A., and Pliskin, W. A,, J . Phys. Chem. 60, 194 (1956). 6 4 . Sachtler, W. M. H., in “Actes du Deuxibme Congrbs Internationale de Catalyse, Paris, 1960,” p. 853. Edition Technip, Paris, 1961. 65. Hirota, K., in “Actes du Deuxibme Congrbs Internationale de Catalyse, Paris, 1960,” p. 856. Edition Technip. Paris, 1961.

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66. Eischens, R. P., Pliskin, W. A., and Francis, S. A,, J. Chem. SOC.22, 1786 (1964). 67. Rienacker, G.,and Bade, H., 2. anorg. u . allgem. Chem. 248,46 (1941). 68. Rienacker, G., Wessing, G., and Trautmann, G., 2. anorg. u. allgem. Chem. 236, 262 (1938). 69. Rienacker, G.,2.Elektrochem. 47, 806 (1941). 70. Schwab, G.-M., Trans. Faraday Soc. 42, Bg8 (1946). 71. Mott, N. F., and Jones, H., “The Theory of Metals and Alloys.” Oxford Univ. Press, London and New York, 1940. 72. Schwab, G.-M., Dechema Monographien 38, 206 (1960). 73. Fuderer-LuetiO, P., and Brihta, I.,Croat. Chem. Acta 31, 75 (1969). 74. Rienacker, G.,Abhandl. deut. Akad. Wiaa. Berlin, K l . Chem., Geol. u. Biol. 3,(1966). 75. Eley, D. D., and Luetid, P., Trans. Faraday SOC.63, 1476 (1967). 76. Rienacker, G., and Hildebrandt, H., 2. anorg. Chem. 248, 62 (1961). 77. Cremer, E., and Kullich, E., Radex Rundachau, p. 176 (1960). 78. Balandin, A. A., et al. Zhur. Piz. Khim. 33, 2476 (1969). 79. Eucken, A., Naturwieaenschaften 36, 48 (1969). 80. Eucken, A.,and Heuer, K., 2.phyaik. Chem. (Leipzig) 196,40 (1960). 81. Schwab, G.-M., and Schwab-Agallidis, E., J . Am. Chem.SOC.71, 1806 (1949). 82. Szab6, Z.G., Solymosi, F., and Batta, I., 2. phyaik. Chem. (Frankfurt) [N.S.] 17, 126 (1958). 83. Stone, F. S., and Evans, P. L., private communication. 84. Mars, P., Unpublished work. 85. Adkins, H., J. Am. Chem.Soc. 45,809(1923). 86. Freidlin, L. Kh., and Levit, A. M., Zhur. Obahehi Khim. 21,1256 (1951). 87. Fischer, J. B., and Sebba, F., in “Actes d u Deuxibme Congrbs Internationale de Catalyse, Paris, 1960,” p. 711. Edition Technip, Paris, 1961. 88. Frachon de Pradel, A., and Imelik, B., J . chim. phya. 56, 1 (1969). 89. Blake, P. G., and Hinshelwood, C., Proc. Roy. SOC.A265, 444 (1960). 90. Menon, P. G., details to be given at the Third International Congress of Catalysis to be held in Amsterdam, 1964. 91. Graeber, E. G., and Cryder, D. S., Ind. Eng. Chem. 27,828 (1936). 92. Komarov, N. A., Chernikova, E. A., Komarov, G. N., and Leonchik, Z . I., Veatnik Leningrad Univ. 16, 120 (1960). 92a. Schwab, G.-M., and Petroutsos, G., J . Phya. Chem. 64, 681 (1960). 93. Szabb, Z. G., and Solymosi, F., 2. Elektrochem. 63, 1177 (1969). 94. Shabrowa, G. M., Doklady Akad. NaukS.S.S.R. 133, 1376 (1960). 95. Royen, P., and Erhard, F., Erddl u . Kohle 6, 196 (1963). 96. Wicke, E., 2. Elektrochem. 63, 279 (1949). 97. Eischens, R.P., Pliskin, W. A,, and Low, M. J. D.,J. Catalyaia 1,180(1962). 98. Holm, N. C . F., and Blue, R. W., Ind. Eng. Chem. 44,107 (1962). 99. Hindin, 5. G., and Weller, S. W., J. Phya. C h m . 60, 1601 (1966). 100. Barhani, H. N., and Clark, L. W.,J. Am. Chem. SOC.73,4638 (1961). 101. Hammett, L. P., “Physical Organic Chemistry.” McGraw-Hill, New York, 1940. 102. Gel’bshtein, A. I., Shcheglova, G. G., and Temkin, M. I., Zhur. Fiz. Khim. 30,2267 (1966). 103. Long, F. A., and Paul, M. A,, Chem. Reua. 67,935 (1967). 104. Ropp, G. A., Weinberger, A. J., Neville, O., J. Am. Chem. SOC.73,6673(1961). 105. Ropp, G. A., J. Am. Chem. Soc. 80,6091 (1968);82,842 (1960). 106. Biegeleisen, J., Hashemeyer, R. H., Wolfsberg, M., and Janknich, P. E., J . Am. Chem. SOC.84, 1813 (1962).

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Schierz, E., and Ward, H. F., J . Am. Chem. SOC.50,3240 (1928). Buiten, J., Dissertation, Delft, p. 89, 1960. Platonov, M.S . , and Tomilov, V. I., Zhur. Obshchey Khim. 8, 346 (1938). Kornienko, V. P., Ukrain. Khim. Zhur. 18,679 (1962). Bircumshaw, L. L., and Edwards, J., J. Chem.SOC.,p. 1800 (1960). Eischens, R. P., and Pliskin, W. A., i n “Actes du Deuxibme Congrbs Internationale de Catalyse, Paris, 1960,” p. 798.Edition Technip, Paris, 1961. 113. Rastrenenko, A. I., and Neumark, I. E., Zhur. Fiz. Khim. 31, 874 (1967). 114. Balandin, A. A., et al. Doklady Akad. NaukS.S.S.R. 118,312 (1968). 115. Balaceanu, J. C., and Jungers, J. C., Bull. aoc. chim. Belg. 80,476 (1961). 116. Freidlin, L. Kh., and Levit, A. M., Izveat. Akad. Nauk S.S.S.R., Otdel. Khim. Nauk. 626 (1961);Chem. Abstr. 48, 3840f (1962). 116a. Noller, H., Hantke, E., and Schwab, G.-M., 2 . physik. Chem. (Frankfurt) [N.S.] 38, 103 (1963). 117. Hirota, K.,Fueki, K., Nakai, Y., and Shindo, K., Bull. Chem. Sac. Japan 81, 780 (1968). 118. Hirota, K., Fueki, K., Shindo, K., and Nakai, Y., Bull. Chem. SOC.Japan 82, 1261 (1969). 119. Imelik, B.,Franpois-Rossetti, J., and Sigli, P., J. chim. phye. 58, 1048 (1969). 120. Whitmore, F. C., J. A m . Chem. SOC.54, 3274 (1932). 121. Day, A. R., “Electronic Mechanism of Organic Reactions,” Am. Book Co., New York, 1960. 122. de Boer, J. H., I d . chim. belge 21, 1169 (1966). 123. Andrianova, T. I., and Bruns, B. P.. Kinelika d Kataliz 1, 440 (1960);Kinetic8 Cataly& (U.S.S.R.),Eng. Tranal. 1, 410 (1960). 124. Rubinshtein, A. M., Doklady Akad. NaukS.S.S.R. 189. 626 (1961). 125. Leugering, H. J., “Heterogener Wmserstoffisotopen Austausch an Al,O, - Oberflachen mit verschiedener Hydroxylgruppenbelegung,” Thesis, Gottingen, 1948. 126. Marec, L. F., and Hahn, D. A., “The Catalytic Oxidations of Organic Compounds in the Vapour Phase,” p. 124.Chem. Catalog., New York, 1932. 127. Pichler, H., Brennatoffchem. 24, 27 (1943). 128. Kornienko, V. P., Ukrain. Khim. Zhur. 18,679;Chem. Abatr. 48,49460 (1964). 129. Keller, A., and Korosy, F., Nature 182, 680 (1948). 130. Beilstein “Handbuch der organischen Chemie,” Vol. 2, System No. 166. Springer, Berlin, 1960. 131. Krogmann, K., 2 . anorg. u . allgem. Chem. 808, 226 (1961). 132. Rienacker, G.,and Toursel, W., 2. anorg. u . allgem. Chem. 807,236 (1961). 133. Zapletal, V., Jedlicke, J., and Ruzicka, V., Collection Czechslov. Chem. Communs. 22, 171 (1967). 134. Ponec, V.,and Danes, V.,CollectionCzechslov. Chem. Communa. 25, 820 (1960). 135. Caujolle, F.,Compt. rend. a d . sci. 204, 446 (1939);Chem. Abstr. 83, 2836 (1939). 136. Bond, G. C., “Catalysis by Metals,” p. 418.Academic Press, New York, 1962. 137, Schuit, G. C. A., van Reijen, L. L., and Sachtler, W. M. H., i n “Actes du Deuxibme Congrbs Internationale de Catalyse, Paris, 1960,” p. 893. Edition Technip, Paris, 1961. 138. Garner, W. E., Stone, F. S., and Tiley, P. F., Proc. Roy. SOC. A211,472 (1962). 139. Courtois, M., and Teichner, S. J., J . Catalysis 1, 121 (1962). 140. Pewny, N. I., i n “Handbuch der Katalyse” (G.-M. Schwab, ed.). Vol. 6,Part 11, p. 341.Springer, Vienna, 1967.

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Application of Spectrophotometry t o t h e Study of Catalytic Systems H. P. LEFTIN

AND

M. C. HOBSON, JR.

Research and Development Department M . W . Kellogg Company, Jersey City, New Jersey

I. Introduction ........................................................ 11. Principles and Theory. ............................................... A. Optical Spectroscopy ............................................. B. Effects of Adsorption on Spectra. ................................... 111. Experimental ....................................................... A. Sample Preparation .............................................. B. Cell Design ...................................................... IV. Surface Structure Studies ............................................ A. Surface Functional Groups ........................................ V. Physisorption Studies. ............................................... A. Silica and Other Oxides.. B. Ionic Salts ...................................................... C. Cracking Catalysts ............................................... VI. Chemisorption Studies.. ............................................. A. SupportedMetals ................................................ B. Oxides and Salts ................................................. C. Cracking Catalysts ............................................... References .........................................................

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115 116 110 118 122 122 124 126 127 130 130 134 139 150 150 157 163 198

I. Introduction Optical absorption spectroscopy is one of the most versatile and informative techniques for investigation of surface chemistry, particularly as related to heterogeneous catalysis. In principle, a t least, it is possible to determine in detail the chemical functionality of a surface, the structure of an adsorbed species and their interactions and interrelationships. Such information, in addition to better defining the nature of surface active sites, is particularly valuable in elucidating the mechanisms of heterogeneous catalysis by identification of the chemisorbed reactive intermediates. Infrared spectra have been very fruitful in studies of relatively simple molecules, such as carbon monoxide chemisorbed on nickel where as many as five modes of chemisorption have been detected. Ultraviolet spectra, being due to electronic transitions, have been most valuable for the characterization of more complex intermediates, such 11u

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as carbonium ions and ion-radicals. The utility of spectroscopic studies of surfaces has recently been broadened by development of experimental techniques for examining a given specimen over both spectral regions, thereby affording information on both the covalent bonding and the electronic states of an adsorhent-adsorbate system. This review covers applications, both experimental and interpretive, of spectroscopy to problems concerning heterogeneous catalysis. Since two excellent reviews ( 1 , 2 )have appeared recently concerning spectral studies in the infrared region, this area is treated rather selectively. Major emphasis is on applications of visible and ultraviolet spectroscopy, since this appears to be the first review of the subject. Although spectroscopy is frequently incapable of providing complete descriptions of complex catalytic systems, it is rapidly providing much new information to extend our insight into the fundamental principles of oatalysis.

II. Principles and Theory A. OPTICALSPECTROSCOPY Optical spectroscopy may be defined as covering the ultraviolet, visible, and infrared regions of the electromagnetic spectrum. In the far infrared (15-25 p ) rotational transitions in the molecule occur. This region has not found wide application in catalysis studies. The infrared (2-15 p ) contains the combined vibration-rotation spectra. Fundamental vibrations, treated as simple harmonic oscillators as a first approximation, produce characteristic absorption bands in this region, and these bands can usually be resolved into a fine structure caused by the superposition of rotational states. The fine structure is sensitive to the type of group the vibrating atoms are in; i.e., the carbon-hydrogen bond vibration shows a distinct difference between methyl and methylene groups in its fine structure. Interpretation of these bands for adsorbed species depends on comparison of the frequencies and relative intensities of the fine structure peaks with the spectra of pure compounds to which assignments have been made. Shifts in the frequency of an identifiable group indicate a change in the electronic environment of the group. In the near infrared (0.8-2.0 p ) are found overtones of the fundamental vibrational frequencies. Weak interactions of a group with its environment can have a pronounced effect on the anharmonicity of group vibrations. Since the relative effects are larger for the overtones than for the fundamental frequencies, these bands are usually very sensitive to environmental factors. Also in cases where

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the fundamental frequencies may be obscured by other strong absorption bands the overtones may not. In the visible-ultraviolet region are found the electronic transitions that may occur in molecules. The chromophoric groups so well known in organic chemistry fall in this region; in fact, the commonly observed absorptions are caused by transitions of electrons not being used in the primary bonds of molecules. For example, pi-electrons in olefinic and aromatic compounds, electron-rich carbonyl and azo groups, carbonium ions, carbanions, free radicals, and many metal ions and organometallic complexes exhibit characteristic absorption bands in this region, While in principle, uv spectra should be quite complex with each band containing a great deal of fine structure from both vibrational and rotational transitions, this fine structure is seldom observed. Although several factors may contribute to this, loss due to collision broadening is probably most important in liquids and solids where the density is so high that interactions destroy the quantization effect. The vibrational and rotational levels overlap to form a continuum within the absorption band. Visible-ultraviolet absorption spectra show a great deal of group character, but are usually lacking in structural features. The combination of infrared and visible-ultraviolet studies of any system has certain obvious advantages. For example, if chemisorbed olefins on a supported metal catalyst are partially hydrogenated to form free radicals, the combined ir and uv can completely define the species. The ir is capable of determining the kind of groups present and the free electron will have a unique absorption band somewhere in the visibleultraviolet region. There are two measurable features of a n absorption band which can contribute information about the system being observed. The frequency a t the point of maximum absorption of a band or of several bands will often serve to identify a molecular species; identification being made by comparison with spectra of known compounds. Usually the identification is not conclusive without other physical or chemical knowledge of the system. Having identified the compound, quantitative information may be obtained by measuring the area under the absorption curve, or the maximum intensity, both of which are a function of the concentration of the absorbing species. This function is usually expressed as Beer’s law,

A = log-I, = CEL

I

(1)

where I , and I are the intensities of the incident and transmitted light, respectively, and log I,/I is called the optical density or absorbance.

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When concentration C , is expressed as moles per liter, E is the molar extinction coefficient. Although the extinction coefficient should be a constant by the above definition, we will see that it is often dependent on the environment.

B. EFFECTS OF ADSORPTIONON SPECTRA 1. Adsorption

Interaction of a molecule with a surface often produces a change in its optical properties. Strong interactions with the surface are generally defined as chemisorption, while interactions involving lesser energies are called physisorption. Definitions of chemisorption and physisorption based upon energies of interaction of the molecules with the surface, although used for many years to describe and distinguish between these types of adsorption, cannot serve as functional definitions for discussions concerning changes in optical absorption spectra. It becomes necessary to redefine chemisorption and physisorption. Physisorption, in terms of spectral changes, is “adsorption which leads to perturbation of the electronic or stereochemical states of the molecule but otherwise leaves the molecule and its entire electron complement intact.” Thus for example, a change in symmetry due to Van der Waals adsorption may cause the appearance or enhancement of a weak band which in the normal state of the molecule might correspond to a theoretically forbidden transition. Spectral changes due to hydrogen-bonding and those which can be attributed to the high polarity of the surface would also fall under our definition of physisorption. Physisorption forces may effect the spectrum of an adsorbed molecule either as a shift in the wavelength (surfatochromism) or as a change in the intensity of an absorption band. Occurrence of new* bands is not usually encountered in physisorption studies as these generally are associated with the formation of a new chemical moiety implying a chemisorption phenomenon. Chemisorption, then, is defined as “adsorption which produces a new chemical species by fragmentation of the molecule or of its electron complement.” Chemisorption will produce spectral bands which are not characteristic of either the adsorbate or of the adsorbent. Observation of new bands can always be taken as an indication of chemisorption; however, absence of such bands is not reliable evidence for the contrary since frequently spectral bands cannot be observed for other reasons.

* Observation of normally forbidden transitions due to symmetry changes which occur upon adsorption are not considered as “new” bands.

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2. Extinction Coeficients

Extinction coefficients of molecules are often significantly altered as a result of adsorption. This effect, aside from any observation of spectral shifts, is of considerable importance to interpretation of spectral data from a catalytic viewpoint since the occurrence of such changes makes it both dangerous and unreliable to attempt in any way to deduce the extent of surface coverage of a catalyst from the observed absorbance of bands due to the adsorbed molecules. These changes cannot be due simply to differences in polarity, since corresponding changes in solvent systems are not known. The change in extinction coefficient may perhaps be the result of peculiar orientations of the molecules and their electronic vectors with respect to the electronic vector of the polar surface. This type of optical anisotropy can lead either to an enhancement or decrease in absorption intensity, depending upon whether the molecule is adsorbed such that its electronic vector is parallel or perpendicular to the electrostatic field of the surface. Although very few quantitative data are available concerning the influence of the surface field on the absorption intensities of ultraviolet and visible bands, several authors have dis. direct and inverse cussed such effects in the infrared region ( 3 , 4 , 5 ) Both variation of extinction coefficient with surface coverage (8) have been observed in the infrared and the ultraviolet regions. The extinction coefficient for chemisorbed ethylene increased with surface coverage on copper oxide whereas it decreased on nickel oxide, indicating that the direction of the variation depends not only upon the adsorbate but also on the nature of the adsorbent ( 6 ) . When several absorption bands are associated with an adsorbed species, these may respond individually to surface coverage. For an aromatic compound adsorbed with its plane parallel to that of the surface the force field normal to the surface can enhance the intensity of in-plane vibrations while the out-of-plane vibrations will be decreased ( 7 ) ,resulting in a nonlinear variation in the relative intensities with surface coverage. Nonlinear variation of the relative intensities of the C =O and C-H frequencies of ketones adsorbed on montmorillonite (5) and of the N-H and C-H frequencies of amines adsorbed on porous glass (8) have been interpreted on the basis of assumed orientations of the adsorbates on the surface. The extinction coefficient of I, adsorbed on calcium fluoride decreased by two orders of magnitude with increasing coverage (9). Molecules adsorbed at low coverages are adsorbed on sites of highest adsorption energy and therefore show the greatest perturbation. Subsequent adsorption involves sites of lower energy and hence these are less perturbed. The resultant molar extinction coefficient, then, a t any coverage is the average over all of the individual molecular extinction

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coefficients. Exponential variation of extinction coefficient with coverage suggests that the extinction coefficient is a function of adsorbateadsorbent interaction similar to the well-known dependence of heats of adsorption on coverage. In principle, the variation of extinction coefficients could be made the basis for determination of a qualitative site energy distribution on the surface. A different site energy distribution could result for each spectral transition, however, depending upon how each transition “sees” the adsorbent. Assuming an exact knowledge of the nature of the electronic transition or vibration associated with each absorption band, and assuming that the effect of the electrostatic field of a surface on these transitions can be calculated reliably, it should be possible, in principle at least, to deduce the orientation of the molecule on the surface from observation of the change in extinction coefficient with coverage. Unfortunately, our knowledge and understanding of optical transition and of the effects of an electrostatic field on these transitions is not sufficiently developed to permit such detailed conclusions. In a practical sense the effects described in the foregoing discussion place a severe limitation on the applicability of spectral studies of adsorbed molecules to the detailed elucidation of the adsorption process and of the stereochemistry involved in surface catalysis. Since the absorption intensity may be either enhanced or decreased as a result of adsorption on a surface, and may either increase or decrease with variation in surface coverage, it becomes very difficult indeed to use spectral data as a measure of the surface concentration of adsorbed species. This is of particular importance when more than one species occupies the surface; e.g., physisorbed and chemisorbed species. In this case the absolute concentration of either species on the surface cannot be measured directly nor can it be reliably inferred from a comparison of the intensity of the bands corresponding to these two species. Moreover, in the identification of an adsorbed species the relative intensities of two or more bands characteristic of that species; e.g., the CH stretching and the CH deformation frequencies for adsorbed hydrocarbons, cannot be used as evidence for the structure of the adsorbed species since the absorption coefficients of the individual bands may change in opposite directions as a function of surface coverage. Thus the relative intensities of such bands cannot be compared to the relative intensities of the same bands observed in solution or in the gas phase. A similar difficulty arises when attempts are made to use the electronic spectra of adsorbed molecules to complement the infrared spectra for identification purposes.

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

121

3. Frequency Shifts Frequency shifts of an absorption band reflect changes in the relative levels of the two energy states of the transition, whether rotational, vibrational, or electronic. Hydrogen-bond formation and changes in the polarity of a, solvent may shift the frequency. Numerous examples and studies of this solvatochromic effect have been reported and severaI theoretical treatments have been made for solvent systems. Similar spectral shifts, caused by

2.

E

W

: Distance from surface

FIG. 1. Potential energy diagram for a bathochromic shift. hv (gas) = ad; hv (ads) = be (10).

both general surface polarity and specific hydrogen-bond formation with a surface, can be expected in the spectra of adsorbed molecules. Shifts may be either in the direction of longer or of shorter wavelength (bathochromic and hypsochromic, respectively), depending primarily upon the nature of the transition responsible for the individual absorption bands. The magnitude of the shift will depend also on the nature of the adsorbent. Both the direction and magnitudes of surfatochromic shifts have been explained (10)in terms of the energy of adsorption by means of potential energy diagrams such as shown in Figs. 1 and 2. For any optical transition in which a molecule in the ground state, A , is promoted to an excited state A’,

hv=AE. (2) Upon adsorption the energy of the transition is modified by the difference in the adsorption energy of the molecule in the ground state and in the excited state. Thus, for the adsorbed molecule Eq. (2) becomes h v = A E +&a -&I, (3)

H. P. LEFTIN A N D M. C. HOBSON, JR.

122

where Q A is equal to the adsorption energy for A and QL. is equal to the adsorption energy for the excited state, A‘. According to Eq. (3) the frequency of the absorption band will depend on the magnitudes There I,. can be envisioned three cases: (a) when of the energies Q A ,& Q A is less than QL8 there is a shift to longer wavelengths; in this case

x.

e

e

W

Distance from surface

FIG.2. Potential energy diagram for a hypsochromic shift. hv (gas) = ad; hv (ads) = be ( 1 0 ) .

the excited state is adsorbed more strongly than the ground state; (b) when &A = & there I. is no shift in the wavelength of the maximum as compared to that of the normal pure compound; and (c) when Q A is greater than Q;. there is a shift to shorter wavelength. Several examples of these types of surfatochromism are described under Section

V, B.

III. Experi mental With the exception of the optical cells and catalyst samples, all of the apparatus and techniques employed in these studies are conventional. Commercial spectrophotometers suffice “as is’’ or with minor alteration. While it is sometimes desirable to obtain high resolution spectra, it is more often the sample rather than the instrument that is the limiting factor in resolving the fine structure of bands. Manipulation of Ramples and materials is carried out in high vacuum systems using the wellknown techniques for handling catalysts and volatile compounds.

A. SAMPLE PREPARATION With the exception of the studies of Pickering and Eckstrom (11) on evaporated metal films in a multiple reflection cell, most of the work

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123

has been done on glass or glass-like materials, or on glass supported catalysts. These may either be in the form of a finely divided powder or of a massive but highly porous plate. 1. Powders

Since light scattering due to finely dispersed materials varies inversely as the fourth power of the wavelength, the difficulties involved in preparing suitably transparent samples for spectral studies increase considerably as we go from the infrared to the shorter wavelength visible-ultraviolet region. I n order to reduce such light scattering to a minimum the particle size of powdered catalysts must be smaller than the wavelength of the light being used. I n the infrared this means the particles should be one micron or less in diameter for most nonabsorbing materials. For metals, which absorb very strongly in the infrared, it is particularly important to have very small particle size. Although, on going towards the ultraviolet, absorption by the metal decreases, losses due to scattering increase so rapidly that it places a severe strain on the sensitivity limits of the spectrophotometer, hence no successful ultraviolet studies of metal catalysts have yet been reported. Attempts to overcome the scattering problem by embedding or suspending the catalyst powder in paraffin or other matrix of suitable refractive index have been made (12-15). Measurements performed on such samples can be only of limited value for catalysis studies for the following reasons: (a) interactions of the catalyst or of the adsorbate with the matrix material either contaminate or otherwise perturb the system, (b) difficulties and uncertainties in determining the extent of adsorption, and (c) inability to define precisely the state of the surface or t o implement variations by pretreatment procedures. Eischens and co-workers ( 1 ) pioneered the use of catalyst powders in the infrared region. Samples were prepared by spreading a paste or slurry of the powder on a calcium fluoride disk and evaporating the volatile liquid. Yang and Garland (16) claim a more uniform sample was obtained when the slurry was sprayed onto the salt plate with an atomizer. Best results were obtained using acetone as the volatile liquid and maintaining the temperature of the salt plate at 70°C. Another modification was used by Fahrenfort and Hazebroek ( 1 7 ) in which the powder was supported on a 200-mesh metal screen. They claim the advantages of a much wider temperature range for sample treatment and no wavelength limitations due to strong absorption by the catalyst. I n some eases powders have been compressed into wafers to overcome manipulation difficulties (18, 19, 20). Vacuum sublimation to obtain thin transparent films (21) has been used for alkaline earth halides.

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H. P. LEFTIN AND M. C. ROBSON, JR.

This technique, however, is not widely applicable to most materials of catalytic interest. Reflectance methods have been used to some extent (22). 2. Glasses Some of the first infrared studies of catalyst surfaces (23) were carried out on porous glass and a number of subsequent investigators have used it (Vycor)as an adsorbent or as a support for metals (6,24,25,26).Transparent massive samples of silica-gel and silica-alumina catalyst have been used by Leftin and Hall (27)and a transparent y-alumina prepared by Peri and Hannan (28).Preparation of these transparent samples is difficult in that the desired result seems to be attainable only through repeated trial and failure. These samples have the advantages of greater ease of manipulation and smaller scattering losses than the powders. Also, since a single sample may be used in studies involving several adsorbates simply by oxidative regeneration between runs, quantitative comparisons can be made. Impregnation of the samples to form supported metal catalysts are limited to less than 10 w t yometal. Otherwise, as pointed out by Eischens and Pliskin ( I ) the metal particle size tends to grow too large, causing excessive light absorption and scatter.

B. CELL DESIGN 1. Infrared Figure 3 is a sketch of an infrared cell used by Eischens and Pliskin

(I).Vertical mounting of the cell, required to prevent loss of the powdered samples, presents one of the disadvantages in using powders since

I

FIG. 3. Cell for catalyst spectre1 studies in the infrared region ( I ) .

SPECTROPHOTOMETRY FOR CATALYTIO SYSTEMS

125

it requires considerable modification in the physical arrangement of the spectrometer. For powdered catalysts which have been compressed into disks and for massive plates the cell need not be mounted vertically. Cells having a furnace wound on a side arm with the sample mounted in a holder that can be moved back and forth from the furnace to the section containing the salt plate windows have been described (19, 28). 2. Visible-Ultraviolet

A cell for the visible-ultraviolet region (29) is shown in Fig. 4. A pyrex extension containing the necessary outlets is added through a graded seal to a commercial fused silica absorption cell having a 1 cm ‘30s

WllO

K

Ouorh to pyrex Absorption cell

(b)

FIQ. 4. Cell for catalyst speotral studies in the visible-ultravioletregion.

path length. The sample is mounted with platinum wire on a rectangular quartz cage that is made to fit snugly into the absorption cell. Samples can be heat-treated in the cell by raising a vertical tube furnace around it. The only instrument modification required is a tall, light-tight cover on the cell compartment to accommodate the cell extension. 3. Combined Infrared and Visible-Ultraviolet By combining the previously described visible-ultraviolet cell backto-back with a gas-type infrared cell containing a recess to hold the rectangular sample mounting cage (Fig. 5) the entire spectral range may be examined on one sample (30))thus eliminating the hazards of not having all the variables identically controlled, as would be the case when different samples and cells are used for the different spectral regions. As pointed out in the discussion of the theoretical aspects of spectroscopy, a wide spectral coverage provides information on both group character and detailed structure of a molecule. This is particularly important in chemisorption studies where reactive intermediates are the species one wishes to detect.

126

H. P. LEFTIN AND M. C. HOBSON, JR.

One disadvantage common to all of the cells is the rather large dead space inherent in their design; this reduces the sensitivity of adsorption measurements, making difficult the measurement of spectral changes as a

NaCl windows--

Sample holder for !ip+ctral rn80surammtf

Graded seal

Sample frame

I c m square silica cell

-t

f

1

--t

FIG. 5 . Cell for catalyst spectral studies in both the infrared and visible-ultraviolet regions.

function of surface coverage. Without adsorption isotherms quantitative comparison between experiments is impossible and even qualitative comparisons are somewhat restricted. There is still room for much improvement in the design of cells.

IV. Surface Structure Studies I n discussing applications of optical spectroscopy to the study of catalytic systems it is convenient to treat separately the principal categories which can be examined; namely, the adsorbent and the adsorbate, the latter being further divided into the physisorbed and chemisorbed states. Since both the adsorbate and the adsorbent are affected in a spectroscopically observable manner during an adsorption process, and since it is difficult to distinguish, let alone separate, chemisorption from physisorption, some overlap between sections is unavoidable.

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

127

A. SURFACE FUNCTIONAL GROUPS 1. Hydration-Dehydration

Most spectroscopic studies of catalyst structure have been concerned with the effects of environment and temperature on surface hydroxyl , (28),titania (33),and some groups of silica ( 8 , 1 9 , 2 4 , 2 5 , 3 1 , 3 2 )alumina clay minerals (18, 34-38). Although several of the investigations have been quite extensive, the properties of the surface hydroxyl have not yet been completely clarified. All of the silica spectra exhibit certain common features; (a) hydroxyl groups on the surface exist in several different states that absorb a t distinct frequencies, (b) the population of these states is dependent on the temperature and gas-phase environment of the pretreatment, (c) hydration-dehydration phenomena are reversible, and (d) physisorption effects depend on hydrogen-bonding ability of the adsorbate. McDonald (31)made an infrared study of the silanol groups on Mallinckrodt Special Bulky Silicic Acid and on Cabosil, two amorphous silicas that represent opposite extremes of surface functionality. Adsorbed water and perturbed hydroxyl groups produced a broad, unresolved band centered around 3400 cm-1 (Fig. 6). On heating the 01

.

,

. , . , . . . . . . , . , . ,., . 3 1

I

I

Frequency in cm-'

FIG.6. Infrared spectrum of Cabosil.

sample in vacuo at 500°C the low-frequency side of the band (28003600 cm-1) decreased and the high-frequency side (3600-3750 cm-l) increased in intensity due to desorption of water that was hydrogenbonded to these surface hydroxyl groups. At 500°C adjacent silanol groups condense, eliminating the 3520 cm-1 band, but intensifying the bands at 3740 and 3750 cm-l. On raising the degassing temperature to 940°C the bands at 3660 and 3740 cm-l were eliminated, while the 3750 cm-1 band, due to free OH, remained. On re-exposure to water vapor

128

H. P. LEFTIN AND M. C. HOBSON, JR.

these effects were essentially reversed, showing that the hydroxyl bands being observed were indeed on the surface and that the isolated silanol groups were very stable, as high as one mole per cent remaining after treatment at 94O0c.Upon physisorption of dimethyl ether the bands a t 3740 and 3750 cm-1 disappeared completely and broad bands at 3320 cm-1 for Cabosil and 3275 cm-l for Bulky Silica appeared. These are the bands for the hydrogen-bonded ether-silanol complex, further demonstrating that the silanol groups were on the surface. The difference between the two silicas appears to be in the relative amounts of hydrogen-bonded hydroxyl groups on the surface. The Mallinckrodt Special Bulky Silicic Acid is more disordered and has more hydrogen-bonded groups than does the Cabosil. At 400°C methanol replaced about 30% of the hydroxyl groups to form methoxyl groups. Failure to displace all of the hydroxyl groups may have been due to small traces of residual water in the reagents. From quantitative analysis of the methoxyl groups by oxidation to CO,, McDonald estimated absorption coefficients as 3 x lo4 cmZ/mole for CH at 2857 and 2959 cm-1, 2 x lo4 cm2/mole for CH a t 2995 cm-l, and about lo6 cmZ/mole for OH a t 3750 cm-1. However, these absorption coefficients may be subject to large error. McDonald interpreted these results to mean; (a) only a few distinct types of silanol groups exist on the surface, and (b) adsorption takes place by hydrogen-bonding to silanol groups and not on siloxy groups. On the second point he was in disagreement with others who have studied the dehydration of silica. I n fact, he commented specifically on a paper by Nikitin, Sidorov, and Karyakin (25) in which a diametrical interpretation was proposed. Folman and Yates (8), in a study of the physisorption of ammonia and methyl chloride on porous glass (Vycor 7930), obtained the data needed to clear up this difference of opinion. They found that for ammonia a t 150°C a plot of the optical density (Fig. 7 ) of the absorption band assigned for hydrogen-bonded OH groups versus a function of surface coverage did not pass through zero. As Yates (39) has pointed out in a recent review, this can be best explained by the initial adsorption occurring on silicon or oxygen atoms in the surface, followed by adsorption on the isolated OH groups. Since isotherms were not obtained in previous studies, the different interpretations may be due entirely to differences in surface coverage. The hydration-dehydration behavior of the hydroxyl groups on alumina (28) and titania (33) is similar in many respects to that on silica. However, there are some important differences. Peri and Hannan (28) found three absorption bands instead of one for “isolated” OH groups after a drastic calcination treatment.

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

129

Adsorption studies demonstrated the independent nature of these three bands. FOTexample, when butene was exposed to a deuterated sample a t 200°C it was found that only the lowest frequency OD group exchanged hydrogen with the butene to form the corresponding lowfrequency OH group. The rate of exchange with deuterium gas to form an equilibrium mixture of deuteroxyl and hydroxyl groups was different for each group. From such evidence Peri and Hannan concluded that the three absorption bands represent three chemically distinct surface hydroxyl groups. .

Volume (cm’/gm) adsorbed

FIG. 7. Peak optical densities as a function of the volume of ammonia adsorbed. Curve (a):band due to perturbed OH groups; curve (b):band at 3370 cm-l due to NH stretch. Points obtained by adsorption 0, by desorption X. (8).

Titania also exhibited some of the characteristics that distinguished alumina from silica. Two distinct “isolated” hydroxyl groups were observed and a possible third type was suggested by the spectra of deuterated samples. The spectra were not as clean-cut as the other two and large differences appeared between the rutile and anatase forms. The relative ease of reduction of titania compared to alumina and silica may contribute to this, Yates (33)suggested that the different OH peaks may be associated with hydroxyls on the surface of different crystal faces, Since these studies have raised more questions than they have answered, much work remains to be done. 2. Adsorption Effects

The frequencies assigned to isolated hydroxyl groups on surfaces are known to be shifted even as a result of the adsorption of nonpolar gases. McDonald (32) attempted to find a correlation between some physical property of the adsorbates and the magnitude of these frequency shifts, He suggested that the polarizability of the adsorbate

130

H. P. LEFTIN AND M. C. HOBSON, JR.

was proportional to SvOH, but the shift caused by nitrogen was about twice as large as that caused by oxygen a t 77°K although there is only about a 10% difference in their polarizabilities. Frohnsdorff and Kington (40)have shown that it is possible to account for the differences between the shifts for nitrogen and oxygen by calculating the contribution of the interaction of the molecular quadrupole moment of nitrogen with the proton of the OH group. Recently, Basila ( 4 1 )in a further study of this question, found that a correlation existed between 8vOH and the ionization potential of the adsorbate, for a series of methyl benzenes and chloromethanes. This correlation was explained in terms of hydrogenbonding being treated as a “charge-transfer” complex. A nonbonding pair of electrons in the adsorbate interacts with the proton, producing an increase in the electron density in the 0 - H bond, and a shift to lower frequencies. The validity of this “charge-transfer’’ model should be reflected in changes in the electronic spectra of the adsorbates.

V. Physisorption Studies Investigations of the effects of physical adsorption on the optical spectra of adsorbates have been limited in number. However, a few observations have been reported on perturbations of molecules by surface electric fields. These involve changes in the symmetry and polarization,

A. SILICAAND OTHER OXIDES 1. Alkanes

I n their now classic study of the effect of surface forces on adsorbed molecules, Sheppard and Yates (26) found that some of the Ramanactive vibrations of methane, ethylene, and hydrogen appeared in the infrared when these materials were adsorbed on silica. The frequency shifts for the molecule on going from the gas phase to the adsorbed phase were all rather small, indicating that no chemical change in the species was brought about by the adsorption. The fact that some Raman-active frequencies became infraredactive on adsorption suggests that high electric fields on the silica surface induced polarization of the adsorbate with subsequent loss of symmetry of the molecule. This loss of sjrmmetry would permit previously forbidden vibrations to appear. They attempted to determine the rotational degrees of freedom of adsorbed methane. On adsorption, methane could have three possible

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

131

configurations. I n case I there would be no free rotation but only torsional oscillations of a rather tightly bound molecule. I n case I1 the methane could be pictured as having three of its hydrogens contacting the surface and rotation could take place only around the axis perpendicular to the surface, resulting in one degree of rotational freedom. I n case I11 all three degrees of rotational freedom persist and the molecule behaves as the usual spherical top. The contours of the absorption bands of each of these models were calculated and compared with the experimentally observed shape of the triply degenerates v 3 vibration (asymmetrical CH stretching). For completely free rotation (case 111) too large a fraction of the over-all intensity would be found in the wings to agree with the experimental curve. Distinction between case I and case I1 was ambiguous. Case I1 seemed to fit the experimental curve a little better on the high-frequency side; however, this may be due to overlap of another band. Because the v 3 vibration is triply degenerate there is another possible cause for the enhanced width of this band; namely, surface forces could cause a slight splitting of the degenerate energy levels. This question can be resolved by studying the spectrum over a wide temperature range. With decreasing temperature the band width due to rotational motion decreases as the q T , while it increases due to splitting of degenerate levels as the molecule becomes more firmly bound to the surface. Because of experimental problems, Sheppard, Mathieu, and Yates (42) studied the temperature effect on the spectrum of methyl bromide instead of methane. Comparison of the double degenerate v4 band of adsorbed methyl bromide with the calculated band for free rotation about the threefold axis at 294°K and 78°K showed a marked sharpening a t the lower temperature and no indication of the formation of a doublet. This ruled out a splitting of degeneracy as a cause of broadening of the band width. Whether the broadening at room temperature was due to a low energy barrier to free rotation or an enhancement of the extinction coefficient caused by interaction with its environment, the fact remained that extrapolation of these results to the methane system permitted the observed band shape for completely free rotation (case 111)of the methane molecule to be reconciled with the calculated contour. As the authors point out, this reopens the question of the rotational degrees of freedom of adsorbed methane. 2. Aromatics

Compounds which may act as proton acceptors in hydrogen-bond formation are strongly adsorbed on silicic acid. The unhindered hydroxyl group on the surface is believed to act as the proton donor in forming a

132

H. P. LEFTIN AND M. C. HOBSON, JR.

hydrogen-bond with some center of high electron density in the adsorbed molecule. Such an acid-base interaction between the “acidic” proton and the pi-electron system has been postulated as being responsible for the selectivity of silica gels toward aromatic systems ( 4 3 ) . The 260 mp transition in benzene provides a means of testing this hypothesis. Due to the sixfold symmetry of the pi-electron system, this transition is theoretically forbidden; however, since the vibrations of the carbon skeleton introduce a degree of asymmetry, the band is observed. If the silicic acid proton were to interact with the benzene pi-electrons at a point not on the sixfold axis it would distort the system, lowering its symmetry, and an enhanced absorption would result. The spectrum of benzene adsorbed on silicic acid did not show an enhancement in the integrated intensity of the 260 mp band (13), indicating that such interactions were, a t best, rather weak. For a series of monosubstituted benzenes adsorbed on silicic acid from cyclohexane solution absorption bands were broadened and shifted to somewhat shorter wavelengths (hypsochromic shift) (13).These shifts amounted to about 1% of the total energy of the transition and are similar to the shift observed for the spectrum of benzene adsorbed on transparent silica gel from the vapor phase ( 4 4 ) .Broadening of spectral bands on adsorption can be interpreted as being due to the statistical distribution of the energy states of the molecule in the field of a polar adsorbent. The spectral shifts may result from hydrogen-bond formation or change in polarity of the medium. Thus, the electronic spectrum of aniline in water as compared to cyclohexane shows a shift to shorter wavelengths by the order of 1000 cm-l (45).This can be interpreted on the basis of hydrogen-bonding with the water acting as a proton donor to the nonbonding electrons on the -NH, group. The hydrogen-bond reduces the extent of interaction between the two Pz electrons and the ring and thus shifts the absorption to shorter wavelengths. Increasing the acidity of the medium by adsorbing the aniline on silicic acid results in a further shift to shorter wavelengths due to the increased protondonating ability of the silicic acid relative to that of water. The extreme case for aniline would be that in an acidic solution, where the principle light-absorbing species is the anilinium ion rather than the aniline molecule. In this case the previously nonbonding electrons are intimately bound in an N-H bond and are thus unavailable for interaction with the ring. The spectrum in this case reverts to that of benzene. From the selected data shown in Table I it can be seen that the extinction coefficients either increase or decrease upon adsorption on silicic acid. These changes cannot be due entirely to the difference in polarity between the cyclohexane and the surface of the silicic acid

133

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

since corresponding changes in solvent systems are not known ; however, they may be the result of peculiar orientations of the molecules on the surface. TABLE I Comparison of Spectra of Compounds i n Cyclohexane Solution and Adsorbed on Silicic Acid ( 1 3 )

Compound

A (mf4 in cyclohexane solution

A

bl-4

adsorbed on silicic acid

E E

(cyclohexane) (silicic acid)

Benzene

261 255 249 243

259 254 248 243

1.7 1.8 2.2 3.5

Chlorobenzene

272 265

271 264

1.3 1.2

N,N-Dimethylaniline

251

237

4.7

4-Nitromiline

323

374-382

3.6

Pyridine

276 261 256 251

261 256 251

0.71 0.72 0.88

280 238

280 234

1 .o 1.1

Thiophenol

-

-

3. Ammonia Variation in the extinction coefficient with surface coverage has been observed by Folman and Yates (8) for the N-H stretching frequencies of ammonia adsorbed on porous glass. A plot (Curve b, Fig. 7 ) of optical density of the N-H stretching band at 3370 cm-l versus surface coverage exhibited a decreasing gradient as surface coverage increased, and on desorption remains well above the adsorption curve. Folman and Yates (8)offeredan explanation of this effect in terms of surface heterogeneity in which the most active sites are silicon or oxygen atoms in the surface and the OH sites have the lower energy of adsorption. The extinction coefficient is high a t low coverages, and decreases as more ammonia

134

H. P. LEFTIN AND M. C. HOBSON, JR.

becomes bound to the OH sites a t the higher coverages. On desorption the ammonia is preferentially removed from the OH sites and the larger extinction coefficient for ammonia on the silicon and oxygen sites causes the hysteresis on the optical density plot.

B. IONIC SALTS DeBoer and co-workers examined the electronic spectra of several compounds adsorbed on vacuum sublimed films of alkaline earth halides. Although this work has been briefly reviewed in Advances in Catalysis ( 4 6 ) ) since it constitutes an important contribution to the subject being considered here it is desirable (even a t the expense of some duplication) to present a detailed description of these studies. The experimental procedures will not be described in detail since they are not widely applicable to many solids of catalytic interest. 1. p-Nitrophenol Para-nitrophenol is strongly and reversibly adsorbed on calcium fluoride. This colorless compound becomes yellow when adsorbed as the result of a bathochromic shift of its spectrum. Decolorization of the surface by exposure to water vapor is reversible upon evacuation, Adsorption occurs through a n ion-dipole interaction between the surface fluoride ion and the dipole of the hydroxyl group. The electrostatic field of the surface interacting with the hydroxyl hydrogen affects the spectrum in the same direction as would be observed for removal of this hydrogen by ionization of the molecule. Variation of this spectral shift with surface coverage was examined in detail ( 4 7 ) . Figure 8 shows the spectra a t various time intervals after exposure of the film to a source of p-nitrophenol maintained at 0°C. At the lowest coverage at which optical measurements could be made the spectrum showed a maximum at 3700 A whereas the pure compound has its maximum absorbance at 3160 A. Equilibrium was attained after 1200 hr (curve 16). Adsorption was then continued at a higher pressure by raising the temperature of the reagent reservoir to 21°C. During the entire course of this adsorption the wavelength of maximum absorbance shifted toward that of pure p-nitrophenol until a second equilibrium was reached. Coverage at the first equilibrium corresponded to 57% of a monolayer, while at the second equilibrium 95% of monolayer coverage was achieved. Each of the observed spectra could be separated into curves caused by two types of adsorption; type a corresponding to the optically perturbed species, and type /I whose spectrum is unperturbed. The result of such a separation into component curves is shown in Fig. 9. Upon

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

135

Wavelength (mp)

FIQ.8. Spectrum of p-nitrophenol adsorbed on CaF, at various time intervals after introduction of the adsorbate ( 4 7 ) .

Wavelength (mpl

FIG. 9. p-Nitrophenol adsorbed on CaF,. Separation of the spectrum into its component curves. A-experimental curve; B-spectrum of type p; C-spectrum of type 01

(47).

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H. P. LEFTIN AND M. C. HOBSON, JR.

further analysis of these data it was shown that a t the first equilibrium 63% of the molecules were absorbed as type 01 and 37% as type p, while a t the second equilibrium they were about equal. Moreover the extinction coefficient of the electrostatically perturbed molecule was about 4 as great as that for the pure compound. 2. Iodine Iodine adsorbed on vacuum sublimed calcium fluoride is brown and its spectrum is shifted toward shorter wavelengths compared to that of ordinary solid iodine (9). On barium fluoride the surface color is more yellow, indicating that the spectrum is even further shifted towards the blue. The visually observed color is due to the combination of the absorption bands for unperturbed iodine and for an electrostatically adsorbed surface species. The spectrum of iodine on calcium fluoride at low coverage is sharp and consists (Fig. 10) of two bands at 3600 A

Wavelength (mp)

FIG. 10. Spectrum of I, adsorbed on CaF',at various degrees of surface coverage (9).

and 2800 A. Since potassium iodine solution shows a similar spectrum, de Boer concluded that I- in solution and F- on a surface have similar effects on the spectrum of iodine. DeBoer ascribed the spectral shift t o polarization of the iodine molecule by the electrostatic field of the negative fluoride ions on the external surface. The close similarity to

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

137

the spectrum of the triiodide ion suggests, however, that perhaps this is more than electrostatic polarization; indeed, chemisorption by formation of a complex trihalide ion might well be consistent with these observations. 3. Alizarin A chemical reaction occurs when alizarin ( 1 ,%dihydroxyanthraquinone) is adsorbed on an alkaline earth halide since, when heated, hydrogen halide is desorbed and the spectrum of the adsorbed alizarin is shifted toward longer wavelengths (48). The electronic spectrum of alizarin adsorbed on barium chloride changes systematically with heating. At 100°C the compound is essentially in the physically adsorbed state as evidenced by the yellow color and by the similarity of its spectrum with that of the pure alizarin. A shift towards longer wavelengths, which is first observed after heating a t 3OO"C, continues with extended heating until the band reaches 4850 A (Fig. l l ) ,the spectrum of alizarin 1.5

P

I .o

if

0.5

400

500 Waveknqth

600

700

(my)

FIG 11. Spectrum of alizarin adsorbed on BaCl,. Curves a and b, in the physisorbed state; curve c, mixture of physisorbed and chemisorbed; curve d, chemisorbed state (48).

in the adsorbed state. Similar behavior is also observed with strontium fluoride and barium fluoride. Here it is interesting to note that the magnitude of the bathochromic shift is dependent also upon the chemical nature of the alkaline earth halide employed. The greatest shift is observed for barium fluoride and lesser shifts for strontium and calcium fluorides. I n the heating process, which leads to the formation of the absorption bands a t longer wavelengths, hydrogen fluoride or hydrogen chloride is evolved and the alizarin is not desorbed. Thus the alizarin on the surface, after heating, is not in a physically adsorbed state but rather is tightly

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H. P. LEFTIN AND M. C. HOBSON, JR.

bound to the surface as the result of a chemical reaction. Further evidence that this is not physisorption is found in the fact that when the calcium fluoride-alizarin system is extracted with anhydrous amyl alcohol no alizarin is obtained in solution. The surface adsorbed species furthermore is not sensitive to alkaline solutions. O n treatment with dilute acid, however, the red-violet color of the adsorbed alizarin is rapidly changed to the yellow color of pure alizarin. Displacement of the alizarin from the surface is shown by the fact that following the acid treatment the adsorbed substance can be removed with amyl alcohol to give a solution whose spectrum is that of pure alizarin. Alizarin which has been heated to an intermediate temperature (curves b and c) responds to treatment with these reagents in a manner which would be expected for a surface containing a mixture of adsorbed species. Washing with alcohol effects a partial removal and is accompanied by a color change corresponding to that obtained by heating to the higher temperatures. Clearly then, a t these intermediate temperatures the surface is occupied both by physisorbed and chemisorbed alizarin. The spectrum under these conditions must be a composite of bands due to both types of adsorbed species. As before, the strongly adsorbed alizarin may be desorbed after a n acid treatment. DeBoer concluded that a reaction occurs at the surface on high-temperature treatment of adsorbed alizarin. I n the case of physisorption, the alizarin molecule is adsorbed on the surface (Structure A) by orientation of the peripheral dipoles of the hydroxyl groups toward the negative halide ion surface layer. On heating, HI? is released and a calcium salt of alizarin is formed on the surface (Structure B). This is further substantiated by the fact that

A0,

*lizarin

I

0

B

ca++

ca++

F-

F-

i

7 F-

b-

Cil++

ca++

F-

F-

(4

(piiGZ-&

./ 0

\

\ 0

(B)

the color of alizarin on barium fluoride, strontium fluoride, and calcium fluoride is qualitatively similar to that exhibited by an alkaline solution of alizarin. Variations in the wavelength of the spectrum of chemisorbed alizarin with the ionic radius of the alkaline earth metal; namely, 4820 A, 4950 A, and 5200 A for the fluorides of calcium, strontium, and barium,

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

139

respectively, is an indication of strong polarization of the alizarin ions by the surface cations. The reaction of alizarin with these fluorides was shown to be limited to the surface layer and each alizarin molecule was shown to take the place of two surface halide ions. This chemical reaction of alizarin with the surface of alkaline earth halides would constitute chemisorption according to our definition.

C. CRACKINGCATALYSTS 1. Visual Color Changes

Visual color changes associated with adsorption play important roles in many commercially and scientifically valuable processes such as chromotography, dying of fibers and photography. Some theoretical conclusions concerning the nature of the surface of the adsorbent have also been drawn from such changes. Qualitative observations of color changes upon the adsorption of a wide variety of neutral, nonpolar, organic compounds and polar dye molecules on the surfaces of silica gel and activated alumina have been reported (49). Among the compounds examined were several derivatives of triphenylchloromethane. Adsorption on an alumina surface caused these molecules to assume colors similar to those associated with their ionization products in strong acid solutions. These results were attributed to ionization of the nonpolar compounds under the influence of the polar surface. Assuming this interpretation to be correct, these observations constitute the first qualitative demonstration of the formation of carbonium ions on the surface of an activated alumina and, as such, may be taken a8 evidence for the acidic nature of the surface of the particular alumina preparation employed. Although ionization may have been involved in some of the observed color changes, it is evident from the results of deBoer that conclusions based solely on visual observations must be viewed with reservation. Of greater interest to catalysis is the application of visual color changes for determination of the acid nature of catalytic surfaces. The inherent acid-like properties of the surface of many solids have long been recognized as being intimately connected with catalytic activity, ion exchange, and adsorption phenomena. Although the chemical effects of acidic sites have been extensively explored, and in fact exploited profitably in several commercial processes, the nature of this acidity is not fully understood. Two measurable quantities may be considered; namely, the surface density of acid sites, and their acid strength. Of

140

H. P. LEFTIN AND M. C. HOBSON, JR.

these, the former has received by far the most attention in the past. Since the literature concerned with various methods of estimation of surface density of acid sites is much too extensive to be cited here, the reader is referred to two excellent reviews on this subject (50, 51). Several major contributions to the estimation of the latter quantity are based on Walling’s (52) extension of the Hammett (53) acidity function, H,, to solid acids. The Hammett acidity function, H,, is a measure of the ability of various media to donate a proton to a neutral base, and is defined as

where aH+ is the activity of the hydrogen ion and the f’s are the activity coefficients for the neutral base and its conjugate acid. Since the theories of acidity functions have been the subject of several excellent reviews (54), further discussion concerning the basis and the justification for Eq. (4)need not be considered a t this time. According to Walling’s extension of the Hammett H , method, the ability of the surface to convert an adsorbed neutral base to its conjugate acid is taken as a measure of the acid strength of the surface. Walling employed five indicator dyes covering apK,rangefrom $4.0 to +0.43. These dyes were adsorbed onto separate samples of catalysts or other solid acid from an isooctane solution. Conversion of the adsorbed dye to the acid form was then visually observed and a relative acid strength was assigned. One of these dyes has been used as an endpoint indicator for the butylamine titration of cracking catalysts suspended in a nonaqueous medium (51, 55), and more recently Benesi (56) has extended this work by employing a series of dyes. By conducting butylamine titrations to the endpoints characteristic of a series of Hammett dyes the apparent distributions of surface acid strengths have been determined for several synthetic cracking catalysts and clay minerals (57, 58). A recent modification permits the method to be used with highly colored solids (59). Although the results of these studies seem to correlate qualitatively with other properties of the surfaces in a n expected manner, quantitative and theoretical conclusions should be viewed with some reservation since the method depends entirely upon visual observation of a qualitative change in the color of an adsorbed molecule. Since both the acid- and base-forms of a dye will be adsorbed simultaneously, visual observation of color changes restricts the method to the use of only those dyes which are much more highly colored in the acid-form. Thus the number of useful indicator dyes is severely limited and the method provides acidstrength values lying between rather broad intervals. Of greater signifi-

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

141

cance is the fact that one cannot be certain that the color changes are actually caused by proton transfer from the surface to the adsorbed dye. As was indicated from the work of deBoer (Section V-B), color changes may also result from spectral shifts of the neutral-form of the dye under the influence of a highly polar surface without a proton being transferred. In his original work Benesi (56)found no evidence to indicate that the observed color reactions were due to anything other than a proton transfer, but similar color changes produced by Lewis acid are known (60).Moreover, since the color of the product of reaction of a basic dye with a proton should not differ greatly from that of its reaction with an electron-deficient center, or Lewis acid, it should be pointed out that no proof is as yet available to show that the titration measures only the number of protons on the surface of the catalysts. Color changes similar to those produced by reactions with a proton are also known to be caused by other agents; e.g., aminoazo-type dyes give acid-like colors in the presence of strong alkali (61).It may reasonably be expected that chemisorbed oxygen and transition metal cations could also modify the color of adsorbed dyes. All that can be said with certainty is that the titration measures the amount of base required to perturb the spectrum of an adsorbed dye molecule. 2. Spectra of Adsorbed Hammett Dyes In view of the uncertainties inherent in Hammett indicator determinations of surface acidity by visual means, a study was made of the spectral behavior of dyes adsorbed on several silica-alumina catalysts and silica gel (62).The effects of catalyst water content, dye concentration, catalyst composition and pretreatment on the spectra of the adsorbed dyes were examined. The Hammett indicator dyes employed and their corresponding p K , values are summarized in Table 11. Reference spectra were determined for the base-form of the dye in iso-octane or methylene chloride solution and for the acid-form in an aqueous sulfuric acid or ethanolic-hydrogen chloride solution. Dyes were adsorbed from isooctane solutions onto thin plates of optically transparent catalysts which were installed in evacuated cells of design similar to that shown in Fig. 4. The catalysts samples were routinely pretreated by calcination in oxygen at 500' to remove any organic contaminants, followed by evacuation at this temperature. To examine the effect of variable water content on the spectra the samples were rehydrated in an atmosphere of water vapor for 24 hr after pretreatment and subsequently evacuated at some lower temperature. Dye solutions were introduced through a side arm. These solutions were suitably dilute so that the absorbance due to dissolved dye was either below the limits of detection or, at

142

H. P . LEFTIN AND M. C. HOBSON, JR.

worst, was quite low and when further diluted by virtue of substantially complete adsorption on the catalyst sample the absorbance of the liquid phase did not differ from that of the pure solvent. Catalysts employed were prepared by hydrolysis of ethyl orthosilicate and aluminum TABLE I1 Hammett Indicator Dyes Dye

PK.4

+ 4.0 + 2.8 + 1.6 + 1.1

Phenylazonaphthylamine Aminoazobenzene Benzeneazodiphenylamine p-Nitroaniline o-Nitroaniline p-Nitrodiphenylamine p-Nitroazobenzene 2,4-Dinitroaniline Benzalacetophenone p-Benzoyldiphenyl Anthraquinone 2,4,6-Trinitroaniline

- 0.2 - 2.4 - 3.3 - 4.4 - 6.6 - 6.2

- 8.1 - 9.3

isopropoxide. Physical properties and compositions are summarized in Table 111. Spectra were scanned periodically over a 24-hr period, after which the intensity of the absorption showed little significant change, indicating that either an equilibrium had been reached or all TABLE I11 Catalysts Used in Spectral Studies

Code DGS-4 SA-11 DSA-1

Description Silica gel Silica-alumina (0.1% Also,) Silica-alumina (12% A1,0,)

B.E.T. surface area (meterllgm) 660

525 280

of the available dye molecules had been adsorbed, The quantity of dye in each case was such that a maximum of 1O1O t o 10l2dye molecules/cm2 of surface could be adsorbed on a sample having a total surface area of 5 square meters. I n general, three situations were obtained corresponding to spectra of dyes which were: (a) converted to their acid-form by

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

143

the surface, (b) adsorbed in the neutral or basic form, and ( c ) both the acid-form and the base-form coexisted on the surface. Figure 12 shows the typical behavior of the absorption intensity a t the wave length of the principal absorption band as a function of time.

FIG. 12. Peak absorbance as a function of time. Benzeneazodiphenylamineadsorbed on silica-alumina DSA-1 from isooctane solution at 26'C (62).

a. Acid-Form. Those indicator dyes whosepK, values are equal to or greater than the H , characteristic of an acid surface will be converted to the acid form when adsorbed. However, if the number of available dye molecules exceeds the number of strong acid sites, the base-form of

00

FIG. 13. Spectrum of phenylazonaphthylamine. Curve A, in isooctane solution; curve B, in ethanolic HCI; curve C, adsorbed on DSA-1; curve D, adsorbed on SA-11 ( 6 2 ) .

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H. P. LEFTIN AND M. C. HOBSON, JR.

the dye will also be adsorbed. I n the examples cited here the concentration of the available dye was maintained sufficiently low to preclude physisorption of the latter type. Figure 13 illustrates the absorption spectra for the adsorption of phenylazonaphthylamine ( p K , = 4.0) on 12% alumina, and 0.1 yoalumina-silicate catalysts, DSA-1, and SA-11, respectively. Curves A and B are the spectra of the base- and acid-forms of this dye as measured in iso-octane solution and in acid aolution, respectively. Clearly, on both catalysts this dye is adsorbed entirely in its acid-form. The bandwidths of the adsorbed species are somewhat broader than those found in solution, as is to be expected. The wavelength of maximum absorbance appears to be shifted to longer wavelengths as compared to that for the acid-form in solution. Figure 14

+

WaveIonpth (A)

FIG. 14. Spectrum of benzeneazodiphenylamine. Curves A, in isooctane solution; curve B, in ethanol solution; curve C, in ethanolic HCl; curve D, adsorbed on DSA-1; curve E, adsorbed on SA-11 (62).

shows similar spectra obtained with benzeneazodiphenylamine ( p K , = +1.5). Here again the dye is adsorbed on the 12% alumina-silicate entirely in the acid-form and without significant shift of the wavelength of the principal band, as compared to that in an acid solution. These data are consistent with the transfer of a surface proton (Bronsted acid) resulting in the formation of the conjugate acids of the dyes which are then held to the surface by electrostatic attraction. This is obviously a case of chemisorption. Another possibility exists which cannot be excluded at this time; namely, that of electron-sharing between a nonbonding electron pair of the dye and an electron-deficient center (Lewis acid) on the surface. It is not to be expected, on the basis of the spectra

SPECTROPIIOTOMETRY FOR CATALYTIC SYSTEMS

145

of the dyes used in this study, that any very striking spectral differences would be found between the reaction products of the dyes with a Bronsted or a Lewis acid. In this respect it may be significant that a 12 m p shift observed in the spectrum of the conjugate acid of 4-aminoazobenzene, formed by reaction with stannic chIoride or with ferric chloride (63), may be taken as an indication that a small but finite spectral difference exists between the conjugate acids of Hammett indicator dyes formed from Bronsted acids and from Lewis acids. On this basis, the 12 mp shift in the spectrum of chemisorbed phenylazonaphthylamine may serve as an indication that the strong acid sites on silicaalumina catalysts are Lewis rather than Bronsted acids. However, in this case, the shift may be due to some less specific effect of the polar surface. Spectra of benzeneazodiphenylamine adsorbed on DSA-I were determined over a range of surface concentrations between 0 and 1 x 1011 dye molecules/cma by exposing the same catalyst specimen to several concentrations of the dye. Figure 15 shows adherence to Beer's

1.5

? 4

-

/ d, ,

1.0

9

0.5-

5r

I/

-

B

0

3

6

9

12

15

FIO.16. Beer's Law plot. Benzeneazodiphenylamine adsorbed on silica-alumina, DSA-I (62).

law over this concentration range, indicating that at these low surface coverages the molecular extinction coefficient is a constant. The calculated molar extinction coefficient for the chemisorbed dye is equal (within experimental error) to that reported for the conjugate acid of this dye. b. Acid- and Base-Forms. Whereas benzeneazodiphenylamine adsorbed on DSA-1 exhibits only the band characteristic of the conjugate acid (Fig. 14, curve D), when this dye is adsorbed on a less acidic

146

H. P. LEFTIN AND M. C. HOBSON, JR.

catalyst (SA-11) a t approximately the same surface coverage the spectrum indicates the coexistence of both the base-form (physisorption) and the acid-form (chemisorption). I n this case the chemisorbed form appears as a shoulder in the 5400 A region while the physisorbed form is evidenced by a very weak shoulder in the 4000 A region. The apparent absorbance maximum shown in curve E a t 4500 A results from overlap of the bands due to the physisorbed and chemisorbed forms. I n principle at least, separation of these bands into the individual component peaks, ooupled with a knowledge of the molar extinction coefficients, should

FIG. 16. Spectrum of o-nitroaniline. Curve A, in isooctane solution; curve B, in water solution; curve C, in 60 wt yo H,SO,; curve D, adsorbed on silica-alumina DSA-1 (62).

permit an evaluation of the indicator ratio required in Eq. (4), and thus afford an exact value of the H , of the surface. While the present data provide an H , value of 1 . 3 for this particular catalyst, it should be rioted that such treatment of the data cannot be justified since the indicator ratio will be strongly dependent upon surface coverage. The only quantitative information which can be obtained from these data is a measure of the amount of dye adsorbed in the form of its conjugate acid, which would be a measure of the total number of acid sites of sufficient strength to convert the dye. I n the present system this amounts to about 0.2 millimole/gm or about 1 x 1 O I 2 acid sites/cm2.It is apparent that the absorption spectrum of the conjugate acid will not obey Beer’s law except a t very low surface coverages where the absorbance at this wavelength is not increased by contributions due to the base-form of the dye. c. Base-Form. Physisorption of the base-form was found in all cases where the surface acidity was too low for chemisorption to occur, or

+

SPECTROPHOTODIETRY FOR CATALYTIC SYSTEMS

147

where the surface coverage exceeded the number of acid sites available. Figures 16 and 17 illustrate two typical examples of the spectra of physisorbed dyes. The most interesting features of these data are the rather large spectral shifts between the dye in solution and in the adsorbed state. This shift amounts to 480 A for the 3720 A band in o-nitroaniline and 960 A for the 3420 A band in p-nitrodiphenylamine as

08

x)

Wavelength (A)

FIG.17. Spectrum of p-nitrodiphenylamine.Curve A, in isooctane solution; curve B, in 80 wt yo H,SO,; curve C, adsorbed on silica-aluminaDSA-1; curve D, adsorbed on silica-alumina SA-11 (62).

compared to the corresponding spectra in isooctane solution. These special shifts are very reminiscent of the observations of deBoer and co-workers mentioned earlier. I n their studies the spectral shifts were often in the direction such that the spectrum of the adsorbed molecule lay intermediate between its spectrum in solution and that of an ionized form of the species so that it would appear as if the adsorbed species were strongly polarized by the surface but not completely ionized. I n the present studies on silica-alumina catalysts it is clear that the polarization by the surface does not result in a species which is partially converted to the acid-form since the wavelength of the physically adsorbed dye shifts in a direction opposite to that exhibited by the conjugate acid. It was found that the magnitude of the shift depends, at least in part, on surface properties. With a given dye; e.g., p-nitrodiphenylamine (Fig. 17), spectral shifts of 960 and 600 A are found for this dye adsorbed on DSA-1 and SA-11, respectively. It should be pointed out that when an absorption band of a base dye borders on the visible, the shift which accompanies physisorption of the dye may be taken falsely

148

H. P. LEFTIN AND M. C. HOBSON, JR.

as an acid reaction or may produce a “wrong color” if this dye is used in the Walling or Benesi procedure. For o-nitroaniline adsorbed on silica-alumina, DSA-1 (Fig. 18), the absorption band shifts to shorter wavelengths with increasing surface coverage over the range of 3 x 10” to 3 x 10l2 molecules/cm.2 Apparently, as the coverage increases, the average perturbation due to the

x)

WnnlrnQth (A)

FIQ. 18. Spectrum of o-nitroaniline on silica-alumina, DSA-I, at several surface coverages (62).

surface field decreases. Adherence to Beer’s law a t the higher coverages (1 - 3 x 1012) is evidenced by the straight-line plot (Fig. 19) which extrapolates through the origin. At the lower coverages, however, a negative deviation from Beer’s Law is observed. While this may be explained on the basis of a variation of the molecular extinction coefficient with coverage, an alternate explanation may be dbtained from a consideration of the characteristic spectrum of this dye. The conjugate acid of o-nitroaniline has its principal absorption band in the ultraviolet at 2740 A, while the base-form exhibits maxima a t 3700 A and 2680 A. Direct measurement of the acid-form of this dye cannot be achieved due to interference from the 2680 A band of the base dye, The band in the 4150-4350 A region for the adsorbed dye corresponds to the 3700 A band of the base dye shifted toward longer wavelength by

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

149

the field of the polar surface. Therefore, any fraction of the dye which is converted to its conjugate acid by interaction with a surface acid site will not contribute to the measured absorption intensity in the 4150-4350 A region. Once all of the surface acid sites have become occupied, subsequent adsorption will be physisorption of the base-form of the dye;

Surface coverage (molecules/cmf) x 10-l'

FIG.19. Beer's Law plot. o-Nitroaniline adsorbed on silica-alumina, DSA-1 (62).

as coverage increases, the relative contribution of the acid-form to the observed absorption intensity will decrease until its percentage contribution becomes lost within the experimental error of the measurements. It is seen that linear extrapolation of the lower portion of the data intersects the concentration axis a t about 2 x 1011/cm2of surface. Accordingly, then, this would correspond to the total number of acid sites on the catalyst which are capable of transforming o-nitroaniline to its conjugate acid. If the spectral shifts observed for physisorbed Hammett dyes results from the fact that the excited states of these molecules are more strongly adsorbed than are the ground states, it is reasonable to expect that the difference between the energies of adsorption of these two levels can account for most of the decrease in energy separation between them. Hence, the energy corresponding to the magnitude of the observed spectral shift will be proportional to the force field of the solid and the dipole moment of the adsorbate. With a given adsorbate, then, it should be possible to differentiate between the field strengths of various solids. That this is reasonable can be seen from the fact that, in all cases, larger spectral shifts are found with a 12% alumina than with a 0.1yoalumina catalyst, reflecting the greater polarity of the surface of the more active

150

H. P . LEFTIN AND M. C. HOBSON, JR.

catalyst. A description of this effect in terms of potential energy diagrams has been given in Section 11,B, 3. Since the proper values of the dipole moments of these dyes are not known, it is not possible to calculate the field strengths of these surfaces; however, order of magnitude agreement with the results of Zettlemoyer (64) are obtained assuming a dipole moment of 1.0 0.

VI. Chemisorption Studies The most extensive portion of the literature dealing with the application of absorption spectroscopy to catalysis is concerned with chemisorption. Changes in molecules produced by chemisorption are deduced from differences between the frequencies and intensities of the spectral bands of the adsorbate and those of the “free” molecule. The structure of the adsorbed species can be reliably ascertained only by direct comparison with spectra of its exact counterparts. When, as is frequently the case, spectra of the reactive intermediates are not available, some structural information can be obtained from changes in the spectrum of the adsorbate that are caused by various physical and chemical treatments; e.g., classification of distinct species as strongly or weakly adsorbed can be made from the effects of temperature and pressure on the adsorption bands.

A. SUPPORTED METALS 1. Hydrogen.

Pliskin and Eischens (65) have observed chemisorbed hydrogen and deuterium on alumina- and silica-supported platinum catalysts. A band a t 21 10 cm-1 was assigned to a weakly bonded form of hydrogen since its intensity could be easily varied by changes in temperature and pressure. A band a t 2058 cm-1 was assigned to a strongly bonded species. Deuterium produced corresponding bands a t 1515 cm-1 and 1479 cm-’, respectively. These assignments were based on spectra of platinum hydride complexes in solution, and on a value of 2083 cm-1 for the hypothetical Pt - H dimer predicted from a plot of the square root of the force constant versus the ionization potential of the metal, for known metal hydrides. Structures of the strongly and weakly bonded forms were proposed from an analysis of the shapes and relative positions of the two absorption bands. The possibility that these bands were due to the symmetric and asymmetric stretching frequencies of a triatomic group made up of two hydrogens attached to one platinum was ruled out on the basis

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

151

of band shape, since both bands would be of the same shape, whereas the 2110 cm-1 band was much sharper than the one a t 2058 cm-l. That the weakly bonded form was an adsorbed molecular ion was ruled out by the absence of a band midway between the hydrogen and deuterium frequencies when a mixture of hydrogen and deuterium was exposed t o the catalyst. The fact that the more strongly bonded hydrogen had the lower absorption frequency ruled out the possibility that the two species were bonded to the surface in the same way but on energetically different sites, for if this were the case, the frequencies would be reversed. Pliskin and Eischens proposed that the strongly bound species was attached to two platinum atoms in either a hydrogen-bond-like structure or by a one electron bond, and the weakly bound hydrogen to one platinum atom by a covalent bond. The only other observation of chemisorbed hydrogen was made by Pickering and Eckstrom (11)on rhodium films. They found a total of 18 absorption bands between 2193 cm-1 and 1416 cm-l which they did not attempt to interpret. It is possible that the bands represent adsorption on different crystal faces of a highly polycrystalline material. 2. Carbon Monoxide

Chemisorption of carbon monoxide on catalysts, mostly metals supported on silica or alumina, has been studied extensively. Eischens and Pliskin have reviewed ( I )a large portion of the work done to date. Supported metal catalysts can be divided into two groups according to the chemisorbed species of carbon monoxide found on them. I n the case of palladium, platinum, rhodium, and nickel two species have been found; one a CO molecule bridged between two metal atoms through the carbon atom, and the other with a single bond between the carbon and the metal. I n the case of copper and iron only the single-bonded species was observed. This first spectral evidence of heterogeneity in chemisorption was obtained by Eischens and Pliskin ( 1 ) .Recently Yates and Garland (66)reported finding five distinct species of carbon monoxide on nickel, as summarized in Table I V and Fig. 20. That these infrared absorption bands represented five distinct surface species was established by observing independent variations in the intensities during changes in the experimental conditions. For example, changing the concentration of the nickel in the catalyst samples produced large changes in the relative intensities of the five absorption bands a t monolayer coverage. For samples containing 1.5'/0 by weight of nickel the bands a t D and E were much more intense than those a t A, B, and C. These low intensity bands became more pronounced when the nickel concentration was increased to 10% and were predominant for samples

152

H. P. LEFTIN AND M. C. HOBSON, JR.

containing 25% nickel. Yates and Garland suggested that these effects were manifestations of the relative amounts of “crystalline” and dispersed nickel present. Catalyst samples poisoned with mercury vapor developed a strong band at D, but lost their ability to produce the bridged species a t A and B and the weakly bonded linear species at E. Structure D was strongly held and desorbed very slowly on evacuation at room temperature. No satisfactory explanation has been offered of the effects of mercury poisoning on the spectrum of adsorbed CO. Surface heterogeneity has also been demonstrated in studies of infrared spectra of nitric oxide adsorbed on transition metals, their salts and oxides. Terenin and Roev (67) found that although NO is more strongly

Frequency (cm-’l

FIG. 20. Infrared spectrum of CO adsorbed OR Nickel at various pressures. Curve 1, 1.7 x 10-4 mm; curve 2, 2.0 x lo-‘ mm; curve 3, 3.6 x lO-*mm; curve 4, 2.3 x lo-* mm; curve 5,1.2 mm; curve 6, 1.2 mm after 12 hr (66).

adsorbed on the oxides, its electroni: configuration and vibrational frequency are not changed as much as on the metals. Both ionic, NO+, and covalent, N=O forms were observed. Adsorption takes place on both metal ions and oxide ions; however, for oxides of nontransition metals NO is bound only on the oxygen sites. Carbon monoxide-metal systems have also been used to demonstrate the effect of carrier material on properties of supported metal catalysts. Eischens and Plisken (1)found that the ratio of concentrations of linear to bridged species was much larger when platinum was supported on silica than when it was supported on alumina. Carbon monoxide on the alumina-supported samples was much more difficult to oxidize than on silica-supported ones. These observations were interpreted as an enhancement of platinum’s ability to donate electrons t o the m e t a l 4 0 bond when alumina is the support. However, the mechanism by which

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

153

TABLE IV Aasignment of Infrared Bands to Surface Species for Carbon Monoxide Chemiaorbed on Supported Nickel Catalyat ( 6 6 )

Band

A

Frequency (cm-* 1

Structure

1915

Relative strength of adsorption

strong

? I

/c\

-Ni-Ni-

B

2035

strong

0

A

I -Ni-

C

1963

0 0 0 I

l

I/

\I

l

medium

c c c -Ni

D

2057

Ni-

0 I C

medium

0 I C

weak

I -Ni-

E

2082

1

-Ni-

154

H. l’. LEFTIN AND M. C. HOBSON, JR.

alumina influences platinum in this fashion remains obscure. Recently O’Neill and Yates (68) investigated the effects of alumina, silica, and titania supports on the nickel-CO system. The ratios of concentrations of linear to bridged species were much larger than for either the silicasupported or alumina-supported platinum samples discussed above. Although this ratio appeared to vary irregularly during adsorption, after evacuation and removal of weakly adsorbed species it reached a constant value of 1.85 for silica samples and 1.04 for alumina samples. Titania samples did not give a characteristic value on evacuation. These ratios cannot be translated into surface coverages since no reliable extinction coefficients are available for these species. That the observed species were adsorbed on the nickel and not the support was demonstrated by Yates (33)who studied the chemisorption of CO on both the rutile and anatase forms of titania. The CO was weakly bound to these surfaces and the absorption bands were at higher frequencies than with the CO-metal systems. 3. Formic Acid

An elegant example of the combination of infrared studies with kinetic and thermodynamic data to determine the overall mechanism of a catalytic reaction is the catalytic decomposition of formic acid. Its decomposition on silica-supported nickel has been investigated by several groups; namely, Fahrenfort and co-workers (1‘7,69),Clarke and Pullin ( 7 0 ) and Hirota, Kuwata, and Nakai ( 7 1 ) . Fahrenfort and Coworkers identified the prominent bands at 1575 and 1360 cm-l rn typical of carboxylate ions by comparison with the spectrum of nickel formate. These bands were assigned to the symmetrical and asymmetrical vibration of the 0 - C - 0 group, respectively. They showed that these bands were absent in the spectrum of formic acid adsorbed on the support. The formation of carboxylate ion (1575 cm-l band) a t room temperature was faster than the response time of the instrument, which was about 10 sec. The proposed mechanism is: HCOOH -------------+ 3 H, + OOCH,,,,

(5)

Mass spectral analysis of the gas phase indicated a small amount of decomposition a t room temperature caused by the side reactions: H C O O H - ............+2 Ha,ls+ GO,

(6)

+ HzO

(7)

HCOOH

.__._._...

-+ GO,,,

The decomposition of the carboxylate ion was followed by observing the change in the intensity of the 1575 cm-1 band. The rate of change of the optical density, dD/dt, of this band paralleled the change in the gas

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

155

pressure in the cell, and the end of the pressure rise coincided with the complete disappearance of the band. Energy of activation values calculated from rate data between 75 and 150°C ranged from 22.7 to 24.4 kcal/mole with the minimum a t 90°C. This agrees well with the value of 22.8 kcal/mole for formic acid decomposition on nickel powder between 125 and 15OOC. The interpretation of the infrared data was further supported by calorimetric measurements. Thus, the heat of adsorption of formic acid at monolayer coverage was 18 kcal/mole which compares favorably with the heat of formation for 4 mole of Ni (OOCH),, 13 kcal. These studies appear to have demonstrated; (a), the presence of the formate ion on the surface, (b) that it is the reactive intermediate in the catalytic decomposition of formic acid, and (c) that the decomposition proceeds from a formate-covered surface. Although the chemisorbed species producing the absorption bands a t 1575 and 1360 cm-1 seems to be well established as the chemisorbed intermediate that is rate controlling in the catalytic decomposition of formic acid, there remains some question of its exact identity. For example, Y a t e ~observed ~~ a similar spectrum for carbon dioxide chemisorbed on titanium dioxide and tentatively assigned it to the .GO, species. The two spectra are compared in Fig. 21. This anionradical is a reasonable possibility since it is isoelectronic with nitrogen dioxide, and thus should be stable. E.P.R. evidence for such species has been observed by Brivati et al. (72), who found a single line with a width of 14 gauss and a g-value of 2.0001 for irradiated lithium formate. Also by diffuse reflectance methods two broad bands were detected with maxima at 3500 and 2700 A. However, corresponding measurements have not yet been made on the catalyst system. 4. Hydrocarbons

Extensive studies of hydrocarbon chemisorption have been made by Eischens and Pliskin ( 1 ) .In a series of studies on olefins and paraffins chemisorbed on silica-supported nickel they were able to show that both associative and dissociative adsorption could occur, depending on catalyst pretreatment. Associative chemisorption of olefins is observed when hydrogen is left on the nickel surface; dissociative absorption, when the hydrogen has been pumped off a t an elevated temperature before chemisorption. The associative mechanism is deduced from the fact that the only absorption bands found when ethylene is added to a hydrogen-covered surface are in the C-H stretching region characteristic of saturated hydrocarbons, and that a C-H deformation band at 1447 cm-1 characteristic of two hydrogens on a carbon is also observed.

H. P. LEFTIN AND M. C. HOBSON, JR.

156

On bare nickel the band intensities are very low, indicating the formation of a surface complex of unknown composition and a low number of hydrogens per carbon. The addition of hydrogen produces a very marked increase in the intensities of the bands and results in a spectrum similar to that of ethylene adsorbed on hydrogen-covered nickel. Formic mid on nlckel

f

I

1800

1700

1600

1500

1400

1300

Frequency (cm?)

FIQ. 21. Infrared spectra. Curve A, formic acid on nickel (69); curve B, carbon dioxide on anatase (33).

Paraffins, of course, cannot have an associative mechanism and do not adsorb on hydrogen-covered nickel. When adsorbed on bare nickel they behave in a manner similar to ethylene. Acetylene, on the other hand, exhibits a unique behavior. Regardless of the pretreatment of the nickel, acetylene produces the spectrum assigned to chemisorbed ethyl radicals. This indicates extensive self-hydrogenation and the formation of a surface carbide. The idea of a surface carbide is supported by a very large increase in spectral intensity when hydrogen is added to chemisorbed acetylene. Eischens obtained some information on the manner of binding of the chemisorbed species by observing the relative intensities of some of the fine structure in the C-H stretching region. In the case of hydrogenated ethylene the intensity of the 2958 cm-l band, due to CH, groups, relative to the intensity of the 2924 cm-l band, due to CH, groups, indicates a 1 to 3 ratio, suggesting existence of an adsorbed butane radical as the principal species. Similar analysis of some of the higher olefins indicates that they are attached to the surface by more than two carbons, perhaps three or four. In view of the rather unpredictable behavior of extinction coefficients with surface coverage (Section 11,B, 3), these systems might stand re-examination. Little, Sheppard, and Yates (73) obtained similar results for acetylene

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157

and ethylene chemisorbed on silica-supported palladium and copper. However, they found an olefinic C-H stretch at 3030 cm-1 for chemisorbed ethylene. A very large increase in intensity on hydrogenation suggests that some olefin is present initially on the bare metal. More extensive self-hydrogenation of ethylene on nickel than on palladium wouldexplainthesedifferences. Little (6)reportsspectrafortheadsorption of these hydrocarbons on the oxides of nickel, copper, palladium, and silver, all supported on porous glass. Ethylene chemisorption results in polymer formation on the surface, as deduced from the observation that much more ethylene adsorbs than can possibly be accounted for by a monolayer; in fact, on nickel oxide the surface species would have to consist of fifteen ethylene molecules for every two nickel atoms. The only identifiable product found is trans-butene-2, I n contrast to the bare metals, there is no change in the spectrum of the adsorbed species on the addition of hydrogen. Moreover, acetylene does not chemisorb on these oxides.

B. OXIDESAND SALTS 1. Acetylenes

A rather extensive study of acetylene and its derivatives adsorbed on alumina and silica has been made by Yates and Lucchesi (74).Their results are explained in terms of acid-base concepts. Strong chemisorption is observed on alumina which has acid sites on its surface, but only physisorption is found on silica. The acetylenic CH vibration of the chemisorbed acetylene produces an absorption band a t 3300 cm-l which lies between the values for the corresponding vibration in gaseous acetylene and in monosubstituted acetylenes; in fact, its value is very close to that of liquid methyl acetylene (3305 cm-1). This implies that the chemisorbed species is held normal to the surface. The C = C frequency (2007 cm-I), of the adsorbate also supports this picture since this value lies between those for monodeuteracetylene (1851 cm-l) and methyl acetylene (2142 cm-1) suggesting that one end of the molecule has an “effective’) mass between 2 and 15. A weakly bound species is also observed with bands a t 3220 and 1950 cm-l which appear a t higher pressures and are rapidly removed by evacuation a t room temperature. The frequency shifts for these vibrations are similar in magnitude to those found on going from gas to liquid, but the relative pressures at which these bands are observed are much smaller than those considered normal for physisorption; therefore, the authors considered that this species is weakly chemisorbed and probably held parallel to the surface.

H. P. LEFTIN AND M. C. HOBSON, JR.

158

Similar results are obtained for methyl acetylene, but dimethylacetylene is completely different. Although the dimethylacetylene is strongly bound to the surface, the CH stretching bands are very close t o their gas phase values and no C = C band is found. This suggests that the species is held parallel to the surface. Furthermore, pre-adsorption of acetylene does not affect the subsequent adsorption of dimethylacetylene. Apparently the adsorption sites for the two species are not the same. I n terms of acid-base concepts the acidic protons on acetylene and methylacetylene react with basic oxygen sites on the surface. The detailed mechanism of this interaction is still an open question. When the protons are replaced by two methyl groups, the “acidic” nature of the acetylene is destroyed and the pi-electrons of the carbon-carbon triple bond then behave as a base and interact with Bronsted acid or Lewis acid sites. 2. Polar Molecules Employing a reflection technique, Kortum and Vogel (75) examined the spectrum of the lactone of malachite green-o-carboxylic acid adsorbed on the surfaces of a series of alkali halides and alkaline earth sulfates. Chemisorption of this compound resulted in formation of a zwitter-ion by ring cleavage [Eq. (S)] as evidenced by the similarity in

fi Y

h N n C - 0

-&Lo

Adsorption

-c----

De sorption

A

B

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

159

the spectrum of the adsorbate with that of malachite green cation (Fig. 2 2 ) . Water vapor slowly caused the displacement of the adsorbed species, accompanied by a bleaching of the color and conversion of the spectrum to that which might be expected for the physisorbed material. From these spectral data it was possible to draw some conclusions concerning the bonding involved in this chemisorption. The similarity of the spectrum of the adsorbate with that of the malachite green cation dictates that the resonance system shown in structure B must remain

.-.-.

FIG. 22. Reflection spectra. X - X X Malachite green lactone in methanol; 0-0-0 Malachite green in methanol; Malachite green lactone adsorbed on sodium chloride (75). ~

intact. Since covalent bonding through the positive center of the zwitterion would destroy this resonance system, the carbonium ion, which constitutes the positively charged half of the zwitter-ion, cannot be bonded covalently to an anion on the crystal surface for if this were the case, the spectrum of the adsorbate would have t o be similar t o that of malachite green carbinol. Further measurements established that a more-or-less polar binding was involved between the carboxylate group and the cations on the crystal. If the chemisorption and concurrent ring cleavage of the lactone occurred through the polarizing action of the cations, then the position of the equilibrium between structures A and B would depend upon the size of the cations and would be shifted to the left with increasing size. Hence, on lithium chloride the lactone would be more strongly chemi-

160

H. P. LEFTIN AND M. C . HOBSON, JR.

sorbed than on the other alkali chlorides. As seen in Fig. 23, the intensity of the band due to the zwitter-ion increased monotonically on going from ceasium chloride t o lithium chloride under otherwise identical conditions. An analogous dependence of the band intensity on the size of the cation was also observed for alkaline earth sulfates; however, the intensities were somewhat greater than for the alkali halides, indicating

FIG.23. Reflection spectra of malachite green lactone adsorbed on dried alkali chlorides. (75).

that the enhanced polarization forces of the doubly charged cations shifted the equilibrium further to the right. This may be partially due to other factors involving the differences between chlorides and sulfates. Even though the intensities were greater on oxides than on the corresponding sulfates, they still decreased with increasing size of the cation. Plots of the maximum intensity of the long wavelength band against mole fraction (Fig. 24) gave Curves (I)and (111) for the chemisorption isotherms on sodium chloride and potassium chloride, respectively. mole fraction, Beer’s law appeared to be obeyed up to about 5 x while a t higher concentrations the isotherms approached a constant value which could not be exceeded. By means of extrapolation, shown by the dotted curve, it was possible t o determine the saturation mole in close agreement fraction of chemisorption to be about 8-9 x with that found previously for butter yellow on barium sulfate (22). Comparison between curves I and I11 shows that the maximum value for the amount chemisorbed was about 25-fold less on potassium chloride than on sodium chloride. This indicates again that the equilibrium on the surface of sodium chloride lies much further to the right. A

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161

plot for the band a t 2700 A characteristic of the unionized lactone is shown in curve 11. Here we are concerned with a physisorption which has no saturation value since the molecules can adsorb upon each other in multilayers. The reflection spectra showed that both types of adsorption occurred concurrently on the surface.

FIG.24. Spectroscopically determined adsorption isotherms for malachite green lactone. Curve I, chemisorbed on NaC1; curve 11,physisorbed on NaC1; curve 111, chemisorbed on KC1 (75).

The adsorption of triphenylchloromethane (22) on barium sulfate was completely analogous. Here the C-C1 bond was so strongly polarized that the spectrum of the adsorbate (Fig. 25) was identical to that of the triphenylcarbonium ion as observed in sulfuric acid solution. On magnesium oxide this compound was not chemisorbed. I n the foregoing examples the spectral data indicated a Lewis acidbase reaction on the surface where the alkali and alkaline earth cations acted as the electron acceptors while the adsorbates were the electron donors. It is quite natural that the reverse situation might be possible; that is, the adsorbent be basic while the adsorbate show acidic properties so that in the chemisorption electron transfer will occur in the reverse direction. Several examples of such adsorption have already been discussed in this chapter. Kortiim (22) found another example in the adsorption of symmetrical trinitrobenzene on magnesia and on alumina. Whereas trinitrobenzene adsorbed on calcium fluoride or silica was colorless, on magnesia it was red with an absorption maximum at 4650A (Fig. 26) and the spectrum of the adsorbed species was very

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H. P. LEBTIN AND M. C. HOBSON, JR.

similar to that obtained in an alkaline ethanol solution of this compound. Elution of the adsorbed trinitrobenzene with methanol provided the unaltered colorless compound. The similarity of the spectra in the alkaline solution and on the krface led to the conclusion that the oxygen ions of magnesium oxide functioned as a basic group for this chemisorption.

-.-.-.-

FIG. 25. Spectrum of triphenylchloromethane: -0-0-0adsorption spectrum in sulfuric acid solution; reflectance spectrum of triphenylchloromethane adsorbed on barium sulfate (22).

V(crn-9

FIG.26. Spectrum of 1,3,5-trinitrobenzene: curve 1, adsorbed on SiO,; curve 2, adsorbed on MgO; curve 3, in alcoholic KOH solution (22).

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

I63

C. CRACKINGCATALYSTS 1. Dyes and Polar Molecules

a. Acid-Base Reactions. Kotov (76) observed only a single band a t 2600 A for aniline adsorbed on a silica-alumina gel a t low coverages. As additional aniline was adsorbed, a second band appeared at 2800 A whose intensity increased with increasing coverage. When capillary condensation occurred the 2800 A band broadened and shifted t o 3000 A. The 2600 A band was identical to that obtained when benzene was adsorbed on this catalyst. The presence of two bands in the spectrum of adsorbed aniline was taken t o indicate the presence of two types of adsorption centers on silica-alumina. Since a band a t 2800 A is displayed by a solution of aniline in methanol, this band was attributed t o aniline adsorbed on surface hydroxyl groups. Interaction with adsorption centers which are specific only to silica-alumina gave rise t o the 2600 A band since the spectrum of aniline adsorbed on silica gel exhibited only the band a t 2800 A. From these data it was concluded that adsorption on the hydroxyl groups occurred only after all of the specific adsorption sites had become occupied. Since the concentration of OH groups on the surface of silica-alumina is considerably greater than that of other acid centers, it is natural that the band with the maximum a t 2800 A should have a much higher intensity than the band a t 2600 A. Conclusions can be drawn concerning the nature of the specific centers on silica-alumina from the similarity of the spectrum of adsorbed aniline with that of benzene. These centers may be of two types, either Bronsted acids or Lewis acids, and may be depicted as shown in structures Si Si I

si-

0 I

0-Al

I

0 I

si

si

A

B

Structure A can easily give up its weakly bound proton and hence function as a Bronsted acid. Structure B possesses electron-accepting properties by virtue of the aluminum atom which contains an empty p-orbital. Reversion of the aniline spectrum to that of benzene will occur when the excess electron density distributed to the ring by the amino group becomes localized on the nitrogen atom. Thus proton addition to the amino group in aniline produces the anilinium ion whose

164

H. P. LEFTIN AND M. C. HOBSON, JR.

spectrum closely resembles that of benzene. Similar behavior is observed when aniline vapors react with anhydrous aluminum chloride. I n this case the electron acceptor binds the unshared electron pair to produce the Lewis acid adduct. Since the changes in the spectrum of aniline adsorbed on the specific centers of silica-alumina can be explained in this manner, it is evident that these centers possess either strong electron-accepting or proton-donating properties. The spectrum of naphthylamine adsorbed on bentonite indicated that the same specific centers were present as in the silica-alumina system. I n order to decide between these two possible structures, Kotov (76) examined the effects of cation exchange and thermal removal of protons on the spectra of adsorbed aniline and naphthylamine. Cation exchange with 1 N sodium chloride solution reduced the number of surface protons by a factor of four, Thermal treatment for 2 hr a t a series of temperatures up to 600°C reduced the number of protons in a regular fashion. The spectra of naphthylamine adsorbed on the base-exchanged and thermally treated catalysts demonstrated that the extent of acid reaction of the amine with the surface was not diminished even after removal of 7504 of the available protons, indicating that on these samples electron-accepting centers on the surface were responsible for the amine adsorption. Since the spectra did not permit one to distinguish between the reaction products of the amine with Bronsted and with Lewis acids, Kotov concluded that on the unheated samples the spectral changes were due to the presence of proton-donating centers and on the heated samples to electron-accepting centers. 6 . One Electron Transfer Processes. While the chemisorption of many basic nitrogen dyes on silica-alumina and other acidic catalysts leads to the expected conjugate acids in which a nonbonding electron pair on the nitrogen is shared with the surface, in several instances the spectral data indicate another type of chemisorption involving the complete transfer of a single electron from the dye molecule to the surface. Close correspondence between the spectra of the adsorbed species and those of the corresponding molecular positive ions was demonstrated by Sidorova and Terenin (77) for diphenylamine and other diarylamines adsorbed on evacuated samples of bentonite and of silica-alumina. Subsequently, Kotov (78) demonstrated a similar behavior for benzidine adsorbed on bentonite and silica-alumina, or on silica gel in the presence of oxygen. In the absence of oxygen, bands a t 3000 A, 4500 A, 7600 A, and 8500 A, were observed in the spectrum of benzidine adsorbed on both silica-alumina and bentonite. An additional band a t 6000 A occurred with bentonite, while with silica gel only a single band was obtained

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166

which was also present in alcoholic solutions of this amine. I n order to promote the ionization of the dye on silica gel, 10 mm of oxygen was introduced into the cell. After 15hr, bandsat 4100 Aand 8500 Aappeared which, by comparison to the 4500 A and 8500 A bands on silica-alumina, were attributed to the benzidine positive ion. Assignment of these bands to a positive molecular ion was verified by comparison with the spectra obtained by Lewis and Lipkin (79) in the photoionization of benzidine in a frozen matrix. The color of benzidine adsorbed on bentonite is reversibly altered by adsorbed water vapor. Vedeneeva (80) found that the intense blue color observed in the prescnce of excess water changed to green, yellow and finally colorless as the adsorbed water was removed. This color sequence was reversed when moisture was added. More recently, Hasegawa (14) reexamined this so-called “benzidine reaction” with both benzidine and tetramethylbenzidine, using benzene as a supporting solvent for the catalyst suspension. Adsorption of benzidine on natural or acid-treated clay produced a greenish-yellow color (curves A and B, Fig. 27) with absorption maxima a t 4400 A, 7800 A, and 8800 A, in close

Wovelength (mp)

FIG. 27. Absorption spectra of adsorbed benzidine;curve A, on natural clay; curve B, on acid-treatedclay; curve C, on alkali-treatedclay ( 2 4 ) .

agreement with that observed (78) under vacuum conditions. These spectra differed from that on alkali-treated clay (Curve C) but were similar to that on a base-exchanged clay where the surface protons had been exchanged by boiling with 5 N sodium chloride. The spectral differences between acid-treated and alkali-treated clays, therefore, do not reside in the replacement of protons with other ions. Since essentially all of the protons were replaced by the base exchange, surface Bronsted acid sites could not have been responsible for production of

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H. P. LEFTIN AND M. C. HORSON, JR.

the bands observed with the natural or acid treated clays. Analogous results were obtained using tetramethylbenzidine. I n order to interpret these spectra, two oxidation products of tetramethylbenzidine were prepared using ferric chloride as an oxidizing agent. These were identified as the quinoidal and meriquinoidal ( 1 : l complex of the base and the quinoid) forms absorbing a t 4700 A and 8800 A, respectively. The spectra were dominated by the quinoidal form on the acid clays and by the meriquinoidal form on the alkali-treated clays. The effect of catalyst heat treatment on the spectra of amines adsorbed on silica-alumina catalysts was examined by Okuda and Tachibana (15) employing the suspension technique. Two bands were observed with benzeneazodiphenylamine. The band at 4400 A was attributed t o the basic form of the dye, while that a t 5380 A to the conjugate acid formed by reaction of the basic dye either with a proton or with an aluminum

Wavelength (mp)

FIQ.28. Absorption spectra of p-phenylenedimine adsorbed on silica-alumina; curve A, not heat-treated;curve B, heat-treated3 hr at 110OC; curve C, heat-treated3 hr at 430'C (15).

atom as a Lewis acid. Increasing heat treatment of the catalyst caused an enhancement in the intensity of the acid-form, indicating that the number of effective acid sites had been increased. To determine if these acid centers were of the Lewis acid type the spectrum of adsorbed pphenylenediamine was examined. This compound, which is a typical electron donor, is suitable for detecting electron-accepting properties. Bands a t 3240 A and 4680 A were observed, The former was attributed t o the free (physisorbed) diamine by comparison with the 3120 A band

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167

from a n alcohol solution. The band a t 4680 A was identified with the semiquinone cation-radical (NH,C,H,NH,)+ which, in aqueous solution, exhibits a maximum a t 4620 A. It is particularly noteworthy that in this case the existence of the cation-radical on the surface was confirmed by the appearance of an electron spin resonance (esr) spectrum. Since neither the spectral band nor the esr signal was obtained from p-phenylenediamine adsorbed on a silica gel containing protonic acid, it was suggested that on the silica-alumina system a Lewis-type acid was responsible for the formation of the cation-radical. Further evidence that the active centers on silica alumina-are Lewis rather than protonic acids was provided by the spectral response to catalyst pretreatment shown in Fig. 28. Here curves A, B, and C represent the spectra of p-phenylenediamine chemisorbed on silica-alumina samples which were heated a t increasing temperatures. The regular increase in the intensity of the 4680 A band, due to the cation radical, was taken as spectroscopic evidence for an increasing number of Lewisacid sites. Although these spectra are in qualitative agreement with the known effect of thermal treatment on the relative abundance of Lewis and Bronsted acidity, quantitative conclusions cannot be drawn since the concurrent increase in intensity a t 3240 A indicates that these measurements were not made under conditions of constant surface coverage. 2. Electronic Spectra of Chemisorbed Hydrocarbons

All of the electronic spectra discussed thus far have involved heteroatomic organic compounds adsorbed on polar surfaces. These data have provided considerable new insight into the nature of the chemisorption process, particularly concerning the electronic properties and character of specific types of adsorption sites. I n the sections which follow we will restrict our attention to spectral studies involving the chemisorption of hydrocarbons on cracking catalysts. These studies, in addition to reinforcing earlier conclusions concerning the nature of the surfaceactive sites, have their principal value in the elucidation of reaction mechanisms by identification of reactive chemisorbed intermediates. Major investigations in this area have been carried out by Webb (20) and by Leftin and his co-workers (27,29, 62, 81, 82, 83). a. Arylalkanes. The demonstration of surface acidity of cracking catalysts (50, 51, 84) in conjunction with product distribution (85) and isotope-exchange studies (86)has led workers (51, 84, 85) to suggest reaction mechanisms which differed in detail but assumed an initial chemisorption process that produced adsorbed carbonium ions. Since the details of this chemisorption process were not known a multi-

H. P. LEFTIN AND M. C. HOBSON, JR.

168

plicity of possible mechanisms for initial carbonium ion formation were proposed. One of the major obstacles to the formulation of an acceptable mechanism for paraffin chemisorption lies in the fact that the chemical nature of the active surface acid sites is not known. Since this acidity may be due to the presence of surface protons, electron-deficient centers, or both, depending upon the hydration state of the catalyst, a t least two possible mechanisms need t o be considered. The first question which must be answered concerning paraffin chemisorption is whether carbonium ions can be formed in this process. With a view toward answering this question the electronic spectra of the compounds listed in Table V adsorbed on transparent platelets of silica-alumina were determined (27, 29, 82) using vacuum techniques.

I

I

Wavelength (mp)

FIG.29. Absorption spectra: curve A, 4*CH chemisorbed on silice-alumina;curve B, 4,COH in conc. H,SO,; curve C, in ether; curve D, 4*C- in ether (29).

As these phenyl-anologs of isobutane all have a single tertiary hydrogen atom, they may be considered as possible precursors of tertiary carboniuni ions. Phenyl derivatives were selected because phenyl groups are known to stabilize carbonium ions and independent existence of these ions have been amply demonstrated by cryoscopic measurements of sulfuric acid solutions of the corresponding carbinols (87, 88) and by conductivity measurements in liquid sulfur dioxide solutions of the corresponding halides and perchlorates (89, 90, 91). Also, the electronic spectra of these ions have been well characterized (92-96). Isobutane is not included in this list since only indirect evidence is available for the independent existence of the tertiary-butyl ion and its electronic spectrum is still in doubt. Studies involving this compound and other aliphatic hydrocarbons will be discussed separately in Section VI, C, 2, c. When a silica-alumina catalyst (previously calcined in 0 and evacu-

169

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

TABLE V Compounds Used in Spectral Studiea

Hydrocarbon

Carbonium ion

Principal absorption band

4040;4320

Triphenylmethane

Triphenylcarbonium ion

4230

1,l -Diphenylethane

Methyldiphenylcarbonium ion

3950

Dimethylphenylcarbonium ion

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H. P. LEFTIN AND M. C. HOBSON, JR.

ated a t 500°C) was exposed to the vapors of triphenylmethane a yellow surface color slowly developed (29). Comparison of the spectrum of the adsorbed species (Fig. 29, Curve A) with the known spectrum (Curve B) of the triphenylcarbonium ion (as determined from a solution of the carbiriol in concentrated sulfuric acid) provided convincing evidence that this ion was formed on the catalyst surface. Comparison with the spectra of the triphenylmethide (97) ion and of the triphenylmethyl radical (97) demonstrated that neither of these species was formed in this chemisorption process.

Wowelenqth (mp)

FIG. 30. Absorption spectrum of triphenylmethane: curve A, in ethanol solution; curve B, adsorbed on silica gel; curve C, B after 4-hr evacuation at 100°C (29).

That the observed spectrum was the result of a chemical reaction between the hydrocarbon and the catalytically active centers of the silica-alumina surface (chemisorption), and not due t o a general surfatochromic spectral shift, was demonstrated from the spectrum of this compound adsorbed on a nonacidic or very weakly acidic silica gel (29). The spectrum (Fig. 30, Curve B) of silica gel exposed to triphenylmethane vapor for 1000 hours at 100°C was identical to the spectrum (Curve A) of an alcoholic solution of this hydrocarbon. The close agreement between these spectra suggested that on silica gel the triphenylmethane was physisorbed. This was further evidenced by the marked loss of spectral intensity (Curve C) attendant to a four hour evacuation a t 100°C. I n contrast, on silica-alumina where the hydrocarbon was chemisorbed as the carbonium ion, no decrease in absorbance was noted even after 48 hr evacuation a t 275OC. These data constituted the first direct demonstration of the formation of carbonium ions as a consequence of ehemisorption of a “tertiary” hydrocarbon on the surface of a cracking catalyst by a reaction involving the rupture of an aliphatic C-H bond. The genemlity of this process of carbonium ion

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

171

formation in the chemisorption of paraffins on silica-alumina was shown from comparison of the spectra of chemisorbed 1,l-diphenylethane and of cumene (Fig. 31) with the spectra of the methyldiphenyl and dimethylphenyl carbonium ions observed from sulfuric acid solutions of the corresponding alcohols. On the basis of the above results,

FIQ. 31. Spectra of carbonium ions: curve A, 1,l-diphenylethaneadsorbed on silicaalumina; curve B, 1,l-diphenylethylenein 98% H,SO,; curve C, cumene adsorbed on silica-alumina; curve D,2-phenyl-2-propanol in H,SO,.

the over-all course of the reaction involved in the chemisorption of tertiary hydrocarbons must be R,CH

Silica-Alumina

R3C+

+ H-

where clearly the hydrocarbon molecule loses a hydride ion (or its equivalent) t o form the observed tertiary carbonium ion. That the carbonium ions do not form covalent carbon-to-surface bonds, as had been suggested previously ( 9 8 ) ,is indicated by the fact that such bonding would destroy the resonance system responsible for the stability and characteristic spectra of these ions. Carbonium ions must be held in the vicinity of the surface by coulombic forces and may be considered to constitute the positively charged half of an electrical double layer. I n the chemisorbed state they may be free to rotate and may even be free t o migrate over limited regions of the surface. Recent nmr data (99) and the isotope exchange resulta of Haldeman and Emmett (100)lend support to some degree of mobility of chemisorbed carbonium ions. According to this view of the chemisorbed state, carbonium ions are readily accessible to attack by oncoming hydrocarbon molecules from the gas phase with which they may undergo the well-known (101)hydride transfer reaction,

H. P. LEFTIN AND M. C. HOBSON, JR.

172

thereby acting as chain initiators for the catalytic reaction. This simple picture, involving the formation of an electrical double layer in the chemisorption process, is also consistent with the theory of Cook et al. (102) for the ion exchange capacity of clay-like minerals and with recent proposals for hydrocarbon reactions over acid catalysts. The nature of the negatively charged half of the chemisorption double layer is closely associated with questions concerning both the mechanism of the chemisorption process and the chemical nature of the surface acid site. Two principal mechanisms shown below were considered :

Mechanivm A (1) RsCH

+ u-+ R P C +

(2) RsCH

+ HA -+ R&+ + A- + H,

Mechanism B

+

R,C = CR,

+ HA+R,CCHR,

R &LCHR p

+ R $33

+

--t

R iC+

+ A-

+ R &H-CHR

a

According to Mechanism A (85),a paraffin molecule may be chemisorbed via direct hydride ion abstraction by the catalyst itself acting either as a Lewis acid center (designated by the symbol u) or via direct reaction with the Bronsted acidity of the catalyst with concomitant evolution of hydrogen. Mechanism B (84) assumes that the only active centers available for carbonium ion formation are Bronsted acids; it requires that only olefins can react directly with the surface to form carbonium ions which then act as initiators for paraffin reactions. Accordingly, paraffins can be converted to carbonium ions only by hydride transfer to these initiating pre-adsorbed carbonium ions. Thermal cracking and traces of olefin or olefin-forming precursors have been suggested as possible sources of these initiators. Using the triphenylcarbonium ion as a probe it was possible to demonstrate (82) that mechanism A-1 could obtain and to exclude A-2 and B from further consideration. It should be noted that if prior olefin adsorption were a necessary prerequisite for paraffin adsorption, then carbonium ions could not be formed from paraffin molecules in a system rigorously freed from olefin or olefin-forming impurities. As triphenylmethane itself is certainly not an olefin precursor, there remained only two possible sources of olefinic impurity to be considered; namely, the catalyst surface and the reagents employed, The former was discounted on the grounds that it was highly improbable that either olefins or carbonium ions on the surface survived the ex-

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SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

haustive 500°C oxidation and evacuation pretreatment employed. Further, since an all-glass, greaseless, system was used and since the only reagent present was the triphenylmethane itself, any such olefin would have to stem from the impurity content of this reagent which had been rigorously purified and shown to have a total impurity level below detectability. To eliminate Mechanism B from further consideration it was only necessary to show that, in any given experiment, more carbonium ions were formed than the maximum possible olefinic impurity level could have provided for. I n the experiment shown in Fig. 32, 1.8 x 10-4 gm of carefully purified triphenylmethane was

3 Wovelength (my)

FIG. 32. Chemisorption of triphenylmethane (1.8 X 10-4gm) on silica-alumina (8.9 metere). Complete chemisorption would correspond to 6 X 10la carbonium ions per square centimeter of surface.

vacuum transferred onto a catalyst sample which had been previously freed of any and all adsorbed olefin or carbonium ions by oxidation and evacuation a t 500°C. The absorbance developed rather slowly at room temperature over a period of 740 hr, however, there was no substantial change after the first 170 hr. Chemisorption of all of the available triphenylmethane would have corresponded to a maximum surface coverage of 5 x 1 0 1 2 triphenylcarbonium ions/cm2 of available surface. Using this value as a n upper limit of the amount chemisorbed would have required that the reagent have had an impurity level of the order of 50 mole yo.In a separate experiment, using a tenfold excess of reagent and determining the amount chemisorbed by an extraction method, a lower limit of the required impurity level (15%) was established. Both of these values were unreasonably high for the reagent employed.

174

H. P. LEFTIN AND &I. C. HOBSON, JR.

The possibility that the Rpectral observations were the result of the action of chemisorbed oxygen was eliminated by showing that identical spectra could be obtained from a catalyst sample which, in addition t o the usual pretreatment, was further treated for a 24-hr period in an atmosphere of hydrogen a t 500°C prior to final evacuation. The foregoing results established for the first time, that it is possible for carbonium ions to form by the chemisorption of parafin molecules on silica-alumina through an intrinsic ability of the catalyst surface t o abstract hydride ions directly from tertiary C-H bonds. It is important to note that these results do not exclude the operability of mechanism B in systems containing olefinic impurities, where it may well be the most favorable path; they merely show that mechanism A is suficient to account for parafin chemisorption.

Wovelength (my)

FIG. 33. Triphenylcarbonium ion formation with supported Lewis and Bronsted acids: curve A, 4,CH adsorbed on silica gel; curve B-1, 500 mm HF added at 26°C; curve B-2, heated 115°C 4 hr; curve C, 4,CH adsorbed on silica gel and B F , added at 26°C; curve D, 4,CH adsorbed on silica-alumina; curve E, 4,COH adsorbed on silica gel and HF added at 26°C.

A choice remained between Mechanisms A-1 and A-2 which requires that the active centers be either Lewis or Bronsted acid, respectively. Since A-2 would lead t o the formation of molecular hydrogen, an attempt was made to detect and measure any hydrogen evolution concomittant with the chemisorption of triphenylmethane. Triphenylmethane was chemisorbed on 28 gm of Houdry S-65 synthetic silicaalumina catalyst in a sealed, evacuated apparatus. When the chemisorption was completed the gas phase, collected using a Sprengle pump,

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

175

was found to contain less than 0.02 cc NTP of hydrogen. Subsequent analysis of the catalyst sample by an extraction technique showed that the amount of triphenylmethane which had been chemisorbed would have provided 11.4 cc (NTP)of hydrogen if Mechanism A-2 obtained. Additional evidence supporting Mechanism A-1 in preference to A-2 was obtained from spectral studies with mounted Bronsted and Lewis acids. Figure 33, Curve A is the spectrum of triphenylmethane physisorbed on silica gel. On exposure of the physisorbed hydrocarbon a t 25°C to 100 rnm of anhydrous BF,, the spectrum changed rapidly to that characteristic of the triphenylcarbonium ion (Curve C) and was similar to that observed on silica-alumina (Curve D) indicating that with this

c 0 c

c n 8

a

Wovelength (mp)

FIG. 34. Effect of water vapor on chemisorbed triphenylmethane: curve 1, 4,CH chemisorbed on silica-alumina; curve 2, after exposure to water vapor 30"C, 25 mm, for 5 min; curve 3, evacuated 1 min at 30°C; curve 4, exposed to water vapor; curve 5, evacuated 1 min at 30°C; curve 6, exposed to water vapor; curve 7, evacuated 5 min at 30°C.

compound, a t least, carbonium ions could be formed by reaction with a known Lewis acid. I n a separate experiment, exposure of physisorbed triphenylmethane to 500 mm of anhydrous HF (Curve B) failed to effect carbonium ion formation even after 48 hr a t lOO"C, indicating that Bronsted acids cannot function in this capacity. The possibility that adsorbed HF could not function as a strong surface acid was excluded by demonstrating carbonium ion formation (Curve E) from preadsorbed triphenylcarbinol which is a stronger base than the corresponding methane. From these data it was concluded that carbonium

176

H. P. LEF'MN AND

M. C. HOBSON, JR.

ions can be formed by direct hydride ion abstraction from the adsorbate by the action of the Lewis-acid sites of the catalyst and that the Bronsted acid sites can not readily react with paraffins under these experimental conditions. The qualitative behavior of the silica-alumina-triphenylcarbonium ion system toward water vapor a t room temperature is shown in Fig. 34. Curve 1 represents the time-invariant spectrum prior to exposure to water. On exposure to water vapor a t 80% relative humidity a t 30°C the characteristic spectrum completely disappeared, as shown by Curves 2 and 6 which coincide with the background curve. Evacuation for only 1 min a t room temperature partially restored the spectrum (Curve 3) and complete restoration was achieved by 5-min evacuation (Curve 7). This phenomenon is a striking demonstration of the ready reversibility of the reaction of water vapor with chemisorbed carbonium ions and demonstrates that the water involved must indeed be rather loosely bound to the surface. Results of a more quantitative examination of this water effect are shown in Fig. 35 where the absorbance a t 4050 A is plotted as a function

FIG.36. Effect of water and of ammonia on chemisorbed triphenylcarbonium ion on silica-alumina.

of the amount of water adsorbed on the sample. The first increment of water caused a small decrease in the absorbance after which there was no further decrease until about 17 water molecules per carbonium ion were taken up by the surface. Subsequent addition caused a very rapid decrease in absorbance until the spectrum was completely bleached. With ammonia, approximately one molecule was required to remove a

SPECTROPROTOMETRY FOR CATALYTIC SYSTEMS

177

carbonium ion from the surface. I n this case, however, the reaction was not reversible and evacuation for extended periods at 100°C caused the desorption of a white solid which condensed on the cold portions of the apparatus and which was shown from its infrared spectrum to be triphenylmethane and not the expected triphenylamine. According to Mechanism A-1 the hydride species removed from the hydrocarbon becomes associated in some way with an electron-deficient center on the surface. If this species exists either as an ionic hydride or as a covalent metal hydride, then reaction with water vapor should afford molecular hydrogen according to the familiar reaction M-H

+ H,O -+

MOH

+ H,

(9)

Simultaneous reaction with the adsorbed carbonium ion to produce the corresponding colorless carbinol ; e.g.,

4&+ -+ H,O 2 +&OH

+ H+

(10)

would be in accord with the observed reversible bleaching of the spectrum. Recent (103) studies have shown that neither of the above reactions describes the system. Exposure of a silica-alumina catalyst, whose surface contained 5 x lo1, triphenylcarbonium ions/cm2, to sufficient water vapor to effect complete bleaching of the carbonium ion color did not afford any noncondensible gases under experimental conditions which, according to Eq. (9), would have provided 40 cc (NTP) of molecular hydrogen, Subsequent extraction of the adsorbed phase followed by infrared analysis of the residue showed that the adsorbed phase consisted entirely of triphenylmethane. There was no evidence of triphenylcarbinol formation according to Eq. (lo), nor of isotope-exchange with the tertiary hydrogen in a similar experiment using D,O. On the basis of these results it was concluded that adsorbed water indirectly reverses the chemisorption. While a detailed picture of the mechanism by which water functions cannot be given at this time, it may be pointed out that these observations are in accord with the known effects of catalyst water content on silica-alumina catalyzed reactions. b. Arylalkenes. Of the three investigations of the electronic spectra of adsorbed aryl-substituted olefins, the earliest was a n undocumented report by Evans (104) that upon addition of 1,l-diphenylethylene (hereafter called DPE) to an oil suspension of acid-treated clay a yellow surface color developed. From the similarity of the spectrum of this clay suspension to that of a sulfuric acid solution of DPE Evans concluded that the carbonium ion, (6,C+CH, had been formed by proton transfer from the clay to the olefin. More extensive investigations by

178

H.P. LEFTIN AND M.

C. HOBSON,

JR.

Webb (20) and by Leftin et al. ( 8 1 , 8 2 , 8 3 , 1 0 5 )revealed that the chemisorption of DPE was considerably more complicated than this simple picture. The spectrum of DPE chemisorbed on silica-alumina (Fig. 36) was found ( 2 0 , 2 7 )to consist of two intense absorption bands in contrast t o the single band found with the clay suspension. Of these, the one at 4230 A ( 6 = 3 x 104) was assigned to the methyldiphenylcarbonium ion, +&3+CH3,by comparison with the spectra obtained from sulfuric

Wovelenqth (mM)

FIG. 36. Absorption spectrum of 1,l-diphenylethylene: curve A, chemisorbed on H,SO,; curve C, in 98% H,SO,. silica-alumina; curve B, in dimethylsulfate

+

acid solutions of either DPE or methyldiphenylcarbinol where this ion has been independently identified by cryoscopic (88) and by nmr data (106). It was not possible to distinguish whether this ion was formed by interaction with a surface proton or with a Lewis acid site since in either case the ions would have closely similar spectra. Assignment of the second major band (A = 6050A) was not as straightforward. Bands a t wavelengths substantially shifted toward the red from those for the expected carbonium ions were first reported by Lavrushin (95) for DPE and for a a-methylstyrene in mixed aceticsulfuric acid solutions. No explanation of the origin of these long wavelength bands was offered, Several years later, Evans and his co-workers (107, 108) rediscovered this phenomenon in the course of their studies on the dimerization of diarylolefins. On the basis of a rather limited process-of-elimination these workers proposed (108) that these bands were caused by nonclassical carbonium ions, or pi-complexes formed between protons and the pi-electron clouds of the olefins. This concept was soon challenged by Grace and Symons (109) who found similar bands from acid solutions containing compounds which had no olefinic pi-electrons; e.g., diphenglmethanol. I n spite of this criticism, Webb (20)

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

179

adopted Evans’ pi-complex hypothesis to explain his spectral data on silica-alumina. Spectra of DPE chemisorbed on silica-alumina are sensitive to the water history of the catalyst. Webb found only the band a t 42308, due to the carbonium ion, when DPE (dried over NaH) was chemisorbed on a dried catalyst (Curve A, Fig. 37). Chemisorption on a hydrated sample afforded a spectrum (Curve B) in which the long wavelength

Wavelength (mu)

FIQ. 37. Spectrum of &C=CH, adsorbed on silica-alumina:curve A, dry catalyst; curve B, hydrated catalyst (20).

band predominated. At intermediate water levels both bands appeared. Webb interpreted these results to indicate that the species responsible for the long wavelength band was a pi-complex (or charge-transfer complex) between the olefin and a surface hydroxyl group formed by the chemisorbed water. Leftin and Hall (83) found that the final spectrum was sensitive t o surface coverage and to ageing, as well as to catalyst water content. Using an all-glass cell in which the catalyst and reagent were in separate compartments isolated by a fragile glass membrane, it was possible t o follow the spectral changes which accompanied increasing surface coverage. At low surface coverages only the band a t 4230 A (carbonium ion) was observed (Fig. 38). The rate of development of this band decreased with time, while as the adsorption continued the band a t 6050 A began to appear and its rate of production increased until its intensity exceeded that of the 4230 A band. I n the presence of an excess of the reagent the spectrum lost resolution, indicating that secondary processes such as dimerization had set in. The carbonium ion band was reversibly bleached by added water vapor (Fig. 39) in a manner similar to that noted earlier with the triphenylcarbonium ion. The band a t 6050 A, on the other hand, while relatively insensitive to water, was

180

H. P. LEFTJX AND M. C. HOBSON, JR.

irreversibly bleached by ammonia. These results suggest that: (a) both bands involve surface acid sites, and (b)these acid sites differ in strength, Adsorption of DPE on the stronger sites to form the carbonium ion will occur first and these can respond to a weak base such as water. When these sites become substantially occupied, adsorption on the weaker sites with formation of the 6050 A species will predominate. These sites, while not sensitive to water, can respond to a stronger base such as ammonia. Curve no. Av. oge (rnin)

I

400

L

500

A

I

600

7' 1 0

Wavelength (m)L)

FIG. 38. Spectra on silica-aluminaat various times following exposure to a source of 1,l-diphenylethylene.

Spectra similar to those on the catalyst surface could be observed in solution of DPE in weak acids such as acetic acid-sulfuric acid or in benzene-trichloroacetic acid mixtures. For the purpose of kinetic measurements the reaction was studied in a ternary solvent composed of chloroacetic, acetic and sulfuric acids. The reaction vessel consisted of two compartments separated by a break-off seal. Solutions containing the olefin in acetic-chloroacetic acid and solutions containing the active acid reagent were separately degassed a t - 195°C and sealed off. After equilibration at 25°C the reaction was initiated by vigorous mixing and the rates were monitored spectrophotometrically. Data from a typical experiment are presented in Fig. 40. Initially the concentrations of the carbonium ion and of the 6050 A species increased simultaneously; after several hours both approached their maximum intensity but while the former remained constant or increased very slowly, the latter decreased in concentration until after about 10 days

181

SPECTROPHOTMETRY FOR CATALYTIC SYSTEMS

the system reached a metastable state. It was evident that in the initial stages of the reaction the 4230 A band increased linearly while the 6050 A band increased in a parabolic fashion, indicating that the initial rate of formation of the latter species was dependent upon the concentration

u

1 a:

I

600

450

750

Wovelength (mjd

FIG. 39. Effect of water and of ammonia on the spectrum of DPE chemisorbed on silica-alumina: Curve 1, DPE chemisorbed on silica-alumina; curves 2 and 4, water added at 25'C; curves 3 and 6, evacuated at 26°C; curve 8, ammonia added.

0

20

40

60

80

100

120

I40

160

180 2 0 0 220

240

Hours

FIQ.40. Kinetics of formation of colored species from DPE in a ternary solvent system; viz., 2.96 M ClCH,COOH in CH,COOH plus 16% H,SO,. Hydrocarbon concentration waa 1.02 X lo-' M in the combined solvent.

182

H. P. LEFTIN AND M. C . HOBSON, JR.

of the carbonium ion. The fall-off in concentration beyond the maximum may correspond to any of a variety of secondary reactions; e.g., dimerization. Rates of formation of both species increased with increasing olefin concentration and with increasing sulfuric acid concentration. Although insufficient data were obtained for a detailed kinetic analysis, it was clearly evident that the reaction was much too slow to be consistent with pi-complex formation. The effects of trace amounts of oxidizing agents on the kinetic behavior was investigated. With the addition of as little as 0.15 wt yo selenic acid the same over-all kinetic pattern was obtained (Fig. 41) but

Minuirr

Fro. 41. Effects of oxidizing agents on the kinetics of formation of the 6060 A species in the ternary solvent. Hydrocarbon concentration was 4.62 x 10-4 M in each run.

the entire process required only about 30 min as compared to over 40 hr in the comparable uncatalyzed system. For direct comparison the uncatalyzed reaction a t the same acid concentration (15% H,SO,) and at a higher acid concentration (22.5%) are shown. Certainly, in view of the modest effect of acid concentration on the rate, the vastly increased rate could not have resulted from the small increase in acid strength of the medium attending the addition of the H,SeO,. Similar results were obtained with peroxydisulfuric acid. More striking, however, was the fact that the addition of as little as 1 x 10-4 moles/liter of K,Fe(CN), also strongly catalyzed the reaction. These results indicated that the species exhibiting the 6050 A band was formed by the oxidation of the olefin. Perhaps the simplest oxidation process which could account for these observations is one which involves the removal of a single electron from the olefin molecule to form a molecular ion analogous to the wellknown (110, 111) cation-radicals found in acid solutions of condensed aromatics. Such ions are characterized by their paramagnetism and by the fact that both the esr and optical spectra of the cation-radical of a

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

183

given hydrocarbon are very closely similar to those for the corresponding anion-radicals obtained by reaction with alkali metals in suitable solvents (112). The possibility that the species responsible for the 6050 A band was a cation-radical was further tested. Electron spin resonance measurements on DPE solutions in mixed acid solvents showed a very weak signal. A single intense resonance line, having a width of 10 gauss and centered a t g = 2.00, was obtained when DPE in CH,Cl, was treated with SbC1,. These solutions were quite unstable but the kinetic behavior appeared to follow the same general pattern as in the mixed acid solutions; i.e., bands a t 4230 A and 6050 A developed rapidly and then faded. A deep blue paramagnetic complex was precipitated when DPE was treated with SbCl, in carbon tetrachloride and a solution of this complex in methylene chloride exhibited both the band a t 6050 A and esr absorption. This behavior is analogous to that reported for condensed aromatics (113) where both optical and esr data have established (111)that in the SbC1, complexes the organic moiety exists as a cation-radical. Early attempts to observe esr absorption for the cation-radical from DPE chemisorbed on silica-alumina either failed or provided very weak signals (81, 83, 114), although ion-radicals of adsorbed anthracene and perylene have been detected (115).More recently (105) the expected esr signal from chemisorbed DPE was observed and was found to be identical to that of the SbCI, complex and similar to that of the alkali metal complex (anion-radical) of this hydrocarbon. The intensity of the signal was found t o be reversibly dependent upon the presence of other adsorbed gases. Adsorbed oxygen decreased the intensity and broadened the line in a manner similar to that observed for the radicals on carbon blacks (116, 117). Hydrogen, argon, and water vapor had an opposite effect, with the latter causing the most pronounced enhancement of the signal. The influence of adsorbed oxygen on the esr signal for adsorbed DPE was opposite t o that reported by Fog0 (118) who observed an enhancement in the signal for adsorbed anthracene on exposure to the atmosphere. On the basis of the pronounced effect of water vapor, it was suggested (105)that Fogo’s observations may have been due to atmospheric moisture rather than to oxygen. A second criterion for assignment of the spectral band a t 6050 A to a cation-radical resides in demonstrating that the corresponding anionradical, having a similar spectrum, can be formed from this olefin. Anionradicals of a number of diaryl-, and tetraarylolefins have been reported from the reaction of these hydrocarbons with alkali metals in diglyme or in tetrahydrofuran (119);however, these studies did not include 1 , l diphenylethylene. I n order to obtain further confirmation of their

184

H. P. LEFTIN AND M. C. HOBSON, JR.

spectral assignment, Leftin and Hall (83) examined the spectra of tetrahydrofuran solutions of the alkali metal adducts of DPE. When a T H F solution of DPE was treated with either lithium, sodium, or potassium an intense blue color (A = 6070 A) a t first developed, then turned to a deep red (A = 4750 A) as more alkali metal reacted. The red or blue solutions were stable, depending upon whether excess metal or excess DPE was present; moreover, they were interconvertible (Fig. 42). Similar behavior has previously been reported for stilbene

3

FIG. 42. Comparisonofspectrafor the cation-radicaland anion-radicalof 1,l-diphenylethylene.

(1,2-diphenylethylene (120))and for other hydrocarbons and has been shown (119)t o be due to the formation of both a paramagnetic mononegative ion-radical and a diamagnetic dianion. I n agreement with these findings, the blue solution of DPE was found to exhibit a strong esr signal (g = 2.003 with a half width about 10 gauss) which was similar to that observed for the SbC1, complex and for the adsorbed species. The close agreement between the spectra of the anion-radical (DPE K) and of the adsorbed DPE provides convincing evidence for the assignment of the 6050 A band t o a cation-radical on the surface. Long wavelength bands due to cation-radicals were also observed when a-methylstyrene (A = 5600 A) and 1,1,4,4-tetraphenylbutediene (A = 5650 A) were adsorbed on silica-alumina. These observations, in conjunction with the demonstration of ion-radical formation from condensed aromatics (115)and from p-phenylenediamine (15)and benzidine (78), may be taken as strong evidence that this type of chemisorption is quite general for silica-alumina and other acidic catalysts. Although, in view of the foregoing discussion, the formation of

+

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

I85

cation-radicals on silica-alumina from the chemisorption of arylalkenes, condensed aromatics, and certain aromatic amines seems reasonably well established, there still exists some question concerning whether this phenomenon is truly an intrinsic property of the catalyst surface or is due to the presence of an extraneous oxidant such as chemisorbed oxygen. I n addition to the esr studies of Leftin, Hobson, and Leigh (105) which seem to dictate against the latter, it is interesting that Roberts, Barter, and Stone (121) noted the formation of a band a t 7500 A when anthracene was adsorbed on silica-alumina under anaerobic conditions. Although not so assigned by these authors, this band is probably that reported (113) for the anthracene positive ion a t 7200 A. Thus, the formation of positive ion-radicals by direct action of silica-alumina

I

500 Wovrlrnplh (mu) 400

1

600

FIG. 43. Spectra of 2-butene adsorbed on silica-alumina: curve A, low coveragedry catalyst; curve B, high coverage-hydrated catalyst; curve C, high coverage-dry catalyst; curve D, A plus heating (20).

appears possible, at least on the aromatic donors considered here. In such an adsorption process electrons would presumably be transferred to the catalyst surface. This offers no difficulty; as it has already been demonstrated that silica-alumina is capable of holding 5 x 1012 hydride ions per om2, there is no apparent reason why this surface should not also be able to accommodate an equivalent number of electrons.

186

H. P. LEFTIN AND M. C. HOBSON, JR.

Since cation-radical formation in the chemisorption of hydrocarbons has not previously been considered in the catalytic literature, the nature, reactions, and mechanism for formation of such species should be of considerable importance t o the elucidation of catalytic reaction mechanisms; particularly in view of the fact that Webb (20)has found spectral evidence for the formation of species other than carbonium ions from butene-2 adsorbed on silica-alumina. It is not possible a t the present time to define either the role of cation-radicals in acid catalysis or the chemical nature of the electrophilic surface sites involved in their formation. c. Aliphatic Hydrocarbons. The two reported investigations of the electronic spectra of aliphatic hydrocarbons chemisorbed on silicaalumina provided results which were contradictory. Webb (20)reported that the electronic spectrum of chemisorbed 2-butene (Fig. 43) consisted of one or more bands depending upon surface coverage, catalyst water content, and heat treatment subsequent t o adsorption. The

FIG. 44. Spectrum of 1-butene chemisorbed on silica-alumina DSA-1: curve A, pumped 1 hr a t 75°C; curve B, heated 1 hr a t 150°C; curve C, heated 1 hr at 200°C; curve D; heated 1 hr at 250°C.

short wavelength band, whose maximum appeared to vary between 3100-3300 A, was observed only a t low surface coverages. Additional bands a t 3800 A, 4650 A, and 5600 A were obtained on heating the sample in a closed cell a t 100°C for 30 min but were not obtained when the sample was simultaneously heated and pumped. The 3100-3300 A band was assigned to a tertiary carbonium ion of a butylene polymer, and the long wavelength bands to a pi-complex or charge-transfer complex on the basis of a similar assignment (now discredited) for adsorbed diphenylethylene.

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

187

I n our studies (122) of chemisorbed aliphatic compounds we were unable to produce the long wavelength bands in the manner described above. The ultraviolet spectrum observed when silica-alumina was exposed t o 2.5 mm of 1-butene for 5 min a t 75°C and subsequently evacuated a t this temperature for one hour (Fig. 44) consisted of a single intense band a t 3000 A. This probably corresponds to the 3100 A-3300 A band observed with 2-butene. Heating and pumping caused an increase in intensity. Further heating in a closed cell for one-hour periods a t 15OoC, 200"C, and 250°C failed to produce the supplementary bands, and it was evident that a time-invariant spectrum had been obtained. I n considering the source of the discrepancy between this work and Webb's, the observations of Jurinak and Volman (123) provided a key. These workers observed that when a catalyst containing a chemisorbed

FIG. 45. The effect of atmospheric contamination on the spectrum of 1-butene chemisorbed on silica-alumina: curve A, 1-butene chemisorbed on silica-alumina; curve B, after exposure to the atmosphere at room temperature.

hydrocarbon was exposed to the atmosphere, a visible lavender color was produced on its surface. Subsequently, it was found that when the sample used to obtain the data shown in Fig. 44 was exposed to the atmosphere a lavender color appeared and the spectrum (Fig. 45) consisted of bands at 3800 A, 4650 A and 5600 A and was closely similar to that observed by M7ebb (Fig. 43). Thus it was concluded that the earlier resuMs were influenced by atmospheric contamination and that the assignment of these bands to ;t pi-complex must be questioned further. It is interesting to note in passing that although the source of these bands is not yet certain, their origin may be similar to that of the long wavelength bands in the diphenylethylene system. Anion-radicals

188

H. P. LEFTIN AND M. C. HOBSON, JR.

(124, 125, 126) have recently been reported for aliphatic systems containing conjugated unsaturation; perhaps cation-radicals can also exist in such systems. Chemisorption of a variety of aliphatic hydrocarbons on silicaalumina gave spectra which were almost indistinguishable and which consisted essentially of a single intense band in the 3000 A region. Figure 46 shows the spectrum of isobutylene and of isobutane. Similar

WoWmqth (A1

FIG. 46. Spectrum of isobutane (---) silica-alumina.

and of isobutylene (-)

chemisorbed on

spectra were obtained (122) with 1,4-pentadiene, and with primary-, secondary-, tertiary- and isobutyl-alcohols. I n all of these spectra the presence of a second band at shorter wavelength may be indicated by the sharply increasing absorbance below 2400 A, however, this region was too close to the short wavelength cutoff of the catalyst samples to permit any firm conclusions concerning such bands. By analogy with the results obtained with the phenyl-substituted analogues, chemisorption of isobutane might be expected to proceed by hydride ion abstraction to form a tertiary-butyl carbonium ion; and it would indeed be convenient to assign the 3000 A band to this ion. Although such alkyl carbonium ions have been useful for the mechanistic description of a variety of reactions of hydrocarbons over heterogeneous acid catalysts, in recent years numerous observations have been reported which apparently are not in accord with this concept. Notable in this respect is the remarkable degree of stereoselectivity associated with the acid-catalyzed double bond isomerization of butenes

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

189

(127,128, 129). This, and other phenomena (130),has led to the postulation of alternative intermediates of a highly speculative nature. Unambiguous spectral assignment of the 3000 A band must be consistent with all chemical, physical, and theoretical data concerning systems in which it occurs. Assignment of this band to a tertiary-alkyl carbonium ion is not entirely consistent with all of the available literature concerning the properties of sulfuric acid solutions of hydrocarbons; nor can i t account for the stereoselective butene isomerization. Bands a t 3000 f.50 A have been reported for sulfuric acid solutions of numerous primary-, secondary-, and tcrtiary-aliphatic alcohols (131, 132); alkyl halides (132),normal and branched chain olefins (133);diolefins, diols, and dihalides (134);and some tertiary hydrocarbons (135).Molar extinction coefficients were remarkably independent of solute structure and were generally of the order of 3000. On the basis of their work with aliphatic alcohols, halides, and hydrocarbons, Rosenbaum and Symons (132)assigned the 3000 A band to tertiary-alkyl carbonium ions. However, these workers (134)later assigned the 3000 A band observed from diolefins, diols, and allylic halides to substituted allylic (alkenyl) carbonium ions. Since it is very unlikely that both saturated alkyl ions and unsaturated alkenyl ions would exhibit their lowest lying electronic transitions at exactly the same frequency, it seems reasonable to assume that while all of these solutions contained different species, they must have had in common a structurally similar chromophore. The possibility that this chromophore was a tertiary-alkyl carbonium ion formed in a rapid isomerization of the other ion precursors in the acid solution cannot be correct for the following reasons: (a) Cryoscopic data for alcohols in sulfuric acid indicate a process more complex than that expected for ionization of these compounds as secondary bases (136); (b)evolution of SO, from sulfuric acid solutions of octene-1 was observed to closely parallel the appearance of the 3000 A band in this system (137), indicating that an oxidation process was involved in forming the chromophore from olefins; (c) the lowest electronic transition for the isoelectronic alkylborines occurs in the 2000 A-2200 A region (138). All of these data may be reconciled with the formation of substituted allylic (alkenyl) carbonium ions. Structures of this type can be derived from any substituted olefin by removal of a hydride ion from a carbon atom adjacent to the double bond (m-hydrogen).Since an olefin may be derived by proton removal from an alkyl carbonium ion, a substituted alkenyl ion may be derived from any alkyl carbonium ion precursor containing a t least four carbon atoms and hence the spectra of their solutions should be very similar, Sulfur dioxide evolution from sulfuric acid solutions of olefins would accompany the formation of the alkenyl

H. P. LEFTIN AND M. C. HOBSON, JR.

190

carbonium ion, since the removal of an a-hydrogen from an olefin must involve an oxidative attack by the sulfuric acid similar, if not identical, to that involved with tertiary hydrocarbons; viz., HzC=CH-CH&Ha

-t 4HaS0, ---+ +

HzC=CHCHCHa

+ SO, + 2H,O+ + 3HSO;

It is interesting to note that the cryoscopic factor of 6 for this reaction is close to the factor of 5.6 obtained (136)for t-butyl alcohol, indicating that with this compound a similar reaction may have occurred subsequent to dehydration of the alcohol. Leftin and Chmeil (139) considered the effecta of solute structure on the rate of formation of the 3000 A band in concentrated sulfuric acid solutions. I n order to avoid secondary reactions, due to the presence of unreacted or undissolved compounds, measurements were restricted to the concentration range between 10-4 and 2 x 10-3 M. Figure 47

-

4

0

5

0I 6 Minutes

0I

J

I

FIG. 47. Influence of solute structure on the rate of formation of the 3000 A absorption band in 98.3% sulfuric acid at 25°C. Solute concentration = 1.25 X moles/liter. Curve A, isoprene; curve B, diisobutylene; curve C, 1-hexene; curve D, tert-butanol; curve E, sec-butanol; curve P, isobutanol; curve G, n-butanol; curve H, isooctane; curve J, cyclohexane.

shows the over-all time dependence of the concentration of the 3000 A species for several types of precursor at 25°C in 98.3% sulfuric acid. The relative rates of appearance of the band depended on precursor structure in the order: diene > olefin > alcohol > paraffin. Among the alcohols, a small structural dependence was observed such that: tertiary > secondary > primary. That formation of the band from an olefin waa not instantaneous, as might be expected if the band were due to an alkyl carbonium ion formed by protonation of the double bond, was demonstrated (Fig. 48) for 1-hexene a t intermediate acid strengths. I n 85% sulfuric acid the rate of reaction of olefin was quite low and was

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

191

very markedly enhanced by the addition of a trace of strong oxidant (selenic acid) indicating that an oxidation process was involved in formation of the chromophore in this case. These observations can be accommodated if the absorbing species is an alkenyl carbonium ion formed according to the following reactions :

Protonation of diolejins

Oxidation of olejins RCH,CH=CH,

8

+ 4H2S04--+

+ SO, + 2H3O+ + 3HSO;

RCH-CH=CHa

(12b)

Dehydration of an alcohol followed by oxidation

+ HaSO4 -+ CHsCHzCH=CHz + H 3 0 + + HSO, + + 4H,S04 -+ CH, CH-CH=CH, + SO, + 2H30+ + 3HSO;

( 1 ) CH3CHaCH,CH,0H

(2) CH3CHzCH=CHz

Oxidation of parafins (1) RCHzCHCH,

+ 4HZSO4

+

(2) R C H , ~ C H , HSO;

I

CH3 (3) RCHaC=CHa

I

CH3

+ 4H,S04

-+

+

+

+ SO, + 2H,O+ + 3HSOa

RCH8C--CH3

I

CH3 RCH,C=CH,

I

+ H,SO, (1.w

CH3 -+

(12~)

R&H-C=CH,

I

+ SO, + 2H30+ + 3HSO-6

CH3

According to these reactions, the relative rates of formation of alkenyl carbonium ions should depend on solute structure in the observed order. Moreover, both the rates and stoichiometry of SO, evolution should show structural dependence. Attempts to measure the rates of SO, evolution from sulfuric acid solutions of hydrocarbons were without success due to the large solubility of the gas. The stoichiometry of these reactions, determined by freezing out the SO, from the solutions after the reactions were completed (Table VI), agreed closely with that predicted for alkenyl carbonium ion formation according to Eq. (12). I n the preceeding discussion data and arguments are presented which provide support for assigning the 3000 A band in sulfuric acid solutions to an alkenyl carbonium ion. It seems reasonable to expect that this would also he a valid assignment for the 3000 A band observed on the catalyst surface; however, in this case it would not be consistent with the infrared spectra reported for adsorbed olefins. Webb (20) cited a

192

H. P. LEFTIN AND M. C. HOBSON, JR.

private communication from Eischens and Pliskin to the effect that the infrared spectrum of isobutylene chemisorbed on silica-alumina was fully consistent with a saturated dimeric or polymeric species. In considering how this can be reconciled with the unsaturated alkenyl carbonium ion dictated by the electronic spectra, it was first necessary to TABLE VI

SulfurDioxide Production at 26°C x lo-* molea/litre of solute)

(98.2% H,SO,, 1

SO, (moles/mole Solute)

Solute

0.0

Isoprene 1-0ctene Diisobutylene t-Butanol a-Butanol Isooctane

0.05

0.96 f 0.05 1.1 1.06 1.1 1.9 & 0.1

ascertain that both the infrared and ultraviolet spectra were indeed due t o the same species. This was established (122)by scanning the spectrum in both regions from the same catalyst preparation under identical experimental conditions using a special cell (Fig. 5) designed for this

0

1

0

2

0

3 0 4 Minutes

0

5

0

6

0

FIG.48. Effect of selenic acid on the rate of formation of the 3000 A band from 1hexene in 84.8% sulfuric acid at 26°C. Solute concentration = 1.60 X 10-8 moles/liter.

purpose. The infrared spectrum corresponding to the 3000 A species for 1-butene on silica-alumina (Fig. 49) consisted of bands at 2960 cm-1 and 2925 cm-1 apparently due to unsymmetrical C-H stretching of methyl and methylene groups (140),respectively; and a less intense

SPECTROPHOTOMETRY FOR CATALYTIC S Y S T E M S

193

band in the 2880 cm-l region which may be due to an unresolved combination of C-H frequencies. This spectrum indeed appeared to be consistent with a fully saturated species; particularly in view of the absence of any appreciable absorbance at frequencies greater than 3000 cm-l normally expected for the C-H stretching associated with olefinic species. Absence of bands in this region was not due to perturbations caused by polar adsorption forces since a distinct band a t 3070 cm-l was observed from 1-butene adsorbed on nonacidic silica gel. On the basis of frequencies assigned from studies involving neutral molecules, these infrared data may be interpreted to indicate that the adsorbed species is a saturated alkyl carbonium ion. Such an interpretation may not be valid, however, since it neglects entirely the effect of the positive charge on the stretching frequencies of C-H bonds in carbonium ions. Since the dipole moments of the bonds formed with charged sp2 carbonium carbon must differ appreciably from those formed with carbon in a normal sp2 state, the positive charge on carbon will cause the C-H stretching to be shifted to lower frequencies. I n a n alkenyl carbonium ion the charge is delocalized over three carbon atoms due to resonance involving two olefinic structures, H H

I I RC=C-CHa d

H H t+

I

I

RC-C=CH,,

e

so that in the hybrid species a discreet double bond does not exist. Therefore, even if we neglect the effect of the positive charge, the stretching bands would be shifted to lower frequencies due solely to delocalization of the double bond. Adding a further shift due to the positive charge makes it entirely possible for the expected olefinic band to be shifted into the region normally attributed to saturated species. Unfortunately it is impossible to estimate these shifts theoretically in any convincing manner, and of course truly reliable interpretations can only be obtained by direct comparison with known spectra of suitable reference compounds containing charged carbon atoms. At the present time infrared spectra of known aliphatic carbonium ions are not available in the literature; however, the recent investigations of Baughan and co-workers (141,142) provided a direct route for obtaining such data, They demonstrated, by cryoscopic and conductivity studies, that stable solutions of aliphatic carbonium ions could be obtained in fused antimony trichloride via ionization of aliphatic chlorides. Conveniently, this solvent is an excellent medium for infrared spectroscopy. In a preliminary study (103) the infrared spectra of the butenyl

194

H. P. LEFTIN AND M. C. HOBSON, J R .

carbonium ion and the secondary-butyl carbonium ion were determined from antimony trichloride solutions of 3-chloro-1-butene and 2-chlorobutane, respectively. The spectra (Fig. 50) differed from each other, as was expected. Of greater significance, however, is the fact that the spectrum of the butenyl carbonium ion (from 3-chloro-1-butene) was found to agree very closely with that observed from chemisorbed 1butene, both in location of the bands and also in respect to the absence of bands above 3000 cm-l. Although the corresponding ultraviolet spectra were not obtained in these systems, due to solvent cutoff, it was noted that the solution containing the butenyl carbonium ion was amber (visually similar to the color of dilute solutions of hydrocarbons in concentrated sulfuric acid) while that containing the butyl carbonium ion was water white. Although much further work will be needed to fully define these systems, these preliminary results appear to support a reinterpretation of the infrared spectrum of chemisorbed 1-butene, and also indicate that considerable caution should be exercized in interpreting infrared spectra of adsorbed charged species on the basis of assignments derived from neutral molecules. Provided the spectral assignment made in the present work is valid, the overall reaction involved in the chemisorption of 1-butene may be described as follows: +

+

@I-

Here a surface Lewis acid (denoted by U)abstracts a hydride ion from the methylene group adjacent to the double bond. This mechanism is in accord with the essential Lewis acid nature of the silica-alumina surface and is consistent with the previously demonstrated ability of this surface t o abstract hydride ions from tertiary hydrocarbons. Since an alkenyl carbonium ion is stabilized by resonance to a greater extent than is a saturated carbonium ion, i t may well be the most stable species which could form in the chemisorption of an aliphatic olefin or its precursor. It seems reasonable, therefore, to presume that such species may be involved in heterogeneous acid catalysis t o a greater extent than has been generally recognized. This chemisorption process does not, of course, exclude the more conventional acid addition to the double bond which may occur under suitable circumstances; but rather, it introduces an alternate path which may well exert a considerable influence on the overall course of catalytic reactions. Thus, for example, since a substituted allylic carbonium ion may be converted to a conjugated diene by loss of a proton, it may be an important intermediate in the formation

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

195

of “coke” on catalyst surfaces. Since several reactions over acidic catalysts have been reported for which saturated carbonium ions cannot offer an adequate explanation, it may be well to consider some of the ways in which this mode of olefin chemisorption could fit into the picture of heterogeneous catalysis.

0.2 -

0.3 -

I

,

I

1

I

It is possible to account for the stereoselectivity in the double bond isomerization of 1-butene over silica-alumina and other acid catalysts by assuming that the chemisorption leads to a butenyl carbonium ion (Fig. 51). According t o this scheme, 1-butene can be chemisorbed to form either a cis (path A) or a trans (path B) butenyl carbonium ion depending upon the conformation of the butene molecule as it approaches the catalyst surface. Since interconversion of the cis and trans forms of this ion is inhibited by the higher bond order of the 2-3 carbon-carbon bond in the resonance hybrid (143),products formed in reactions involving this ion will have the same geometric configuration as that of the ionic intermediate. This conclusion finds support from the observed retention of geometric configuration in the hydrolysis of cis-crotyl chloride (144). If kinetic control of the reaction resides in the chemisorption step, then the observed cis-trans ratios imply that for the chemisorption of 1-butene 13, > k,. Evidently, the steric influence of the methyl group in the

196

H. P. LEFTIN AND M. C. HOBSON, JR.

trans conformation effectively interferes with the approach of the allylic hydrogen to the surface. Another feature of the reaction which must be considered is that the stereoselectivity decreases at high conversion levels. Although there may be several explanations of this effect, it can be accommodated by the present mechanism on the basis of a slow, k,, interconversion of the ionic configurations. 0

-.-.-_ A

.'

.-*+---

\ 0.1

-

/

i .i ,/

em-'

FIG. 50. Vibrational spectra of carbonium ions. Curve A, sec-butyl carbonium ion in fused SbC1, a t 80°C (generated from 2-chlorobutane); curve B, butenyl carboniurn ion in fused SbC1, a t 80°C (generated from 3-chloro-I-butene);curve C, carbonium ion fornied in the chemisorption of 1-butene on silica-alumina.

On this basis, then, it appears that the stereoselectivity of olefin isomerization can be adequately explained within the framework of classical carbonium ion theory without invoking such dubious species as the pi-complex or cyclic carbonium ion. The attractiveness of this mechanism is further enhanced by the fact that it is similar to the allyl carbanion mechanism proposed for the base catalyzed isomerization (145). Perhaps it would not be too great an extrapolation to presume that a similar mechanism, involving allyl radicals, may obtain over reduced metal catalysts.

SPECTROPHOTOMETRY FOR CATALYTIC SYSTEMS

197

Condon (130),in discussing the formation of 2,5-dimethylhexane from the reaction of isobutane with AlCl,, postulates a protonated cyclopropane intermediate to explain the skeletal isomerization of the 2,2,4-trimethylpentyl ion, which he presumes was formed by attack of t-butyl ion on isobutylene. I n this reaction the isobutylene was produced by removal of a proton from the initially formed t-butyl ion. The cyclic intermediate can be avoided if part of this isobutylene was first converted to butenyl ions which then added to the remaining isobutylene t o provide the 2,5-dimethylhexyl skeleton directly. A

(qouchrl

B

(trans)

ti (cis)

FIQ.51. Proposed mechanistic scheme for the stereoselective double bond isomerization of 1-butene over silica-alumina.

It would be beyond the scope of this article to attempt to explore all possible applications of alkenyl carbonium ions to catalytic reaction mechanisms. The several examples which have been discussed serve to illustrate the potential utility of optical absorption spectroscopy to problems in heterogeneous catalysis. ACRNOWLEDQY~NT The authors are very pleased to express their appreciation to Dr. Heinz Heinemann for his interest and encouragement, to Dr. Ernest Solomon for helpful suggestions, and to Miss Mathilda Black for her review of the manuscript. One of us (M.C.H.) is, in part, indebted to the laboratory for Research on the Structure of Matter, University of Pennsylvania, supported by the Advanced Research Projects Agency of the Department of Defense for financial support.

198

H. P. LEFTIN AND M. C. HOBSON, JR.

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134. Rosenbaum, J., and Symons, M. C. R . , J . Chem. SOC.p. 1 (1961). 135. Lavrushin, V. F., Kursanov, D. N., and Setkina, V. N., Doklady Akad. NaukS.S.S. R . 97, 256 (1954). 136. Newman, M. S., Craig, R. A., and Garrett, A. B., J . Am. Chem. SOC. 71, 869 (1949). 137. Gonzalez-Vidal, J., Kohn, E., and Matsen, F. A., J . Chem. Phya. 25, 181 (1956). 138. Davies, A. G., Hare, D. G., and Larkworthy, L. F., Chem & Ind. (London) 48, 1519 (1959). 139. Leftin, H. P., and Chmeil, J., Unpublished research, M. W. Kellogg Co., 1961. 140. Bellamy, L. J., “The Infra-red Spectra of Complex Molecules,” Chapters 1-3. Wiley, New York, 1958. 141. Porter, G. B., and Baughan, E. C., J. Chem. SOC.p. 744 (1958). 142. Davies, A. G., and Baughan, E. C., J. Chem. SOC.p. 1711 (1961). 143. Eliel, E. L., i n “Sterio Effects in Organic Chemistry” (M. S. Newman, ed.),pp. 94,95. Wiley, New York, 1956. 144. Young, W. G., Sherman, S. H., and Winstein, S.,J. Am. Chem.Soc. 82, 1376 (1960). 145. H a g , W. O., and Pines, H., J. Am. Chem. SOC.82,387 (1960).

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Hydrogenation of Pyridines and Quinolines MORRIS FREIFELDER Organic Chemistry Department. Research Diwiawn. Abbott Laboratories. North Chicago. Illinoia

. .

I Introduction .................................................... I1 Catalysts ........................................................ I11. Occurrence of Side Reactions ......................................

IV . V VI VII . VIII I X. X XI XI1. XI11. XIV

. . . . .

. XVI .

xv .

. . xx.

XVII XVIII . XIX X X I. X X I I. XXIII XXIV .

. xxv.

203 204 206 Catalyst Poisoning ................................................ 208 Reduction of Pyridine Salts ........................................ 209 Effect of Substituents on Hydrogenation ............................ 210 Pyridinecarboxylic Acids .......................................... 213 Pyridylalkanoic Acids ............................................ 216 Esters and Amides in the Pyridine Series ............................ 216 Hydroxypyridines and Derivatives ................................. 217 Pyridylalkanols .................................................. 220 Aminopyridines .................................................. 220 Pyridylalkylamines ............................................... 222 Dipyridyls and Related Compounds ................................ 223 Quaternary Compounds ........................................... 224 Some Selective Reductions. ....................................... 226 227 A Reduction in the Presence of Ketone Function .................... B. Reduction in the Presence of Nitrile Function ..................... 234 C. Reduction in the Presence of a Benzenoid Ring .................... 238 D . Selectivity in the Presence of Other Heterocycles .................. 237 Quinolines ...................................................... 238 Isoquinolines .................................................... 239 Effect of Substituents ............................................ 241 Quinoline and Isoquinolinecarboxylic Acids .......................... 242 Esters and Amides ............................................... 243 Hydroxyquinolines and Related Compounds ........................ 244 Amino Compounds ............................................... 245 Quinolylketones .................................................. 246 Concluding Remarks ............................................. 248 References ...................................................... 248

.

1. Introduction Catalytic hydrogenation of the pyridinoid nucleus has been carried out under a variety of reaction conditions by many investigators in the past fifty years . Unfortunately. much of the investigation has been conducted on the basis of superiority of one catalyst over another while 203

204

MORRIS FREIFELDER

ignoring the effect of substituents on the rate and extent of reaction. Another factor contributing to the gaps in knowledge is that a number of the reports on ring reduction are isolated preparations of piperidines needed only as intermediates. Only by piecing together bits of information in a review can we see whether there is a need for further study of catalytic hydrogenation of pyridines and related compounds.

II. Catalysts The discovery by Sabatier (1)of the reducing power of finely divided nickel, although unsuccessful in the reduction of pyridine to piperidine, led to the investigation of other nickel catalysts (prepared from nickel salts or oxides) for the same purpose. The development of Raney nickel catalyst (2) gave to the chemist a more active form of this metal. Its activity in the hydrogenation of the pyridine ring was first studied by Adkins (3) and his co-workers at the University of Wisconsin. Since this investigation, it has enjoyed wide use in the catalytic reduction of many pyridines. I n general, nickel in its various forms requires elevated temperature and pressure conditions for the catalytic reduction of pyridines. Hydrogenations with nickel on keiselguhr (4)or nickel chromite (5), for example, employ similar rigorous conditions. Copper chromite ( 6 ) (copper chromium oxide) has also been investigated. With this catalyst temperature conditions are usually higher than with nickel catalysts. There is a report of the use of a palladium catalyst in the reduction of some 2-(/3hydroxyalkyl) pyridines at 130" and 200 atmospheres pressure (7). Recently hydrogenation of pyridines has been reported with ruthenium dioxide (8). While high pressure is necessary, reaction conditions are much milder than with other catalysts used at superatmospheric pressure, and reaction time at 90" and 70 atm pressure is generally short. The temperatures necessary for successful conversion of many pyridines under high pressure conditions with most catalysts (except ruthenium) also favor side reactions. Indeed, in the hydrogenation of pyridine with copper chromite at 220" arid 100-150 atm ( 6 )only 50% of piperidine was obtained along with an equal amount of a high boiling condensation product. Reduction, therefore, under more moderate conditions is desirahle. The work of Skita with colloidal platinum (9) led the way toward hydrogenation of the pyridine ring under low pressure conditions. The use of Adams catalyst (lo),of far greater value, gave further impetus to the investigation of low pressure conversions of the ring. The latter

HYDROGENATION OF PYRIDINES AND QUINOLINES

205

catalyst, despite certain disadvantages inherent in the conditions necessary for reduction, has been widely used up to the present day. A highly active form of Raney nickel has been described and its use noted in a low pressure hydrogenation of the pyridine ring of quinoline (11).2-Dodecylpyridine and 2-pentadecylpyridine have also been converted to the corresponding piperidines with this active catalyst at low pressure (12).Unfortunately it does not retain its high activity for too long a period and as a result has not found widespread use. There is another report of a reduction of the pyridine nucleus with normally active Raney nickel under a few atmospheres pressure, but as much catalyst as compound is used (13). There is considerable literature on nickel catalysts promoted with platinum, but these have not been applied to the hydrogenation of the pyridine ring under low pressure or high pressure conditions. The advent of rhodium on a support ( 1 4 )led to its use by other workers from the same laboratory who reduced pyridine to piperidine using an equal weight of 5% rhodium on alumina (15).Investigation of rhodium on carbon in the hydrogenation of a number of pyridines (16)showed it to be effective in a neutral medium when the catalyst ratio is high enough. Each catalyst has advantages and disadvantages. Reduction of pyridine compounds with nickel requires conditions which lead to side reactions, although there are circumstances where some nickel catalysts can be used to advantage. Platinum oxide, except for isolated instances, is only effective under acidic Conditions (or when the pyridine nitrogen is neutralized, as a hydrochloride salt for example). Rhodium on a carrier, probably the best catalyst for use at low pressure, is effective in neutral solvent when the ratio of metal to compound is 1 to 2%. It is, however, adversely affected by strong nitrogen bases. As a catalyst in the presence of an acid it is more active when the acid is an organic one. Ruthenium dioxide, probably the most active catalyst for conversion of the pyridine ring, may be the least selective, particularly in the presence of a benzene ring, Copper chromite, which requires higher temperature conditions than most catalysts for hydrogenation in general, is of questionable value for this system. There are too few reports on the use of palladium catalysts in the reduction of pyridines to draw any concrete conclusions. There is a report (17) which shows that it is capable of promoting ring reduction at high pressure. Another report (18) indicates that it may be more useful than platinum oxide when hydrogenation of certain pyridines is carried out in acetic acid at 70-80' and 3-4 atm pressure. It has been found (19)that pyridine can be converted to piperidineacetate by reduction in alcohol with pal-

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ladium on carbon at 60" and less than 3 atm pressure, if an equivalent of acetic acid is present.

I l l , Occurrence of Side Reactions The main side reactions occurring during the catalytic hydrogenation of pyridines consist of the following: (1) condensation, (2) rupture of the carbon nitrogen bond, and (3) alkylation of the piperidino nitrogen. Other reactions may also occur, such as hydrogenolysis of certain groups or reduction of benzene rings, but only the above three will be covered in this section. Self condensation of products of reduction or interaction between resultant piperidines and starting materials or intermediates in the course of hydrogenation may take place under moderate conditions, but are favored under more forceful conditions. Adkins (20) points out that conditions necessary for successful conversion of many pyridines with nickel catalysts also favor side reactions ( l ) ,(2), and (3). They are noted with other catalysts where temperature and pressure conditions are high. Rupture of the carbon nitrogen bond is probably an effect of temperature, but may also be dependent on the catalyst. Sabatier (I)reported that in passing pyridine and hydrogen over nickel, amylamine was obtained. It is likely that piperidine was the intermediate and that decomposition or cracking occurred as the temperature rose.

The experiments of Zelinsky and Borisov (21) appear to substantiate that piperidine is the intermediate. A high yield of it was obtained in a flow type hydrogenation of pyridine in the presence of palladized or platinized asbestos, A similar flow type reaction with the same catalysts (22) gave not only piperidine, but side products which the investigators suggest might be 1,2,3,4-tetrahydropyridineand 1,2,5,6-tetrahydropyridine (2-piperideine). The observation, however, needs further confirmation. Reinvestigation of the products of reaction with modern

HYDROGENATION O F PYRIDINES AND QUINOLINES

207

analytical tools could give an insight to the mechanism by which ring reduction takes place. I n further study with other catalysts, Sadikov and Mikhailov (23) report isolation of two condensation products in addition to piperidine. One of these is 1-methyl-3-propylpiperidine(N-methylconiine), the 1'-dipiperidyl (C16H3,,Nz). The authors' other 2,2'-dimethyl-3-propyl-l, suggested reaction scheme to obtain these products is based on conversion of pyridine to piperidine and subsequent rupture of the piperidine ring. The latter of the two products may have also been obtained in the hydrogenation of pyridine with reduced copper a t 220" and 160atm (24). I n this instance 50% of piperidine was obtained in addition to a product with the same empirical formula. The succession of reactions that follows is probably the pathway by which N-alkylation takes place during the high temperature hydrogenation of pyridines in alcohols: (a) RCH,OH

(b) RCHO

-H , 1 ,RCHO

3 3-5 =I>+

+ HN

-+

R Y - - N 3

OH

(c)

RCII-N

I

OH

RCH,--N

A II,O

1%

R = H or ally1

Reaction (a)is, of course, reversible. However, under high temperature in the presence of hydrogenation catalysts, dehydrogenation is favored. Addition in reaction (b) then takes place. Reaction (c) follows yielding B by hydrogenolysis of the - OH group. An example of the occurrence of N-alkylation is shown in the following experiment.

l-EthyZ-2,3-dirnethyZpiperidine.A solution of 107 gm (1.0 mole of 2,3-dimethylpyridine in ethyl alcohol was hydrogenated a t 150°C and 200 atm pressure in the presence of 20 gm of Raney nickel. After hydrogenation was complete (12-15 hr) the cooled reaction mixture was filtered from the catalyst. The solution was concentrated and the residue distilled. A fraction was collected a t 155-165" (750 mm). It weighed 44 gm and corresponded to 33% of N-alkylated product. Elemental analysis confirmed its structure.

208

MORRIS FREIFELDER

N-alkylation has been observed with the use of nickel catalyst which has been stored in ethyl alcohol (25). A hydrogenation of pyridine a t 200" and 200 atm with such a catalyst gave 9Ol0 of 1-ethylpiperidine in addition to the desired product. Sawa (26)points out that 1-ethylpiperidine is always obtained in the high pressure hydrogenation of pyridine in ethyl alcohol with nickel catalyst. He makes use of this propensity toward N-alkylation in hydrogenating pyridine with nickel on kieselguhr in the presence of a variety of alcohols to obtain good yields of 1-alkylpiperidines from c, to C*@ N-alkylation is not necessarily limited to nickel catalysts. 1-Butylpiperidine wm obtained in a reduction of pyridine in n-butyl alcohol in the presence of copper chromite (27, a ) . However, when hydrogenation was carried out in the presence of tertiary amyl alcohol, only piperidine was obtained. Little N-alkylation occurs during hydrogenation of pyridine in the presence of polyols (27, b). Substituted pyridines undergo little Nalkylation in the presence of a branched chain alcohol. An example of this is seen in the hydrogenation of 2-picoline in isopropyl alcohol a t 230-240". Only 7.3% of 1-isopropyl-2-pipecolinewas obtained (28, a ) . With more sterically hindered pyridines chain length of the alcohol is a factor (28, b ) . I n the hydrogenation of 2,6-lutidine in various alcohols N-methylation occurred to the extent of 56.6% a t 250°C. A dramatic drop occurred when ethyl alcohol was used giving only 19% of the 1-ethyl derivative. Less than 10% N-alkylation took place with npropyl and n-butyl alcohols. While N-alkylation usually takes place a t elevated temperatures, there is a possibility that it can occur at lower temperature. This is suggested by literature reports of the nickel catalyzed alkylation of aromatic amines in refluxing alcohols (29). The mechanism of action is similar to the scheme described on a previous page of this article. However, it has not been observed with the moderate temperatures used in the high pressure hydrogenation of pyridines in the presence of ruthenium catalyst (16),nor in the reductions a t 50-60" with rhodium catalyst at low pressure (8).

IV. Catalyst Poisoning Maxted in his study on the poisoning of metallic catalysts (30) has suggested that this phenomenon occurs because the unshielded structure of the poisoning atom allows strong bonding (chemisorption) between it and the surface of the catalyst. I n another publication (31)

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209

he has shown that pure dry pyridine is capable of acting as a poison toward platinum catalyst in the hydrogenation of cyclohexene. Since it has the ability to inhibit other reductions, he reasons it would be expected to be self-poisoning in its own reduction. Indeed, in the hydrogenation of pyridine in a neutral solvent in the presence of platinum oxide uptake of hydrogen is extremely slow or does not take place a t all (10). Maxted has stated that the slow hydrogenation rate of pyridine is due to the obstructively long adsorbed life of piperidine on the catalyst surface. This has been suggested by others. However, no proof has been offered as to which of these compounds causes poisoning. Each contains a nitrogen atom with an unshared pair of electrons and therefore either, according to Maxted’s theory, is capable of inhibiting reduction. Accordingly, a modification of the experiment of Maxted and Walker (31)was carried out in this laboratory in an attempt to find the answer.

Experimental. A solution of 8.2 gm (0.1 mole) of cyclohexene in 50 cc of absolute ethyl alcohol was hydrogenated under 2 atm pressure in the presence of 0.1 gm of platinum oxide. Uptake of hydrogen was complete in twenty minutes. This was taken as the standard experiment. The same low pressure Parr hydrogenation apparatus was used throughout. The equipment was thoroughly cleaned after each run and all reductions were run in duplicate. The same batch of catalyst, solvent and cyclohexene was used in each experiment. I n a second set of experiments 1 cc of pure distilled pyridine was added to the standard hydrogenation experiment. Uptake of hydrogen was observed. I n 20 min it was 40% of theory, a t 30 min 52%, at 55 min 60%. It was 62.5% of theory at 75 min and only 65% at the end of 420 min. I n the third set 1 cc of redistilled piperidine was added to the standard experiment. Uptake in 20 min was negligible. Shaking was continued for 420 min during which time no uptake of hydrogen was observed. These results indicate the piperidine is indeed the offending compound.

V. Reduction of Pyridine Salts When pyridine is reduced catalytically in the form of a salt or in acid solution, uptake of hydrogen proceeds rapidly in contrast to its reduction as base in a neutral solvent (10).This falls in line with Maxted’s concept of the toxic or unshielded nitrogen atom in both pyridine and piperidine bases and the nontoxic or shielded salt form (30). Calculation of the

210

MORRIS FREIFELDER

amount of acid used in the hydrogenation of pyridine with colloidal platinum (9) leads to the assumption that its presence is the secret of success. Maxted’s study of the poisoning effect of the pyridine and piperidine nitrogen and the subsequent shielding by acid is limited to platinum oxide. However, other metallic catalysts are similarly affected. Earlier in this article it was mentioned that it was possible to convert pyridine to piperidine a t low pressure in the presence of a palladium catalyst when an equivalent of acid was used and the reaction carried out at 60” (19).In the absence of acid no uptake of hydrogen occurred. Walker (18)attaches great importance to the reaction temperature in his report. However, in all reductions but one, acetic acid was used as solvent and it (the acid) tied up the ring nitrogen allowing hydrogen uptake to proceed. The lone reduction carried out in ethyl acetate was a 2substituted product. The use of acid medium in the hydrogenation of some pyridines with basic side chains with rhodium on alumina has been reported (32). A patent describes the preparation of various salts of piperidine from the reduction of pyridine in the presence of an equivalent of an organic or inorganic acid and rhodium catalyst (33).The acids range in strength from acetic to phosphoric and hydrochloric acids. Of interest is the more rapid reduction obtained with rhodium on alumina compared with an equivalent amount of platinum oxide (based on metal content) under the described conditions. There are disadvantages in carrying out the hydrogenation of pyridines as salts. When the corresponding piperidine bases are desired, isolation of some of them in good yield is difficult because of their extreme water solubility. The isolation of pure piperidinecarboxylic and piperidinealkanoic acids after reduction as hydrochloride salts is troublesome and often extremely difficult. The reduction of the corresponding pyridine acids now no longer requires the presence of an acid. This will be discussed in a subsequent section.

VI. Effect of Substituents on Hydrogenation Steric effects also play a part in the hydrogenation of the pyridine ring. Adkins and co-workers ( 3 ) first commented on the favorable effect of substituents in 2- and 6- positions. 2,3-Dicarbethoxypyridinewas hydrogenated at a lower temperature and in a shorter reaction period than 3-carbethoxypyridine in the presence of either Raney nickel or nickel on kieselguhr. This suggests “physical shielding” by a group or groups adjacent to the nitrogen atom. The fact that 2-methylpyridine

HYDROGENATION O F PYRIDINES AND QUINOLINES

21 1

(2-picoline) and 2-benzylpyridine are reduced under milder conditions than pyridine itself and the short reaction time necessary for the conversion of 2, 6-dicarbethoxypyridine adds further credence to the effect of 2- and 6- substituents. This effect, however, is discounted by Smith ( 3 4 ) who suggests that initial hydrogen pressure, amount of catalyst and type of solvent may have had some influence on reaction time. I n some other work by Adkins (35) there was little significant difference between the reduction of 2- and 3-methylpyridines. Similarly, hydrogenation in the presence of a nickel silica gel catalyst shows more rapid conversion for pyridine (5 hr) than for 2-methylpyridine (7 hr) (36) On the other hand, comparative reductions of pyridine and 2-methylpyridine in the presence of ruthenium dioxides (2% ratio of catalyst to compound) show that the 2-substituted compound is reduced more readily. Indeed, with a lower catalyst ratio the difference is more significant. With a 0.1 yo ratio of ruthenium dioxide catalyst 2-picoline was completely converted to 2-methylpiperidine in 20 min a t 200' and 70 atm pressure. The hydrogenation of pyridine under similar conditions took 90 min. Unpublished work from the same laboratory (37) on the reduction of 2-picoline a t room temperature and 2 atm pressure in the presence of a 20% weight ratio of 5 % rhodium on alumina shows complete hydrogen uptake in 4 hr as against 6 hr for analytical reagent redistilled pyridine under the same conditions. In view of the conflicting reports on the effect of ortho substituents on reduction of the pyridine ring, a series of hydrogenations modeled after the one in Section I V was carried out in this laboratory. It had been shown that piperidine caused immediate poisoning in the reaction. A comparison of the effect of 2-methylpiperidine should have a bearing on the validity of "physical shielding." 2,6-Dimethylpyridine and 2,6-dimethylpiperidine were also included. Compound

Time (min)

Standard Pyridine

20 20 (40% uptake) 420 (65% uptake) complete poisoning-no 28-30' 70b 20 20

Piperidine 2-Picoline 2-Methylpiperidine 2,6-Dimethylpyridine 2,6-Dimethylpiperidine

"SS%, 23 min. b60%, 10 min; 76%, 35 min; 86%, 45 min.

uptake

212

MORRIS FREIFELDER

Experimental. Cyclohexene was hydrogenated in the described manner and time of hydrogen uptake noted. Then 1 cc of piperidine was added t o a standard experiment and observation made, Thereafter in subsequent experiments 1 cc each of the remaining compounds listed in the table was added to a standard experiment and the results recorded. The results indicate that piperidine is a far more potent poison than 2-methylpiperidine. Therefore it can be assumed that as the hydrogenation of pyridine proceeds, more poisoning will occur as piperidine is formed and uptake of hydrogen should be slower than when 2-picoline is reduced. The results substantiate the protective effect of substituents adjacent to the ring nitrogen. This is even more pronounced with the 2,6-derivatives, It must be pointed out, however, that “physical shielding” takes place only when the nitrogen atom is not in an ionized form. Hydrogenations of pyridine and 2-substituted pyridines when carried out in acid solution or as salts show that pyridine is reduced first in comparison with other pyridines. For example, under the conditions used by Skita (9) pyridine was reduced more readily than 2-picoline. The same order of reducibility was observed in hydrogenations in acetic acid in the presence of Adams catalyst (38).I n another comparison of the reduction of mixed hydrochloride salts (39) i t was found that pyridine reduced first. This is further amplified in the reduction of equimolecular binary mixtures of hydrochloride salts in the presence of platinum oxide (40). The latter reference is of interest since it appears to indicate that, once the nitrogen atom is in an ionic form, not only should substituents on the ring produce steric effects when in 2-position but the size of the substituent should likewise affect the rate of hydrogenation. This is seen in the following order of hydrogenation: pyridine > 2-picoline > 2-ethylpyridine > 2-propylpyridine > 2-(2-methyl-2-hydroxypropyl)pyridine > 2-phenylpyridine. 2-( 2-Hydroxypropyl)pyridine, also mentioned, may be anomalous. It was hydrogenated at a slower rate than pyridine but its rate was equal to that of 2-picoline. The size of the constituent in 2-position may have some bearing on hydrogenation in a neutral medium. A single instance involving the comparison of the reductions of 2 4 2-hydroxyethyl) and 2-(3-hydroxypropyl) pyridine (16) shows the latter compound being converted more rapidly to the corresponding piperidine. However, many more comparisons are necessary before this can become a valid conclusion. Smith (34) has stated that as the methyl group moves away from the

RYDROC4ENATION OF PYRIDINES AND QUINOLINES

213

ring nitrogen reduction becomes increasingly difficult. This probably applies to groups larger than methyl. Not enough work has been carried out to show the difference between 3- and 4- substituents. There are, however, enough examples in the literature to show the difference between 2- and the other monosubstituents. Several are shown in the hydrogenation of substituted pyridines with rhodium on carbon (16). I n a comparison of 2- and 4-dodecylpyridine a t low pressure, 2-dodecylpiperidine was obtained while the 4-derivative did not reduce (12). The successful conversion of 2-( 2-hydroxyethy1)pyridine with Raney nickel at 150”and 130 atm ( 4 1 )and the unsuccessful attempt by another investigator with the isomeric 4-substituted product may be another example (42). An example of the effect of a 2-substituent on the reduction of the pyridine ring can be seen in the hydrogenation of the isomeric 3-(2-, 3and 4-pyridylmethyl)-2-indolones(18). Conversion of the 3- and 4isomers waa carried out in glacial acetic acid, while ethyl acetate was used for the %isomer. Despite the fact that reduction of the %isomer was carried out in a neutral medium, uptake of hydrogen proceeded readily, This is an indication of how “physical shielding” can inactivate the potential catalyst poisoning power of the unshielded nitrogen atom. The difference in reducibility of 2-substituted pyridines compared to similar 3- and 4-compound in the absence of any acidic agent is no doubt due t o physical shielding by that group. It is unfortunate that in another series ( 4 3 ) )observations of the reduction times were not made during the conversion of 2- and 4-isopropyl and 2,4-, 2,5- and 2,6-dimethylpyridines to the corresponding piperidines. The information might have added to the knowledge of the effect of substituents on hydrogenation of the ring.

VI I. Pyridi necar boxylic Acids I n one of the earliest reports among this group of compounds pyridine 2,6-dicarboxylic acid was hydrogenated with colloidal platinum (44) to give the corresponding piperidine derivative as the hydrochloride salt. Later Hess (45) investigated the reduction of the isomeric pyridine carboxylic acids in acetic acid with the same catalyst isolating the products as hydrochloride salts in good yield. Other hydrogenations in this series appeared where reaction of the hydrochloride salt was carried out in aqueous solution (46)with a high catalyst ratio (platinum black). A later report described the reduction of nicotinic acid hydrochloride in aqueous solution with 2% by weight of platinum oxide ( 4 7 ) . This procedure has been in general use up to the present day. The

214

MORRIS FREIFELDER

drawback, of course, is the problem of obtaining the free acids. Several other methods have been reported which should supplant the McElvairi and Adams process. I n each the need of an acidic agent to overcome catalyst poisoning has been eliminated and by each method the free acid is obtained. They will be discussed subsequently. Sorm ( 4 8 ) has described the hydrogenation of picolinic acid in glacial acetic acid with platinum oxide. He was able to obtain the corresponding piperidine carboxylic acid by concentration of the acetic acid solution. Since piperidine-2-carboxylic acid and acetic acid are probably of the same order of acidity, it is likely then that acetic acid is not strong enough to break open the zwitterion form or the hydrogen bonded form of the piperidine acid:

C=O

I

1-0

However, when isonicotinic acid (pyridine-4-carboxylic acid) was reduced under the same conditions the resultant piperidinecarboxylic acid was not isolated per se but as the hydrochloride salt. I n the conversion of nicotinic acid Sorm reported considerable decarboxylation yielding only 50% of nipecotic acid along with piperidine. He postulates that the hydrogenation of nicotinic acid takes two paths:

H

H B

He considers A a disguised form of a &aldehyde acid and as such readily loses carbon dioxide. When nicotinic acid was reduced in dilute hydrochloride acid only 10% of piperidine was obtained, indicating a small amount of decarboxylation. Decarboxylation has been observed during the hydrogenation of nicotinic acid in aqueous solution in the presence of ruthenium (8))rhodium (16), and platinum oxide (49).It has been

HYDROGENATION OF PYRIDINES AND QUINOLINES

215

prevented by reducing a solution of the sodium salt (8, 50). I n another report (51) not only has decarboxylation been avoided but the free piperidine-3-carboxylic has been obtained. Reduction was carried out in aqueous solution containing a slight excess of ammonia in the presence of rhodium on alumina. After the reaction was complete the free acid was obtained by concentration of the solution to dryness. Decarboxylation has also been observed in the hydrogenation of the betaine, trigonelline (48): 0

0' -

COOH

LO(-,

I

CH,

Q+lNy CH, I

AH3

The ease of direct isolation and high yields of the isomeric 2- and 4piperidinecarboxylic acids obtained after hydrogenation of the corresponding pyridine compounds with ruthenium dioxide a t 90" and 70 atm (8) and a t low pressures with rhodium (16) and platinum oxide (49) should make any of the three procedures one of choice dependent on available equipment or catalyst. The work with platinum oxide is of particular interest since it is the first report of success with it in the pyridine series, in the absence of an inhibitor of the toxic ring nitrogen. I n the absence of such an inhibitor poisoning takes place (30).Rhodium catalysts are also affected by nitrogen bases (52) but the authors theorized that the carboxyl group should neutralize the toxic effect of the ring nitrogen and allow hydrogen uptake to proceed with either catalyst. The results confirmed their reasoning. The difference in the ease of reducibility of the 2-derivative compared to the 4-acid is of interest and showed up in the hydrogenations with each catalyst. Palladium on carbon has been used successfully to obtain piperidine-2-carboxylic in almost quantitative yield after reduction of picolinic acid in aqueous solution a t 60" and 3 atm pressure (53). The time of reduction, however, is much longer than with rhodium or platinum catalyst. Since there is always a difference in reduction time between the 2- and 4-derivatives, hydrogenation of isonicotinic acid should be prohibitively long a t low pressure with palladium. However, if pressure equipment is available, reaction time should be decreased considerably. It is of interest that decarboxylation has not been observed in any hydrogenation of picolinic or isonicotinic acids.

216

MORRIS FREIFELDER

VIII. Pyridylalkanoic Acids There are only a few reports of catalytic reductions among this group of compounds. Most of the work has been confined to the #3-pyridylpropionic acids. Recently hydrogenation of 3-pyridylacetic acid in aqueous solution has been described in the presence of platinum oxide, rhodium on carbon or rhodium on alumina ( 5 4 ) .Here again, as in the reductions of picolinic and isonicotinic acids, it is shown that reaction will take place without the aid of acidic agents. The isomeric #3-piperidylpropionic acids have been obtained as salts by the platinum catalyzed hydrogenation of the corresponding pyridylacrylic acids in acetic acid (55) or dilute hydrochloric acid ( 5 6 , 5 7 ) .A few other piperidinealkanoic acids have obtained by reduction of the pyridine acids under similar conditions. Only Hall ( 5 8 ) carried out a hydrogenation which gave the free acid. He obtained an excellent yield of #3-3-piperidylpropionic acid after reduction of an aqueous solution of /3-3-pyridylacrylic acid at elevated pressure with ruthenium dioxide. The corresponding p-4-pyridylacrylic acid was reduced in aqueous ammoniacal solution with Raney nickel at 190' and 40 atm ( 5 9 ) . The product of reduction was not identified per se but was converted to the ethyl ester. Freifelder found that both js-3- and 4-pyridylpropionic acids could be reduced at low pressure in aqueous suspension with platinum oxide or rhodium on a carrier ( 5 4 ) ,but uptake of hydrogen was slow because of poor solubility. It was concluded that the piperidyl acids should act like amino acids, therefore hydrogenation of the pyridyl acids in weakly ammoniacal solution should not only lead to more rapid uptake but the free acids should be obtained after concentration. The reductions carried out in this manner (54) proved these points. 4-Piperidylacetic acid was also prepared in the same way. The catalyst used in these reductions in weakly ammoniacal solution was 5% rhodium on alumina or 5% rhodium on carbon. It is interesting, although not unexpected, that platinum oxide is ineffective under these conditions. Indeed, from the excellent yields and the simplicity of work up, this low pressure method could be a general one for the hydrogenation of not only pyridylalkanoic acids, but also any pyridinoid compound having a carboxyl group.

IX. Esters and Arnides in the Pyridine Series Hydrogenation of these compounds does not pose much of a problem. Nickel under elevated pressure and temperature has been used with success (3, 60). Platinum catalyzed reductions in the presence of an

HYDROGENATION OF PYRIDINES AND QUINOLINES

217

acid are satisfactory. Rhodium on carbon gives excellent results. Ruthenium dioxide where pressure equipment is available may well be the catalyst of choice in reduction among this group. A case in point may be seen in the hydrogenation of N, N-diethylnicotinamide. Uptake of hydrogen in the presence of rhodium on carbon at 60" and 2.7 atm was very slow (16).Reduction at 90" and 70 atm in the presence of ruthenium dioxide was complete in one hour (8). The difficulty in reducing this compound at low pressure is shown in another example (61) where uptake of hydrogen with platinum oxide in aqueous acid took 34-39 hr a t 50". Some esters may undergo cyclization if the side chain is long enough and it is in 2-position. The following is an example of such a reaction (62). COOC,H,

When the same pyridyl compound was reduced under low pressure conditions with platinum oxide in acetic acid, the intermediate piperidyl ester was obtained (63). It was, however, converted to the same ring closed compound on distillation.

X. Hydroxypyridines and Derivatives Wheland (64) describes 3-hydroxypyridine as a true phenol. Since phenols are the easiest of the substituted benzenes to hydrogenate catalytically (65), it should be of interest to examine some of the reductions of 3-hydroxypyridine for comparison. Cavallito (66) reports that neither it nor 3-benzoxypyridine was hydrogenated in the presence of Willstatter's palladium sponge catalyst (67). Under low pressure conditions in ether or dioxane Raney nickel and platinum oxide were ineffective. However, other examples show that reduction takes place readily under a variety of conditions. Biel used Raney nickel at 125" and 50 atm (68),excellent yield of 6-propyl3-hydroxypiperidine resulted from reduction of the pyridine in acetic acid with platinum oxide (69). Ruthenium in the conversion of 3hydroxypyridine in aqueous solution gave very high yield of the corre-

218

MORRIS FREIFELDER

sponding piperidine (70). Of the catalysts used in these reductions ruthenium is probably the most useful. Since rhodium catalysts have been shown t o be effective in pyridine reductions (16),it would be of interest to study their use in this group of compounds. I n the reduction of 1-phenyl-3-hydroxypyridiniumchloride considerable hydrogenolysis took place with the formation of l-phenylpiperidine (34%) along with 31yo of the desired 1-phenyl-3-hydroxypiperidine (71). While Kao (72) reported very good yield of 3-hydroxypiperidine, Biel and his co-workers (73) found that piperidine (50% yield) was obtained in addition to 30-40”,10 of 3-hydroxypiperidine. They postulate that an intermediate “ally1 type alcohol” is formed which is more prone to hydrogenolysis than 3-hydroxypiperidine:

;’fLNJ H,o

0’”’ +

H

H

3-Hydroxypiperidine was recovered unchanged when subjected to similar reaction conditions. When 2-hydroxypyridine, generally considered as 2-pyridone (64) is hydrogenated only two molecular equivalents of hydrogen are taken up. Grave ( 7 4 )in some early work suggested that the N-methyl derivative, l-methyl-2-pyridone, is a true dihydropyridine and should require only two equivalents. Ruzicka (75) in an earlier publication describes 2-piperidone as a d-acid lactam, not a true ketone. When either 2pyridone or any 1-substituted-2-pyridoneis subjected to catalytic reduction the resultant product is either a cyclic amide or an N-substituted one. Of all the groups amenable t o catalytic hydrogenation the amide is the most difficult to reduce. 2-Piperidone as an aminoketone should be reduced further to 2-hydroxypiperidine. Furthermore, the nitrogen atom of a 2-pyridone or piperidone should not be basic enough to cause catalyst poisoning. Cavallito (66) was able to hydrogenate 2pyridone with palladium catalyst in neutral solution, but not 3-hydroxypyridine. I n neutral medium the unshielded nitrogen atom of the 3-derivative inhibits the activity of the catalyst. Platinum catalyst was used to reduce 1-methyl and 1-ethyl-2-pyridones in neutral solution ( 7 6 ) .When Raney nickel is used in pyridine reductions tempera-

HYDROGENATION OF PYRIDINES AND QUINOLINES

219

ture and pressure conditions must be elevated in order to overcome inactivation by the resultant piperidines. Yet Gautier was able to convert N-substituted pyridones to the corresponding piperidones (77, 78) under conditions that are too moderate for the hydrogenation of most pyridines with it. In only a few instances has enough hydrogen been absorbed to carry the reaction beyond the 2-piperidone stage. Galinovsky (79) obtained 1-methylpiperidine from the reduction of 1 -methyl-2-pyridone with a 50% ratio of platinum oxide for 42 hr. 3-Methyl-2-pyridone was converted to 3-methyl-2-piperidone in 80% yield by reduction in methyl alcohol in the presence of Raney nickel a t 200-240" and 120 atm (80). When the reaction temperature was raised to 280" some of the cyclic amide was reduced to 3-methylpiperidine and further alkylated to 1,3-dirnethylpiperidine. Structurally, 4-hydroxypyridine or 4-pyridone resembles the 2isomer. However, when it is subjected to catalytic hydrogenation, three molar equivalents of hydrogen are absorbed as in the reduction of the 3-isomer, giving 4-hydroxypiperidine, There is no record in the literature where only two equivalents of hydrogen are taken up in a catalytic hydrogenation. A number of unsuccessful hydrogenations are reported. Cavallito (66) could not convert the 4-derivative, but found that 4-hydroxypiperidine could be obtained by hydrogenation of 4-benzoxypyridine. Some 4-pyridones have been reduced in glacial acetic acid in the presence of platinum catalyst a t low pressure (81). Others found both palladium and platinum catalysts ineffective (75, 82, 83). When the 4-derivative was reduced in ethyl alcohol in the presence of Raney nickel, uptake of hydrogen did not take place below 200 "and 100atm pressure. Under these conditions, N-alkylation occurred, yielding 1-ethyl-4-hydroxypiperidine. In aqueous solution at 175" and 65 atm, yield was poor and much starting material was recovered (83). In contrast, the N-methyl derivatives were reduced a t 150" and an initial pressure of about 100 atm (84). In addition to l-methyl-4-hydroxypiperidineabout 10% of failed l-methylpiperidine was obtained. 1,2,6-Trimethyl-4-pyridone to undergo hydrogenation with platinum or palladium in glacial acetic acid (85), but in the presence of Raney nickel at 125" and 100 atm or nickel on silica, reduction was successful. Several other 1-methyl-substituted 4-hydroxypiperidines were similarly prepared. I n the hydrogenation of alkoxypyridines, 2-methoxypyridine was found to undergo hydrogenolysis (74). It was reported that the ring reduced first, then the ether group was cleaved with the total uptake of four equivalents of hydrogen. When the reaction was interrupted

220

MORRIS FREIFELDER

after the uptake of these equivalents the products of reduction were piperidine and 20 yo of 2-methoxypyridine. Hydrogenolysis did not take place during the reduction of 3- or 5-methoxy-4-pyridones (85). The stability of 3-alkoxypyridines toward reduction has been reported by others (86). 4-Alkoxypyridines hydrogenated in the presence of ruthenium dioxide a t 140" and 150 atm have been converted to the . the other hand, corresponding alkoxypiperidines in good yield ( 8 6 ~ )On conversion of 3-methyl-4-propoxypyridinea t 150" and 150 atm in the presence of Raney nickel gave only a 25% yield of the corresponding piperidine.

XI. Pyridylal kanols Hydrogenation of this group of pyridines requires little comment. A number of catalysts were employed under a variety of conditions. Ruthenium dioxide is favored in this laboratory in the reduction of these compounds because of the speed of reaction and the absence of side products. 2-( 2-Hydroxyethy1)piperidinehas been obtained in excellent yield after the hydrogenation of the corresponding pyridine (8). I n contrast, a reaction in the presence of Raney nickel a t 150" and 130 atm gave less than 50% of reduction product plus di- [2-(2-piperidyl)-ethyl] ether (87). 4-(2-Hydroxyethyl)pyridine has been converted t o the . and Eldred piperidine by means of ruthenium catalyst ( 8 6 ~ )Brown reported that there was no hydrogen uptake using nickel catalysts a t 200" and 200 atm (87a). I n attempting to reduce pyridine-2-aldehyde in one step t o Z-hydroxymethyl-piperidine with ruthenium] uptake of hydrogen stopped after one equivalent was absorbed. It was then found necessary to complete the hydrogenation with rhodium catalyst (87). 3-(2-Piperidy1)propanol is readily prepared by reduction with Raney nickel a t 125" and 170 atm. However, when the temperature is raised to 200°, indolizidine (octahydropyrrocoline) is formed in 78% yield (88).

XII. Aminopyridines Aminopyridines are similar to hydroxypyridines in that the 2- and 4aminopyridines can exist in several forms. 3-Aminopyridine is considered as an aromatic amine related to aniline and behaves so on hydrogenation. Nienburg (89) found that uptake of hydrogen stopped sharply with the absorption of three molar equivalents. 3-Aminopiperidine was obtained in good yield. Another investigation confirmed his results (90).

HYDROGENATION O F PYRIDINES AND QUINOLINES

22 1

While Wheland (91) suggests that 4-aminopyridine is an aromatic amine little success has been achieved in hydrogenation. Orthner (92) cites failure with platinum catalysts under a variety of conditions. Very low yield (16.5%) resulted from reduction of the hydrochloride salt with platinum catalyst under 80 atm pressure (93). I n the reduction of N-(4-pyridyl)morpholine where the 4-amino nitrogen atom is tertiary, hydrogenation in alcohol was successful with ruthenium (8),but unsuccessful with rhodium on alumina or platinum oxide in acetic acid. It is possible that the pressure conditions used for reduction with ruthenium catalysts may be conducive to conversioii of 4-aminopyridine, since these catalysts are less inhibited by strong nitrogen bases. In the hydrogenation of 2-aminopyridine under acidic conditions two equivalents of hydrogen were absorbed, giving tetrahydroaminopyridine as the hydrochloride salt, which on further reduction was converted to piperidine hydrochloride and ammonia ( 7 4 ) . When 2-aminopyridine hydrochloride was hydrogenated in aqueous solution four equivalents of hydrogen were absorbed, yielding piperidine hydrochloride and ammonia.

H

Kirsanov and Ivashchenko (94)in a study on the existence of 2-aminopiperidine hydrogenated 2-aminopyridine in the presence of acetic anhydride with platinum oxide. They observed that the required amount of hydrogen was absorbed and obtained the 1,2-diacetyl derivative. They also reduced 2- [N,N-diphenyllaminopyridine.I n each case the product was identified by nitrogen analysis only, which may cast some doubt on the structure. Birkhofer (95) in an attempt to hydrogenolyze the benzyl group in 1-benzyl-2-iminopyridine found that nuclear reduction occurred yielding 2-amino-3,4,5,6-tetrahydropyridine and l-benzyl-2-iminopiperidine.2-Benzylaminopyridine was converted to 2-benzylamino-3,4,5,6-tetrahydropyridineinstead of undergoing debenzylation. The uptake of two equivalents of hydrogen is in line with some unpublished work by the author, who in carrying

222

MORRIS FREIFELDER

out the hydrogenation of 2-diethylaminopyridine found that only two equivalents of hydrogen were absorbed.

1-(2-Hydroxy-2-phenylethy1)-2-iminopyridine as the hydrochloride salt was converted to the corresponding piperidine ifi the presence of rhodium on carbon with the absorption of three equivalents of hydrogen (96).When the free base was hydrogenated deamination occurred.

I

CH,CHC,H,

I

OH

The protection against deamination when reduction is carried out as a hydrochloride salt is in contrast to the work of Grave ( 7 4 ) with 2aminopyridine. Nevertheless, it seems to emphasize a combination of effects which includes that of a 1-substituent. These effects suggest the possibility of hydrogenating a 1-substituted-2-iminopyridine in acid solution with subsequent removal of the protecting group to obtain 2-aminopiperidine.

XIII. Pyridylalkylamines I n the most commonly employed procedures for the hydrogenation of this group of compounds reaction is carried out under acidic conditions with platinum catalyst, usually a t slightly elevated temperature for increased speed of reaction. Yields generally are good except in those cases where it is necessary to isolate extremely water soluble bases. Since piperidines cause catalyst poisoning it would be expected that more basic compounds resulting from reduction of pyridylalkylamines or related compounds would be even more potent poisons. The failure

HYDROGENATION OF PYRIDINES AND QUINOLINES

223

of Profft to reduce 2-(butylaminomethy1)pyridine with Raney nickel (97) is not surprising. However, the mild conditions used t o convert nicotine to hexahydronicotine (98)is of interest. Another hydrogenation in methanol in the presence of nickel required conditions (70 atm initial pressure and 170”) which led to N-alkylation and partial hydrogenolysis (99).

Reductions in the presence of rhodium on carbon a t low pressure showed that pyridylalkylamines could be hydrogenated in the absence of any acidic agent (16).I n the hydrogenation of 4-(2-benzylaminoethyl)pyridine some hydrogenolysis occurred yielding 4-(2-aminoethy1)piperidine in addition to the desired product; debenzylation appeared to take place after ring reduction. Ruthenium dioxide a t 70 atm and 90-100” may be the catalyst most suitable for this group of strong bases, since it appears to be less poisoned by them than other catalysts (8). It is affected, however, when the amino group in the side chain is primary.

XIV. Dipyridyls and Related Compounds The simplest and most successful method for the hydrogenation of the isomeric dipyridyls has been reported by Smith (100). Reaction is carried out under low pressure conditions in water or alcohol with a moderate excess of hydrochloric acid in the presence of platinum oxide. Rhodium on a carrier has been tried by this author in the conversion of 2,2’-dipyridyl but the excess acid has an inhibiting effect on the catalyst. When acetic acid was used in place of hydrochloric acid reduction with rhodium was very slow. I n the absence of acid no uptake of hydrogen took place. High pressure hydrogenation of 4,4’-dipyridyl with Raney nickel in ethyl alcohol resulted in N-alkylation (201), giving l,l’-diethy1-4,4’dipiperidyl in 74% yield. However, under similar conditions ( 15O-180”, 100-150 atm) 2,2”-dipyridyl yielded a mixture of 2,2-dipiperidyl and 1-ethyl-:!, 2’-dipiperidyl. If hydrogenation of dipyridyls is to be carried under higher pressure conditions it might be preferable to consider ruthenium catalyst since N-alkylation has not been observed with its use (8). A partially hydrogenated dipyridyl, l,l’-diacetyl[4H,4‘H]-4,4‘-

224

MORRIS FREIFELDER

dipyridyl, was readily reduced in the presence of platinum oxide in acetic acid at low pressure, but resisted reduction with Raney nickel in alcohol or dioxane a t 25-70" and 3-100 atm pressure (102).

XV. Quaternary Compounds The hydrogenation of alkyl, aryl, or aralkylpyridinium salts is a moderately uncomplicated means of obtaining N-substituted piperidines. The quaternary salts reduce more readily than the salts of the parent pyridines ( l o )and yield of N-substituted piperidines is usually better than obtained by reaction of the piperidine with the appropriate halide. The reduction may be carried out under a variety of conditions of pressure and temperature, in the presence or absence of base, in water, alcohol, or acetic acid. Raney nickel, palladium on a carrier, rhodium on a carrier and platinum oxide have all been used, with platinum oxide enjoying the widest use. Reaction in general is straightforward. There are, however, several examples in which the isolation of tetrahydropyridines are reported. The reduction of the methiodide of methyl isonicotinate in methyl alcohol with platinum oxide gives methyl 1 -methylpiperidine-4-carboxylate in good yield (103). When the reduction is carried out in a shorter reaction period, l-methyl-l,2,5,6-tetrahydroisonicotinateis formed. It is claimed that length of reaction time determines the product. A tetrahydro compound was obtained by using a lesser amount of catalyst for the hydrogenation (104).Lyle and co-workers (105)conclude that formation of the tetrahydro product results from inactivation of the catalyst. They were able to obtain the tetrahydro compound in 71% yield by reducing crude quaternary iodide in the presence of previously used catalyst. When the corresponding methyl-pyridinium bromide was hydrogenated only the completely saturated compound was obtained. The hydrogenation of the methiodide of ethyl nicotinate in alcohol in the presence of Raney nickel and triethylamine, a t 120" and 120 atm for eight hours gave over 80% of ethyl-1-methylpiperidine-3-carboxylate (106).When the reduction was run for 4 hr a t 80" and the same pressure, 54% of the hexahydro compound and 38% of a tetrahydro product were obtained. The results of these hydrogenations suggest a possible poisoning effect of iodide ion, either in the form of starting quaternary compound or when partial hydrogenation takes place and the hydriodide salt is formed. However, there are enough examples of the hydrogenation of pyridine quaternary iodides to show that the effect is not general, but

HYDROGENATION OB PYRIDINES AND QUINOLINES

225

may be limited only t o pyridinium compounds containing electron withdrawing groups conjugated with the ring. I n a private communication, Dr. R. E. Lyle has pointed out that tetrahydro derivatives were obtained when 3-cyanopyridine quaternary compounds were hydrogenated, He believes that most pyridine quaternary salts containing an unsaturated electron withdrawing group in the 3-position are capable of undergoing stepwise addition of hydrogen to yield tetrahydropyridines. The work of Grob and Ostermayer (foe),in which tetrahydro compound was obtained, lends support to his suggestion. I n contrast to this, Kuznetsov and Libman (107)obtained the completely hydrogenated compound, although it must be pointed out that no attempt was made to get any other product. Since the reported u values (108)of the 4-COOC,H5 and 3-COOC,H5 groups show that the former has a higher electron withdrawing effect, then the results of hydrogenations with the 4-COOCH3 derivative (103,lU4, 105) are not unexpected if the thesis of the effect of conjugated electron withdrawing groups has any creditability. The formation of derivatives from quaternary compounds containing an ester function could take place in the following manner:

64 COOR

I

I

CH, Ie

Uptake of the first molar equivalent of hydrogen gives the bracketed intermediate. Although the 5,g-double bond is in conjugation with the 3,4-bond it is a much weaker system than that of the 3,4-bond and the carbonyl system. It could be expected that the 5,6-bond would be preferentially attacked as if it were an isolated double bond. 1,4-Addition of hydrogen followed by migration of the 2,3-double bond would also give the same intermediate shown above. However, since the 4-position is occupied, 1,2-addition of hydrogen seems more logical. Pyridine quaternary iodides having an amide or substituted amide group in various positions have been reduced catalytically (61,109,110), but partial hydrogenation has not been observed or attempted. The amide nitrogen probably causes sufficient change in the electron withdrawing power of the group to decrease the tendency for hydrogenation to stop a t the tetrahydro stage. Lyle (105) suggested that catalyst poisoning aids in stopping the

226

MORRIS FREIFELDER

reduction a t the tetrahydro stage. Grob and Ostermayer (106) found that when triethylamine was used some tetrahydro compound was obtained even a t high pressure. I n this latter instance the base neutralized the shielding effect of hydrogen iodide on the ring nitrogen after the first equivalent of hydrogen was absorbed. As a result this leaves the intermediate dihydropyridine

with an unshared pair of electrons. Catalyst inhibition will occur (30). This coupled with the suggested effect of the electron withdrawing groups then causes the reduction to stop a t the tetrahydro stage. The partial hydrogenation seen in the preceding example (106) leads to the thought that catalytic reduction of the whole series seen below should be investigated in the presence of sufficient base to neutralize the resultant hydrogen halide. Low pressure conditions should be employed.

I

R'

R

=

IG

I

R'

2-, 3-, 4-nitrile, ester, amide; R'

= alkyl.

The results could supply an answer to the validity of the effect of unsaturated or conjugated electron withdrawing groups on the entire system. Indeed the investigation could be expanded t o include R = carboxyl and to substitute other ions for iodide ion.

XVI. Some Selective Reductions I n general one can assume that groups capable of being reduced will undergo preferential attack in the presence of an aromatic system. One can predict, with some degree of certainty, that a functional group will be selectively hydrogenated in the benzenoid series. With the more readily reducible groups this also can be done in the pyridine series. Among substituents which undergo hydrogenolysis it is impossible to retain nuclear chlorine, bromine or iodine when reducing the pyridine ring. Others like 0- and N-benzyl attached to a side chain may not be

HYDROGENATION OF PYRIDINES AND QUINOLINES

227

affected during ring reduction. Examples can be seen in the hydrogenation of 4-ethoxycarbonyl-4-(2-pyridy1)butyl benzyl ether ( I l l ) , and of 4-(2-benzylaminoethyl)pyridine(16). I n the latter case some debenzylation did take place when uptake of hydrogen was well over theoretical, although it appeared that it occurred after ring reduction. Baker and McEvoy reduced l-benzyl-l,2-dihydro-2-pyridylideneacetophenone in alcohol with Raney nickel under 2-3 atm pressure and noted a sharp break after two equivalents of hydrogen were absorbed (112).The corresponding piperidylidene compound was obtained in 60% yield. When platinum oxide was used under similar conditions two equivalents of hydrogen were quickly absorbed with a n unsharp but definite break and more hydrogen was slowly taken up. However, in alcohol containing hydrochloric acid, hydrogenation in the presence of platinum showed a very sharp break a t three equivalents of hydrogen. If uptake was allowed to continue a total of five equivalents resulted.

The difference in selectivity may be explained. The ease with which the pyridylidene compound was reduced with Raney nickel a t low pressure is an indication of its non-aromaticity. The selectivity then is a preferential one of two nonaromatic bonds over one highly conjugated double bond. The selectivity was likely aided by the basic unshielded nitrogen which inhibited the activity of the catalyst and prevented further hydrogen uptake. In the presence of hydrochloric acid the inhibiting effect was removed and uptake proceeded quickly with platinum oxide to 1-benzyl-2-phenacylpiperidine.

A. REDUCTION IN

THE

PRESENCE OF KETONEFUNCTION

The selective reduction of functional groups in the pyridine series over ring hydrogenation is probably due to the effect of the nitrogen atom. It is conceivable that this selectivity could be changed, particularly

228

MORRIS FREIFELDER

in the case of less readily reducible functional groups, by converting the ring nitrogen to an ionic form and thereby removing its poisoning effect on the catalyst (30).An examination of the reduction of certain pyridines containing ketone groups should therefore be of interest. I n the hydrogenation of 2-phenacylpyridine with Raney nickel a t room temperature and pressure 2-(2-hydroxy-2-phenethyl)pyridine was obtained (113).Reduction in the presence of platinum oxide gave the same product (114).2,6-bis-Phenacylpyridinegave the corresponding bis-carbinol. However, when these investigators hydrogenated 2phenacylpyridine with platinum in acetic acid four equivalents of hydrogen were absorbed and 2-(2-hydroxy-2-phenethy1)piperidinewas obtained. Similarly, 2,6-bis-phenacylpyridinehydrochloride yielded the corresponding bis-carbinol of the piperidine. Howton and Colding (115) reduced methyl 2-phenacylpyridinium bromide in methanol with platinum oxide and obtained an 88% yield of I-methyl-2-phenacylpiperidine hydrobromide. I n another reduction they allowed uptake to proceed further and claimed that 20% of 1methyl-2-(hexahydrophenacy1)piperidine was formed. I n the hydrogenation of the methobromide and metho-p-toluene sulfonate of 2-phenacylpyridine 1-methyl-2-phenacylpiperidinewas obtained (116). Vl'hen the methiodide compound was reduced about 8% of l-methyl-2(2-hydroxy-2-phenethy1)piperidine was also obtained. A study of these reductions shows that three reaction routes can be followed, leading to the pyridyl carbinol, piperidyl carbinol, and piperidyl ketone:

0 / 0

CH,CHO HC,H

CH,COC,H,

\

( )'

QcH,CHoHc6Hs

H

Route (A) is the expected reaction route in neutral solvent. I n the preparation of the piperidyl carbinol, in the presence of hydrochloric acid, reaction follows route (B),which may involve concurrent reduction of the ring and carbonyl group or may actually follow (A) or (C) with

HYDROGENATION O F PYRIDINES AND QUINOLINES

229

subsequent further hydrogenation of each until a total of our equivalents is absorbed. Only when the quaternary compound is reduced is (C) followed. An examination of the hydrogenation of 2-phenacylpyridines with the ring nitrogen in a shielded or ionic form may indicate how reaction takes place. 2-Phenacylpyridine hydrohalide in solution can (or may) exist in equilibrium between the two species seen below:

q+j CH,COC,H,

+ HX

0)

‘CH,COC,H, Xe (El

I n the described reductions only one equivalent of acid (HX) was used. Because of this, only partial protonation may occur leaving the unprotonated base to follow route (A) and the protonated species (conjugate acid) to follow (C). With the quaternary compounds no equilibrium corresponding to that of protonation can occur so that the species is present in only a single completely ionic form (F).

0

/N CH,

‘cH,COC,H, Xe

(F)

The examples show that form (F)leads to preferential ring reduction. The reduction of (E)should follow the same pathway as (F)because each is in a completely ionic form and the poisoning effect of the nitrogen atom will be neutralized (30). If the proposal that incomplete protonation is due to the insufficiency of acid, further or even complete protonation might be achieved by the addition of excess acid, By this means, unless the catalyst would be adversely affected, 2-phenacylpiperidine could be obtained directly from 2-phenacylpyridine. It is possible that this selectivity could be extended in the presence of other functional groups. The validity of the hypothesis that quaternization favors ring reduction over ketone function is given support in the work on l-phenacylpyridinium halides. Krohnke and Faslod (117) found that in water

230

MORRIS FREIFELDER

1-phenacylpyridinium bromide underwent ring reduction without affecting the ketone. At higher pressure, however, a 30% yield of 1-(2-hydroxy-2-phenethy1)pyridiniumbromide resulted (118).A series of 1-(4-nitrophenacyl)-4-alkylpyridinium bromides was converted t o the corresponding 4-aminophenacylpiperidines (119). It was found that with 1-phenacyl-4-alkylpyridiniumbromides ring reduction was favored if the 4-alkyl group was methyl or ethyl (120). When the 4-alkyl group was larger the corresponding piperidine carbinol formed. Examination of molecular models of phenacylpyridines shows that they can be readily rotated so that the carbonyl group can be moved away from the surface of the catalyst. Selective ring reduction is therefore possible under the proper conditions. However, removal of the methylene bridge between the pyridine ring and the carbonyl group causes a distinct difference in movement of the molecule. For example, with models of the isomeric benzoylpyridines rotation seems t o be restricted to those positions which suggest that both the pyridine ring and the carbonyl group should make good contact with the catalyst. Lack of selectivity, therefore, could be expected. Crook and McElvain (121) found that in the reduction of 2-, 3- and 4-benzoylpyridine in water containing an equivalent of hydrochloric acid four equivalents of hydrogen were absorbed. Only the corresponding piperidyl phenyl carbinols were obtained. Lyle and his associates (122) reported that they could not convert 4-benzoylpridyine or the methylpyridinium iodide or bromide to the corresponding piperidines without reducing the carbonyl group. Lyle and co-workers (123) followed the hydrogenation of methyl 4-benzoylpyridinium bromide by means of ultraviolet absorption spectra and found that the carbonyl group was attacked first. They were able to stop the reduction and isolate the pyridinium carbinol. Hydrogenation of some acetylpyridines shows that they cannot be converted to the corresponding piperidyl methyl ketones. When 2acetylpyridine or its methosulfate was reduced selective ring conversion could not be accomplished (124). Homoveratryl3-acetyl-6-methylpyridinium chloride when hydrogenated in glacial acetic acid in the presence of palladium on carbon yielded only the piperidyl carbinol(l25). Luke6 and his associates (126) found that methyl 4-acetylpyridinium bromide gave the piperidyl carbinol under very mild conditions (0.5 atm and 19'). I n contrast t o previously published work (121),Lyle and his associates (123, 127) found that when 3-benzoylppidine hydrochloride was hydrogenated 3-benzoylpiperidine was the major product. When methyl 3-benzoylpyridinum bromide was reduced, 1-methyl-3-benzoylpiperidine hydrochloride was obtained in 40% yield (127). When this hydrogenation was re-examined, the products of reduction were

HYDROGENATION OF PYRIDINES AND QUINOLINES

231

shown to be 1-methyl-3-benzoylpiperidine(40%)) a pure diastereoisomer a 10% mixture of the diastereoisomers, of the piperidyl carbinol (€!yo), and 15% of 1-methyl-3-benzylpiperidine. The approximately equal amounts of piperidyl ketone (40%) and piperidyl carbinols and 1-methyl-3-benzylpiperidine (33%) suggest that two competing reactions may take place. I n one reaction, the carbonyl group is attacked first and the resultant pyridyl carbinol is further reduced to the piperidyl carbinol. The presence of the 3-benzyl compound may be due to hydrogenolysis of the pyridyl carbinol and subsequent ring reduction or hydrogenolysis of the piperidyl carbinol. I n the second competing reaction, selective reduction of the pyridine ring takes place. The ultraviolet absorption spectra obtained by Lyle seems to suggest that reaction indeed is selective. Ring reduction takes place first, with subsequent reduction of the carbonyl group and later hydrogenolysis of the hydroxyl group to give the 3-benzyl derivative. Much more interesting, however, is the indication of stepwise hydrogenation of the pyridine ring. Lyle observed in following the ultraviolet absorption spectra during reduction that the strong band for the starting material decreased in intensity as the reaction proceeded. Simultaneously, another strong band of much longer wavelength (315-320 mu) appeared. This band decreased in intensity as the reduction proceeded to completion, a t which time a band, identified as one for the benzoylpiperidine, was seen in addition to the one a t 320 mu. This band was assumed to be one for l-methyl-3-benzoyl-1,4,5,6-tetrahydropyridine. This compound was never isolated per se but its identifying ultraviolet spectrum was shown to be similar to that of 1-(2-phenethyl)-3-benzoyl-l,4,5,6-tetrahydropyridine, prepared in low yield by catalytic hydrogenation of 1-( 2-phenethyl)-3-benzoyl-1,4dihydropyridine in neutral solution. From Lyle’s results, it might be assumed that the reduction could proceed in the following manner by 1,4 addition since the 4-position is unoccupied:

232

MORRIS FREIFELDER

The addition of hydrogen to give the intermediate dihydro compound is likely a fast reaction, while further conversion to the tetrahydro derivative is much slower. If the 1,4-dihydropyridine is indeed the true intermediate, then selective reduction of the 5,6 bond leading to the tetrahydro compound is obvious. While Lyle was not able to isolate this particular tetrahydropyridine, he did show that i t was possible t o obtain such a compound from a corresponding dihydropyridine. It would be of interest to attempt a hydrogenation of methyl 3-benzoylpyridinium iodide in the presence of a base (106)and in the presence of an inactivated catalyst (105).It is doubtful that any dihydropyridine could be isolated but the combination of the ring nitrogen no longer shielded after uptake of the first equivalent of hydrogen and the less active catalyst could inhibit reaction sufficiently to stop it a t the tetrahydro stage. A comparison of the effect of the 3-benzoyl group and the reported ineffectiveness of the 3-acetyl group, of equally strong electron withdrawing power, on selective reduction of the pyridine ring of pyridine quaternary compounds causes speculation. It would be of interest to determine whether the effect is steric by comparison with a pyridine quaternary salt containing the 3-COC(CH3),group. It has been shown in Chapter XV that pyridine quaternary iodides containing either the 3- or 4-methoxycarbonyl group can be partially hydrogenated. It is of interest that there is such a significant difference when the 3- and 4-benzoylpyridine quaternary compounds undergo hydrogenation. The reasons for this are not clear and make one hesitate t o offer an explanation. That the use of excess acid may favor ring reduction may be seen in

OH J)

OH (K)

HYDROGENATION O F PYRIDINES AND QUINOLINES

233

the work of Pitha and Ernest (128).They hydrogenated 5-(2-pyridyl)2-pentanone in a 60% excess of dilute hydrochloric acid and obtained 4-methylquinolizidine. I n a similar reaction 2-keto-5-(2-pyridyl)-lpentanol was converted to 4-hydroxymethylquinolizidine.There are two ways in which cyclization could take place and these are shown at bottom of facing page. Ifreduction of (G) would proceed through (J)to (K)it should be difficult to split out water in order to obtain (I)as readily as the authors did, A further example of the use of excess acid favoring ring reduction may be seen in the conversion of 1-(2-pyridy1)-2-propanoneto the corresponding piperidylketone (isopelletierine) (129). I n this reaction the ketone is hydrogenated in glacial acetic acid in the presence of platinum oxide. Since such a large excess of glacial acetic acid is present, form (M) in the equation below could be favored, which would give the desired product.

When the methosulfate of 1-(2-pyridyl)-2-propanone was hydrogenated the pyridine ring was preferentially reduced to give N-methylisopelletierine (130). Additional evidence of the effect of excess acid may be seen in some work of Boekelheide and co-workers. They showed that when phenyl 3-( 2-pyridyl)-propyl ketone was hydrogenated in aqueous alcohol containing more than two molar equivalents of hydrochloric acid, 4-phenylquinolizidine was obtained (131).From the reduction of other pyridylketones in the same medium or in aqueous dilute acid (132)they obtained a mixture of quinolizidines. The results of these reductions bear out that ring conversion is the first step.

I

R

R = CH, or C,H,

"

234

MORRIS FREIFELDER

(P) was readily converted to (0).It is most likely the intermediate formed by the reaction shown below:

I n the presence of excess acid the base (N) can become more or less completely converted to ( Q ) ,which will lead to (S),the piperidyl ketone. An intermediate addition product (T) may be formed and water split out to give (P),or (S) can go directly t o (P)by condensation and loss of water. The same investigators (132)showed that when compound (N),where R =CH,, was hydrogenated in neutral solvent under pressure in the presence of Raney nickel only the corresponding piperidylcarbinol was formed. They concluded that the carbonyl group was reduced prior to the ring.

B. REDUCTION IN THE PRESENCE OF NITRILEFUNCTION 2-(2-Cyanoethyl)pyridine was reductively cyclized t o indolizidine (octahydropyrrocolinej by carrying out the reaction in aqueous alcohol containing two molar equivalents of hydrochloric acid in the presence of platinum oxide catalyst (133). The authors suggest that preferential ring reduction is essential for cyclization and propose the following reaction scheme:

235

HYDROGENATION OF PYRIDINES AND QUINOLINES

acHa \

CH,

bHC,H,

111

N (V)

The formation of (VI) and (IX) was cited as corroboration that preferential ring reduction is necessary for subsequent cyclization. The formation of (VII) and (X)was said to indicate that selective conversion of the nitrile group precluded cyclization. The pathway to indolizidine from (I)may be explained by another mechanism in’the scheme that follows this discussion. I n Section XVI,

236

MORRIS FREIFELDER

of this article, it was suggested that conversion of the pyridine nitrogen to an ionic form could foster preferential hydrogenation of the ring in the presence of ketonic function. The same principle could apply to the reduction of pyridylnitriles with the formation of (IIa). I n the catalytic hydrogenation of nitriles an imino compound is the first intermediate (134).This will account for (IIb). The precedent for the conversion of a nitrile to an aldehyde through the intermediate imino compound in the presence of water and hydrochloric acid has been established by Kuhn and Kirschenlohr (135), and confirmed in this laboratory. Others. report on similar reductions (136, 137, 138). The reductive cyclization of (IIc) to (IV) probably goes through either (IId) or (IIe). When Winterfeld and associates (139)reduced 2-( 1,3-dicyanopropyl)pyridine in alcoholic hydrogen chloride in the presence of platinum oxide or in acetic acid in the presence of palladium, cyclization did not take place. From a review of the reductions in this section it appears that not only is an excess of hydrochloric acid necessary (other strong acids might work as well) but water also must be present for cyclization to take place. The yield of (IV) (43%) may be due to incomplete formation of (Ia) because not enough acid is present to force the reaction to go in the desired direction, as proposed in the alternate scheme. A reinvestigation of the conversion of (I) to (IV) under Bockelheide’s condition in the presence of a larger excess of acid and an examination of the products of reaction could give a better insight to the mechanism of action.

c. REDUCTIONI N

THE PRESENCE OF A

BENZENOID RING

The selective hydrogenation of the ring in the presence of a benzenoid ring requires little comment. Adkins states that pyridinoid compounds are easier to reduce than benzenoid derivatives (10). It is of interest that Overhoff and Wibaut (38) reported that they could not hydrogenate 4-phenylpyridine under acidic conditions with platinum oxide while 2-phenylpyridine reduced readily. Molecular models fail to show why it cannot be reduced. Both 2- and 4-phenylpiperidine have been prepared by reduction under pressure with nickel catalysts ( 3 ) . They also reported that in the hydrogenation of 2,6-diphenylpyridine all three rings were affected while with 2,4,6-triphenylpyridine only the benzene rings appeared to be converted. Molecular models of 2- and 4-phenylpiperidine suggest that once the pyridine ring is reduced conversion of the benzene rings should be

HYDROGENATION O F PYRIDINES AND QUINOLINES

237

difficult. The addition of a methylene group between the two rings makes reduction of the benzene nucleus much less difficult. Veer and Goldschmidt (140) found that in the conversion of 4-benzylpyridine with platinum catalyst in acetic acid some 4-cyclohexylmethylpiperidine was also obtained. In contrast, with rhodium in neutral solvent the only product was 4-benzylpiperidine, while during the conversion of the 2-derivative 8% of the completely reduced compound was obtained (16).Under the reaction conditions with ruthenium there was a tendency toward reduction of both rings with each compound ( 8 ) . The lack of selectivity with ruthenium was also observed in the reduction of 4-(2benzylaminoethy1)pyridine which gave 4-(2-cyclohexylmethylaminoethy1)piperidineplus a small amount of 4 42-aminoethy1)piperidine (8). The reduction of phenylpyridyl carbinols in acetic acid or as hydrochloride salts with platinum oxide is widely reported. Rhodium in acetic acid has given good results in this laboratory. Walker (18) reports considerable success with palladium on carbon when reducing the acetate salts of such compounds. The only disadvantage of this method is possible accompanying hydrogenolysis of the OH group.

D. SELECTIVITY IN THE PRESENCE OF OTHER HETEROCYCLES Before discussing the reduction of the ring in the presence of other heterocycles it might be of interest to mention the selective reduction of one pyridine ring in the presence of another. McCarthy and his associates (141) hydrogenated a-phenyl-a-(2-pyridyl)-2-pyridinemethanol and obtained a 75% yield of a-phenyl-a-(2-pyridyl)-2-piperidinemethanol. The same group showed they could selectively reduce a-(2-furyl)or-phenyl-2-pyridinemethanol,as hydrochloride salt in alcohol or as base in acetic acid, to the corresponding piperidine, in the presence of a platinum catalyst, without affecting the furan ring. Bader and Oroshnik (142) hydrogenated a-(3-indolyl)-2-pyridinemethanol in the presence of platinum oxide in alcohol solution containing a considerable excess of acetic acid. They obtained 3.5% of the piperidine carbinol and 51% of 2-skatylpiperidine [2-(3-indolylmethy1)piperidine] showing that the indole ring remained intact. The same authors (143) showed that 4-skatylpyridine was converted to 4skatylpiperidine. Overhoff and Wibaut (38) showed that the pyridine ring reduced first when the pyrrole nitrogen was unsubstituted. They selectively hydrogenated the pyridine rings of 2-( 2-pyridyl) and 3-(2-pyridy1)pyrrole in alcohol containing an equivalent of hydrochloric acid with platinum

238

MORRIS FREIFELDER

oxide catalyst. When I-( 2-pyridy1)pyrrole was reduced cleavage to piperidine and pyrrolidine resulted. With 1-methyl-2-(2-pyridy1)pyrrole both rings were reduced. 1 -Methyl-2-(3-pyridy1)pyrrole underwent hydrolytic cleavage to yield 3-(4-aminobutyl)piperidine.This latter reduction is of interest because Spath and Kuffner (144) found that reduction in acetic acid with palladium on carbon yielded only 1-methyl-%(3-pyridy1)pyrrolidine. Ochiai and his group point out that the pyrrole ring is more difficult to reduce than the pyridine ring but note that substituents on the pyrrole ring have a bearing on the course of reduction (145).When they hydrogenated some 3-substituted-2-methyl-4-pyridylpyrrolesin acetic acid with platinum the resultant piperidines were the only products obtained. However, with %methyl-&(2- and 3-pyridy1)pyrroles both piperidy lpyrroles and piperidylpyrrolidines were found.

XVII. Quinolines It is well known that the pyridine ring of quinoline is hydrogenated before the benzene portion of the molecule, although substituents do exert an effect. Hamilton and Adams ( 1 0 ) converted quinoline hydrochloride to 1,2,3,4-tetrahydroquinolinein alcohol in the presence of platinum oxide a t low pressure. Adkins and Billica using a highly active nickel catalyst obtained the 1,2,3,4-tetrahydro derivative readily in a reduction a t 1-3 atm ( 1 1 ) . Among the earliest workers, Darzens (146) used a reduced nickel catalyst to obtain the same product a t elevated temperature. Conversion of quinoline to the completely saturated decahydro compound requires a longer reaction period and more drastic reaction conditions. Ipatief (147) carried out a reduction in the presence of nickel oxide a t 240" and 110 atm pressure for 12-14 hr to obtain decahydroquinoline. Adkins and Cramer ( 3 5 ) with nickel on keiselguhr obtained the tetrahydro derivative a t 150" and 160 atm after 1-4 hr, but found it necessary to raise the temperature to 175" and 175 atm for 6-8 hr to obtain a mixture of cis and trans decahydroquinolines. Even more striking is the difference shown by Tsushima and Sudzuki (148).Using a reduced nickel catalyst they obtained 71-73% yields of 1,2,3,4-tetrahydroquinolinea t 70-100" and 70-61 atm but had to raise the temperature to 210" a t 70 atm to obtain 62% of trans decahydroquinoline. Skita and Meyer (149) describe the ease of converting quinoline to decahydroquinoline with colloidal platinum in acetic acid. However, a large amount of catalyst was used to induce complete reduction in nine hours while conversion t o the tetrahydro derivative took a much shorter time. Overhoff and Wibaut (38) in

HYDROGENATION OF PYRIDINES AND QUINOLINES

239

carrying out a hydrogenation in acetic acid in the presence of platinum oxide also found that a long reaction period was necessary for conversion to the completely saturated compound. The conditions used to prepare decahydroquinoline have a n effect on the amount of cis and trans isomers that are formed. As might be expected, high temperature should lead to a preponderance of trans product. This was born out in the work of Sugino and Mizuguchi (24). They reduced quinoline to 1,2,3,4-tetrahydroquinolinein the presence of copper a t 130" and 140 atm. I n order to obtain the decahydro compound it was necessary to raise the temperature to 260". Trans decahydroquinoline resulted. Further confirmation may be seen with the use of nickel catalyst by Tsushima and Sudzuki (148).Huckel and Stepf (150) made a study of the hydrogenation of quinoline using Skita's method (149).They found that the formation of cis and trans isomers was dependent on the acidity of the solution. Reduction in acetic acid containing excess concentrated hydrochloric acid yielded 65% of the cis product and 35% of trans, while reaction of quinoline oxalate in acetic acid gave 80% trans and 20% cis decahydro compounds. Reduction in neutral solution gave only 1,2,3,4-tetrahydroquinoline. Overhoff and Wibaut (38) obtained mostly trans product in the hydrogenation of quinoline with platinum oxide in acetic acid, while reduction of the hydrochloride salt in alcohol gave a mixture of cis and trans products.

XVIII. lsoq u i no1i nes Addition of hydrogen to isoquinoline also occurs first in the pyridine ring. However, reduction of the benzene portion of the molecule is far more difficult than with quinoline. Helfer (151) could not hydrogenate isoquinoline beyond the tetrahydro stage with colloidal platinum. Skita (152) could convert it to decahydroisoquinoline by the addition of a large amount of chlorplatinic acid after the tetrahydro compound had formed or by adding sufficient chlorplatinic acid a t the beginning of the reduction so that the total platinum content was over 20% by weight of the starting isoquinoline. Witkop (153) could obtain the decahydro derivative by low pressure reduction in acetic acid solution containing excess sulfuric acid only if a large amount of platinum oxide was used. He pointed out that when o of catalyst to compound was present or the sulfuric less than a 1 0 0 ~ratio acid was omitted, reduction stopped a t the tetrahydro stage. It is of interest that similar to the work of Hiickel and Stepf with quinolines (150) the presence of strong acid resulted in a predominance of cis product. Witkop in some later work showed that high pressure re-

240

MORRIS FREIFELDER

duction with nickel (18O0, 270 atm, 15-18 hr) led t o an 80% yield of trans compound (154). The refractory behavior of isoquinoline toward reduction beyond the tetrahydro stage may be compared to the difficulty in the hydrogenation of /3-phenethylamines to the corresponding cyclohexylalkylamines. Metayer (155) has reported on the problem of reducing phenylalkylamines where the amino group is beta to the aromatic ring. Successful reduction with nickel catalyst required high temperatures that were conducive t o side reaction. Zenitz and his co-workers (156) in their study with platinum oxide in acetic acid a t low pressure often found it necessary to make several additions of catalyst t o induce complete ring hydrogenation. Kindler (157) used excessively high catalyst ratios (25% by weight of palladium black) to obtain complete uptake of hydrogen. Gray and Heitmeier suggest the possibility of transannular interaction between the basic nitrogen and the aromatic ring as a cause of the difficulty in reducing the benzene ring of phenethylamines and tetrahydroisoquinoline (158).However, an examination of molecular models seems to point to poisoning caused by good contact between the basic nitrogen atom with its unshared pair of electrons and the catalyst surface. Further credence as to this possibility may be seen in the work of Woodward and Doering (159)who found that they could not convert 7-hydroxy-8-methylisoquinoIinebeyond the corresponding 1,2,3,4tetrahydro derivative with platinum oxide. However, when they subjected the N-acetylated tetrahydro compound t o the same hydrogenation conditions, they did obtain the decahydro compound although there was some concomitant loss of the hydroxyl group. A more interesting example is seen in the work of Rapala and his associates (160) who showed that 1,2,3,4-tetraisoquinoline-3-carboxylic acid could be converted to the corresponding decahydro compound by carrying out reduction in neutral solvent in the presence of rhodium on alumina a t 50-100" and 15 atm pressure. I n this instance the carboxyl group neutralized the poisoning effect of the basic ring nitrogen and allowed hydrogenation to go to completion. Evidence of this may be seen when the same workers could not induce similar ring conversion of ethyl 1,2,3,4-tetrahydroquinoline-3-carboxylate. This author had shown that ruthenium dioxide was useful in reducing p-phenethylamines to the corresponding cyclic compounds (161). The relationship of these compounds to isoquinoline led to an investigation of the use of ruthenium in the hydrogenation of isoquinoline. It was found that under the moderate high pressure reduction conditions (go", 70 atm) when the reaction waa interrupted after the uptake of two

HYDROGENATION OF PYRIDINES AND QUINOLINES

24 1

equivalents of hydrogen decahydroisoquinoline was always found as a contaminant. By raising the temperature and increasing the catalyst level decahydroisoquinoline could be obtained in good yield.

XIX. Effect of Substituents The effect of substituents on reduction in the quinoline series in the presence of nickel catalyst has been studied by von Braun and his associates. The first work was on the effect of substitution in the benzene ring and the 2- and 3-positions (162).They pointed out that when the benzene ring was substituted or there was a substituent in 2-position reduction of the pyridine ring takes place more readily than quinoline itself and that 1,2,3,4-tetrahydro compounds are obtained. When 3-substituents are present mixtures of 1,2,3,4- and 5,6,7,8-tetrahydro compounds result. A second study on methylated quinolines (163))in which all but 5-methylquinoline were investigated, confirmed that substitution on the benzene ring leads exclusively to 1,2,3,4-tetrahydro compounds. It was further seen that loading of the pyridine ring favors 5,6,7,8-tetrahydro formation. 3- and 4-Methylquinolines lead to mixtures, while 2-methylquinoline gives 96% 1,2,3,4-tetrahydro and 4% of the 5,6,7,8-product. The size of the alkyl group in 2-position also had some influence (164). When 2-propylquinoline was hydrogenated, 65% of 1,2,3,4-compound, and 35% of the 5,6,7,8-tetrahydro derivative resulted. However, when the group in 2-position was phenyl (162) only 2-phenyl-l,2,3,4-tetrahydroquinoline was obtained. There are a number of examples in the literature confirming this observation. Oldham and Johns (165) reduced 2-phenylquinoline in the presence of platinum on zirconium oxide and found that the resultant 2-phenyl1,2,3,4-tetrahydroquinolinecould not be hydrogenated aay further. However, when the hydrogenation was carried out with platinum oxide, 2-cyclohexyldecahydroquinoline was obtained. 2-Phenylquinoline-4carboxylic acid was reduced in the presence of colloidal platinum to the tetrahydro derivative (9). When more catalyst was used, a decahydro compound was obtained. This product was shown to be 2-cyclohexyl1,2,3,4-tetrahydroquinoline-4-carboxylic acid (166), which on further reduction was converted to the cyclohexyldecahydro compound. The investigators found that 2-phenylquinoline was first converted to the tetrahydro compound when two equivalents of hydrogen were absorbed. With the uptake of five equivalents only 2-cyclohexyl-1,2,3,4-tetrahydroquinoline was obtained. Molecular models seem to show that the phenyl group in the 2-position will make better contact with the catalyst than the benzenoid portion of the quinoline ring.

242

MORRIS FREIFELDER

Campbell and co-workers suggest that the ease of nuclear reduction appears to be a general property of 8-substituted quinolines leading to conversion of the pyridine portion of the molecule (167).The reductions of 8-acetyl, 8-carboxy, and 8-carbethoxyquinolines with platinum oxide were investigated. There are examples of other 8-substituted quinolines in the literature supporting their work. I n view of Campbell’s work with 8-quinoline-carboxylic acid, the hydrogenation of S-hydroxyquinoline8-carboxylic acid is of interest. I n this instance the 5,6,7,8-tetrahydro compound was obtained (168).While von Braun (162, 163) has shown that 3-substituents lead to mixtures of 1,2,3,4- and 5,6,7,8-tetrahydro derivatives, and Campbell has shown that quinoline-8-carboxylic gave only 1,2,3,4-tetrahydro product, this finding of Ochiai suggests a very profound effect of the 3-hydroxyl group. Further study of its effect on ring reduction is in order. Isogai (169) reported that when 5-dichloromethyl-5-methyl-8(5H)quinolone was hydrogenated in neutral solvent with either palladium on carbon or platinum oxide, 5-dichloromethyl-5-methyl-8-hydroxy5,6,7,8-tetrahydroquinolinewas obtained. I n view of von Braun’s reports (162. 163) that substitution in the benzene portion of the molecule favors reduction of the pyridine ring, Isogai’s results may be surprising. However, one must consider that in this instance the benzene portion is not truly aromatic. The hydrogenation could be viewed as that of a conjugated carbonyl system selectively reduced in the presence of a pyridine ring. Since the reduction was carried out in neutral medium the pyridine portion of the molecule should be less favored. There has not been enough investigation of substituted isoquinolines to draw any definite conclusions about the effect of substituents on either ring. Woodward and Doering (159) reported the conversion of 7-hydroxy-8-methylisoquinolineto the 1,2,3,4-tetrahydro product in glacial acetic acid in the presence of platinum oxide. They were unable, however, t o reduce 7-hydroxy-8-piperidinomethylisoquinoline in alcohol and other solvents with the same catalyst. An examination of a molecular model does not show any apparent hindering effect because of the size of the group in 8-position, but rather tends to suggest poisoning from both the unshielded nitrogen atoms.

XX. Qui noli ne and lsoqu ino1i necarboxylic Acids Since all the quinolinecarboxylic acids are known, it is unfortunate that so little work has been carried out on catalytic reduction among this group of compounds. From the work of von Braun (162,163)on substitution in the benzene

HYDROGENATION O F PYRIDINES AND QUINOLINES

243

ring, it could be assumed that presence of a carboxyl group might have a similar effect and favor 1,2,3,4-tetrahydro formation. The only reduction of this type was reported by Campbell (167) who showed that 8-quinolinecarboxylic acid was converted to the 1,2,3,4-tetrahydro compound, Among the quinolines with the acid group in the pyridine portion of the molecule, the 2-acid has been reduced to 1,2,3,4-tetrahydroquinoline-2-carboxylic acid in the presence of platinum oxide in acetic acid (170). The 3-acid as sodium salt has been hydrogenated with Raney nickel a t high pressure t o give a 1,2,3,4-tetrahydro derivative (171). The hydrogenation of 2-phenylquinoline-4-carboxylicacid (9,166)t o the 1,2,3,4-derivative shows that a carboxyl group in 2-, 3-, or 4-position should, in general, result in conversion of the pyridine portion unless some strong directing group causes a profound change (168). Among isoquinolinecarboxylic acids little is found concerning reduction. One reference reports a hydrogenation of isoquinoline- 1-carboxylic acid to the 1,2,3,4-tetrahydro compound which required thirteen days for completion of the uptake of two molar equivalents of hydrogen when platinum oxide was used in acetic acid (172).It would be of interest to see whether uptake of hydrogen would proceed a t a more rapid rate in a neutral solvent in the same manner as described for pyridinecarboxylic acids (16, 49). The only report of successful reduction of any quinoline or isoquinolinecarboxylic acid in neutral solvent may be seen in the work of Campbell (167). It should therefore be of interest to investigate reduction of all the acids in neutral solution in order to learn the effect of p H and position of the acid group on reducibility.

XXI. Esters and Amides There are a few reductions among this group which require some comment. I n view of von Braun's studies on the loading of one ring which leads to hydrogenation of the other ring (162),the reduction of methyl 2,3-dimethylquinoline-4-carboxylategave surprising results (173).Despite complete substitution of the pyridine ring, only 24% of 5,6,7,8-tetrahydro compound was obtained along with 70% of 1,2,3,4tetrahydro derivative. Chiaverelli (171)found that if methyl quinoline3-carboxylate was reduced with Raney nickel at 100' and 110 atm, the corresponding 1,2,3,4-tetrahydro compound was obtained. With palladium on carbon at 60-65" and 90 atm a dihydro product resulted, which was assumed to be a 1,2-dihydroquinoline. It may be another example of the effect of electron withdrawing groups on reduction of the ring as proposed in Chapter XV. The dihydro compound was readily converted to the tetrahydroquinoline. Hydrogenation of the corre-

244

MORRIS FREIFELDER

sponding 3-diethylamide in the presence of palladium on carbon a t 60" and 90 atm resulted in formation of the py-tetrahydro product. When ethyl quinoline-2-carboxylate was hydrogenated in the presence of palladium oxide, not only was the desired tetrahydro compound obtained, but in addition, a dimeric tetrahydro compound bridged at &position was found (170). When platinum oxide was substituted, a quantitative yield of 1,2,3,4-tetrahydro derivative resulted.

XXII. Hydroxyquinolines and Related Compounds Since it is well known that phenols are the easiest of the substituted benzenes to hydrogenate, an examination of the reduction of hydroxyquinolines and isoquinolines is in order. It should be of interest t o note the effect of the phenolic hydroxyl group on ring hydrogenation. The work of Cavallito and Haskell(66) with a palladium sponge catalyst (67) in neutral solution showed that, while the introduction of the hydroxyl group permitted more rapid reduction than with quinoline itself, ring hydrogenation occurred in the pyridine ring. 3-Hydroxyquinoline was reduced more slowly than the 5 - , 6-,7-,or 8-derivatives but it, too, yielded the Py-tetrahydro form. 2-Hydroxyquinoline (carbostyril) and 1-hydroxyisoquinoline (isocarbostyril) gave the corresponding 3,4dihydroquinolines. This was to be expected since the two isomeric compounds are not phenolic but behave like amides. The behavior of the 4-hydroxy compound is of interest since it was not hydrogenated. However, this might be expected since the authors reported 4-hydroxypyridine was likewise unaffected by reduction under the same conditions. 4-Methyl-2-hydroxyquinolinefailed to reduce probably because of steric effects. A series of benzoate esters were also subjected t o hydrogenation. The were obtained as benzoates of 3-,5-,6-,7-OH-Py-tetrahydroquinolines expected. 2-Benzoxyquinoline and 1-benzoxyisoquinoline underwent hydrogenolysis and gave, for the most part, toluene and the corresponding 3,4-dihydro compounds. 8-Renzoxyquinoline was converted to the 1,2,3,4-tetrahydro compound with accompanying migration of the benzoyl group to the ring nitrogen. This migration is not an uncommon occurrence, often noted during the reduction of esters of nitrophenols. The 4-benzoyl derivative was hydrogenolyzed to toluene and 4-hydroxyquinoline. Cavallito and Haskell reported that Raney nickel and platinum oxide were ineffective catalysts under their conditions. Under more drastic conditions, 6-hydroxyquinoline was reduced by means of a nickel catalyst a t 140" and 60 atm (174) and in the presence of nickel on kieselguhr a t 140' and 80-100 atm (174, to 6-hydroxy-1,2,3,4-

HYDROQENATION OF PYRIDINES AND QUINOLINES

246

tetrahydroquinoline. When the temperature was raised to 180' the decahydro compound was obtained in low yield. Ochiai and Kawazoe (176)found that the hydrogenation of isocarbostyril in glacial acetic acid in the presence of platinum oxide gave l-hydroxy-5,6,7,8-tetrahydroisoquinoline. l-Ethoxy and 1-isopropoxyisoquinolinewhen hydrogenated under similar conditions also gave the bz-tetrahydro compound. 5-Hydroxyisoquinoline was readily converted to the corresponding 1,2,3,4-tetrahydro derivatives under low pressure conditions by means of platinum oxide (176~). When an attempt was made to prepare the corresponding decahydro compound by carrying out the reduction in alcohol in the presence of an active nickel catalyst (1.2)at 125' and 200 atm only 2-ethyl-5-hydroxydecahydroisoquinoline was obtained. It was necessary first to acetylate the tetrahydro compound and then subject it to further hydrogenation with the same nickel catalyst in order to avoid N-alkylation.

XXIII. Amino Compounds There are only a few reported hydrogenations of amino compounds where the amino group is attached to the pyridine ring. The reduction of 3-aminoquinoline with Raney nickel at 55' and 90 atm (171)gave a good yield of 3-amino-l,2,3,4-tetrahydroquinoline plus a bridged compound, 3,3'-iminobis (1,2,3,4-tetrahydroquinoline).The hydrogenation of l-aminoisoquinolinewith platinum oxide in acetic acid plus excess sulfuric acid gave 65% yield of l-amino-5,6,7,8-tetrahydroisoquinoline in addition to 1-aminodecahydroiaoquinoline (167). The structure of the bz-tetrahydro compound was proved by subsequent conversion to 5,6,7,8-tetrahydroisoquinoline. Some amino compounds with the amino group on the benzene ring have also been reduced, but most of them had an additional substituent on the benzene ring and as expected gave Py-tetrahydro derivatives.

XXIV. Quinolyl Ketones The hydrogenation of quinolyl ketones is worthy of considerable discussion because of the variety of results. Selective reduction, aa compared with pyridyl ketones reported in Section XVI, B, is complicated by the benzenoid portion of the molecule. Let us first consider phenacyl compounds. There are too few examples of reductions among them. However, the few reported examples are straightforward. Phenacylisoquinolinium bromide contains a nitrogen atom existing in a completely ionized or nontoxic form (30).It should

246

MORRIS FREIFELDER

therefore follow the pattern set forth in section XVI, B and lead to selective reduction of the pyridine portion over ketone function. The reported hydrogenation does indeed follow the expected path, yielding 2-phenacyl-l,2,3,4-tetrahydroisoquinoline(177). I n the same section it was suggested that if complete protonation could be induced in acid solution, then the ring nitrogen would again be present in a nontoxic form and the ring reduction would be the favored reaction. I n the hydrogenation of 2-phenacylquinoline an alcoholic solution containing hydrochloric acid in the presence of platinum oxide both 2-phenacyl1,2,3,4-tetrahydroquinoline and 2-(2-hydroxy-2-phenethyl)quinoline were obtained (178). The results could be an indication of the existence of equilibrium shown below, each side leading to the expected product.

H

Cle

1

Ha

The reduction of quinolyl ketones of the following type produces results that require some explanation: R

\ a

q\" R

e

R =acetyl, benzoyl, or substituted benzoyl, R' =H

R=H, R'=acetyl or benzoyl All of the hydrogenations to be discussed were carried out in neutral solution. Campbell and his associates (167,179)reported on the reduction of 2-,4-, and 8-acetylquinolines. They reported that when 2- and 4acetylquinoline were hydrogenated at low pressure, 2- and 4-quinolyl methyl carbinol were obtained. This was to be expected since the unshielded or toxic nitrogen atom should inactivate the catalyst (30) for ring reduction. When a little hydrochloric acid was present during the reduction of 2-acetylquinoline, the corresponding pinacol was the main product (179). When 8-acetylquinoline was hydrogenated in neutral

HYDROGENATION O F PYRIDINES AND QUINOLINES

247

solution (167) the authors pointed out that two molar equivalents of hydrogen, instead of one, were absorbed without a break and uptake stopped. They stated that the product obtained appeared to be a dihydroquinolyl carbinol which was not identified except that a benzoate was prepared. Campbell found that when ferrous sulfate was added, according to a procedure of Tuley and Adams used to induce selective carbonyl reduction (180),three equivalents of hydrogen were absorbed. A mixture was obtained but not identified. The difference in reaction between 2- and 8-acetylquinoline can be ascribed to the fact that the 8-acetyl group is closer to the ring nitrogen than the 2-acetyl group. Therefore it can shield the ring nitrogen sufficiently to prevent catalyst inactivation and allow the pyridine portion of the molecule to either reduce before or concurrently with the carbonyl group. This effect in the 8-position is more pronounced with the benzoyl group. Woodward and Kornfeld reported partial hydrogenation of 3-acetylquinoline (181). The structure, 3-acetyl-1,4-dihydroquinoline,was 0 3 - c O c H s

HS Raney Nickel

~

09-cocHs \

‘N

H

proposed for the dihydro compound. Failure to proceed beyond the dihydro state, as well as lack of carbonyl reactivity, was concluded to be due to interaction between the imino and carbonyl group. The hydrogenation of benzoylquinolines is incompletely covered. In the work of deDiesbach (182) 2-, 4-, 6-, and 8-benzoylquinolines were investigated. At room temperature 2-benzoylquinoline was converted to 2-phenylquinolylcarbinol. At elevated temperature and pressure (70-100’ and 50-70 atm), 2- and 4-benzoylquinoline were converted to the corresponding tetrahydroquinolyl carbinols. When 8-benzoylquinoline was hydrogenated a t room temperature and low pressure, the tetrahydro carbinol was obtained. When molecular models of 2- and 8- benzoylquinoline were examined, it appeared that the 8-benzoyl group had a far greater shielding effect on the ring nitrogen. It is likely that the 8-benzoyl group has the same effect on the nitrogen atom of quinoline as 2- and 6- substituents have on the pyridine nitrogen as suggested by Adkins (3), and that for this reason, ring reduction takes place so readily. On the basis of catalyst inhibition by the quinoline nitrogen, it might be expected that, in the case of 6-benzoylquinoline, the carbonyl group would be reduced selectively. However, in deDiesbach’s investigation (182)the ring was reduced without affecting the ketone in every in-

248

MORRIS FREIFELDER

stance. In some later work with di[quinolyl-(6)]-ketone (183)’di-[1,2,3, 4-tetrahydroquinolyl-(6)I-ketone was obtained, again showing selective ring reduction with a ketonic substituent in 6-position.

XXV. Concluding Remarks It was the author’s purpose not to prepare a comprehensive review but rather to examine catalytic hydrogenations among this group of compounds in order to determine how further knowledge may be gained. Accordingly, some pertinent references may not have been included. If some of the suggestions made in the various chapters of this text induces further investigation of reduction by catalytic means among pyridines and quinolines, then the author’s purpose will indeed be amply served.

ACKNOWLEDGMENTS The author is indebted to Dr. Warren J. Close, Head of the Organic Chemistry Department of this laboratory for the critical editing of this study and for the opportunity of discussing some of the reactions. We are particularly grateful to Dr. John Tadanier, also of this laboratory, for some of the ideas gained from him on reviewing possible mechanisms of action.

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W.,Connor, R.. and Adkins, H., J . Am. Chem. Soc. 54,1661 (1932). 5. Lazier, W. A., U.S. Patent 2,018,680 (1936). 6. Adkins, H., and Connor, R., J. Am. Chem. SOC. 58, 1081 (1931). 7. Lavignino, E. R., Chauvette, R. R., Cannon, W. N., and Kornfeld, E. C., J. Am. Chem. SOC.82. 2609 (1960). 8. Freifelder, M., and Stone, a.R., J. Org. Chem. $6,3806 (1961). 9. Skita, A.. and Brunner, W., Ber. 49, 1697 (1916). 10. Hamilton, T. S., and Adams, R., J. Am. Chem. SOC. 50,2260 (1028). 11. Adkins, H., and Billica, H. R., J. Am. Chem. SOC. 70,696 (1948). 12. Ames, D. E., and Bowman, R. E., J. Ghem. SOC.p. 1067 (1948). 13. Levy, J., U.S. Patent 2, 813,100 (1967).

HYDROGENATION OF PYRIDINES AND QUINOLINES

14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24.

249

Rosenblatt, E. F., U.S. Patent 2,675,390 (1954). Gilman, G., and Cohn, G., Advancea i n Catal. 9,733 (1957). Freifelder, M., Robinson, R. M., and Stone, G. R., J . Org. Chem. 27,284 (1962). Rylander, P. N., and Steele, D. R., Engelhard Induatriea Tech. Bull. 2, 19 (1962). Walker, G. N., J. Org. Chem. 27, 2966 (1962). Freifelder, M., Unpublished results. Adkins, H., “Reactions of Hydrogen,” p. 64. Univ. of Wisconsin Press, Madison, Wisconsin, 1937. Zelinsky, N. D., and Borisov, P., Ber. 57, 150 (1924). Sadikov, V . S., and Mikhailov, A. K., J. Ruaa. Phya. Chem. SOC.58, 527 (1927); Chem. Abstr. 21, 3364 (1927). Sadikov, V . S., and Mikheilov, A. K., Ber. 61,421 (1928). Sugino, K., and Mizuguchi, J., J. Chem. SOC.Japan 59, 867 (1938); Chem. Abatr. 82, 9095 (1938).

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28.aSawa, Y., and Inoue, K., Ann. Rept. Shwrwgi Research Lab. 4, 26 (1954); b. Ibid. F. 31; Chem. Abalr. SO, 15529 (1956). 29. Rice, R. G., and Kohn, E. J., J. Am. Chem. SOC.77, 4952 (1955); Ainsworth, C., ibid. 78, 1635 (1956). 30. Maxted, E. B., AdvanceainCatal. 3, 129 (1951). 31. Maxted, E. B., and Walker, A. G., J. Chem. SOC.p. 1093 (1948). 32. Brady, L. E., Freifelder, M., and Stone, G. R., J . Org. Chem. 26,4757 (1961). 33. Freifelder, M., and Robinson, R. M., French Patent 1,277,823 (1961). 34. Smith, H. A., in “Catalysis” (P. H . Emmett, ed.), Vol. V, p. 175. Reinhold, New York, 1957. 35. Adkins, H., and Cramer, H. I., J. Am. Chem. SOC.5% 4349 (1930). 36. Ushakov, M. I., and Bronevskii, A. I., Zhur. Obehchel Khim. 7, 750 (1937); Chem. Abatr. 31, 5799 (1937). 37. Freifelder, M., Unpublished work. 38. Overhoff, J., and Wibaut, J. P., Rec. trav. chim. 50, 957 (1931). 39. Ushakov, M. I., and Yakovleva, E. V., Zhur. Obahchei Khim. 7, 753 (1937); Chem. Abatr. 31, 5799 (1937). 40. Ushakov, M. I . , and Promyslov, M. Sh., Zhur. Obahche‘l Khim. 17, 1016 (1947); J . Oen.Chem. U.S.S.R. (Engl. Tranal.)17,1022, (1947);Chem. Abstr.42,4583 (1948). 41. Norton, T. R., Seibert, R. A,, Benson, A. A., and Bergstrom, F. W., J. Am. Chem. SOC.68, 1572 (1946). 42. Brown, H. C . , and Eldred, N. R., J. Am. Chem.Soc. 71,449 (1949). 43. McElvain, S. M., and Carney, T. P., J. Am. Chem. SOC.68,2592 (1946). 44. Hess, K., and Wissing, F., Ber. 48, 1907 (1915). 45. Hess, K., and Liebbrandt, F., Ber. 50, 385 (1917). 46. Vorhees, V., and Adams, R., J. Am. Chem. SOC.44, 1397 (1922). 47. McElvain, S . S., and Adams, R., J. Am. Chem.SOC.45,2738 (1923). 48. Sorm. F., Collection Czechalov. Chem. Commum. 18, 57 (1948). 49. Freifelder, M., J. Org. Chem. 27, 4046 (1962). 50. Raasch, M. S., J . Org. Chem. 27, 1406 (1962). 51. Freifelder, M., J . Org. Chem. 28, 1135, (1963).

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Freifelder, M., J. Org. Chem. 26, 1835 (1961). Freifelder, M., Unpublished work. Freifelder, M., J. Org. Chem. 28, 603, (1963). King, J. A., Hofman, V., and McMillan, F. H., J . Org. Chem. 16, 1100 (1951 Galinovsky, F., Vogl, O., and Weiser, R., Monatsh. 83, 114 (1952). Kleiman, M., and Weinhouse, S.,J. Org. Chern. 10,562 (1945). Hall, H. K., Jr., J . A m . Chem. SOC.82, 1209 (1960). Webb, J. L., and Convin, A. H., J . A m . Chem. SOC. 66, 1456 (1944). Grob, C. A., and Itenk, E., Helv. Chim. Acta 37, 1672 (1954). Lasslo, A,, Marine, W. M., and Waller, P. D., J . Org. Chem. 21, 958 (1956). 62. Ratusky, J., and Sorm, F., Collection Czechoslov. Chem. Communs. 19, 340 (1954). 63. Ratusky, J., Reiser, A., and Sorm, F., Collection Czechoslov. Chem. Communa. 20, 801 (1955). 64. Wheland, G. W., “Advanced Organic Chemistry,” 3rd ed., Sec. 14.9 p. 706. Wiley, New York, 1960. 65. Adkins, H., “Reactions of Hydrogen,” p. 60. Univ. of Wisconsin Press, Madison, Wisconsin, 1937. 66. Cavallito, C. J., and Haskell, T. H., J . A m . Chem. SOC.66, 1166 (1944). 67. Willstatter, R., and Waldschmidt-Leitz, E., Ber. 54, 113 (1921). 68. Biel, J., U.S. Patent 2,802,007 (1957). 69. Marion, L., and Cockborn, W. F., J . A m . Chem. SOC.71, 3402 (1949). 70. Hall, H. K., Jr.,J. A m . Chem. SOC.80,6412 (1958). 71. Koelsch, C. F., and Carney, J. J.,J . A m . Chem. SOC.72, 2285 (1960). 72. Kao, C. H . , J . Chem. Eng. China 15,80 (1948);Chem. Abstr. 44,3993 (1950). 73. Biel, J. H., Friedman, H. L., Leiser, H. A., and Sprengler, E. P., J . A m . Chem. SOC. 74, 1485 (1952). 74. Grave, T. B., J . A m . Chem. SOC. 46,1460 (1924). 75. Ruzicka, L., and Fornasir, V., Helv. Chim. Acta 3,806 (1920). 76. Ruzicka, L., Helv. Chim. Acta 4, 472 (1921). 77. Gautier, J., Compt. rend. acad. sci. 205, 614 (1937). 78. Gautier, J., Compt. rend. acad. sci. 225, 880 (1947). 7 9 . Galinovsky, F., and Stern, E., Ber. 77, 132 (1944). 80. Takahashi, T., andKariyone, K.,Chem. Pharm. Bull. 8, 1106 (1960). 81. Emmert, B., German Patent 292,456 (1914). 82. Armit, J. W., and Nolan, T. J., J . Chem. SOC.p. 3023 (1931). 83. Bowden, K., and Green, P. N., J.Chem. SOC.p. 1795 (1954). 84. Ishii, T.,J. Pharm. Soc. Japan 71, 1097 (1951);Chem. Abstr. 46,5042 (1952). 85. Campbell, K. N., Ackerman, J. F., and Campbell, B. K., J . Org. Chem. 15, 337 (1950). 86. Paul, R., and Tchelitcheff, S., Bull. SOC. chim. France p. 341 (1947); Baker, B. R., and McEvoy, F. J., J . Org. Chem. 20, 136 (1955). 86a. Stach, K., Thiel, M., and Bickelhaupt, F., Monatsh. 93, 1090 (1962). 87. Freifelder, M., Unpublished work on (hydroxyalkyl) piperidines. 87a. Brown, H. C . , and Eldred, N. R.,J. A m . Chem. SOC.71,445 (1949). 88. Boekelheide, V., and Rothchild, S., J. A m . Chem.Soc. 70,864 (1948). 89. Nienburg, H., Ber. 70, 635 (1937). 90. Walker, G. N., and Moore, M. A., J. Org. Chem. 26,432 (1961). 91. Wheland, G. W., “Advanced Organic Chemistry,” 3rd ed., Sec. 14.9 p. 709. Wiley, New York, 1960. 92. Orthner, L., Ann. 456, 226 (1927).

52. 53. 54. 55. 56. 57. 58. 59. 60. 61.

HYDROGENATION OF PYRIDINES AND QUINOLINES

25 1

93. Yakhontov, L. N., Yatsenko, S. V., and Rubstov, M. V., Zhur. Obshchell Khim. 28, 3115 (1958); Chem. Abstr. 53, 10205 (1959). 94. Kirsanov, A. V., and Ivashchenko, Y. N., Bull. soc. chim. France p. 2279 (1936). 95. Birkhofer, L., Ber. 75, 429 (1942). 96. Shapiro, S. L., Soloway, H., andFreedman, L . , J . Org. Chem. 26, 818 (1961). 97. Profft, E., Chem. Tech. (Berlin)6 , 4 8 4 (1954);Chem. Abstr. 50,332 (1955). 98. Britton, E . C., and Horsley, L. H., U.S. Patent 2,834,784 (1958). 99. Sperber, N., Papa, D., Schwenk, E., Sherlock, M., and Fricano, R., J . A m . Chem. SOC.73, 5752 (1951). 100. Smith, C. R . , J . Am. Chem. SOC.50,1936 (1928). 101. Beyerman, H. C., and Bontekoe, J. S., Rec. trav. chim. 74, 1395 (1955). 102. Frank, R. L., Pelletier, F., and Starks, F. W., J . Am. Chem. SOC.70, 1767 (1948). 103. Supniewski, J. V., and Serafinovna, M., Arch. chem. farm. 3, 109 (1936); Chem. Abstr. 33, 7301 (1939). 104. Lyle, R. E., and Lyle, G. G., J . Am. Chem. SOC.76,3536 (1954). 105. Lyle, R. E., Perlowski, E. F., Troscianiec, H. J., and Lyle, G. G., J . Org. Chem. 20, 1761 (1955). 106. Grob, C. A., and Ostermayer, F., Helv. Chim. Acta 45, 1119 (1962). 107. Kuznetsov, S. G., and Libman, N. M., J . aen. Chem. (U.S.S.R.)(Engl. Transl.) 29, 2385 (1959). 108. Taft, R. W., Jr., in “Steric Effects in Organic Chemistry” (M. S. Newman, ed.), Chapter XIII, p. 595. Wiley, New York, 1956. 109. Tomita, K., J . Pharm. SOC. Japan 71,220 (1951) (in English). 110. Sugasawa, S., and Degueli, Y., J . Pharm. SOC.Japan 76, 968 (1956); Chem. Abstr. 51, 2771 (1957). 111. Winterfeld, K., and Augstein, J., Ber. 90, 868 (1957). 112. Baker, B. R., and McEvoy, F. J.,J. Org. Chem. 20,118 (1955). 113. Boehringer, A., British Patent.312,919 (1930). 114. Scheuing, G., and Winterhalder, L., Ann. 473, 126 (1929). 115. Howton, D. R., and Golding, D. R. V., J . Org. Chem. 15, 1 (1950). 116. Luke& R., Kova?, J., and Blaha, K., Collection Czechslov. Chem. Communs. 21, 1475 (1956). 117. Krohnke, F., and Faslod, K., Ber. 67, 656 (1934). 118. Riegel, B., and Wittcoff, H., J . Am. Chem. SOC.68, 1805 (1946). 119. Truitt, P., Buttram, J. R., and Herd, R., J . Am. Chem. SOC. 74, 5448 (1952). 120. Truitt, P., Bryant, B., Goode, W. E., and Arnwine, B., J . Am. Chem. SOC.74, 2179 (1952). 121. Crook, K. E., and McElvain, S. M., J . A m . Chem. SOC.52, 4006 (1930). 122. Lyle, R. E., Leone, S. A., Troscianiec, H. J., and Warner, G. H., J . Org. Chem. 24, 330 (1959). 123. Lyle, R. E., Warner, G. H., and Nelson, D. A., Bol. 8oc. quim. Peru. To be published during 1962-63. 124. Clemo, G. R., Raper, R., and Vipond, H. J., J . Chem. SOC.p. 2095 (1949). 125. Hohenlohe;Oehringen, K., Monatsh. 93, 586 (1962). 126. Luke& R., Strouf, O., and Ferles, M., Chem. Listy 51, 923 (1957); Chem. Abstr. 51, 14724 (1957). 127. Lyle, R. E., and Warner, G. H.,J. Med. Pharm. Chem. 3,597 (1961). 128. Pitha, J., and Ernest, I.,Collection Czechoslov.Chem. Communs. 24,2632 (1959). 129. Wibaut, J. P., Kloppenburg, C. C., and Beets, M. G. J., Rec. trav. chim. 63, 134 (1944).

262 130. 131. 132. 133.

MORRIS FREIFELDER Wibaut, J. P., and Kloppenburg, C. C., Rec. trav. Chim. 65, 100 (1946). Boekelheide, V., and Agnello, E. J.,J . Am. Chem. SOC.72, 6006 (1960). Boekelheide, V., and Rothchild, S., J . Am. Chem.SOC.71,879 (1949). Boekelheide, V., Linn, W. J.,O’Grady, P., and Lamborg, M.. J . Am. Chem. SOC.75, 3243 (1963).

134. 135. 136. 137. 138. 139. 140. 141.

von Braun, J., Blessing, G., and Zobel, F., Ber. 56, 1988 (1923). Kuhn, R., and Kirschenlohr, W., Ann. 600,116 and 126 (1966). Umio, S., J . Pharm. SOC.Japan 78, 1072 (1968). Kallischnigg, R., and Krastinat, W., German Patent 1,032,262 (1968). Umio, S., J . Phrm.Soc. Japan 79, 1133 (1969). Winterfeld, K., Hiiring, W., and Schmialek, P., Arch. Pharm. 295, 24 (1962). Veer, W. L. C., and Goldschmidt, St., Rec. trav. chim. 65,793 (1946). McCarthy, F. J., Tilford, C. H., and Van Campen, M. G., Jr., J . Am. Chem. SOC.79, 472 (1967).

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Bader, H., and Oroshnik, W., J . Am. Chem.Soc. 79,6686 (1967). Bader, H., and Oroshnik, W., J . Am. Chem. SOC.81, 163 (1969). Spiith, E., and Kuffner, F., Ber. 68,494 (1936). Ochiai, E., Tsuda, K., and Ikuma, S., Ber. 69, 2238 (1936). Danens, a.,Compt. rend. acud. aci. 149, 1001 (1909). Ipatief, V., Ber. 41, 991 (1908). Tsushima, S., and Sudzuki, J., J . Chem. SOC.Japan 64, 1296 (1943); Chem. Abatr. 41, 3801 (1947).

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Skita, A., andMeyer, W. A., Ber. 45,3689 (1912). Huckel, W., and Stepf, F., Ann. 458, 163 (1927). Helfer, L., Helv. Chim. Acta 6 , 786 (1923). Skita, A., Ber. 57,1977 (1924). Witkop, B., J . Am. Chem. SOC.70,2617 (1948). Witkop, B., J . Am. Chem. Soc. 71,2659 (1949). Metayer, M., Bull. aoc. chim. France p. 276 (1962). Zenitz, B. L., Mach, E. B., andMoore, M. L . , J . Am. Chem.Soc. 69,1117 (1947). Kindler, K., Hedemann, B., and Scharfe, E., Ann. 560,216 (1948). Gray, A. P., and Heitmeier, D. E., J . Am. Chem. SOC.80,6274 (1958). Woodward, R. B., and Doering, W. E., J . Am. Chem. SOC.67,860 (1946). Rapala, R. T., Lavignino, E. R., Shepard, A. R., and Farkas, E., J . Am. Chem. SOC.

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Freifelder, M., and Stone, G. R., J . Am. Chem. SOC. 50,6270 (1968). von Braun, J., Petzold, A., and Seeman, J., Ber. 55, 3779 (1922). von Braun, J., Gmelii, W., and Schultheiss, A., Ber. 56,1338 (1923). von Braun, J., Gmelin, W., and Petzold, A., Ber. 57,382 (1924). Oldham, W., and Johns, I. B., J . Am. Chem.SOC.61,3289 (1939). Skita, A., end Wiilff, C., Ber. 59,2683 (1926). Campbell, K. N., Kerwin, J. F., La Forge, R. A., and Campbell, B. K., J . Am. Chem. SOC.68, 1844 (1946). Ochiai, E., Kaneko, C., Shimada, I., Murata, Y.,Kosuye, T., and Kawasaki, C., Chem. P h r m . Bull. Tokyo 8,126 (1960). Isogai, K., J . Chem. SOC. Japan 81, 1694 (1960); Chem. Abatr. 56,2420 (1962). Wieland, H., Hettche, O., and Hoshina, T., Ber. 61,2371 (1928). Chiaverelli, S., and Marini-Bettal, G. B., Uazz. chim. Gal. 82,86 (1962); Chem. Ah & .

79, 3770 (1957).

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47, 6944 (1963). 172. Solomon, W., J . Chem. SOC.p. 129 (1947).

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173. von Braun, J., and Lemke, G., Ann. 478,176 (1930). 174. Miyamoto, T., and Kataoka, A., J . Phamn. SOC. Japan 58,478 (1939); Chem. Abatr. aa, 9306 (1939). 175. Miyaki, K., and Kataoka, A., J . Phatm. SOC. Japan 59,222 (1939); C L m . Ab&. 84, 7910 (1940). 176. Ochiai, E., and Kawezoe, Y . , Pharm. Bull. (Tokyo)5, 606 (1957). 176a. Georgian, V., Harrison, R. J., and Skaletzky, L.L., J . Org. Chem. 27, 4571 (1962). 177. Goode, W. E., J . Am. Chem. SOC. 70, 3946 (1948). 178. Wright, P. E., and McEwen, W. E., J . Am. Chem. SOC. 76,4540 (1964). 179. Campbell, K. N., Helbing, C. H., and Kerwin, J. F., J . Am. Chem. SOC.68, 1840 (1946). 180. Tuley, W. F., and Adams, R., J . Am, Chem.Soc. 47,3061 (1925). 181. Woodward, R. B., and Kornfeld, E. C., J . Am. Chem. SOC. 70,2508 (1948). 182. de Diesbach, H., Pugin, A., Morard, F., Nowaczinski, W., and Dessibourg, J., Helv. Chim. Acta 35, 2322 (1952). 183. Kiihnis, H., and de Diesbach, H., Helv. Chim. Acta 41, 894 (1958).

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Modern Methods in Surface Kinetics* Flash Desorption. Field Emission Microscopy. and Ultrahigh Vacuum Techniques

GERT EHRLICH General Electric Reseawh Laboratory. Schenectady. New York

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I Macroscopic Rate Studies .............................................

256 A . Adsorption Measurements ......................................... 256 B. Flash Desorption .................................................. 271 C. Energy Parameters in Macroscopic Adsorption Experiments ............ 304 I1. Field Electron Microscopy ............................................ 311 A . Physical Principles ................................................ 312 B Experimental Considerations ....................................... 330 C. Comparison of Techniques-The Interaction of Rare Gases with Metals ... 334 I11 Field Ion Microscopy ................................................ 347 A . Physical Principles ................................................ 349 B . High Field Effects in Practice ....................................... 360 C Direct Observation of Adatoms ..................................... 367 D . Practical Implementation ......................................... 386 391 E . Prospects ........................................................ 391 IV Ultrahigh Vacuum Techniques ........................................ A . Are Ultrahigh Vacuum Techniques Necessary? ........................ 392 B The Golden Rule .................................................. 393 C. Vacuum Systems ................................................. 393 401 D . System Operation ................................................. E Pressure Measurements ............................................ 404 F Gas Handling .................................................... 414 G Sources ofEquipment ............................................. 421 424 References ........................................................

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The outstanding feature of the past decade has been the development and perfection of a variety of experimental techniques that allow a new and deeper insight into the structure of surfaces. and into the elementary atomic events occurring on them . I n many instances these techniques have merely confirmed the correctness of previous proposals; in others. entirely new understanding of the interaction of gases and solids has been achieved . What is important. however. is that experiments. properly executed and interpreted. can now remove from the realm of

* Supported in part by the Directorate of Chemical Sciences. Air Force Office of Scientific Research . 256

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speculation most of the elementary facts dealing with interplay of gases and an initially clean surface. This tremendous advance has, in large measure, been made possible by the reduction of ultrahigh vacuum techniques to a matter of complete routine. For the first time it has become easy t o specify exactly the environment at an interface and to perform meaningful measurements on samples with a surface area of less than a square centimeter. As a consequence, considerable progress has been made in apprehending the relation between surface structure and the elementary rate processes important in adsorption. A proper judgment of the validity of these findings, as well as any extension of such work, must rest upon a detailed appreciation of the experiments involved. It is the aim of this article to review the experimental methods upon which these advances have been based-the flash filament technique, flash desorption, field emission and field ion microscopy, and the use of ultrahigh vacuum procedures. Flash desorption, although still dependent upon macroscopic wire samples, has made it possible to quantitatively measure rate processes involving the transfer of molecules between the gas phase and the solid. I n principle, even the dependence of surface kinetics on atomic structure could be established by studies on macroscopic samples. The specification of surface features below lOOA is difficult on such samples. For this purpose measurements in the field emission and ion microscope are more convenient and powerful-they afford a view of the surface on a scale approaching atomic dimensions. However, such work can only be properly carried out against a background of detailed macroscopic information, and it is this sequence from macroscopic measurements to direct observation of atoms that will be followed here. Results and insights obtained by modern experimental techniques have been outlined elsewhere (1)--they will be presented here only to the extent that the findings illuminate experimental problems. The theme to be remembered throughout is that dependence upon a single technique, even though a powerful one, is dangerous, particularly when the execution of a variety of modern surface experiments is straightforward.

I. Macroscopic Rate Studies A. ADSORPTION MEASUREMENTS 1. Historical

The initial process in any interaction of gas with a solid involves the adsorption of molecules from the gas phase. The kinetics of this event

MODERN METHODS IN SURFACE KINETICS

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are conveniently specified by the sticking coefficient k,, that is by the probability that a molecule striking a surface will actually be adsorbed on it. Classically the amount of material adsorbed on a solid exposed to a gas is measured by ascertaining the pressure diminution in the system during adsorption. I n principle, the kinetics of this adsorption step can be followed by measuring the pressure as a function of time. I n practice this has not been easy to do. With clean metal surfaces the rate of adsorption is of the same order of magnitude as the rate of supply from the gas phase; this process can therefore only be observed if pressures in the range below 1 p can be measured quickly. More precisely, if adsorption on a sparsely populated interface (fractional 1/10) is of interest, then the pressure p (in mm) and coverage 9 the length of the experiment T (in sec) must satisfy the condition p x 7 10-6. I n general, pressure detectors capable of such performance have become available only recently, in the form of the inverted (BayardAlpert) ionization gauge (described in Section IVYE, 1). As early as the 192OYs,however, Langmuir and his co-workers (2) did have the means for detecting pressures as low as 10-15 mm for one particular vapor, cesium, and were therefore able to lay the foundations for modern studies of adsorption kinetics. Cesium, on striking a tungsten wire maintained a t a temperature P ! > 1200°K evaporates as an ion (3). By surrounding this wire with a cylindrical electrode a t a negative potential the rate a t which the wire is struck, and therefore the pressure, can be measured as an ion current to the cylindrical collector. The amount of cesium adsorbed on a tungsten surface can be similarly determined by heating (flashing) the surface; the cesium evaporates as a, positive ion and the surface cover is therefore directly registered as a current burst. This technique is quite limited, however. The rate a t which atoms evaporate is related to hLi, the rate of ion evolution by N

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where 4 is the electron work function of the metal surface and I the ionization potential of the gas phase atom colliding with the surface; wi and w, denote the statistical weighting factors of ion and atom, respectively. Only for I << 4 will ion evolution dominate. This technique is therefore confined to the alkali metals, such as cesium and rubidium. There is yet a further limitation. As the surface concentration rises, the work function 4 drops. At higher coverages an in-

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creasing fraction of the desorbed material is therefore made up of neutral atoms, and is not detected. This limitation can be overcome with a separate sample and detector, as was first suggested by Lewi Tonks and carried into practice by Taylor and Langmuir ( 3 ) .Tonks' suggestion was to mount a second wire (D) parallel to the sample (S)as shown in Fig. 1. The detector D is kept a t a

-

FIQ. 1. Surface ionization gauge for flash desorption of Cs ( 3 ) . Cs atoms, evaporating when sample S is flashed, are ionized on the detector wire D (T 1200°K) and counted as ions at the collector cylinder C,. G,, G,-guard rings; A-galvanometer. Sample 4 1 volts negative to D, 4 volts positive to G,, G,, C,. potential N

high temperature (T > 1200" K ) and is positively biased with respect to the collecting electrode. Cesium striking the detector wire is ionized, and the current to the cylinder therefore serves to measure the ambient pressure. When the sample has been coated by this ambient, its temperature is raised to -1800" K , with the potentials adjusted so that the cesium evaporates as atoms. A fixed fraction of the evaporating atoms, determined by the geometry of the system, is intercepted by the detector and ionized. This current burst therefore gives a measure of the amount desorbed. By these methods it was possible t o measure surface coverages from 0 10-4 into the multilayer range. Despite its tremendous power, Tonks' arrangement is still limited to the study of easily ionizable metal atoms, inasmuch as it depends on a surface ionization detector. This limitation was removed by Apker ( 4 ) who substituted a hot filament ionization gauge as a detector and was therefore, for the first time, able to measure the evolution of electronegaN

MODERN METHODS IN SURFACE KINETICS

259

tive gases as well. With this development the stage was set for the systematic application of flash filament techniques to the study of the interaction of gases with solids. The only further technical advance, beyond refinements of experimental methods, occurred a decade later with the introduction of the mass spectrometer to analyze the composition of the evolved gas ( 5 , 6 ) . 2. Determination of Adsorption Kinetics

The Bayard-Alpert gauge measures pressures down to 10-lo mm; as shown in Fig. 2, the evolution of less than 1012 molecules occurring over

tFIG. 2. Desorption of Xe from W using Bayard-Alpert ionization gauge (7). N-gas density; t-time after heating sample, Sample area = 1.3 cm2,cell volume = 650 cmS.

a time span of milliseconds can be determined routinely. Even the most rapid adsorption process may therefore be followed quantitatively by techniques now standard. Rates of adsorption measured by the flash-filament technique are summarized in Fig. 3. These data are based on a knowledge of three related parameters: A(dn/dt),the absolute rate of adsorption occurring on a sample of geometrical area A , p the pressure in the system, and n the instantaneous surface concentration. Procedures for determining these parameters are outlined in the next few pages. a. Closed Systems. Consider a system of volume V containing the adsorption sample and a gauge arranged so as to measure the pressure a t the sample. The sample having been properly prepared, a fixed quantity of the gas whose adsorption kinetics are of interest is introduced a t time t = to. I n an ideal experiment, only the sample surface

interacts strongly with the gas, and gas admission is accomplished in a time interval short compared to any process of physical interest. The change in gas density N after admission is complete is then dictated by

1.0-

' ' ' ' "

-

I

1

1

i .

IO-" a n , molccu,les/em* SURFACE CONCENTRATION

FIQ.3. Sticking coefficients determined by flash desorption, as a function of surface concentration n. Data: Cs (3); N, ( 8 ) ;CO (9); Xe (7).

competition between the pumping speed S, of the surface and reevolution from the sample, occurring at a rate F,. This competition process is described by*

Immediately following the introduction of gas adsorption predominates. As indicated in Fig. 4 the pressure in the system drops and continues It is convenient to express these conservation conditions in terms of the density N , rather than the pressure p , since the ion gauge actually responds to the former. The two terms will be used interchangeably, hopefully without causing confusion.

261

MODERN METHODS I N SURFACE KINETICS

t o do so until evaporation finally compensates. S,, the pumping speed, is related to the sticking coefficient k, through

V being the mean speed of the molecules in the gas phase, known from kinetic theory. The sticking coefficient can therefore be determined

- 10"

-IO'O -

E

-

a

-

3

v, v)

-

'I

NO EVAPORATION

n E

Q

-3

-lo9 =

- cJi

- z c3

-108 AMOUNT ADSORBED =

- CLOSED SYSTEM

-10'

FIG. 4. Pressure cycle for adsorption experiments in a closed system on sample of unit area. t,-start of adsorption; t,--start of heating; t,---end of desorption. Amount adsorbed = V ( N o - N 1 ) = amount desorbed.

from the slope of the density vs time curve, provided reevaporation, F,, is negligible. Even if this criterion is not satisfied, the rate constants may still be deduced by measuring d N / d t , starting a t different initial

values of N . A plot of the net rate of adsorption ( N S , - F p ) against N , keeping the amount adsorbed nA constant then yields a straight line (10) of slope S,, arid a y-intercept equal to the evaporation rate. An exampIe of this procedure is given in Fig. 5 for the physical adsorption of Xe on W. I

n =52x lo'*

25 -

120110l2 MOLECULESlcm:

-

20-

FF/A

52x 10" 69 86

-5-

i,

Ib

A

120 137

3b

410

sb

40

E,

-

molehm$nim kml/mnk 0

-

.0094x10n .31 .67 2.02 4.68

m

eb

6.5 6.0 59 5.8 5.7

-

40,n:12Ox10~4

-

FIG. 5. Net rate of adsorption as a function of impingement rate. Xe at T, = 300°K on tungsten at 80°K (7). Slope = sticking coefficient k,; intercept = evaporationrate = F,/A.

The net quantity of gas held on the surface a t any instant t , must just equal the decrease in the number of molecules in the gas phase, or, more formally t

l0

If a t time t , the sample is heated to a temperature a t which evolution is complete, then the starting conditions are restored. The amount desorbed in the interval ( t , - t J is obviously % ( t i )A

=

j: tl

(FR

- N X p ) dt

=

V (No - N l ) ,

(4)

MODERN METHODS I N SURFACE KINETICS

263

The return to the starting pressure on heating insures that there are no extraneous side reactions occurring. I n such an ideal system the entire kinetics of adsorption can therefore be deduced very simply from the variation of pressure with time. b. Flow Systems. In a flow system the pressure is dictated by competition between the rate a t which gas is admitted to the cell, FA, and the rate at which it is adsorbed on the surface, or else removed by flow through a controlled leak to the pumps NS,. The equation for mass balance is therefore

If the rate of gas admission F A and the leak rate 8, are known, then the rate of adsorption on the sample ( N S , - FF) can be deduced from the pressure vs time curve, just as in the closed system. At the start of the experiment FA the flow of gas into the system, as well as the leak rate S , are adjusted to achieve the desired starting concentration N o . This is accomplished with the sampIe surface at a high temperature a t which no net adsorption occurs. A steady state is established in which the flow of gas into the cell balances the leak rate out; that is FA = NOSE. Once the sample filament is allowed to cool, it acts as a sink for gas and the pressure drops. This diminution continues only long enough for a new steady state to be re-established, as indicated in Fig. 6. The net flow of gas into the system ( N o - N ) SE is now balanced by the net loss of gas to the sample (FF- NS,). I n a closed system, a steady state is established only when the rate of evaporation of the adsorbed layer becomes comparable with the rate of adsorption on the sample. In a flow system, the steady state pressure achieved with the sample adsorbing can be adjusted by altering the flow of gas into the cell. Once a new steady state is reached, VdNI dt = 0. The net rate of adsorption is just equal to the net flow into the cell, and is therefore immediately known, provided the exit speed S , has been previously determined. The number of molecules adsorbed a t time t , after the start of the experiment is given by

QERT EHRLICH

264

That is, the amount adsorbed is now equal to the difference in the amount of gas in the container a t time to and a t time t , , plus the net inflow Qi. This quantity is equal to the product of the exit speed and the shaded area in Pig. 6. Just as in the closed system, a record of the pressure diminution in the cell suffices to determine all the rate parameters.

FLOW SYSTEM

h

TIME (SECONDS)

FIG.6. Pressure cycle in flow system during adsorption on sample of unit area. Qi = amount of gas added to cell during adsorption interval il-t,,; Qo = amount lost during desorption interval t , - t i ; S , = exit speed; P A = rate of gas entry from reservoir.

The concentration of adsorbed material may also be found by flashing the surface to an appropriate high temperature in the interval t to t 2. I n an obvious extension of Eq. (4), the amount evolved is then

If the heating is done so rapidly that desorption is substantially complete before appreciable amounts of gas can leave the cell, the gas

MODERN METHODS I N SURFACE KINETICS

265

density rises to a peak value ( N , - N , ) . This then corresponds to the number of molecules adsorbed on the filament up to time t , , corrected for a small net loss during the heating. The peak will of course exceed the original steady state value No, inasmuch as gas is liberated that entered the cell from the reservoir during the adsorption interval, Such rapid heating is not a prerequisite for determining the amount adsorbed. If the heating is continued until the original steady state is restored N a = N,. The amount of gas added to the cell in the adsorption must then balance the loss of material during the desorption interval, Qi, step, Qo. The shaded area must therefore be equal to that crosshatched in Fig. 6. This return to the starting conditions also insures that the total number of molecules adsorbed equals the number desorbed in the interval ( t , - t , ) . The identity of these quantities serves as a convenient check on experiment. In a static system, knowledge of the amount adsorbed and desorbed depends upon the same experimental parameter, the density difference (3, - N).,This need no longer hold true in a flow system. During the adsorption interval ( t l - .to), the net flow into the cell in Eq. (6) usually determines the amount adsorbed. This quantity therefore depends critically on the value of the exit speed S,, as well as upon the density vs time curve. I n desorption, conditions can be arranged to make the flow term small. The amount desorbed will then be determined by the peak height ( N , - N , ) , with only a minor correction involving the pumping speed. I n a flow system, the rate parameters can therefore be determined in two essentially independent ways, (1) From the pressure drop as a function of time, as indicated on the previous page. (2) By measuring the amount of gas desorbed after different adsorption intervals A t . Differentiating the curve of n vs At obtained under reproducible flow conditions then yields the net rate of adsorption. This can again be separated into contributions from adsorption and desorption rates by determining the pressure dependence. Only adsorption and desorption have been specifically considered here. It is obvious that any surface reaction resulting in a pressure or concentration change in the gas phase is susceptible to study by this technique. 3 . Experimental Implementation a. Closed Systems. The ideal systems described so far are difficult to achieve in practice. In a real closed vessel, adsorption occurs on the container walls and gas is also removed by the gauge. The envelope, coated with gas a t the starting pressure, may further reemit gas as the

266

OERT EHRLICH

pressure in the cell drops. The pressure in a real closed system is therefore dictated by the relation.

where F A and S, now designate gas flow or pumping of any type except that due to the sample. Pumping by gauge and walls is difficult to suppress. Their effectiveness can usually be measured and appropriate corrections made. Gas evolution from the walls during a run presents a more difficult problem.

B.

FIG. 7. Flow system for adsorption studies, with independent control of entrance and exit speed. F-adsorption sample; I-ionization gauge. Flow out of gas reservoir (B) and into cell controlled by valve V,; gas flow out of cell regulated by adjustable exit port E.

Operation in closed systems is also limited since, for any initial pressure, only a small range of surface concentrations can be conveniently explored. The rapid establishment of the initial gas concentration is hard to achieve. Two ways are apparent. (1) Gas can be admitted through a valve, keeping the sample hot enough to prevent adsorption. (2) Gas may be introduced into the cell by breaking ampules after the filament is cold. In the former, precautions are necessary to prevent inflow of gas even when the valve is closed. Even then, in order to achieve high coverages a high initial pressure is necessary, and the sample will therefore pick up appreciable amounts of gas while still warm; that is, under ill-defined conditions. With gas ampules, on the other hand, the number of introductions is limited and impurities from crushing the glass seals may cause spurious results.

MODERN METHODS IN SURFACE KINETICS

267

Inasmuch as some gas flow through the cell appears inevitable, it is convenient t o design experiments to take full advantage of a flow system. Indeed, no rate studies have been carried out in a closed system using modern techniques. b. Flow Systems. The principles and practice of adsorption measurements in flow systems were first outlined by Molnar and Hartman (II), and almost a t the same time by Wagener (12).The techniques introduced by these investigators differ only in one technical detail-the method of admitting gas-and are otherwise completely identical. A schematic of the equipment for adsorption measurements in a flow system is shown for ready reference in Fig. 7. The flow of gas into the cell is controlled by valve, V A .The rate a t whichgas leaksout of the cell can be adjusted with the magnetically operated ground glass valve E, leading to the pumping system.* With the sample hot, so that no adsorption occurs on it, the rate FA a t which gas enters from the reservoir is just equal to the rate a t which gas is transferred to the pumps, NS,, and

Knowing S,, we therefore know FA. The leak rate S , can be determined very simply by introducing a burst of gas, either by opening the valve or by flashing a filament, and then following the pressure drop, that is, from 1 dN

- = s,.

( N o - N ) dt

*4lternatively, starting with the exit wide open, the port can be put into position suddenly, with the inflow at its appropriate setting. The logarithmic derivative of the pressure rise then also yields the pumping speed of the exit hole. It is convenient to operate with the gas reservoir a t a pressure much higher than that in the cell. F A ,the flow into the cell, is then dependent only on conditions in the reservoir; these are unaffected by the escape of molecules into the cell and remain constant throughout a run, provided the temperature is maintained. The net rate a t which gas enters the cell may also be controlled by interposing a capillary between gas reservoir and cell. This is the method adopted by Wagener and is illustrated in Fig. 8. Both entry and exit of gas occurs through the same capillary, instead of through separate valves. A t the start of the experiment, with no pumping by the sample, the gas density in the reservoir, No, is the same as that in the cell. As

* A more detailed description of these parts is given in Section IV.

268

OERT EHRLICH

adsorption does occur, the pressure in the cell diminishes and a net flow equal to ( N o - N ) S , is established. This is precisely the action in the flow system already described. The use of the capillary has one disadvantage compared with the double valve method: The flow through the cell cannot be as readily altered, since this now requires introduction of an entirely new capillary.

TO PUMPS

a TRAPS

FIG. 8. Flow system for adsorption studies a la Wagener (12).F-sample surface; B-gas bulb. Gas density in reservoir section measured by ion gauge I,, and controlled at a constant setting N o by adjusting valve V. Density N in cell measured by gauge I,. ( N o - N ) 8, = net flow into cell determined by capillary A, with pumping speed S,.

The ability to regulate the flow through the cell is very important. Consider two extreme situations: (1) An unreactive impurity constitutes a fraction zo of the main constituent in the gas bottle. Assuming that there is no fractionation during admission, then in the steady state with the sample hot the ratio of impurity to reactant in the cell will be the same as in the reservoir. However, only the main component can be adsorbed on the sample once this cools, and its concentration is then diminished. The fractional impurity concentration during the adsorption step thus increases to = z0 ( l sI?/sE)*

+

(2) The gas in the cell can interact with the pressure measuring device to produce contamination at the rate* NS,. Assuming that this contaminant adsorbs on the sample a t the same rate as the main constituent, z, the ratio of contaminant to main component is given by z =

sc SF

+

sE

* This is an important effect with all but the most inert gasee.

MODERN METHODS IN SURFACE KINETICS

269

and can be minimized by making 8, large. However, the ratio of the steady state pressure with the filament hot (and therefore inactive) to the steady state pressure when adsorption is occurring on the filament is, in the absence of re-evolution of gas from either the walls or the sample, given by

A large exit speed 5, >> S p establishes an almost constant pressure, This in turn insures that the gas cover of the walls remains constant and that the impurity content of the gas in the cell will not exceed that of the reservoir. At the same time it is clear that such a setting makes accurate rate measurements difficult. The sticking coefficient is determined by [NOIN- 11 and for N,/N ,- 1 this ratio would have to be known with great accuracy. I n the other extreme, when SF >> S, we can approximate a closed system. This maximizes the pressure drop in the cell on adsorption, but brings with it all the evils of a closed system. I n practice, the adjustment of S,/S, must therefore be varied to suit each system. The ratio S,/S, also affects the lowest surface concentration that can be examined by flow techniques. During the initial period of adsorption the sample is not at the temperature of the surroundings. Over the time interval T in which the surface changes from the temperature necessary for cleaning to that of the experiment, adsorption occurs under ill-defined conditions. Significant rate measurements can only be taken with the sample a t the bath temperature. The amount adsorbed during this interval 7 can be minimized by operating a t low steady state pressures, or a t low settings of S,. The first adjustment restricts the number of molecules initially present in the system that may collide with the sample. The second lowers the amount of gas flowing into the cell. The limits over which the exit speed can be changed have already been discussed. Similarly, the starting pressure cannot be lowered a t will. At very low pressures ( p << 10-9 mm) the background of residuals may contribute importantly to the total gas composition. Even if this can be kept in bounds, the usual ionization gauge becomes difficult to handle reliably at p 10-10 mm. The time interval 7 can also be controlled by arranging the sample to maximize heat conduction to the bath. Thus as indicated in Fig. 9 for a 10 mil tungsten filament 16 cm long, the cooling period is on the order of 3 min. On such a sample adsorption rates have N

270

GERT EHRLICH

been accurately determined a t concentrations as low as 10 x 10l2 molecules/cm2. So far it has been implicitly assumed that the density in the cell is uniform throughout and has the same value a t the sample as a t the 6

I

I

-I

I

I

5

I\

AMBIENT TEMPERATURE = 298’K

4

v)

t z

3

$3 4

a t m c 4

2

iI I

-

I UNIT =6.15 x 10-lOmm

I

SAMPLE FILAMENT: 10 mil TUNGSTEN, 16 cm LONG

I

I

I

I

I I

I I

I 2 t-lo, MINUTES

I

I

3

4

-

FIG. 9. Sample temperature and nitrogen pressure following high temperature cleaning (T 22OO0K).Adsorption takes place a t bath temperatures only for ( t - t o ) > 3 min (13).

gauge. When the sample is adsorbing a pressure gradient will be established, however. It is therefore vital in carrying out such measurements to bring gauge and filament as close together as possible. The only practical limitation is the need to shield the sample from the thermal radiation of the gauge, as well as from stray electrons. The order of

MODERN METHODS I N SURFACE KINETICS

27 1

magnitude of these gradients in practical systems will be presented in Section IV, E, b. Conceptually, the determination of adsorption kinetics in flow systems is straightforward. The actual execution of the experiments is, however, beset by numerous small difficulties. Above all else, persistent attention to detail is therefore required to insure success. Information on adsorption rates is still sparse, and despite its difficulties such work is therefore urgently needed.

B. FLASH DESORPTION I n heating the sample a t the end of the adsorption interval, it is not only possible to count the total number of molecules evolved. As is clear from Fig. 10, a qualitative picture of the nature of the binding energies

Fro. 10. Multiple binding states detected by flash desorption. N-gas density, t-time after start of heating. Maximum flash temperature T = 2200°K. Time scale: 1 unit = 100 msec. Surface concentration-N,; ny = 50 X 10l2molecules/cm2, na = 11 X 10l2, = 136 X 10" (8); CO: 12, = 108 X lo'', no, = 206 X lo", no1 = 202 X lo", nos = 62 x 10'2 (9).

of the adsorbed species can be obtained; it is in this fashion that a detailed view of the formation of intermediate binding states in adsorption on metals has been attained ( I ) . Under some circumstances, the quantitative kinetics of the desorption can be measured as well, and rate laws obtained in this fashion are summarized in Table I. Such rate measurements under changing tempera-

212

GERT EHRLICH

tures are certainly not novel. As early as 1930 Urbach (15) studied the well depth for electrons in phosphors by observing the luminescent intensity (proportional to the rate a t which electrons escape from their traps) under continuous heating. Many of the analytical techniques for handling the data originated with him. Since these original studies, rate measurements with temperature as a continuous function of time have been found useful in examining heterogeneous chemical reactions (16))oxidation (17))annealing of defects in solids (I&),as well as chemical transformation in minerals (19). TABLE I Quantitative Flaah Deaorption Measurements from Tungeten

Desorbing species Na

(B)

Evaporation rate (molecules om-' sec-l) na x 1.4

Refecence

x lo-* exp [-81 x 10*/RT]

6

exp [-76 X 10a/RT]

9

co ( P I )

n X 3 X lo1*

Ha (B)

na x 6

exp [-31 x 10a/RT]

14

H

n X 2.2 X 1Ola exp [-67 X 10a/RT]

14

x

The fact that temperature is not maintained constant must not be construed as a disadvantage. Isothermal conditions are difficult to achieve experimentally. More than that, when several processes with different rate constants occur simultaneously, it is difficult to resolve the component steps by isothermal measurements. At temperatures a t which all the processes take place the rate behavior will be confused by the overlapping of individual reactions. I n isothermal measurements these can be resolved only by separate determinations at temperatures at which only one reaction is important-either so low that processes with a high activation energy are inhibited, or else so high that reactions with a low activation energy have gone to completion. By continuously changing the temperature, however, a single experiment can span the appropriate temperature regions. The technique of varying the temperature is therefore particularly well adapted for examining complex, multiple step processes. The adaptation of these methods t o simple surface processes (20) occurred as soon as the use of filamentary adsorption samples became widespread. The temperature of these samples can be readily adjusted by resistive heating and the ionization gauge makes it possible to analyze the desorption process in detail.

MODERN METHODS I N SURFACE KINETICS

273

When at the end of the adsorption interval the sample surface is heated,* the pressure change in the cell is dictated by Eq. (5). That is, as the temperature is raised, the pressure changes in the competition between gas evolution from the surface F , and removal by pumping through the port and gauge NS,. To this must be added the usually smaller contributions of gas inflow from the reservoir FA, balanced by readsorption on the sample NS,. The way in which the gas density changes as the sample is heated therefore depends upon the course of the evolution, the heating curve for the sample, the pumping speed out of the cell, and the rate law for adsorption. In order to establish a basis for deducing quantitative rate parameters from the experimental data, the form of the pressure vs time curve will be analyzed in the next few paragraphs, on the assumption that readsorption is negligible. 1. Qwclitative Analysis

a. Reaction Order and Heating Schedule. The order of the desorption can be readily determined from the pressure-time curve, which differs markedly for gas evolution obeying first- and second-order rate laws. This is most conveniently demonstrated when the rate of loss of gas from the cell is negligible, Provided the pumping speed is so small that throughout the time interval during which P, has a significant value the condition 1 >> SBt / V is satisfied, Eq. (4) simplifies to

tl

If we can represent the rate of evolution from a surface of area A by an Arrhenius equation

- -an _ -- F~ dt A

,

= 12% vz exp

x=l,2;

(8)

then

AN- - 1 - e x ~ [ - X I AN,

, x=

1 ;

(9)

* Sections 1 and 2 are primarily based on material in ref. 21.

274

GERT EHRLICH

For a hyperbolic heating curve ( 6 )in which 1/T = a $- bt

X

=

=

[

Rvz exp - b ED

-

~~

221

(exp

[ - E+]bt

- 1)

C {exp [Bt] - 13

For a linear heating curve (ZZ), T

=a

+ bt

where y = E , / R T , and Ei (y) is the exponential integral. For desorptions starting a t low temperatures, for which the condition E,/RT > 1 is always satisfied, this expression reduces to

Evolution curves for first- and second-order desorption with the same activation energy and heating curve ( 1 / T linear in time) are shown in Fig. 1 1 . They differ qualitatively in shape and in their dependence upon the initial concentration of adsorbed gas, n (tl). The second order curve is shifted to lower times (that is, to lower temperatures) as the initial concentration is increased, whereas the first order desorption is, of course, independent of initial concentration. The second order curve also has a characteristic s shape not found from first-order evolution. The activation energy for desorption may itself be a function of the concentration; this then removes the independence of concentration which holds for a simple first-order desorption. However, only the first-order evaporation terminates at the same time regardless of the initial conditions and of the detailed shape of the heating curve. A distinction between first- and second-order reactions can therefore always be accomplished by this criterion, even when the desorption energy is concentration dependent. That the shape of the evolution curve is also sensitively dependent upon the heating curve is apparent from Fig. 1 2 , in which the same desorption is followed for different temperature schedules. However, despite these variations, for a fixed heating curve the appearance of first and second order curves is distinct enough to permit identification. b. Dependence on Heating Rate and Pumping. In any real system the

MODERN METHODS IN SURFACE KINETICS

275

gas density in the cell, in the absence of readsorption on the sample, is dictated by the competition between evolution and loss of gas by pumping of one sort or another. The effects of pumping and heating rate are most easily apprehended for first order desorptions, following a heating curve such that 1/T remains linear in time.* As appears from the curves calculated in Fig. 13 increasing the pumping speed, keeping the rate of temperature rise constant, causes a considerable reduction in the pressure peak.? For a system in which significant pumping occurs, not only through the exit port but also in the cell itself, the pumping speed may be difficult t o determine quantitatively. Accurate analysis of the desorption curve under such conditions is not feasible. This loss due to pumping can be countered by increasing the rate of temperature rise of the sample as shown in Fig. 14. However, the higher heating rate not only causes an increase in the peak height but also shifts the gas evolution to higher temperatures. Experimentally, such increased heating rates can be realized by diminishing the sample diameter or by increasing the potential across the sample. Mechanical instability, as well as the danger of overheating this specimen, limit the range of accessible heating rates, however. When the pumping speed S E / V of the system is high, the pressure rise becomes a small difference between the rate of evolution and the total flow out of the cell. Under these conditions, the shape of the

* Provided that the rate of adsorption N S , as well as the flow from the reservoir, FA, are negligible during the time interval over which the filament temperature is raised, it follows from Eq. (7)that the density N as a function of time is given by

tl

For a first order desorption under a heating schedule giving a linear variation of 1/T with time, this assumes the form

where X , B , and C are defined in Eq. (10). The integral can be approximated by the tabulated values of the incomplete function, rx ( p 1 ) = Jf X p exp [-XId X or for small values of p = S/, V B , through the expansion XP N 1 plnX.1.

+

+

t A pumping constant S,/V = 4.8 sec-l, such as controls one of the curves shown there, corresponds in a one liter cell to complete loss of all molecules striking a n area of 0.4 cme.

OERT EHRLXCE

276

L

o

90-

o

.

l

l

1

I

I

I

I

I

I

I

a

mmm-

& m-

1

I, YlLLlSECOllW

Fro. 11. Dependence of evolution curve on rate law for desorption. (a) First- vereue second-order desorption; E D = 80 kcal/mole, v1 = 3.64 x 10" gec-1, v s = 8 x molecules-' sec-l em*. For variable heat, E D = E i - v n ; EL = 80 keal/mole, 7 = 0.3 kcal/mole per 10'3 molecules/cma. (b) First-order desorption with concentration dependent desorption energy. E L = 110 kcel/mole, 7 = 0.3 kcal/mole per 101' molecules/ cm*, v1 = 3.64 x l O l a sec-l. Heating schedule: 1/T = a bt; a = 9.96 X lo-' (OK)-', - b = 1.192 X lo-* ( O K SIX)-'.

+

MODERN METHODS IN SURFACE KINETICS

1

HEATING CURVES: HYPERBOLICf a t bt

277

/4

1,'K mWSTbNT WLTAGE

yx)

350

400

450

Mo

%a

1

I, MILLISECONDS

FIQ. 12. Effect of heating schedule on evolution curve. First- and second-order desorption (with rate pararileters as in Fig. 11) are shown in (a)for heating curves indicated in (b). Second-order reaction maintains a shape.

pressure vs time curve resembles that of the evolution rate, since in the limit dNldt = 0, N = F,/S,. The desorption rate itself goes through a maximum. As shown by Urbach (15), the requirement d/dt ( - d n / d t ) = 0, relates T,, the temperature of the peak rate to the desorption

278

GERT EHRLICH

energy E , through

Here x again indicates the order of the desorption reaction. For a simple first-order evolution this maximum is obviously independent of the starting concentration regardless of the form of the heating curve. When the desorption energy is a function of concentration, or desorption I

1

I

I

I

I

1,YILLISECONOS

FIG. 13. Effect of exit speeds, on first-order desorption. Rate parameters and heating curve as in Fig. 11. Peak height diminishes with increasing S,, and shifts to lower temperatures.

occurs as a second-order reaction, then at higher starting concentrations the maximum is shifted to lower temperatures. The behavior of the peak in the desorption rate for different initial coverages can therefore serve to distinguish the order of the reaction.*

* This method has been favored by Redhead (23).

MODERN METHODS I N SURFACE KINETICS

279

At the rate maximum, that is at time t,, (d2n/dt2) = 0 and it therefore follows from differentiating Eq. ( 5 ) that

v dZN(t,)

-

-dN(t,) _

s,

at

at=

*

At the maximum of the pressure vs time curve dNldt = 0. These temperatures can be brought close to one another by a large exit speed S, and a slow heating rate. For qualitative purposes, such as the determination of the reaction order, the pressure vs time curve at high pumping speeds can be substituted for a plot of the temporal evolution of the rate.

.90/

1.192

,596

.I99 I

200 1200

I

I

440

I440

I

280

240 520

1680

'0°

I

320 1920

1 , MILLISECONDS

I

I

-bid

-

I

360 -1.192 680 -596 2160 .I99

FIG. 14. Variation of evolution curve with rate of heating at a fixed exit speed (S,/V = (OK)-'. Peak 4.8 sec-l). Rate parameters as in Fig. 11. 1/T = a bt, a = 9.95 shifts to higher temperatures at higher heating rates.

+

+

2. Quantitative Analysis of Desorption

The problem now is the derivation of quantitative rate data. Given are the experimentally determined pressure time trace [corrected for any loss of gas by pumping in accord with Eq. (7)], together with a temperature time trace for the sample. The pressure rise in the cell may be brought about under two limiting conditions ( 6 , 2 0 ) .If the temperature change of the sample is slow com-

280

OERT EHRLICH

pared with the rates of evolution and adsorption processes, the steady state surface concentration can be maintained at any temperature. That is, the net rate of change of the gas density, d N / d t , is negligibly small a t any given temperature, and the change in density with time which is experimentally observed arises from a displacement of the steady state as the temperature of the sample is raised." I n the opposite extreme, to which we have so far limited ourselves, the temperature of the sample is raised rapidly compared to the rate processes determining the concentration of adsorbed material. The system a t any instant is far removed from the steady state, and the rate of the evolution reaction determines the surface concentration. a. Evolution Kinetics Rate Determining. The instantaneous slope of the curve of surface concentration n vs time t obtained from the experimental pressure-time trace by correcting for the net loss of gas is just

-dn/dt = nz vz exp [ -E,/RT]

= nx

kg) ( T ) .

(13)

For the simplest possible systems, in which ED and v are both independent of temperature and n, the slope of a single evolution curve at different times (temperatures) yields the rate law of desorption. The correct concentration dependence (x = 1 or 2) may be ascertained by plotting In n-zdnldt vs 11RT. A linear relation with slope - E, will be obtained for the correct value of the order of the reaction. The preexponential can then be determined by substituting EBandnzin Eq. (13). Even for such a simple system it must be established that the slope of the experimental curve is properly represented by Eq. (13). At a fixed temperature and surface concentration, the rate of desorption must be independent of the way in which the sample is heated. This independence must be checked experimentally by determining d?z/dt at different heating rates. The experimental evolution curve arises from the competition with readsorption NS, and gas flow ( N - NJS,. The more rapid the temperature change, the further away from the steady state the system is a t any instant, thereby minimizing the contribution of adsorption. Also, the higher the heating rate the smaller the time interval during which the gas flow must be considered. Only when independence of heating rate has been established can we conclude that the evolution rate alone is being measured. This method of analysis is valid for any arbitrary heating curve. It has the added advantage that it can be applied to systems in which the heat

* This is analogous to the situation considered in the previous section. However, there the steady state was dictated by evolution from the sample and removal by pumping out the port. Here the gas is re-adsorbedon the sample surface.

MODERN METHODS I N SURFACE KINETICS

281

of adsorption is itself a function of the surface concentration. When this situation prevails the semilogarithmic plot of dnln' against 1/T will not lead to a straight line for reasonable values of x. I n its stead, a whole family of desorption traces must be obtained a t different initial concentrations, arid possibly also a t different heating rates. Each curve of this family then yields a value of the slope dnldt, for a fixed surface population n, but a t different temperature. A plot of In (dnldt)against 1IRT gives, as usual, the heat of a desorption a t a fixed surface concentration without, however, any indication of the order of the reaction. This can be deduced from a given desorption trace, using the previously attained heat value, by fitting the slopes a t different temperatures with the appropriate concentration term nz. The heat of desorption can be deduced more simply from the maximum in the rate of evolution. The maximum rate as well as the surface concentration a t the maximum, both of which enter into Eq. (12), can be obtained directly from the curve of evolution rate vs time or temperature. The graphical integration required for n can be avoided in simple first-order reactions by adopting a procedure customary in the study of thermo-luminescence. On substituting for d In n/dt from Eq. (9), Eq. (12) rearranges to

and the desorption energy can therefore be written as

As already pointed out by Randall and Wilkins (24),the term in ED/RTk on the right-hand side amounts to a small correction for the rate parameters of interest; the desorption energy can therefore be adequately related to the temperature T , of the rate maximum by

For semiquantitative studies, it is permissible to assume a value of v( 5 x 1 0 1 2 sec-I), and only T , is required for evaluating ED. This temperature can, moreover, be deduced from the peak pressure at low heating rates and high pumping speeds, without knowledge of the rate and does not demand any special heating curve. I n practice, values of the preexponential may vary by two to three orders of magnitude as a surface becomes filled (I),introducing errors of more than 20% in N

282

OERT EHRLICH

the desorption energy. Inasmuch. as any change in the rate law is of interest, this procedure for evaluating the desorption curve must be abandoned in a quantitative analysis. The need t o assume a value for the preexponential term can, however, be eliminated by a further refinement, due to Booth (25). If the temperature of the maximum rate is determined a t two values of the heating rate b, then from Eq. (14) it follows that

For a desorption process with E D = 40 kcal/mole, a 100-fold change in the heating rate (from 1 to 100"K/sec) raises T, from 673°K t o 783°K. This is therefore an entirely feasible method for checking desorption energies. When the heating schedule conforms to one of the simple forms, such as l / T or T linear in time, the rate parameters can be derived from the variation of the surface concentration with temperature in yet a different way ( 6 ) .For a temperature-time curve in which l / T is linear in time, the amount of gas evolved is given aa a function of time by Eq. (10). The instantaneous surface concentration can therefore be written as

Here

and

The order of the desorption reaction, as well as its activation energy, can be determined from a logarithmic plot of either (n(tl)- %)/enn(tJ or (l/~)ln[n(t,)/n]against l / R T . A straight line with slope - ELI will be obtained for the plot of the correct concentration dependence. The correction factor E itself depends upon ED,but for a temperature range from 1000-2000°K and a heat on the order of 50 kcal/mole it is small, and can be conveniently accounted for by successive approximations.

MODERN METHODS IN SURFACE KINETICS

283

This was the method employed in the first quantitative study of flash desorption. When the sample temperature is raised proportionally to time t , the desorption follows Eq. ( 1 1). This can generally be approximated by

Straight lines for 1st and 2nd order reactions are obtained on a semilogarithmic plot of (l/TZ)In [n(t,)/n],or (1/T2) [n(tl)- n ] / [ n n(t l )] against l / R T with slope equal to the desorption energy. Neglect of the last term in Eq. (18) introduces an error of < 1 % for chemisorbed systems; this can again be corrected by successive approximations for very small values of the adsorption energy, where this term may become important. For either heating curve, it must be established that the concentration a t a fixed temperature is linear in the time required to term JdnT/nZ reach that temperature. This test is equivalent to proving that the slope of the evolution curve is independent of heating rate. It should be noted that of all these methods and their variations, the direct determination of the desorption rate as a function of concentration and temperature is the most powerful and straightforward. This method is not limited to a simple heating curve or to a simple form of the rate expression. It can, if necessary, encompass concentration dependence in both the heat and the preexponential. The other techniques, such as the use of integrated rate expression, or of the maximum evolution rate are, however, very useful for rapid evaluation of the desorption energy in simple systems. b. Temperature Displacement Determining. For an evolution curve determined by the rate of desorption, the slope of experimental pressuretime plot must be independent of heating rate. On the other hand, in order to prove that evolution occurs through a displacement of this steady state surface concentration, we must establish just the opposite, namely, that dn/dt = (dn/dT) (dT/dt). The slope of the evolution curve at any given temperature must be a linear function of the slope of the temperature-time curve. If this condition is obeyed, then the surface concentration is related to the gas density through

N

= AnZ(vZ/S,) exp

( - E,/ RT ).

Just as in the analysis of the desorption rate, the parameters defining the steady state can in general be derived from a family of evolution curves. A semilogarithmic plot of the gas density N a t a constant surface

284

OERT EHRLICH

concentration n against 11RT yields

The rate constant for adsorption on metals is generally high and endowed with a small activation energy; the last term can, under most circumstances, be safely neglected, and this procedure therefore results in an approximation to the energy of desorption. When desorption occurs as a simple first- or second-order reaction, the order can be deduced by plotting ln(N/n”) against 1IRT. A straight line is obtained if the exponent x correctly corresponds to the order of the desorption, and the slope of this curve then yields again the energy of the desorption reaction. It should be noted that by changing the rate of the temperature displacement it is possible to move from one desorption regime to another. Even starting with a steady state surface concentration at the adsorption temperature, the rate of desorption can subsequently be made to dominate, merely by imposing upon the sample a square temperature pulse with a rise time short compared t o the half life of the desorption at the higher temperature. Conversely, a temperature-time curve with a monotonically diminishing slope may result in a change from a rate limiting to a temperature displacement limiting evolution during the course of a given pressure burst. 3. A Practical Example-Evolution of N , from W . The application of these general principles to practical problems is best illustrated by an analysis for an actual system-the desorption of nitrogen from tungsten. Of the three chemisorbed states a, fl, and y the most strongly bound, p, will be considered first. a. High ~ e ~ p e r u Desorption. ~~re That the f i state is not evolved in a simple first-order reaction follows immediately from the behavior of the desorption curve with increasing concentration. It is clear from Fig. 15 that the evolution begins a t lower times, and therefore lower temperatures, as the surface concentration increases. A concentration shift means either a first-order desorption, with the desorption energy diminishing a t higher concentrations, or else a second-order reaction. This point can be settled by inspection. A first-order desorption, even when complicated, will terminate a t the same time regardless of the initial concentration, provided the heating curve remains unchanged. I n Fig. 15 the entire evolution is shifted toward the left, toward lower temperatures, and this signifies a desorption of the second order with respect to the adsorbed material.

MODERN METHODS IN SURFACE KINETICS

285

A more rigorous, quantitative analysis of the pressure time curve for /3 nitrogen proceeds as follows (6). The nature of the processes involved is revealed by the family of curves in Fig. 16 taken at different heating rates. It is apparent that fixing the temperature does not fix the population of adsorbed entities or the number already evolved. The smaller the rate of heating (that is, the smaller the heating current) the greater the number of adsorbed entities evolved at lower temperatures. Clearly the steady state approximation does not apply to these systems: The pressure rise is dictated by the rate of desorption.

FIG. 16. Adsorption of N, on W at T = 243"K, showing growth of a and /3 states. Time scale: 1 unit = 100 msec. (a) Adsorption interval dt = 4 min, R , =6.5 X 10l2 molecules/cm', no = 39 X 10'2; (b) At = 12, n, = 16 X 10l2, np = 114 X 10l2; ( c ) At = 21, R , = 14 x lo", ~p = 204 X 1012; (d) dt = 26, R , = 13 X lo", n,g = 233 X 10la. With increasing concentration, evolution of /3 shifts to lower t.

By resorting to the bridge circuit shown in Fig. 27 the sample temperature can be raised so that over the range of p desorption the reciprocal temperature is proportional to the heating time. The desorption can then be simply analyzed according to Eq. ( 1 7 ) . In Fig. 17 the data from a typical evolution curve are plotted according to both a first- and a second-order rate law. From the fact that only the latter gives a satisfactory straight line we can tentatively conclude that the desorption of /3 nitrogen is a second order reaction. That this tentative conclusion is not obviated by competition from adsorption is indicated by linearity

OERT EHRLICH

286

of the plot of (n ( t l ) - %)/en a t fixed temperatures against the time of reaching these temperatures, shown in Fig. 18. If readsorption were occurring, deviations from a straight line would appear a t low heating rates (at low values of - I / b ) . 1

1

I

1

I

1

1

-

80 -

-

m\ v)

W

n(tfl=85x i0l2 MOLECULES/CM2

6c-

-

+ 40c

Y

30

-

-

TEMPERATURE PK)

FIG. 16. Amount of N, desorbed up to a fixed temperature T a t different heating rates. Heating rate ( - b ) = [ - 0.66 x lo-' + 2.06 x Iz](OK sec)-l. Fixing T alone does not determine surface concentration.

There still remains the possibility that the heat of desorption depends upon the concentration of adsorbed nitrogen, and that the simple analysis used here might be inapplicable. I n a plot of the type shown in Fig. 17 the concentration on the sample filament is continually decreasing as high temperatures are reached. If E D increased with concentration, the magnitude of the integral

would diminish below that indicated on the graph a t low concentrations (at low values of l/T). A straight line would not prevail for either of

MODERN METHODS I N SURFACE KINETICS

287

the simple rate laws, inasmuch aa this criterion was established on the assumption that E, f E , (n).However, if the desorption is plotted for increasing initial concentrations of nitrogen we find, as shown in Fig. 19, that straight lines are obtained for each value of n ( t l ) ;these are parallel to each other if the data are represented as a second-order desorption. A given temperature now corresponds to quite a different value of n on different curves, and for a concentration dependent heat of desorption parallel lines could not be expected.

FIG. 17. Analysis of pressure-temperature curves for 4, nitrogen, indicating secondorder desorption kinetics.

The criterion for the insignificance of the adsorption reaction was also derived under the assumption of a constant E D .However, the curves in Fig. 20, obtained at an intermediate coverage but over a range of heating rates favoring the reverse reaction, still have the same slopes as their counterparts at higher ( - b ) and show as good a fit to a second-order

288

QERT EHRLICN

desorption. Under the conditions of these experiments the evolution of fl nitrogen must therefore be dictated by the desorption reaction. This desorption is satisfactorily represented by a second-order rate law, with a constant heat of desorption over the concentration range 7 < n < 250 x 10l2 molecules/cm2.

0.a

0.7

0.6

0.5 C

C

0.3

0.2

0.I

:

FIG.I8. Dependence of nitrogen desorption kinetics on heating rate. Linearity indicates Eq. (17) is obeyed, and that readsorption is negligible.

The variation of the slope of the desorption plot with concentration of p nitrogen is shown in Fig. 21. The heat of desorption is seen to be independent of cover within the limit of error of the measurements f 4 kcal/mole, having a value of 81 kcal/mole up to np= 250 x 1012

MODERN METHODS I N SURFACE KINETICS

289

molecules/cma. The rate law for desorption can be evaluated from Eq. ( 1 7 ) by substituting the experimentally determined heat of desorption ED and the heating rate ( - b ) , and is found to be

[

-1 dn = 1.4 x 10-2 exp n2 dt

fi-x1O9] cmZ/molecules sec. RT

The quantitative flash desorption measurements therefore establish that at the temperature of evolution ,8 nitrogen is present on the surfaces as atoms. Desorption then involves the collision of two energetic adatoms with one another and their evaporation as a molecule.

Fro. 19. Desorption plots and desorption energy E D a t increasing initial coverages of /3 nitrogen on W. Both appear constant with concentration.

b. Low Temperature States. Quantitative desorption measurements are not yet available for either the 01 or the y state. Qualitative observations (8) reveal considerable complexity in the behavior of these forms of nitrogen.

290

OERT EHRLICH

From the maximum concentration achieved a t 298°K we can estimate a binding energy of -20 kcal/mole for nitrogen held on 01 sites. Just as for the /3 state, the desorption is shifted towards lower temperatures as the length of the adsorption interval is increased. This is evident from

-

-

-

0.1 -

- EFFECT

-

OF HEATING RATE

6.00

I

7.00

I

6 .00

1 0 ' T

I

7.00

(.K)-'

FIG. 20. Desorption energy of nitrogen (at a fixed surface concentration)for different rates of heating. E , independent of heating curve.

flashes shown in Fig. 22. The concentration of 01 itself goes through a maximum. At the same initial concentration of a, evolution occurs a t lower temperatures for the higher concentration of /3. This suggests

MODERN METHODS IN SURFACE KINETICS

291

that the desorption energy for the CY state is diminished by the presence of large numbers of nitrogen atoms in the fi state. For the y state, evolution terminates a t the same time even for a 20fold increase in concentration. This is therefore a first-order reaction. However, it appears from Fig. 23 that the onset of desorption shifts

FIG.21. Experimental desorption energy of /3 nitrogen as a function of surface concentration.

t o lower temperatures, unlike a simple first-order desorption with constant heat. This shift could result from a drop in the desorption energy a t higher fi concentrations. However, the y peak itself is seen in Fig. 23 t o be a composite. A substate with a lower binding energy forms a t higher concentrations, complicating a quantitative analysis. The desorption traces do a t least show that part of the y state evolves by simple first-order kinetics. This part, amounting to 1/4 of the total, must therefore be present on the surface as molecules, or in the form of isolated atom pairs. The temperature range of the desorption and the absence of a significant pressure dependence for the rate of formation a t 115°K indicate a binding energy on the order of 9 kcal/mole for this molecularly bound state. Even these qualitative observations therefore help t o define the atomic state of the adsorbed material.

-

4. Experimental Implementation

The details of executing successful flash desorption experiments must be adapted to suit the needs of each individual problem under study. There are, however, a few general points of technique that deserve brief mention. a. Resolution of Multiple Peaks. The analysis of desorption occurring in distinct stages follows along the lines already indicated, except for such differences as arise from overlapping of the evolution curves or from conversion of different entities to one another during desorption. Consider, for example, two entities n , and n2evolving independently

292

GERT EHRLICH

FIG. 22. Evolution of a nitrogen from W at increasing concentrations of the j3 state, after adsorption a t T = 200°K. Temperature range for desorption: 200 < T < 600°K. T and p scales uniform throughout. (a) na = 53 x 10l2 molecules/cm*; (b) 117 x 10l2; (c) 146 X 10l2; (d) 220 x 10".

MODERN METHODS I N SURFACE KINETICS

293

GERT EHRLICH

294

of one another with kinetics -

dn,/dt = ny v y ) exp [ - Eg’/RT]

- dn,/dt

= n: v r ) exp [

,

- Eg)/RT] ,

E p >Ey.

For desorption with different heats, the extent to which simultaneous evolution of both species occurs will be greater the higher the rate of temperature change. Thus under a step function temperature pulse, n,

:115

a 10”

298 a 10” noltcultr/cm2

t

N

tFIG 23. Desorption of y nitrogen from W at reduced heating rate. Time scale: 1 unit = 0.2 sec. Adsorption intervals at T 115°K. Left: At = 20 min; Right: dt = 35. Arrow indicates substate evolving with first-orderkinetics. N

both the material with a low and a high desorption energy evolves simultaneously, despite the fact that completely independent evaporation, without apparent overlapping, is possible. This can be accomplished by allowing the evaporation of the weakly bound component to proceed a t a low temperature, a t which the rate for the higher is negligible, and by raising the temperature once this has come close to completion. Coalescence of the evolution curves a t higher heating rates is quite a general effect that holds for linear as well as hyperbolic temperature curves. A specific example is given in Fig. 24, for two independent first order desorptions a t heating rates differing by a factor of 100. The overlapping of the evolution curve also changes the qualitative features of the pressure time trace, which are significant for the recognition of the existence of separate adsorbed entities. I n general, the increase in heating rate can be compensated by increase in the sweep rate of the detector. Here, however, noise may introduce a limit. More than that, the evolution curves are compressed a t higher heating rates, making the observation of the inflection point between two peaks more difficult. For qualitative detection of this inflection it is therefore

296

MODERN METHODS I N SURFACE KINETICS

advantageous to operate a t lower heating rates and at higher pumping speeds. Under these conditions an estimate of the desorption energy is most readily accomplished by just noting the temperature corresponding to the pressure peak. I

I

I

1

,b

(OK

s(C

1

I

1

I

i':1.192x 16'

1

1.192 x

16'

RESOLUTION OF DUAL PEAK

.20I

260

I

280

I

300

I

1

I

1

320 340 360 1, MILLISECONDS

380

FIG. 24. Resolution of dual peaks: Effect of heating Two components, with the same initial concentration, kinetics with: v(:) = 3.54 x sec-l, E 2 ) = 80 sec-l, E;) = 97.5 kcal/mole. Heating curve: 1/T = a

+

I

1

400

420

rate ( - b ) and exit speed (BB). evolve according to first-order kcal/mole; d,"' = 4.17 X l O l 3 bt, a = 9.95 X lo-' (OK)-'.

I n a real system the evolution of different entities from the surface may not occur independently, and conversion of one to the other during the evaporation may occur with an activation energy E , smaller than that required for removal to the gas phase. For a system in which species 1 can convert to 2 a t a rate

296

GERT EHRLICH

the ratio of the rate of conversion v, to that of desorption vD is just

d ~ ) / v= ~ )n:-Z (vJv(2)) exp [ - ( E , - E $ ) ) / R T ] . This is greater the lower the temperature, and for the gas evolution to faithfully describe the initial distribution on the surface a high heating rate, which displaces the evolution to higher temperatures, is therefore desirable. This may, however, bring about overlapping of the individual desorption peaks.

FIG.26. Flash desorption cell and ultrahigh vacuum system. C-desorption cell, with sample F, mounted on glass covered tungsten leads L, and gauge G. Volume of cell is defined by ground glass port D and E. A, H-ion gauges for leak checking and pressure monitoring; B-Granville-Phillips valve; I-liquid nitrogen trap; S---selective getter bulb.

b. Detection of Desorbed Gas. The Bayard-Alpert gauge, coupled with a high-speed iod current detector, provides an instantaneous record of gas density. However, in the actual flash filament cell the density is not measured in the immediate vicinity of the sample. Because of the size of the gauge and the desire to minimize electrical interference, the two should be mounted in separate containers connected by large bore tubing, as indicated in Fig. 25. The possibility of density gradients within the cell during desorption must therefore be examined. When the

MODERN METHODS IN SURFACE KINETICS

297

sample filament is cooled or heated during the adsorption cycle and the local gas density undergoes sudden changes, a steady state distribution is established only after a time interval long by comparison with E, the average transit time of a molecule from sample to gauge. If the geometrical arrangement of the cell is idealized as a long tube of diameter d , in which a gas source with a cosine distribution (the sample filament) is at a distance L from the density detector (a gauge) then (26)

Here U is the average gas velocity and T the mean lifetime on the surface of molecules colliding with the tube walls. I n a properly designed cell (typical dimensions L = 13 cm, d = 4 cm), the transit time for nitrogen is only of the order of a millisecond. For the evolution of chemisorbed gas which occurs over an interval of a t least 150 msec and for which T is small this delay is negligible. I n the desorption of a low temperature state which is complete in 20-40 msec, or for carbon monoxide which has an appreciable life-time while adsorbed on the glass walls, the gauge may not give an instantaneous indication of the processes occurring in the vicinity of the sample. This difficulty is particularly pronounced in studying physical adsorption. With xenon, for example, gas from the desorption peak is adsorbed on the leads and glass parts, which are of necessity cooled when the sample is brought to the temperature (T< 115°K) a t which adsorption occurs. To compensate for this, the desorption must be carried out very rapidly, over an interval of -10 msec or less. Since the delay time is -1 msec even discounting adsorption, quantitative determination of the desorption kinetics is difficult. I n desorption, the cell does not behave as an isothermal system. Gas evolved from a hot filament may be as much as 1000" above the temperature of the rest of the system. However, even if temperature equilibrium is not established by collisions with the cell walls before the gas enters the gauge, the change in ion current is still proportional to the increase in the gas density, and a measure of the number of evolved molecules can be obtained without correcting for their temperature. The kinetic energy of the particles evolved from the sample does affect the processes (such as readsorption on the sample, as well as escape from the cell itself) which dictate the gas density in the cell. Under the usual experimental conditions, even these effects are small. Assuming a linear rise in gas density with time after the start of the desorption, the average number of collisions that a molecule in the cell has made with the walls at time t is wA,/8tV. For estimating a lower limit, w can be approximated by the ordinary, average speed a t the cell temperature.

298

GERT EHRLICH

The ratio A J V is of the order of unity, and thus even for desorption intervals as short as 5 msec the molecules make on the order of 30 collisions with the walls and should achieve the wall temperature. G. Control and Measurement of Temperature. Control of the sample temperature during the course of evolution is a t the heart of flash desorption measurements. For slow heating rates (on the order of 2000"/min) the temperature can be varied by adjusting the voltage across the sample with a mechanical drive. For rates an order of magnitude higher this adjustment is most conveniently accomplished electronically. The exact form of the temperature vs time curve is not vital to success, however, since adequate analysis of the data can be carried out for an arbitrary heating curve. Measurement of the instantaneous temperature is readily accomplished for filamentary samples by following the resistance. The filament thus serves both as a sample and a resistance monitor. The first qualitative determinations of the desorption spectrum were carried out using a bridge circuit, shown in Fig. 26, designed by J . P. Molnar [see Hagstrum ION GAUGE

FLASH FILAMENT

R?

FIQ. 26. Desorption spectrometer using dc bridge (27). R, = 50R; R, = 5OOOi2; R , = 75a; R,. << R , (for flash filament hot or cold). dc power supply: 0-112 volts, 4 amp.

(27)l.The dc potential drop developed across the sample by the heating current gives a signal proportional to the filament resistance, provided t h a t the current through the filament remains constant as it warms up. This circuit is suitable only for short, large diameter samples. For fine wires ( 6 mil or less with W) or for high heating rates a dc power supply with output voltage considerably in excess of 125 volts would be required. This limitation can be overcome by separating the heating and measuring function as in the desorption spectrometer, shown in Fig. 27, which was suggested by Rodbell and employed in the 1st quantitative flash desorption studies ( 6 ) . Here the sample is heated by direct current,

MODERN METHODS I N SURFACE KINETICS

289

but the resistance is determined on a 10 kc impedance bridge. The impedance Z of the arms is adjusted so that 2, >> Z,, and Z, >> 2,. The unbalance of the bridge when the filament is flashed results in a signal proportional to filament resistance. This is amplified by a vacuum tube voltmeter and a pulse transformer, detected, passed through a

125 V DC

FIG. 27. Desorption spectrometerusing 10 kc bridge (6). R,,R,,R,are part of standard impedance bridge; C , = 30 p j , C, =.0-100 ppf.

filter to remove 60 cycle noise, and displayed on one axis of a n oscilloscope. To calibrate the scope deflections, R,, a compensated decade resistor, may be substituted for the sample filament, which is isolated from the direct current supply by the blocking capacitor C,. The scope deflection is found tovary linearly with calibrationresistance to within 1%. The change in resistance of a 0.0245 cm diameter tungsten filament during the course of the flash, as obtained with this bridge circuit, is shown in Fig. 28. The transient is caused by the initial current surge, but is of negligible duration compared to the total time necessary t o raise the temperature to 2200'K. This surge does not affect measurements above 1000°K. It may, however, make it difficult to examine phenomena that occur immediately a t the start of the heating. This difficulty can be avoided by abandoning bridge techniques entirely, and by simultaneously measuring the potential drop across the sample and across standard resistors in series with it. If these emfs are recorded on a scope, even rapid temperature changes occurring on a

300

GERT EHRLICH

millisecond scale can be determined as indicated in Fig. 29. This method lends itself to essentially any form of heating curve and has the advantage of simplicity as well as of flexibility. d . Temperature Effects. I n converting the filament resistance into a measure of the sample temperature, it must be borne in mind that the

FIG. 28. Change in filament resistance during desorption with ac spectrometer. Sample filament: 10 mil W. Steady state current I = 6 amp. Resistance markers: 0.lQ; time scale: 1 unit = 0.5 sec.

temperature is not uniform over the entire sample. For an electrically heated filament, such as customarily employed in flash desorption, temperature variations over the length of the sample arise (1) from nonuniformities in the cross sectional area or the impurity content, and (2) as a consequence of energy transfer to the surroundings. Elimination of the first source of difficulty presents only trivial problems. Such nonuniformities can be detected by cinematographic examination of the incandescent filament during a temperature flash. Temperature nonuniformities of the second type always exist in an electrically heated wire near the supporting leads, when these are in good contact with thermostated environment. The current-voltage characteristics of very short tungsten filaments show a negative differential resistance-the current in a low impedance circuit, including the filament, diminishes with increasing voltage. This effect arises whenever the fractional increase in energy dissipation on increasing the temperature is smaller than the fractional increase in the resistance of the sample (28). Under these conditions a steady state with a filament a t a uniform temperature is impossible whether the energy losses occur

MODERN METHODS IN SURFACE KINETICS

IOV

301

a

FIQ.29. Pressure-temperaturemeasurementduring rapid desorption (7). (a)Schematic of circuit for simultaneous recording of ion current and of potentials across sample filament and standard resistor. (b) Oscilloscope display for desorption of Xe. Time scale: 1 unit = 1.48msec.Temperaturerange:t = 0,T = 116"K;t = 13.32 msec, T = 254°K. Total Xe concentration: 150 X 10leatoms/cm2.

302

QERT EHRLICH

primarily by conduction through the leads or through the gas surrounding the filament. As a measure of the temperature distribution over a typical sample when brought to a steady state by resistive heating there is plotted in Fig, 30 the resistance between equally spaced points on the wire (labeled 1 through 6 ) for different heating currents. Temperature differences

POTENTUL PROBE

wILwT!AL PROBE

FIQ.30. Variation of temperature over resistively heated filament. Tungsten sample: 15.6 om long, 2.62 x 10V cm diameter; leads at 77'K. Potential for temperature flash: 10 volts. Distance between probes: 1-2, 1.19 cm; 2-3, 2.19 cm; 3-4,2.23 cm; 4-5,3.04 cm; 6-6, 6.95 cm.

over the length of the filament amount to more than 100"K, even a t the lowest steady state currents, corresponding to an average filament temperature T , , = 200°K. I n passing through the range of the Busch anomaly (which appears only as constant current independent of the applied voltage a t a central filament temperature T,, 200"K), the resistance gradient, and therefore also the temperature gradient, becomes much steeper. The resistance distribution during a temperature flash for over-all filament resistances corresponding to those in the steady state are also shown in Fig. 30. Throughout the range covered (220 < T , , < 1480°K) the temperature is significantly more uniform during the flash than in the steady state; for temperatures a t the filament center of T,,< 370"K, the nonuniformities are negligible, never exceeding 40" over the entire N

MODERN METHODS I N SURFACE KINETICS

303

length. For qualitative purposes, such as the isolation of different binding states, such a filament can therefore be considered to be at a uniform temperature. Even for quantitative rate determinations a correction will be required only above room temperature. This is one of the significant advantages of following desorptions under continuously increasing temperatures, as contrasted with measurements obtained for a steady state distribution. The more rapid the rate of heating the smaller the correction for end losses. For current pulses below 1 msec however, appropriate corrections to the resistivity have to be made in order to account for the skin effect in filaments exceeding 0.1 cm in diameter. End losses can also be minimized by operating with long, fine wire samples. Inasmuch as quantitative corrections for temperature nonuniformity are difficult i t is desirable to design the experiment in order to minimize such effects. I n the course of an experimental cycle the sample is first cleaned by resistive heating to 2200°K; it is then allowed to cool to the temperature of the surrounding bath a t which the desorption is studied. At the end of the adsorption interval the sample is rapidly heated in order to count the number of molecules evolved. Once the heating current is interrupted the temperature of the filament only slowly approaches that of the surroundings. As already pointed out, significant measurements can therefore only be achieved at the end of this cooling period, which is normally on the order of 3 min. The cycling of the sample temperature also changes the temperature of the cell walls even in the thermostated enclosure, and this in turn may affect the concentration of adsorbed entities on the walls. During the high temperature cleaning of the sample the inner walls may be depleted of their gas layer which is then restored during the adsorption interval. Conversely, after a desorption flash, gas may again be evolved from the walls surrounding the filament. Effects of this type may under some circumstances be appreciable. For nitrogen at a pressure of 3 x lo-'mm we estimate a surface concentration n of 1 .li X lolo molecules/cm2 and for carbon monoxide a t p = 7 x lo-' mm, n = 1.4 x 1012 molecules/cmz. This leads to lifetimes of the order of 2 x and 2 x 10-2 sec, respectively. With carbon monoxide difficulties due to adsorption on the glass must therefore be expected both during the adsorption and desorption cycle. Qualitative observations for oxygen also indicate lifetimes on glass comparable to those of carbon monoxide. Despite some of these difficulties flash desorption is an exceedingly powerful technique in that it gives a direct, numerical count of adsorbed material. Flash desorption has already yielded considerable qualitative

304

GERT EHRLICH

information on the presence of adsorbed gases in different binding states, and on the nature of the desorption kinetics, as well as quantitative data on the surface populations in various levels and on the absolute kinetics of their decay. The chief advantage of the technique is its experimental simplicity. I n principle, some of the same information could be obtained isothermally, by imposing a square temperature pulse on the sample. I n practice this is difficult to accomplish.

c. ENERGYPARAMETERS I N MACROSCOPICADSORPTIONEXPERIMENTS Flash desorption, as well as other kinetic measurements, are fruitful sources of information on the energetics of surface processes. Indeed, for some systems, especially those in which processes occur a t high temperatures, the traditional techniques such as calorimetry and isotherm determinations are difficult to execute and interpret. I n order to compare the results obtained by flash desorption with kinetic and equilibrium measurements by more standard techniques, a sketch of the interrelations between energy parameters (29) is in order. 1. Calorimetry

a. Reversible and Irreversible Gas Introduction. Calorimetry on gas adsorption has without exception been done under adiabatic conditions. When only van der Waals forces are operative, reversible adsorption of dm moles of gas can be accomplished by letting a frictionless piston diminish the volume of the calorimeter by d V . The heat of adsorption qnrdetermined from the temperature rise of the calorimeter and its contents is then given by

Here uQ - u1 is the difference in the partial molar internal energy of the gas G and the adsorbed layer 1, and - p ( a V / b n ) sis the adiabatic differential heat of compression. Reversible adsorption can only occur at temperatures such that diffusion of the adsorbed gas is fast enough t o establish an equilibrium distribution a t the surface during the time of the measurements. This criterion holds for chemisorption as well, but implies that observations are made at correspondingly higher temperatures. For chemisorption a t room temperature, equilibrium pressures are small and reversible compression by standard techniques is difficult. More than that, to permit clean operating conditions, calorimeters

MODERN METHODS IN SURFACE KINETICS

306

with a cooling time of minutes are generally employed (31, 32) and adsorption must be completed in an interval small compared with this decay time. It is therefore customary to introduce gas irreversibly by opening a valve to a high pressure reservoir ( p 10-1mm). The calorimetric heat determined in this way is just

-

qaz. = ug - u1.

(22)

This differs from the value obtained under adiabatic reversible conditions only through the heat of compression. The latter will be of the order of 2RT, R being the gas constant, and differences between qapand qai for chemisorption are therefore well within the limit of the expected experimental error. b. Equilibrium and Nonequilibrium Measurements. In calorimetric experiments, several related processes with rather different relaxation times are involved in the approach to an equilibrium surface layer. An atom or molecule is bound by the surface if, on colliding with the surface from the gas phase, the atom gives up its translational energy. Such a chemisorbing atom achieves its final equilibrium state only after a series of additional energy transfers to the lattice. The efficiency of this transfer is as yet not quantitatively established. Model calculations indicate that 98% of the heat of adsorption is lost from the adatomsurface bond in only a couple of collisions (33). This process should therefore reach equilibrium during the time of the calorimetric determination. A more severe requirement is posed by the demand that the sites at which the adsorbed atoms are finally bound be equilibrium sites. On a heterogeneous surface, for example, an atom must be able to sample the different sites available in order to discriminate against those on which the chemical potential is high. During the time T of the measurement, the root mean square distance covered by an adsorbed atom is X = ?!(WT)''',

where t i s the jump distance and w the jump frequency. The distance x must exceed a characteristic separation on the surface, such as the average separation of the different types of sites, for an equilibrium distribution to be established. Equivalent criteria have to be applied when the attainment of equilibrium is limited by other processes such as surface rearrangement or slow transformation of one binding state into another. When the temperature of adsorption is so low that processes like diffusion or rearrangement cannot establish equilibrium during the

306

CERT EHRLICH

time r of the experiment, then the heat value obtained experimentally is a number average,

dm, now denotes the number of adsorbed species held a t surface features

a t which A u may differ, either because of the adsorbed environment, or because of intrinsic properties of the surface itself. The number dmi is dictated by the kinetics of the adsorption step; the nature of the surface distribution must be properly specified to permit any detailed interpretation. 2. Adsorption Isotherms

a. Isosteric Heats. The temperature dependence of the equilibrium pressure (34)necessary to maintain r,the ratio of adsorbed gas to solid, yields the isosteric heat qr. according to

where h, and h, represent the partial molar enthalpy of gas and adsorbed material. Underlying Eq. (23) are the assumptions that the partial molar volume of the gas far exceeds that of the adsorbate and that the perfect gas approximation is justified. The isosteric heat is therefore related to the calorimetric value by

+

(24) qai RT * The differences between these quantities are again small, and of the order of the error in the measurements, assuming of course that the calorimetric heat qai is determined under equilibrium conditions. b. Temperature Dependence. The isosteric heat is itself a function of temperature and the explicit dependence of the isosteric heat upon this variable is given by Qst

=

T

where qo is the heat of adsorption a t absolute zero and Ac, the difference in the partial molar heat capacities of gas and adsorbate a t constant pressure and composition This temperature dependence can be quite significant in comparing measurements obtained in different temperature ranges. Consider, for example, a molecular gas which adsorbs atomically on a uniform surface.

r.

MODERN METHODS IN SURFACE KINETICS

307

I n the gas phase, each molecule has three degrees of translational plus two of rotational freedom so that cp $R,plus a small contribution from vibration which will increase to R a t higher temperatures. At low temperatures the atoms in the adsorbed layer will be localized and vibrationally unexcited. I n this temperature range the isosteric heat therefore increases as QR.As the temperature is raised, however, and surface vibrations begin to contribute, the specific heat of the adatoms will approach that of the gas and finally exceed it, causing a diminution of the heat. This trend reverses once the adsorbed layer approaches the two dimensional gas. Presumably the vibration perpendicular to the surface contributes somewhat before the vibrations in the gas phase become important. The specific heat of the adsorbed layer will therefore continue to exceed that of the gas by a quantity of the order of +R, until all vibrational degrees of freedom are excited and the gas again dominates. The minimum in Ac, occurs when the adsorbed gas is fully excited vibrationally and the unbound entities are not, giving a value of Ac, - 4 R. A diminution in the isosteric heat by 5 kcal/mole can therefore be expected on raising the temperature from the laboratory environment t o 1000° K.* N

-

N

3. Reaction Studies a. Rates. Kinetic measurements are preferred over thermodynamic ones because of the additional atomistic information they may provide. The activation energy E derived from the logarithmic temperature dependence of the rate constant k,

can be related to the equilibrium heats as follows: For an adatom layer formed from a molecular gas a t a pressure p , the condition of detailed balance which must be met at equilibrium can be written as p (27r m,kT)-1'2kA ( T ,n) = n2kD( T ,n ) . (26)

k, is a sticking coefficient for gas colliding with the surface a t a rate p (27 m,kT)-'12, and kD the rate constant for molecular evaporation. Both kA and kD can be considered functions of the temperature T and the surface concentration n. This concentration, moreover, is always

* A more quantitative estimate of the temperature dependence of the heat can be made by adopting a model for the variation in the potential energy of a n adatom over the surface such as has, for example, been proposed by Hill (35).

GERT EHRLICH

308

given in terms of the bulk or geometrical surface area and as such is directly related to the ratio of adsorbed to solid atoms which enters in Eq. (23). Equating the logarithmic derivative of p obtained from Eq. (23) to the same derivative obtained from Eq. (26) it immediately follows that the activation energy for desorption E D is simply related to the isosteric heat by

r

All energy parameters in Eq. (27) refer to equilibrium. For metals a t least, this imposes no special restrictions since the diffusion of adatoms usually is rapid enough to establish an equilibrium distribution over the surface regardless of the supply of atoms from the gas phase (36). Provided the activation energy for adsorption is negligibly small, the desorption energy is therefore equivalent to the isosteric heat. Two matters require further attention, however. (1) Approach to equilibrium. I n flash desorption the activation energy should be completely equivalent to that found by standard methods. The extent t o which an equilibrium distribution of adatoms over the surface can be obtained may, however, remain in question. I n an adsorbed layer formed a t low temperatures, a t which low surface mobility prevents any approach to equilibrium on the time scale of the experiment, a redistribution of the adatoms can only occur between the temperature marking the onset of surface diffusion and the act of evaporation. If this temperature interval is swept through in a time t such that an atom adsorbed on a given site is able t o travel a distance larger than some characteristic separation of equivalent sites, then an equilibrium distribution will be established before desorption takes place. The root mean square distance < x > traversed by an evaporating molecule is specified through the relations n

I

<x >=

(__J n(t-1 1

\

-

\

xZ(t)dn

)

l/Z '

,

)

-' &t*)

f,

w ( T )l a( T )d r ,

x2(t)=

(28)

t,

dn

= nx kD@'dt = n'v,

exp [ - E , / R T ] dt .

Here n ( t l ) is the surface concentration prior to evaporation, v, the preexponential term in the rate constant for evaporation, and we now explicitly introduce the temperature dependence of both the jump frequency W , and the jump distance l.For example, in the desorption

MODERN METHODS IN SURFACE KINETICS

309

of /?nitrogen from W described in Section I, B, 3, a, the mean distance traversed by an evaporating molecule amounts to a t least 7.5 x 103A; equilibrium over different sites will therefore be established provided distinct surface features, on which adsorption behavior may be different, do not exceed < x > in one critical dimension. Even if this condition for equilibrium is met, a desorption energy characteristic of the equilibrium state of the surface a t the desorption temperature may not be obtained. At temperatures such that lattice atoms can diffuse, rearrangement of the surface to a configuration minimizing the free surface energy in the presence of adsorbed atoms may take place. The time interval before the onset of desorption also has to be long on the time scale of this surface rearrangement in order to obtain true equilibrium of the surface layer. (2) Identi$cation of mechanisms. Implicit in comparing isosteric or calorimetric heats with desorption energies is the presumption that the rate process under study is the important one in establishing the distribution prevailing in the thermodynamic measurements. Thus, in Eq. (26) it is assumed that only molecules are present in the gas phase and that adsorption is entirely dissociative. At high temperatures and low pressures the gas phase will, at equilibrium, consist partly of atoms. Rate processes involving these atoms may under some circumstances predominate and have to be accounted for. An extreme example of the care required in making comparisons between the different energy parameters occurs in the dissociative adsorption of a molecular gas, in which the atoms are more strongly bound to the surface than to each other. For such a layer, gas evolution occurs predominantly as atoms. The activation energy derived from desorption experiments therefore refers to the transfer of atoms from the surface to the gas phase, whereas calorimetric measurements indicate the energy change in transferring molecules from the gas phase. I n addition to the usual temperature dependence discussed in Section I, C, 2, b, and to the corrections implied by Eqs. (24) and (27), the two quantities therefore also differ because of the differences in the state of association of the gas phase entities. Such comparisons can therefore only be reliable when the detailed kinetics are clear. Another example of physical interest arises in flash desorption. There the population and desorption energy of molecules held in different states of binding can be determined in detail. To compare these measurements with the results of calorimetric and isotherm studies, the desorption energy for each state must first be properly weighted by its population. This is illustrated for the adsorption of CO on tungsten in Fig. 31. There a diminution of the differential heat of adsorption

310

GERT EHRLICH

occurs (although the desorption energies of all states are maintained constant) through an increase in the contribution of more weakly bonded entities. An analogous effect may also be of importance in desorption. At high coverage and low temperatures gas evolution will come entirely from such weak states. Even if the desorption energy for these entities is correctly determined, the frequency factor for the desorption will be underestimated, unless the concentration is directly measured, as it is in flash desorption.

e

FRACTIONAL SURFACE COVER

FIG.31. Diminution of differential heat of adsorption through addition of molecules t o weakly held substates. Heat is estimated for instantaneous population a t 298'K as determined by flash desorption. Contribution of individual states a, PI,pa, Ps based on constant energy of 20, 53, 75 and 100 kcal/mole respectively.

b. Steady State Determinations. Closely related to the energy quantities derived from reaction rates are the heat values obtained from measurements under steady state rather than equilibrium conditions. Such steady state experiments are customary in examining atomization a t hot surfaces and in other situations where the difficulty of maintaining a complete container a t a uniform high temperature becomes significant. For a steady state between adatoms a t a surface a t temperature T,

311

MODERN METHODS I N SURFACE KINETICS

surrounded by molecular gas at a lower temperature T,, Eq. ( 2 6 ) is transformed to p (27r m2kTQ)-1'2 k,(T, T,, n ) = n2kD( T ,n )

.

k D , the rate constant for evaporation, is here recognized as unchanged from that at equilibrium. On metals, surface processes should suffice to establish local equilibrium. The temperature derivative of the steady state pressure which is to be compared with Eq. 23 now gives us

The rate constant for adsorption in metals is generally high and endowed with a small activation energy. Under most circumstances, the last term can therefore be safely neglected. Deviations from the equilibrium value of the sticking coefficient must be expected, however, if there is an activation energy opposing adsorption. Under such conditions the cold gas can be excited into the higher energy levels necessary for a reaction only on the surface. The solid thus just acts as an energy reservoir. Provided that the surface temperature is much above that necessary for reaction, further changes in temperature will have little effect upon the efficiency of the excitation process. Even under these conditions we therefore obtain a gocd approximation to the desorption energy from the temperature derivative of the steady state pressure. The energy parameters for chemisorption derived from different experimental measurements are therefore comparable, provided that for nonequilibrium determinations both in calorimetric and in rate studies the molecular processes are properly identified. Here the detailed insights achievable with flash desorption methods are particularly important.

II. Field Electron Microscopy The power of the flash desorption technique lies in its ability to give a rapid, direct count of the number of adsorbed species on a surface. However, information on the properties of the adsorbed layer is obtained only indirectly, by deduction from adsorption and desorption measurements. To supplement these indirect studies, there are needed techniques that yield information on the properties of the adsorbed material by direct observation of the gas layer. Such methods have been available for many years. For example, the change in thermionic emission of electrons after adsorption of a,

312

CERT EHRLICH

gas was used as early as 1923 to follow changes in surface concentration (37).Not too many years thereafter, in 1927, low-energy electron diffraction was first employed to examine the structure of the gasmetal interface (38). In the early 1930's, information on the atomic arrangement of crystal surfaces was also deduced from the diffraction of atomic beams (39). These techniques were, however, limited either by difficulties of execution or of interpretation. In low-energy electron diffraction one of the crucial problems has been the identification of the entities from which scattering occurs. Even as late as 1961, diffraction measurements by themselves have not been able to distinguish between a clean and contaminated surface. On the other hand, changes in electron emission are not readily correlated with changes in surface concentration. The techniques for carrying out work function measurements are also far from straight forward, as is indicated by the large discrepancy in the results obtained by different investigators (40, 41). Finally, standard determinations of the work function measure an average quantity, into which the nature of the sample as well as the special techniques enter in a significant way. An important advance in the ability to observe the surface layer directly was made by Muller who in 1936 invented the field emission microscope (42). This instrument approached for the first time the ideal of being able to scan the surface on a scale that approaches the realm of atomic dimensions and to follow rapid changes in the surface layer. Most important of all, it gave a direct indication of the cleanliness of metal surfaces. The field emission microscope has already been discussed at some length in this series (43).Here the primary aim will be to establish a perspective between the different techniques for examining surface kinetics. Consideration will therefore be limited to selected topics* that bear directly on the observation of adsorption and on the determination of the properties of the adsorbed layer.

A. PHYSICAL PRINCIPLES 1. Emission Intensity and Resolution

As the name implies, observations in the field emission microscope are based on the emission of electrons from a metal into a high electric field. The electron current i in a one dimensional system, in which the free electron approximation is used to describe the metal, is given by the

* For review of the foundationsand general problems of field emission measurement8 see ref. ( 4 4 ) .

MODERN METHODS I N SURFACE KINETICS

313

Fowler-Nordheim equation

.

t & =

;

(1.54 _____ 1O1O) A __ T exp

[ -0.68

f] .

(30)

A is the emitting area, 4 the work function (in ev), and F the electric

field expressed in volts/A; the dimensionless quantities f and g are very slowly varying functions of x = 3.79 FlIz/+ (which is just the fractional change in the work function induced by the applied field). The field is related to the applied voltage V and the radius r of the emitter

D

FIQ.32. Field emission microscope for adsorption studies. A-gas bottle; B-break off seal; C-inverted ionization gauge (also serves as selective getter); D-GranvillePhillips valve; E-ionization gauge; F-grounding rings; G-double Dewar ; H-emitter assembly (tip mounted on hairpin support wire, equipped with potential leads for memuring resistance); I-anode terminal; J-willemite screen settled onto tin-oxide conductive coating; K-ground glass port; L-trap.

314

CERT EHRLICH

through F = K V / r , so that the F(ow1er)-N(ordheim)equation can be abbreviated in the form

i/V 2 = LY. exp [ -p43’2V ] .

(31)

The intensity of emission depends upon both the work function and the field a t a particular point on the surface. The course of adsorption can therefore be readily followed: Addition of n molecules per unit area, which in the adsorbed state have an effective dipole moment of M , Debye units, changes the work function by*

-A+

= 1.2

x 1O1*mMn .

(32)

FIQ.33. Hard-sphere model of body-centered cubic emitter, [110] oriented. Number of nearest neighbor atoms around surface atom: White-4; Dark Grey-5; Light Grey-6, Black-7.

A layer of 1014 atoms/cm2,each with a dipole of only 0.1 D changes the work function by 0.04 ev. The significant improvement over standard emission observations,

* The dipole M , is defined as positive for adatoms with an excess of positive charge away from the surface.

MODERN METHODS I N SURFACE KINETICS

315

however, lies in the magnification and resolution inherent in this microscope. I n order to conveniently attain the high fields (- 0.3 volts/A) necessary for appreciable emission, the sample is shaped into a needlelike point with a radius of -2000A or less. As indicated in Fig. 32 this is mounted in a bulb of radius p . Electrons tunneling out of the tip are accelerated toward the fluorescent screen and there produce an image of the emitter a t a magnification c p / r where c is a compression factor, on the order of 0.6. The resolution is limited to -20A by the tangential velocity of the electrons in the free electron gas. Their kinetic energy

FIG.34. Field emission from clean tungsten surface in standard configuration,Orientation of the crystal surface is indicated by the orthographic projection (b).

of motion parallel to the interface amounts t o -0.1 ev, and remains unchanged in the act of tunneling resulting in a rms velocity of 2 x lo7 cmlsec. It is thus not possible to detect individual adsorbed atoms but

316

OERT EHRLICH

only larger aggregates. The resolution is sufficient, however, to specify emission changes occurring on regions of known crystal orientation. As indicated in the hard-sphere model in Fig. 33 a hemispherical emitter surface is made up of a variety of crystal planes. The closer packed of these, such as the (110), (211),and (100) in the body centered cubic lattice, have a higher work function; they therefore appear as dark spots in the more brightly emitting areas corresponding to stepped regions. With the aid of a standard orthographic projection (Fig. 34). the orientation of the emitter and the identity of the planes can thus be deduced from the symmetry of the pattern. 2. Determination of the Work Function Increment The properties of the adsorbed layer become apparent only through their effect on electron emission. The correct evaluation of the work function is therefore of primary importance in the interpretation of observations in the field emission microscope. The Fowler-Nordheim equation describing the electron flux in terms of the electric field and the work function of the surface can be applied in two ways to the problem of determining the emission properties of the gas covered surface: (1) Keeping the field constant, the change in emission current during adsorption can be related to changes in the work function. (2) Provided the work function $ and the terms a and p are insensitive to changes in the applied field, the slope of the FN plot, In (i/Vz) versus 1/ V, is given by

The change in this slope on adsorption can be used to arrive at a value of

A d , the work function increment. The first method has the advantage of experimental simplicity. More than that, the field acting on the adsorbed layer is maintained constant, so that serious effects arising from this source can be avoided. However, this method of deducing At$ rests on the assumption that the FowlerNordheim equation quantitatively describes the emission process. The potential diagram (Fig. 35) on which this equation is based is that of clean surface, to which an electron is attracted by an image force. When an atom is adsorbed on such a surface, an additional well must be added in front of the surface, Inasmuch as this change only occurs a t a few spots on the surface where atoms are adsorbed, the emission can no longer be treated as a one-dimensional problem. The Fowler-

MODERN METHODS IN SURFACE KINETICS

317

Nordheim equation therefore cannot be expected to give an accurate description of the observed emission current. The second method is more complicated to implement. I n order to follow the work function change during adsorption, the current-voltage relation must be determined rapidly at a number of applied fields. If the applied field exerts an important effect upon the polarization of the adsorbed material, then changing the field may also change the state of the adsorbed layer. The important advantage of this procedure, however, lies in the fact that it is insensitive to the shape of the barrier to electron emission.

ZERO FIELD 0

SCHOTTKY SADDLE: /3.79F " 2 e ~

I, A DISTANCE FROM IMAGE PLANE

FIG. 36. Potential energy of electron at a clean metal surface. Energy levels appropriate for high index plane of tungsten. %,-position of Schottky saddle; E,-Fermi energy.

Three differently shaped barriers (44 b, 46, 47) are shown in Fig. 36. A t the same value of the field F and the work function 4, the emission densities differ by more than an order of magnitude; still the slopes of the FN plots are all proportional to +3/2 to within the limit of accuracy of the usual measurements (-1 %). Provided the applied field is not too high, so that the thickness of the barrier still exceeds .a few atomic diameters, the changes in the area under the potential curve with applied fields are concentrated in those regions where the potential is linear in distance. The details of the barrier close to the surface thus will not affect the determination of the work function. The field dependence of the factor f in Eq. (30) is so slight that over the range of fields in a normal FN plot the variation in the slope is less than *%. The emitting area A , as well as the term g are also dependent

318

GERT EHRLICH

upon the field. Since the exponential factor dominates, these variations do not affect the experimentally determined slope. Polarization by the applied field appears to be a more serious matter. The dipole moment (in Debye units) induced by the field F (volts/A) in an adatom of polarizability a0 (A7 can be approximated by M = a. F / 3 . Values for the work function change in a monolayer of gas induced 4.

I

2-

I0STEP- FUNCTION -I -

---. (v

u.

0

-2-

-3-

COMPOUND BARRIER

-4 -

-6 -7

-

-Q

i

4

4

I/F

5

A f A

9

AIV

FIG.36. Electron emission from metals with differently shaped surface barriers. j-current density in amp/cm2, F-field in volts/A. Calculations for 4 = 4%= 4.5 ev, 4, = 6.5 ev; width of conduction band = 6.4 ev; d = 2.5A. Slopes of F-N curves identical to better than 1%.

by the applied field, on the assumption that polarizability of the adsorbed entities is the same as in the gas phase, are listed in Table 11. These changes are quite appreciable, on the order of 0.2 ev, for nonmetallic adsorbates. Once this dependence of the work function on the applied field is

319

MODERN METHODS IN SURFACE KINETICS

recognized, the EN slope becomes

However, expanding $312

g3I2around its zero field value 4 ~ (01, ~~’ =

(0) +

$312

and d In (i/V 2 ) d (1iV)

-

3/2

dF

- p g 3 ” 2 (0)

(

0

+ ...

.

(34)

The slope of the FN plot therefore gives an approximation to the value of the work function a t low fields. This is of course analogous to the TABLE I1 Field Induced Work Punelion Increases (at F = 0.3 volts/A)

Adatom

aOQ(A*)

Reference

N 0 H Xe Ba

1.13 0.77 0.66 4.01 75

48 48 49 49

a

50

Monolayer concentration Reference ( 1014 molecules/cm2) 5.60 -10.0 9.76 2.00 3.60

8

14

7 51

A P (ev) 0.21 0.24 0.21 0.27 1.81

Polarizability of free atom.

’Estimated using Topping’s depolarization relation (52).

situation in thermionic emission (53).There the slope of the Richardson ) plotted against l / T ,approaches the value equation, in which In ( i / T 2 is of the work function a t T = 0. The work function changes determined by field emission, listed in Table 111, are in good agreement with the values obtained a t low fields. The differences between the two sets are of the same order of magnitude as differences in the work function determined by a single technique. The high fields necessary in field emission do not appear t o seriously affect work function measurements. The polarization of the adsorbed layer by the applied field does, however, introduce an additional factor in the preexponential term 01. This further complicates the derivation of work function changes using the variation in the emission current wit.h adsorption a t a fixed field.

320

QERT EHRLICH

The field dependence affords a much more reliable method of evaluating work function changes since it is not tied to an oversimplified picture of the emission process. I n this, there is again a close analogy to the determination of activation energies in chemical reactions. Using TABLE 111 Work Function Increments on W from Field Emiseion and Lcnv Field Techniques (at monolayer cover)

-

System

A d (ev) Field Emission

N (clean) N (dirty) 0

H Xe

B&

A d (ev) Low Field

Reference

- 0.1-0

41 41 56 57 44a 45 44d 58

0.60 1.8-1.9 0.43 0.40 - 1.38 - 1.3 - 2.42

0.20 (310 plane) 0.50 1.90 0.48

Reference 55 54 54 54

- 1.14

54

- 2.64

59

transition state theory, it is possible to write a theoretical expression for the preexponential in the rate equation. An experimental determination of the absolute value of the rate can then be used to evaluate the activation energy. A more reliable and more generally accepted value can be attained from the temperature dependence of the rate. 3. Interpretation of Work Function Changes Even a clean, single crystal emitter surface is not made up of a single plane alone; for such a surface the work function, determined from the Fowler-Nordheim plot, is a n average quantity. The total emission from a point is the sum of currents from the microscopic regions of uniform characteristics, which will be denoted by the subscript i. I n the high applied fields, each contributes independently, and

i

= $ ii = I

Vz $ aiexp [ -&$3’2/

V] ,

The slope of the experimental plot now becomes d In ( i /V z ) d (1/V)

1

-

d

aiexp [

-~

mi

exp [ -Pi+i8”1

d

(l/V

V]

MODERN METHODS I N SURFACE KINETICS

321

The FN plot therefore yields the $ power of the intensity average ( 4 5 , 6 0 )of ( f l ~ 4 ~ ’When ~ ) . adsorption occurs, each individual adatom acts just like an area of different emitting properties. If emission is lowered by adsorption, those spots on which adsorption occurs will not contribute to the work function, which may therefore remain initially unaffected. The only change noticeable may be in the total emitting area. Conversely, if adsorption raises emission, then the areas on which adsorption occurs may dominate, even on a single, well-defined plane. I n field emission the work function therefore need not be linearly related to surface concentration. There is one further difficulty in deducing the effect of a single adsorbed atom or molecule on the electron work function of the surface. The concentration itself is not amenable to direct measurement in the field emission microscope. For example, if the work function drops in adsorption and one part of the crystal region dominates the emission pattern, then this can in principle arise because the concentration of adsorbed material is higher there, because the work function increment per adatom is higher, or through a combination of both effects. A quantitative resolution of these factors has so far not been carried out, and even a qualitative separation can only be accomplished in a few isolated instances. Low field or contact potential measurements on well-defined macroscopic surfaces have a n advantage here. The total amount of adsorbed material can be measured separately by flash desorption. Moreover, the contact potential A+ corresponds to an area average, which is also approached in low field emission measurements. The change in the contact potential in adsorption can therefore be unequivocally related to the dipole moment per adatom through Eq. (32). The difficulty in this approach lies in the preparation of a truly uniform surface of macroscopic size, which has not as yet been accomplished. 4. Interpretation of Image Changes

The problems in interpreting the changes in a field emission pattern brought about by adsorption are similar to those encountered in interpreting work function changes. Consider, for example, the sequence A clean tungsten observed in the adsorption of nitrogen on tungsten (41). tip in the standard [ 1101 orientation is shown in Fig. 34. Such a tungsten emitter was exposed to a low pressure stream of nitrogen and the emission patterns were photographed as a gas layer built up. Upon adsorption of the first few molecules (< 60 x 1OI2 molecules/cm2) the pattern in Fig. 37 changes significantly as the regions around the (100) poles become more prominently emitting. Once the surface is more heavily

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GERT EHRLICH

covered, the region around the central ( 1 10) grows darker but the (1 1 l} planes, as well as the immediate surroundings of the {211}s, become more prominently emitting. Although these photographs provide a detailed map of the emission properties of the partly covered surface they do not, without additional information, allow us to map the distribution of adsorbed material

FIG. 37. Field emission patterns from tungsten after nitrogen adsorption at T = 300°K. (a) Adsorption interval At = 6 min, n < 60 X 10la molecules/cma, (b) At = 10, n < 100 x 1012, ( c ) At = 16, (d) A t = 31, (e) At = 71, ( f ) A t = 106. Final pressure 8 x lo-" mm.

-

over the surface. The enhancement of the {1 1 lls, for example, is a relative effect. Adsorption occurring on the (331) and on planes around the (11 l} reduces the emission from these by comparison with the (111). This itself need not necessarily arise from differences in surface concentration. The variation in intensity may, conceivably, come about entirely through differences in the work function increment per adatom on these different planes. However, by careful analysis, it can at least be demonstrated that adsorption does indeed take place on most of the planes exposed. The

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changes in intensity distribution are accompanied by startling effects on the emission properties. As appears from the FN plots taken very rapidly during the adsorption and displayed in Fig. 38 adsorption lowers the emission and after the first few (,- 10) minutes the field current at constant applied voltage drops monotonically with increasing

- lo9 a

665 6.62 6.64659 599 6.605.64

106 71

At =

-14.4 I

I

I

I

I

I

l

31 16 6 10 ,

I

I

I

0 min ,

-

-

I

-$, VOLTS -' FIG.38. Fowler-Nordheim plots during adsorption of N, on W at T = 300°K. Adsorption intervals At identical with those in Fig. 37. Emission current i in amps, applied potential V in volts.

nitrogen cover. On the other hand, the work function deduced from these plots initially diminishes from 4.5 to 4.21 ev and then rises slowly 4.4 ev as saturation is approached. The disappearance of reaching the (100)s as dark regions is accompanied by a reduction in the work function. Inasmuch as the emission from the (100) is comparable to that of its surroundings, it should contribute to the intensity average making up the experimentally determined value of the work function. The clean (100) of tungsten has a work function 4 = 4.71 ev. The experimental value of 4.21 ev for the partly covered point therefore N

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GERT EHRLICH

suggests that adsorption and a lowering of the work function has occurred on this plane. A similar argument can be applied to most of the other planes except the (1 lo} and (211). Even when clean, the work function of these surfaces exceeds by far that of the surrounding higher index a r e w On completion of adsorption they still remain unemitting relative to the other crystal areas. This can again be interpreted in two ways. These planes may have remained bare when exposed to nitrogen; on the other hand, if adsorption did occur, it did not result in an enhancement of electron emission relative to that of the surroundings. Whatever may have taken place on these low index planes is shrouded in darkness. Although the field emission pattern really only gives direct information on the relative emission from different areas, it is possible to draw qualitative conclusions about surface concentrations if the experiment is properly designed. I n general, this identification is accomplished by taking advantage of surface diffusion. A deposit can be formed on the surface by evaporation from a directional source. If the emitter is maintained at a low enough temperature, this deposit will be localized on one side of the tip. This is demonstrated in Fig. 39, where a beam of N, (at T = 300°K) was allowed to impinge upon a tungsten emitter at T 20°K; the shadow side remains clean, Once the surface is heated 40°K migration occurs. The motion of the adsorbed gas into the to T initially clean regions can be readily followed. Inasmuch as a known gradient is established a t the start of the experiment, the usual difficulty of disentangling concentrations and work function increments is avoided. This technique is therefore one of the best for studying the structural dependence of surface processes. This same end can sometimes be accomplished even without shadowing (as will be demonstrated in Section 11,C, 2, a). At low temperatures, for example, a uniform distribution may be formed in the adsorption process. At temperatures at which diffusion and therefore equilibration are possible, rearrangement may then occur. From differences in these emission patterns it may then be possible to deduce relative equilibrium concentrations on different crystal planes. Only little definitive information on structural specificity has been obtained from the interpretation of emission pattern formed in ordinary adsorption experiments. Diffusion studies, however, are extremely powerful. Not only is surface migration an extremely interesting topic in its own right; it is largely through the observation of diffusion that unequivocal information has been derived on the variation of rates with variations of structure a t the atomic level. One of the most important advantages of the field emission microN N

MODERN METHODS IN SURFACE KINETICS

326

scope is the ability to ascertain the cleanliness of a surface. I n most techniques, such as flash desorption or low energy electron diffraction, the state of the substrate can only be inferred indirectly from the pretreatment of the surface or from the subsequent adsorption behavior.

FIG. 39. Surface migration of nitrogen deposited on tungsten at 20°K.Beam direction indicated by arrow. (a) Deposit at T N 20aK, localized around (OlO), (121), and (111). (c)3minat40°K,lmin (b) T N 40"Kfor 1 min. Motionoverregionsaround(OlO)and(lll). at T N 45°K. Boundary motion toward (100). (d)T N 45°K for 1) min. (e) T = 66OK. Boundary advance toward (100) haa stopped. (f) Room temperature for 1) min. (100) pole disappears, (010) emerges. (8) Room temperature for 1) min. Room temperature adsorption form of {lll}s and (211)s develops. (h) T 700°K for 1 min. Boundary in (100) region begins to fade, (111)s and {211}s change to high temperature form. (i) T N 700°K for 1 min. Equalization of (100)and (010) poles starts.

-

For metals at least, the appearance of a clean field emission pattern has been established, using as a guide the behavior of ,tungsten and similar refractory metals. With tungsten, the conditions under which a clean surface can be attained have long been known. The emission

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pattern (Fig. 34) of such a clean surface is distinguished by the gradual variation of emission intensity with direction. The Wulff plot for a metal has cusps only a t the directions perpendicular to the close packed plane. Only in these directions is the surface flat, and relatively nonemitting. I n going from one low index plane t o an adjacent one, there should be no ridges and therefore no sharp variations in emission intensity. On a contaminated surface in contrast, impurities are usually held a t specific planes and result in a sharp, highly detailed pattern, readily distinguished from the clean emitter. This should be apparent from the comparison of the clean Ni surface in Fig. 40, with one contaminated by oxygen. As a general rule, the more interesting the appearance of a field emission pattern, the dirtier the surface.

Determination of Rate Parameters Even without being able to relate emission changes to changes in surface concentration directly, energy parameters can be derived in the field emission microscope. For surface diffusion (57) the procedure is straightforward: The diffusion coefficient D, can usually be simply determined by observation of boundary motion. x,, the distance traversed during time t is given by x, = d/o,t; the diffusion coefficient can be written as D, = P V, exp [ - E,/RT] . (36) 5.

Here t is the average jump distance and V, the preexponential for the rate of jumping, both assumed temperature-independent. Em,the barrier to motion over the surface, is obtained as usual from the derivative

For desorption, the analysis is more difficult. The time interval AT that elapses between the change from one fixed emission current value to another, or for a fixed change in the emission pattern, is in the limit AT --f 0 inversely proportional to the rate of the desorption reaction. A semilogarithmic plot of this time interval AT against 1/T therefore corresponds to a plot of In (dnldt) vs 1/T a t constant concentration, and yields the desorption energy according to Eq. (25). The nature of this energy parameter was already discussed in Section I, C, 3, a. For metals, diffusion should be sufficiently rapid to maintain surface equilibrum during evaporation even when the surface is energetically heterogeneous. If changes in the total emission are observed in fixing AT, then the observed energy corresponds to the isosteric heat, that is, to the change in partial molar enthalpy in going from the gas

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FIG.40. Effect of contamination on field emission pattern (61). (a) Clean Ni surface, showing smooth variation of emission intensity with orientation. (b) [loo] oriented Ni tip, after exposure to oxygen ( p 5 x mm) and heat treating at T 900°K.

-

-

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phase to the adsorbed layer. Such measurements are not necessarily more powerful than standard methods for determining isosteric heats. Indeed, field emission may fail to reveal the detailed distribution of adsorbed material in different binding states that is immediately apparent in flash desorption. However, if the field of observation can be limited to a small portion of the surface, to a single crystal plane, for example, the energy parameter will correspond to the activation energy for the changes occurring on this particular area (in this given environment), In this respect, emission microscopy has a tremendous advantage over more macroscopic measurements of desorption-the effect of structural heterogeneity on the desorption energy can be apprehended. Because of this detail it is both more difficult and more important to properly identify the mechanism by which changes occur. A change in the emission pattern of a surface partly covered by adsorbed gas may indicate the onset of surface diffusion, evaporation or rearrangement of the substrate. The identification of the appropriate mechanism by ordinary field emission measurements may be difficult.

0

b

FIG.541. Changes in field emission from nitrogen covered tungsten after heating. (e)W tip efter adsorption of NPfor 11 min at N 6 x 10-8mm. (b) Pattern after heating to 370°K for 2 min without further adsorption.

Consider, for example, the sequence in Fig. 41. There the temperature of a tungsten tip, covered by a small amount of nitrogen, was raised from 300 to 370°K. The most significant changes occur in the region around the (111) and (211) planes. I n shadowing experiments (such as indicated in Fig. 39), no redistribution of nitrogen was noticed at these low temperatures. Diffusion therefore does not appear responsible. Flash desorption, however, indicated the evolution of the a state in this temperature range (Figs. 10, 22). With the aid of supporting information from separate diffusion and evolution experiments the

MODERN METHODS IN SURFACE KINETICS

320

changes in field emission on heating can therefore be properly assigned to the removal of nitrogen held in a weakly bound state." There is one further practical difficulty in deducing rate information from observations in the field emission microscope. Changes in the patterns are usually not large enough to permit identification of con0.1. Such measurements therefore centration differences less than A 0 correspond not to the rate do/&, but rather to an integral change ABlAt. For simple desorption reactions the rate constant k, can be evaluated immediately from

-

- A In 8 1

Ae

1

1st order =

kDAt 2nd order

.

(38)

Special precautions (62)are necessary when the rate constant is itself dependent upon surface concentration. Consider an extreme case; evaporation in which the rate is determined by the variation of the heat of desorption with cover. For a rate law

vD = -dO/dt

= vexp

RT

it was already pointed out by Langmuir (63) that the time increment for a change from O1 to a lower concentration O2 is given by

Because of the exponential dependence upon cover, the lower of these rates will determine the time interval for the usual concentration intervals and therefore

When the desorption energy decreases with increasing concentration, the rate constant as well as the desorption energy determined from experiment corresponds t o the lower concentration. Especially in estimating values of the frequency factor it is therefore important to establish the exact form of the rate law, by testing the concentration dependence over a wide range of 8.

* That this is held preferentially on the {lll)s again does not follow unequivocally from the fact that pattern changes are concentrated in this region. Such specificity is. however, confirmed by flash desorption from ribbons with small contributions from [ l l l ] oriented grains.

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These difficulties can be avoided by measuring the time interval required for a given emission current increment a t a fixed voltage. The emission varies exponentially with work function, and therefore gives a far more sensitive indication of change than the voltage, so that the true rate of change of surface concentration may be apprehended directly. This method has one added advantage. Current changes at amp can be easily measured. I n contrast detection of pattern changes requires a total emission a t least one order of magnitude higher, and therefore also a higher field. The applied field lowers the chemical potential of the adsorbed layer, by an amount that can be approximated as 0.623 ( F / 3 )[BN,,(F/3) + M,] ev. For adatoms with a dipole moment of - 2 0 , and a polarizability 0.7A3, a field of 0.3 volts/A lowers the potential by 0.13 ev. Such changes may become important in both equilibrium and rate studies. However, in any field emission measurement it is desirable to minimize exposure to the field, and to establish that the act of observation has not perturbed the system. N

B. EXPERIMENTAL CONSIDERATIQNS Field emission microscopy is by now quite a standardized operation, described in some detail in the literature. Only two developments that considerably simplify experimentation will be mentioned here. 1. Observation of Surface Difiusion

The full power of the field emission microscope is only revealed in measurements of surface diffusion. However, shadowing the emitter with gas presents quite a problem in studying the migration of permanent gases. This demands a directed gas source in the system, as well as provisions to trap excess gas that is not deposited on the field emission point. Gomer ( 4 4 4 has accomplished this very ingeniously by immersing his apparatus entirely in liquid helium. The cold walls trap any colliding gas molecules, with the exception of helium and hydrogen. Gas can then be deposited on the emitter by activating a suitable source (discussed in Section IV, F, 1, a ) built into the field emission tube. However, this technique requires sizeable quantities of liquid helium and observation of surface events occurs through the coolant, and is therefore tricky. The requirements for shadowing can be more conveniently met (41) by the microscope shown in Fig. 42, in which localized deposits of any gas desired for study can be readily formed on one side of the field emission tip. I n this tube, patterned after one of Muller’s low temperature

MODERN METHODS IN SURFACE KINETICS

331

field ion microscopes, the stem (1) bearing the field emitter assembly (point, supporting filaments and potential leads for temperature determination) is developed into a Dewar vessel and is itself shielded by two separate concentric Dewars (2 and 3). These are pierced on a level with the emitter tip. Through this opening a beam of gas can be shot a t the emitter from a small diameter tube (B) leading t o the separately

"I

1:

I

\/" 1

w V t

N,

1

Pumps and Traps

7

and Traps

FIG. 42. Field emission microscope for surface diffusion studies. A-anode terminals; B-beam tube; C-tin oxide conductive coating; I-inverted ionization gauge; Nliquid nitrogen cooled trap; P-magnetically operated ground glass port; T-thermocouple gauge; V-Granville-Phillips valve (type C ) ; W-willemite screen. l - e m i t t e r assembly (point mounted on supporting filament carrying potential leads) developed into Dewar; 2-trapping Dewar filled with liquid hydrogen; 3-shielding Dewar for liquid nitrogen.

pumped gas handling system. The duration and intensity of this beam are regulated by a bakeable valve (V). I n operation the second Dewar is cooled with liquid hydrogen, and gas molecules not condensed on the field emission assembly are trapped on colliding with the glass walls. Provided the coolant in the emitter stem (1) has been properly chosen

UERT EHRLICH

332

to bring the tip temperature below that at which surface migration can occur, only one side of the emitter is coated with gas and the spreading of this deposit over the point can then be examined on raising the emitter temperature. This design incorporates several desirable features. ( 1) The microscope is continuously pumped ao that a variety of experiments can be carried out in it, (2) Only minimal amounts of liquid hydrogen (or helium) are required (70 cc for the central Dewar and 150 cc for Dewar #2). (3) Surface diffusion of any gas which can be trapped on a liquid helium cooled surface can be studied, merely by changing the gas bottle.

Work Functiom Measurements of the current voltage relation in field emission have in the past been so lengthy that work functions derived from the FN plots 2. Rapid Determination of

9'"

Emission Anode

ir

*-I

T

I I 5 V AC

-.-

500 M

._.

t

euchlnq Voltoqe

2

Qlomv Recorder 0

FIQ.43. Circuit for continuous recording of field emieaion voltages (41). Electronic relay shorts recorder input when bucking volteges exceeds voltage across R.

have generally not been taken during the adsorption step itself. This is, of course, a serious limitation, since interpretation of changes in the emission current at fixed fields has been demonstrated as unreliable. By automatic recording of both the emission current and the applied voltage, work functions can be rapidly determined as follows. The electron current in the interval 5 x 10-8 to 2 x 10-7 amps is measured by a do microammeter, the output of which is displayed on a recorder. The potentisl across a small portion (R in Fig. 43) of the current limiting

MODERN METHODS M SURFACE KINETICS

333

resistor separating the anode from ground is plotted on a separate 1800A, field emission currents recorder. For points with a radius r only become appreciable at anode potentials V > 2800 volts, and the voltage range covered in the FN plot therefore amounts to but a small fraction, approximately 7 yo, of this minimum potential. To measure changes in anode voltage more accurately, a part of the potential drop across the resistor R is therefore balanced out using the circuit sketched N

-

FIQ. 44. Effect of ion bombardment on field emitter. Bombardment of W at 79'K by Ar at average pressure of lo-' mm. Voltage V required for 2 x lo-' amp emission. Clean emitter, V = 3020 volts, (A) V = 3020, (B)2910, (C) 2830, (D),2710. Total duration of bombardment, 108 min.

in Fig. 43 so that a full scale displacement on the recorder corresponds 1000 volts. With these modifications, all data for a Fowlerto A V Nordheim plot can be taken in 30 sec, with an average deviation of less than 0.5% in the slope. Significant work function measurements during the course of adsorption are therefore feasible. N

334

OERT EHRLICH

This is important, not only for properly interpreting emission changes, but also to minimize the effect of ion bombardment. Electrons traveling toward the screen may ionize gas molecules present in the tube. These ions are accelerated toward the tip, where they can significantly alter the emission properties. This already occurs a t moderately low pressures. Thus the sequence of photographs (45) in Fig. 44 was obtained with argon a t an average pressure of lo-' mm. At 79°K there should be essentially no adsorption (n < lo6molecules/cm2)of argon on a tungsten surface. The changes in the pattern arise entirely from bombardment during the intervals (amounting to a total of 4 min in going from A t o B, and from B t o C) during which the field is on while the emission is periodically checked. At lower temperatures (T 20°K) bombardment would be superposed on effects arising from adsorption and would complicate the interpretation of the pattern changes. Such ion bombardment can be minimized, however, by minimizing the product iNAr, where i is the emission current, N the gas density and A7 the interval of measurement. Automatic recording of the current voltage relation accomplishes just this.

-

C. COMPARISONOF TECHNIQUES-THE INTERACTION OF RAREGASES WITH METALS The relative merits of flash desorption and field emission, as well as the problems encountered in the study of adsorption phenomena by each can best be appreciated by sketching the information attained on one particular system. Here we will consider the data on the interaction of xenon with W, obtained by both methods ( 7 , 4 5 ) . 1. Flash Desorption

a. Experimental Observations. Adsorption of gases through van der Waals forces is by far the simplest and most readily amenable to theoretical analysis, but little work has been done to characterize these interactions by modern methods. This is particularly true for flash desorption studies. These depend upon adsorption on a small, well defined sample, in preference to adsorption on the container walls of much larger surface area. I n van der Waals interactions, the driving force for adsorption on the metal sample differs little relative to that on the glass. Accurate counting of the amount of gas taken up and given off by the sample therefore involves serious corrections for competing adsorption on the cell wall. Both the difficulty of executing measurements, as well as their interest, is demonstrated by Pig. 45. On rapidly heating the sample above the

MODERN METHODS I N SURFACE KINETICS

335

adsorption temperature (82"K), the gas pressure rises discontinuously. After an initial evolution of gas, denoted by 01 in the desorption trace, the pressure remains essentially constant. Only after the temperature has risen further does gas evolution resume, resulting in another peak.

N

t

-

-

FIQ.46. Flash desorption of Xe from W after adsorption at T 80°K (7). Time scale: 1 unit = 60 msec. Total surface concentration 6 x 10l2 molecules/cm2.Break in curve is indicative of 2 binding states.

This simple observation proves the existence of different states of binding for xenon held on a polycrystalline tungsten surface. The 01 and /3 peaks on the pressure vs time curve are indicative of xenon evolving a t different rates from the surface-in these experiments impurities were excluded by the usual ultrahigh vacuum techniques, supplemented by selective gettering (as outlined in Section IV, F, 2). The nature of the sites on which these states form does not emerge from these measurements. However, it is clear that the /3 state is the more stable, evolving a t higher temperatures. Upon reaching a maximum value, the pressure in Fig. 45 drops rapidly as the evolved gases readsorb on the cold container walls. This 20 msec. I n this system pumping occurs with a typical half-life of 50 liters/sec. By contrast, by adsorption on the walls amounts to flash desorption of N, and CO (as in Fig. 10) was carried out a t low pumping speeds (S, < 0.2 liters/sec) and a t a heating rate such that loss of gas during the period of desorption constituted a small correction, amounting to less than 3%. For these two gases the peak height N

N

336

OERT EHRLICH

therefore gives a measure of the amount desorbed. I n the evolution of Xe, l/6 of the molecules evolved up to the time the pressure peaks are lost by pumping, and even a t the peak the evolution may be only 90% complete, Consequently, the amount desorbed must be calculated by estimating the pumping losses over the entire heating cycle [Eq. (7)]. Despite difficulties in measuring this pumping speed accurately, the course of the adsorption can be followed as a function of time. The results of such experiments, the first on the rate of physical adsorption on a well defined metal surface, are showq in Fig. 3 as a plot of sticking probability vs surface concentration. The rate of adsorption has two interesting features. Even a t low 0.4; not every xenon atom concentrations the sticking coefficient is colliding with the surface condenses on it. Furthermore, the rate of adsorption is insensitive to the amount adsorbed. The sticking coefficient rises only slightly over the range of surface concentrations for which measurements can be accurately made. This still corresponds to a number of adsorbed atoms small compared to the total number of atom sites on the surface, which can be approximated by 10'5 per cmz. The total amount of xenon adsorbed under the conditions of these 10-7 mm, only adsorption experiments is also small: At 82°K and p 200 x 1012 molecules/cm2 are held by the tungsten surface. Changes in the interaction of the adsorbed material with the surface as the amount adsorbed increases also can be deduced by flash desorption. Desorption measurements similar to those of Fig. 45, taken at ever increasing concentrations, indicate that the relative concentration of the more tightly bound /3 state diminishes as the over-all surface population increases. A measure of the stability of the adsorbed layer is given by the evaporation rate which can be deduced from the plots of the adsorption rate against pressure, as in Fig. 5. As the surface concentration increases, so does the rate constant for evaporation. The energies of desorption at increasing surface coverages can be estimated, assuming a first-order desorption with a frequency factor of 5 x 10l2sec-1. The values in Fig. 5 have only semiquantitative significance because of the difficulties inherent in such measurements but do follow the trend expected from flash desorption-at higher concentrations the desorption energy diminishes. Moreover the magnitude of the binding energy, 6 kcal/mole, corresponds to that expected for van der Waals interactions. b. Interpretation. The adsorption rate deduced from these measurements does not conform to the expected. Langmuir (64) originally envisioned that the rate of physical adsorption should be proportional to 1 - 8, the fraction of the total number of surface sites unoccupied. N

N

-

MODERN METHODS IN SURFACE KINETICS

337

On this picture, an atom on striking a localized gas layer would be reflected if it collided with a filled site, and bound if it struck an empty one. This is not at all the behavior found for the adsorption of xenon which rather resembles that of chemisorbing systems. The sticking coefficient does not vary significantly as the surface concentration is 112 the increased. Despite this, even at low concentrations, only total number of atoms striking the surface is captured. The slight concentration dependence can be accounted for on the assumption that in a collision a xenon atom has a finite probability of remaining at the surface regardless of whether or not the struck site is filled. In a collision with a filled site, that is with a xenon adatom, the incoming atom may still be able to give up its kinetic energy and may be bound a t least transiently. Thereafter, this atom can either jump to another site or else evaporate. Assuming that at the temperature of the experiment xenon atoms on the average make n such jumps before evaporating from a complete monolayer, and that the behavior of the layer does not change with cover, then the sticking coefficient would on this model be proportional to 1 - (O).+l, and would indeed vary little with increasing surface concentration (65). The fact that even the initial value of the sticking coefficient is less than unity suggests that the processes actually occurring at the surface are more elaborate. Physical adsorption should not involve any activation energy, but despite this simplification, two basically different effects may participate: (1)An atom colliding with the surface may not be able to give up its energy of motion perpendicular to the interface, and may therefore be rescattered into the gas phase. The over-all adsorption rate is, on this picture, limited by the efficiency of energy transfer to the lattice. (2) The surface may be energetically heterogeneous. On certain crystal faces the xenon atom may be more weakly held than on others, and under the experimental conditions the equilibrium concentration may be negligibly small on the poorer planes. A xenon atom colliding with one of the latter will then be permanently adsorbed only if it can jump over this patch and reach a location at which the binding is stronger. On this model, the adsorption rate is limited by the probability that such an adsorbed atom will be able to migrate away from a poor patch before reevaporation. Flash desorption does not provide information on surface diffusion, or on the crystal dependence of the binding energy; a decision on the mechanism of adsorption based solely on this one type of measurement is therefore not possible. The interpretation of the data on evaporation is again equivocal. Although it is clear that the binding of xenon on a tungsten surface N

338

GERT EHRLICH

occurs in at least two states, the nature of these cannot be specified. It is reasonable to assume that differences in the binding energy stem from differences in the surface structure; the fl state is observed a t very low coverages a t which interaction with neighboring atoms should be negligible. However, the particular atomic arrangement that enhances binding remains unknown. The decrease in the desorption energy also cannot be interpreted unequivocally. This could arise entirely from repulsive interactions between adsorbed atoms, or else from intrinsically different binding on different parts of the surface. This is, of course, the classical difficulty of interpreting macroscopic heat measurements. Despite this, these observations are of interest. They establish a lower bound on the efficiency with which the kinetic energy is transferred in a collision. Moreover, desorption proves directly the heterogeneity of the tungsten surface involved. 2. Field Emission

I n the field emission microscope observations are made on the properties of the adsorbed layer itself. This technique therefore should not suffer from the difficulties experienced in flash desorption of xenon; it should, moreover, yield detailed information on the structural dependence of the interaction. Two problems might prevent useful results : (1) The field required for emission may seriously affect the surface distribution of a weakly adsorbed gas. The importance of such field effects can only be estimated experimentally. (2) Adsorption of a rare gas may not affect the electron emission sufficientlyto permit significant observations. Previous work on evaporated films by Mignolet ( 5 4 ) established quite significant work function changes in physical adsorption and makes it clear that lack of sensitivity should not prove a limitation. a. Experimental Observations. ( 1 ) Adsorption of Xenon at T > 79°K. When xenon atoms a t 300°K strike a tungsten tip* maintained at 79°K the field emission pattern changes as in Fig. 46 and 47. Most significant is a marked structural specificity. The increase in emission, indicating a lowering of the work function, appears to be concentrated in the vicinity of the (100) poles. Only a t the very end of the adsorption, Fig. 47d, does the emission spread to other surface regions. Here the (111) planes appear as small but well-defined triangles; still the pattern is dominated by the emission from the ring around the (loo}, segmented only by the dark 100 zone line. With continuing adsorption the voltage necessary to maintain a constant emission current diminishes monotonically and a work function of 3.12 ev is obtained (from a Fowler-Nordheim

* Observations on the adsorption of Xe on Mo, quite siniilar to those on W, are to be found in ref. 66.

339

MODERN METHODS I N SURFACE KINETICS

plot) for the saturated point compared with 4.5 ev with the clean emitter . If adsorption is carried out with the emitter at 85"K, the sequence of patterns as well as the changes in the voltage are initially the same as a t 79°K. However, saturation occurs at a much lower coverage which corresponds to a work function of 3.75 ev and to a pattern resembling

FIG. 46. Field emission from tungsten at 79°K during adsorption of Xe. Pattern changes are concentrated around (100). Voltage ( V ) required for emission of 5 x 10-0 amp: clean emitter, V , = 4770. (A)A t = 12 min, V = 4720; (B)A t = 3, V = 4670; (C)At = 4, V = 4600;(D) At = 5, V = 4500.

that in Fig. 47c. Adsorption around the (100) poles can be observed a t temperatures as high as 104°K. However, xenon is completely removed within 30 sec after raising the temperature to 114°K. From the temperature dependence of the rate of change of the field (required to maintain 100°K constant emission) on raising the tip temperature from 77' to we estimate a heat of desorption of 8-9 kcal/mole. The disappearance of emission from the (111) planes on raising the temperature from 79"

-

-

340

QERT EHRLICH

to 85'K suggests a desorption energy of -5 kcal/mole for the material visible there. (2) Adsorption at T < 79°K. As the temperature of the tungsten surface is lowered the intensity distribution over the emitter becomes quite different. The first change appears at 72°K. If adsorption is

FIU.47. Continuation of X e adsorption on W at T = 79°K. (A)At = 8 min, V = 4280; (B) At = 12, V = 3930; (C)At = 22, V = 3300; (D)At = 47, V = 2900. Fiml Xe pressure lo-' mm.

-

allowed to occur below this temperature and above N 65"K,the sequence of patterns shown in Fig. 48 is observed. These are distinguished by a brightly emitting border around the lower edges of the emitting cap and leading from the edges of the (1 lo} plane to the zone line and the edge of the (211) plane. After the initial stages (in Fig. 48) the patterns become almost indistinguishable from those obtained a t the higher temperatures. Such behavior is to be expected if above 72°K adsorbed xenon is mobile on all planes and adsorption results in an equilibrium distribution

MODERN METHODS I N SURFACE KINETICS

341

of adatoms over the surface, and if below this temperature the mobility of the adlayer around the (100) is limited. The appearance of the low temperature pattern, as well as the concentration of the gas, are on this model dictated by the rate of supply of gas to the cap. This rate is greatest around the borders of the (110) and (211) planes, inasmuch as these regions receive xenon both directly from the gas, as well as by surface diffusion across the regions of the (1 lo} and (1 1 l} planes, which apparently hold gas only weakly at these temperatures.

FIG.48. Adsorption of X e on W at 66"K, showing boundary formation along [ l l l ] zone lines. Clean emitter required 4770 volts for 6 x loeQamp emission. (A) At = 144 min, V = 4710; (B) At = 20&, V = 4680; (C) At = 26, V = 4620; (D) At = 31, V = 4510. Pressure 10-lOmm.

-

This explanation assumes that the absence of any pronounced initial changes in the electron emission from the region around the (111) indicates a lack of adsorption. It is conceivable, however, that this interpretation may not be justified, that adsorption in this crystallographic region does not affect the surface dipole and therefore does not

342

QERT EHRLICR

become visible. The course of the adsorption a t still lower temperatures must be examined to decide between these alternatives. At 20°K the mobility of adsorbed xenon should be limited, and differences in binding energy on the surface thus should no longer influence the distribution of adsorbate over the emitter. As is apparent from Fig. 49, adsorption on a liquid hydrogen cooled emitter initially raises the emission everywhere except for the high index planes, which

Fro. 49. Field emission from W at 20°K after adsorption of Xe. Initid adsorption raises emission from {ill} as well as around {loo}. Clean emitter, V = 4600 for 5 x amp. (A) At = 9 t min, V = 4270; (B) A t = 26, V = 4000; (C) A t = 39, V = 3500; (D) At = 138, V = 2770.

remain dark. This indicates that adsorption around the (111) poles does significantly affect the surface dipoles, and that the relative darkness of these areas a t higher temperatures was interpreted correctly as a lack of adsorption. (3) Field effects. One more variable, besides pressure and temperature, may influence the stability of the adsorbed layer and the appearance

MODERN METHODS IN SURFACE KINETICS

343

of the field emission patterns-the electric field. This is applied only during brief intervals (- 10 sec) necessary for observing the patterns recorded here. Under the conditions of these experiments, a t temperatures below 79"K, the field does not appear t o affect any significant changes. However, a t this temperature it is possible a t saturation to diminish the emission from the (1 1 1); a t a somewhat higher temperature it can be eliminated completely by quickly raising the field. This is, of course, the expected behavior. Each individual xenon adatom constitutes an effective dipole, with its positive end pointing away from the surface. The application of the field necessary to allow electron tunneling raises the potential energy of a dipole thus oriented; this transforms the adsorbed layer. Such a rearrangement should be significant only for the adsorbed entities in equilibrium with the gas, for which a small change in driving free energy results in a displacement of the equilibrium. A possible change in the adsorbed gas a t very low fields is more difficult to examine. However, when a rearrangement is observed (at the (111)s) it is found to be slow compared to the time necessary to raise the field sufficiently to observe the emission pattern (- 1 sec). Moreover, the work function changes determined by field emission are in qualitative agreement with contact potential changes, measured by Mignolet in the absence of a high field (Table 111).I n field emission, the work function average is weighted in favor of regions of high emission in accord with Eq. (35); a larger diminution of the work function must be expected than in contact potential measurements, and this is shown by the experimental data which result in an area average. The field can therefore be eliminated as being responsible for the observed distribution of adsorbed xenon. b. Interpretation. The most striking feature of the changes observed in the field emission microscope is the pronounced dependence upon the structural features of the surface. That this specificity is a result of differences in the binding of gas atoms a t various lattice positions, and does not arise solely because of differences in the work function increment per adatom on different crystal planes, has been established by experiment. The changes in emission therefore indicate quantitative differences in the concentration of adatoms over different crystal regions. It follows from the sequence in which xenon is adsorbed on a tungsten tip, as well as from the border formation, that the binding energy of xenon is greatest in the regions around the (100) poles; the binding energy must be considerably lower on the (1 lo), as well as on the (1111, and on the regions immediately around these and on the (211)s. For

344

OERT EHRLICH

-

desorption from the region around the (loo), the activation energy 5 from the (111). The amounts to 8-9 kcal/mole, compared with detailed order in which the binding energy varies over the lattice planes cannot be established from the observations, but in Table I V the values are listed qualitatively as high, H, and low, L. Comparison of the arrangement of Xe on a bcc lattice model (as in Fig. 50) with the experimentally observed distribution of Xe over the

FIG. 60. Xe atoms adsorbed on hard-spheremodel of a tungsten emitter. Number of Xe contacts is largest at (310),(611) planes.

field emitter indicates that binding is strongest when the adatom is surrounded most intimately by lattice atoms. That the forces can be approximated as being pairwise additive is suggested by the agreement of the qualitative order of stability in Table I V with that estimated on the assumption of van der Waals forces between individual atom pairs, This correlation also holds for surface diffusion. From the rate of disappearance of the boundaries observed in Fig. 48, the activation

MODERN METHODS IN SURFACE KINETICS

345

energy for migration in the atomically rough regions between the (110) and (100) planes can be estimated at 3.8 kcal/mole; this compares with 1.5-3.4 kcal/mole of diffusion over the smoother (111)s. This structural specificity is not only of interest for its bearing on the nature of physical adsorption in general; it also permits a more thorough analysis of the flash filament experiments. TABLE IV

Van der WuaZs Interactions for Xe om W

Plane

v (ev)

116 123 130 120 112 100 122 111 110

4.30 4.52 4.31 4.34 4.65-4.88 4.71 4.35 4.39 6.99

ELJIB, (111)

Calculated’

E D Experimental

1.38 1.38 1.29 1.29 1.25 1.08 1.07 1.00 0.89

Data (44b). Attractive pair interactions only (45).

3. Reexamination o j Macroscopic Measurements The observations in the field emission microscope establish that under the conditions of the adsorption studies of Section 11, C, 1, a, xenon was mobile, and that the binding energy varies with the crystal orientation. The structure of the filament surface is not known in detail; however, this is a polycrystalline specimen, the cylindrical surfaces of which are made up of planes with the orientation (hhk). The condensation coeficient must therefore be an average quantity, and energy transfer may thus not be the limiting step. I n the adsorption experiments, only a small portion of the surface was filled when brought to a steady state a t p = lo-’ mm. If this concentration were to correspond to saturation of the entire surface the xenon 50A2, more than three adsorbed would have a cross-sectional area of times the gas kinetic value. A more reasonable view is that only the rougher, stepped planes are filled. From the calculated binding energy ratios which appear in fair agreement with experimental results found by field emission for the (111) and {130}, we can estimate for Xe a heat of 5.4 kcal/mole on the (100) and 4.5 kcallmole for the (110) plane.

-

346

GERT EHRLICH

Under the conditions of the rate experiments evaporation was found t o be appreciable for xenon held with a binding energy of 6 kcal/mole. During a run ( p < 5 x 10-9 mm) the (110) planes should, therefore, have a cover of less than 1011 atoms/cniZ, the (1 11}’s less than 5 x 10l2. This is also consistent with the direct observation of the (111) planes in the field emission microscope-at pressures comparable to those in the adsorption experiments the (111) was occupied a t 79”K, but not at 85’K. The picture appropriate t o the adsorption experiments is therefore the following. Xenon atoms, striking high index planes, will, if they are able to give up sufficient energy, be bound. Atoms colliding with smoother facets not only have to lose their kinetic energy, they will be saved from subsequent reevaporation only if they can escape to the rougher areas. Assuming pair-wise interactions, diffusion on the (1 10) should occur 0.5 kcal/mole and on the (11 l} over a barrier of 1 over a barrier of kcal/mole. The mean square distance covered by a n atom prior to evaporation (67) is

-

N

Under the experimental conditions atoms on the (110) and (111) could therefore cover a distance more than 5 x 105A.If the smooth facets on the sample are of this order of magnitude or smaller, then collisional energy transfer will still limit the over-all rate: Even though these planes do not act as sinks, they serve as highly efficient catch basins. Evaporation from such smooth planes will contribute to a lowering of the adsorption rate if the size of these planes exceeds < x >. Facets wider than < x > (that is, on the order of 0 . 1 mm) are not too likely, however, and reevaporation of xenon probably does not contribute significantly to the deviation of the sticking coefficient from unity. The presence of multiple states in flash desorption, which in Section 11, C, 1, b, was assigned to a structural effect, has been confirmed by the observation of significant variations in the binding energy of Xe over the surface. More than that, the diminution of the heat of adsorption, deduced from macroscopic measurements, likewise appears to be dependent upon the structure of the surface-in the field emission microscope, the {I 1I > planes were found to fill in last and with a lower binding energy than typical of rougher planes. An even closer look at the atomic details of the interaction of xenon with tungsten cannot be achieved; the binding is so weak that observation in the field ion microscope does not appear feasible. However, flash desorption measurements extended to crystals with a single surface

MODERN METHODS IN SURFACE KINETICS

347

orientation and carried out over a wider temperature range should provide an unequivocal answer to the efficiency of condensation. It is clear even from this single, limited example that the field emission microscope and flash desorption supplement each other in an important way. Field emission microscopy can, if properly executed, provide detailed information on the distribution of adsorbed material over a surface and on the rates of atomic processes occurring in the adlayer itself. The fields necessary for such observations do not appear as a significant obstacle. Flash desorption, on the other hand, provides straightforward data on the kinetics of molecular transfer between the gas phase and the surface. Not only is this vital to the understanding of the kinetics of heterogeneous reactions, but from such data the detailed energetics of gas binding at a surface can be readily deduced. Without field emission data it is difficult to relate this information to the structure and properties of the surface layer. Conversely, without the information available from macroscopic measurements, the nature of the rate processes responsible for the changes observed in the field emission microscope is difficult to identify unequivocally. Finally it should be noted that although the theory of field emission is intricate and not yet fully developed, the experiments are simple and straightforward. Difficulties arise only in the interpretation of emission changes in terms of the chemical processes occurring on the surface. I n contrast, the conceptual basis of macroscopic adsorption measurements in flow systems is ludicrously simple; however, there the execution of reliable experiments is not.

I l l . Field Ion Microscopy Although the field emission microscope is capable of yielding detailed information on adsorption phenomena, its resolution is insufficient to reveal the location of individual adatoms. Since the limitation lies in the high tangential velocity of electrons, the (in retrospect) obvious solution t o the problem of visualizing atoms is t o form an image with particles initially a t thermal velocities. This step was taken by Miiller (68) in 1951. I n his ion microscope, as usually operated, the sample is surrounded by helium. An image is then formed by imposing a field sufficient to bring about ionization of the helium at the surface, but not in free space. Ions thus formed are accelerated along the lines of force toward the fluorescent screen, on which are mapped those places on the surface a t which the rate of ionization is highest. The increase in resolution over the field emission microscope attainable by this method is demonstrated in Fig. 5 1. Whereas the field emission image only reveals

348

GERT EHRLICH

variations in intensity with orientation, the ion image locates the more protruding of the individual surface atoms. Despite its amazing resolution, the utility of the ion microscope in probing adsorption is quite uncertain inasmuch as ( 1 ) the conditions under which adsorbed gas atoms and molecules become visible have have not been established; (2) in observing atomic details with helium ions the adsorbed gas may be removed from the surface; (3) even if the

-

-

FIQ. 61. Ion microscopeview of tungsten tip in standard configuration. Surface formed by field evaporation at 15.4 kv; He image at T 20°K. Tip radius 650A.

adlayer is visible the surface may be altered from its normal state under the double influence of adsorption and applied field. I n this section we will explore the use of the ion microscope for studying adsorption phenomena, with special emphasis on the ability t o detect single gas atoms. So far the microscope has been applied principally to the examination of the surface structure of metals and of imperfections in solids. These, as well as other bulk phenomena, will not be touched upon.*

* A thorough review of the ion microscope, its operation and applications, has been given by Mdler (69).

349

MODERN METHODS I N SURFACE KINETICS

A. PHYSICAL PRINCIPLES The physical basis of the low temperature helium ion microscope is straightforward even though all the details of image formation are not yet understood. The sample, in the shape of a fine needle, with a tip radius of a few hundred angstroms, is mounted in a field emission tube. Helium ia admitted at a pressure of -1p at which the mean free path of the ion is just comparable to the tip to screen distance. At higher pressures, ion-atom scattering would blur the image. The screen is then made negative with respect to the emitter, which is cooled by liquid hydrogen. When the applied field at the tip approaches 4 volts/A, helium atoms a t the surface may lose a valence electron to the metal by tunneling. The resulting ion is accelerated toward the fluorescent screen and we thus obtain a photographic record of those spots on the emitter at which ionization occurs most frequently, that is, of the more protruding atomic positions at which field and the helium supply is highest.

-

1.

Field Ionization of Helium

The energy relations that must be obeyed to make field ionization possible are indicated schematicallyin Fig. 52a.In free space, the potential well of an atom placed in a uniform field is distorted symmetrically. At high fields (with helium, F > 4.5 volts/A) the barrier behind which the electrons are trapped (shaded in the illustration) is sufficiently thinned and electrons can tunnel through. The rate of tunneling has been evaluated explicitly for hydrogen atoms (70, 71) and hydrogen molecule ions, H i (72);for the former, the rate constant for field ionization can be written as

k,

= 0.9438

F

x 1Ol6 p exp

[ -0.684

-

F

I

sec-1

.

(41)

The exponential represents the probability of an electron escaping into the field on colliding with the core potential. It has the form familiar from the Fowler-Nordheim relation [Eq. (30)], with the ionization potential replacing the work function as a measure of barrier height. The frequency of collisions with the core potential is given by the preexponential factor. A t a field of 2 volts/A, this amounts to 0.51 x 10l6 sec-1, compared with 0.66 x 10ls sec-1 for the Bohr frequency, yielding a rate constant k p = 2.12 x lo7sec-1. For other atoms the form of the transmission coefficient should remain unchanged, the tunneling probability varying exponentially with the ratio I3I2/F;in the absence of detailed calculations it is convenient to retain the preexponential unchanged as well.

350

QERT EHRLICH

At a surface, these relations are slightly altered. At the very same applied field the overlapping of the Coulomb potential of the ion core and the image potential of the metal indicated in Fig. 52a diminish the barrier t o tunneling. Field ionization a t the surface is therefore possible under conditions a t which ionization in free space is still negligible. As an atom approaches the surface in a fixed applied field the barrier confronting a

fv

e1;

F=4.5V/A

I

/

F: 4.5VIA

1

,He Free Space

Surface

I

I

1

+-2A

D istonce PIC. 52a. Electronic potential for He in a high electric field. Dotted curve at right repeats potcntial in free space. +-work function of mctal; I-ionization potential ; z,distancc from image plane a t which electron level in atom lines up with Fermi level.

valence electron diminishes. The probability of tunneling rises monotonically as x the distance to the metal becomes smaller, until a critical separation x, is achieved. Here, the electron level is lined up with the Fermi level; closer to the surface tunneling can occur only if energy is put into the system. These relations become more transparent on examining the energy change (44 d ) in the system as a whole. I n the absence of an electric field, the potential energy of a helium atom, shown in Fig. 52b as a function of its separation from the surface is

MODERN METHODS IN SURFACE KINETICS

351

essentially a straight line. There is a slight dip, close to the surface, as weak van der Waals forces come into play; this is followed by a sharp rise, dictated by the Pauli exclusion principle. The helium ion, however, experiences a long range interaction with its image in the metal and is therefore subject to a force* e2/4x2, where - e is just the charge on the

i c3 wr (

ZERO FIELD

x,A

OISTANCE FROM IMAGE PLANE FIG. 62b. Potential curves for field ionization of He at a metal with 4 = 4.50 ev. I n an applied field of 4.5 volts/A, isoenergetic switch from atomic (A) to ionic (A+)curve can occur at 2, N 4.3A. a,-polarizability of atom; a+-polarizability of ion. Repulsion of ionic and atomic curves is not indicated in this or following diagrams.

electron. To create the ion in free space, an electron must be removed a t the cost of the ionization potential I. The electron can be accommodated in the metal, thereby lowering the total energy required by the value of the work function 4. The total energy increase in going from an atom to an ion consequently is I - 4.

* Electrostatic units will be used throughout this section for greater clarity, except when specifically indicated. For practical units, the following will prove convenient: Potential-volts; Field-volts/A; Energy-ev; Length-A; Polarizability-AS; Dipole moment-D.

352

GERT EHRLICH

We now impose a field on the metal surface. This lowers the potential energy of the atom by 4 ol,F2. As indicated in Fig. 52b, the polarization by the field changes the energy but little. I n contrast, the energy of the ion is lowered by Fex, in addition to the small polarization term 4 a +F2 . There is one further effect: The field penetrates into the metal, depleting a thin surface layer of electrons. This shifts the image plane, from which distances are measured, and the coordinate x therefore itself becomes a weak function of the field. In a field F these two potential curves would now intersect a t a distance x , from the surface. This is, in general, prevented by the combination between wave functions for the ionic and the atomic state, which leads to a rounding and therefore a separation of the curves. Consider now a helium atom approaching the surface from the right along curve A . A switch to the ionic state can be achieved a t constant energy on reaching the separation x,. A t distances larger than xc, the transformation atom --f ion lowers the total potential energy. This would result in the formation of an electron with excess kinetic energy. Maintaining helium in the atomic state closer to the surface than x, would, however, raise the potential and requires energy input. Ionization could, of course, still be accomplished by raising an electron from one of the occupied levels up to the Fermi energy. The distance xc a t which ionization can occur by tunneling without energy change is now obtained by equating the energy of the atomic and ionic states. At x, and in a field F ,the ionic state* which in free space was I - 4 above the atomic has been lowered by ( F e z 8 a+ P 2 e2/ 4x) The . atomic level, however, is only depressed by 4 a,F2 and therefore the condition on x, becomes

+

+

Fex, = I - + -

e2 -

4xc

+ +Fz(ao- a,).

(42)

With helium the forbidden zone starts 4.25A from a tungsten surface when the field is 4.5 volts/A. However, the probability that a helium atom accelerating toward a tip will be ionized is small, even at fields of 4 volts/A and a t a distance of only 5A . Such an atom traveling toward the surface has a chance of only one in 100 to be ionized in traversing a distance of one angstrom. I n order to understand the formation of the ion image, it is therefore necessary to examine the motion of the helium atom subsequent to colliding with the surface.

* We shall throughout consider only singly ionized atoms. The charge q will therefore not be specified in this or subsequent equations. Similarly, corrections for resonance between the ionic and atomic states will not be explicitly included anywhere.

MODERN METHODS IN SURFACE KINETICS

353

2. Formation of the Ion Image

The sample surface immersed in a dilute helium atmosphere at temperature T , is bombarded by atoms a t a rate p(2mnkT,) -'/' per cm2. When a field F is applied to the point, the potential energy of the helium atoms is lowered by 4 a,F2. In the vicinity of the point they are therefore accelerated perpendicularly to the surface (reaching a kinetic 0.15 ev). The trajectories of atoms that would riot normally energy of collide with the surface are bent toward the tip, which in the presence of a field acts as if it had an effective radius r' given by * N

At fields such that the probability of ionization in space, and therefore also blurring is small, most of the atoms intersecting the sphere of capture penetrate closer than the critical distance x, and rebound. Even without energy exchange the direction of the reflected atom will be partly randomized. The energy of motion perpendicular to the surface is therefore a t least partly converted to motion parallel to the surface. Atoms undergoing this randomization are temporarily trapped by the field gradient and thus must make a number of collisions with the surface. The probability of ionization while traveling parallel to the surface therefore increases; however, the resolution of an image thus -~ formed would be poor. The time of flight of the ion is T = p d m / 2 e V , and the diameter of the scattering disc at the screen ~ T V , , where , v,,, the velocity of the atoms parallel to the surface, is related to the translational energy E through v = d v m . The resolution, 8, which is just this diameter divided by the magnification cplr, is simply

-

(44)

C

assuming that it is the transverse velocity of the ion and not the uncertainty principle that is limiting. For a liquid hydrogen cooled tip of 400A radius, the resolution with He as image gas without energy accom8A. modation would be limited to If the atom on colliding with the lattice very quickly achieves thermal equilibrium, that is, if the average translation energy is lowered to $ kT,then the resoltion becomes

-

) units. 40 *See footnote on p. 366.

practical

J

(45)

354

GERT EHRLICH

On a 400A point, the tangential velocity of helium ions in thermal equilibrium with the metal establishes a lower limit of 1A for the resolving power; the limitation imposed by the uncertainty principle is of the same magnitude. A truly quantitative theory of energy accommodation in lattice collisions is not yet available. However, model calculations have been carried out by McCarroll (73), and are summarized in Fig. 53. These

.01I

I

I

I

I

I I

IIII

10

1

I

,

I 1 1 1 1 1

I

I

,

100 Calk

1

1

1

1

1

1

lo00

I I , I I

0000

("KI

TRANSLATIONAL ENERGY FIG.53. Accommodation coefficient for He and Ne atoms colliding with W at a specific energy en. Estimates for a 1-dimensionalharmonic lattice (73).

suggest that at incident energies in excess of 0.01 ev the accommodation coefficient for He already approaches a constant value of -0.02, independent of the energy of the colliding particle. Sufficient energy is therefore exchanged even in the initial impact to trap most of the atoms in the inhomogeneous field, provided the image gas is a t a low enough gas temperature ( T~ 2 0 ' K ) .Subsequent collisions t'hen can bring about complete thermal equilibrium with the lattice. As long as the kinetic energy of thermalized or rebounding atoms is less than &,F2,they must stay close to the surface, executing short hops into the inhomogeneous field surrounding the emitter. Escape into space can take place in only two ways: (1) I n hopping to heights larger than x, the atom may become ionized; a t low fields, this will only occur rapidly close to x,. The ion is then accelerated toward the screen and is recorded there. (2) The hopping atoms may wander toward the

MODERN METHODS I N SURFACE KINETICS

355

shank of the emitter, into a region in which the potential gradients perpendicular to the surface are small. There they may evaporate. Hopping atoms penetrate into a region of higher potential energy and are opposed by a force a0F dF/dr. The atom is turned back a t the height at which the gain in potential energy just equals the initial kinetic energy. Inasmuch as the field drops off as F , = Po (ro/ry, 1 < x < 2, the distance traversed is proportional to kTro/aoF,2,where F , is the field a t the surface itself. The atomic processes in the ion microscope can be most easily understood by starting with the emitter immersed in the image gas, but without any field applied. As the field is turned on, the rate of supply of gas t o the tip rises.*The concentration of imagegas in the tip region increases for three reasons. The capture cross section of the tip is enlarged in the field, the distance covered by hopping into the field diminishes, and the depth of the trapping well (equal to the polarization energy 4 a0F2) increases. This increase in the surface concentration continues until the field becomes high enough to bring about ionization. This will occur first a t those places on the surface a t which the field is highest. These areas have three advantages: (1) The supply from the gas is highest there, as atoms are accelerated into the region of greatest field inhomogeneity. (2) The loss by diffusion is least. (3)The barrier to tunneling is more transparent. As the field is raised further still, both the hop height and the edge of the ionization zone are brought closer to the surface; the former is affected more strongly, however (varying as l / F 2 compared to 1 / F ) and the fraction of the hopping atoms that are ionized a t these spots therefore diminishes. Other areas with a larger radius of curvature may, a t this point, become effective for ionization, and draw upon the supply of trapped helium atoms. This will necessarily affect the course of ionization elsewhere, Iny lowering the atom concentration a t the surface. At even higher fields, the incoming image gas may undergo ionization before colliding with the surface, a t some distance from x,. At this stage, the ionization process loses its dependence upon the immediate atomic structure, and the image becomes blurred. This qualitative account serves to underline the complexity of the physical processes involved in forming the ion image, and the difficulty of arriving a t a detailed interpretation. Neither the theory of atom scattering a t surfaces, nor the description of field ionization are sufficiently advanced to allow a quantitative account of image intensities. However, such a quantitative evaluation is not really necessary in

* I have profited here from discussions with D. G. Brandon, also Brit. J. Appl. Phys, 1963.

Cambridge University. See

356

OERT EHRLICH

order to derive incisive information on surface processes, and the more qualitative aspect of ion microscopy will therefore be stressed henceforth. 3. Field Evaporation and Desorption The ionization of a helium atom a t the surface in the presence of a high electric field has been outlined in Section 111, A, 1. As appears from Fig. 54, the field desorption of more strongly bound electronegative

XC I

,

,

,

FIG.54. Field desorption of electronegative adatom A ( I - 4 > > 0). V (r)-atomic potential at distance r ; E,f,-activation energy for field desorption. Ionization occurs at refor adatoms sufficiently energetic to reach this point.

atoms, for which I - 4 is also high, occurs in much the same manner ( 4 4 4 . The only real difference is introduced by the greater binding energy. At fields so low that x, is larger than the range of the forces binding the atoms to the surface, the adatom must first be thermally desorbed. Only then can ionization occur, as it does for helium. Under these circumstances the rate of desorption is just proportional to ] , the activation energy E$u equal to the normal exp [ - E ~ D / R T with desorption energy ED.However, ionization can be accomplished without

357

MODERN METHODS IN SURFACE KINETICS

thermal activation, by increasing the field so that the ionic curve is depressed almost to the bottom of the well for the atomic level. An isoenergetic conversion from atomic to the ionic state can now take place a t

Fez

=I

-4 -

e2

4x

+ 4 F 2 (a,, - a,)

-

V

(2)

,

(46)

where V(x)is the atomic potential (a negative quantity) at the point x, Whereas with helium only the ion could interact strongly with the surface, now there is strong binding [ V(x)] for the atomic state as well. This tends to increase the fields required for ionization, thereby counteracting the generally lower values of the ionization potential. The barrier E,, facing the atom can be adjusted arbitrarily, by applying different electric fields. From the rate of desorption as a function of T and F , the atomic potential V can then in principle be deduced experimentally.

x,A FIQ.55. Field desorption of ion. x gives position of ionic ground state at surface relative to vacuum; e ( Pe)”2-Schottky saddle; x<,-equilibrium position of ion. Field desorption at T = 0°K occurs when e (Fe)’’2= E D I - 4 Feze; E D I - 4 = x ( 0 ) ,E D = desorption energy for adatom.

+

+

+

Quite a different situation appears to prevail when the ionization potential of the atom is close to the work function of the surface. The energetics are clarified by first considering a hypothetical system, in which an ion is held on the surface a t an equilibrium distance x,. For such an ion, the attractive part of the potential energy curve in the absence of a field (Fig. 55) can be closely approximated by the potential e2/4x(or, in practical units 3.601~ev) arising from the interaction of the

358

QERT EHRLICH

charge with its image in the metal. The energy of the ground state on the surface relative to that in the vacuum is given by x. When a high field is now superposed, the resulting curve shows a maximum e(Fe)”‘ [equal to 3.79 F1Ieev] below the zero field vacuum level, a t a distance 5, = (3.60/F)””.This is the well known Schottky saddle (53, 58). If the field is raised further still, the saddle can be lowered sufficiently to coincide with x, the ground state level of the ion. Evaporation will then occur, even a t T = O’K, a t a field

x2 (F).

4]F fj?; Pe3

=

(471

As indicated in Eq. (47) the energy level of the adsorbed ion is itself affected by the field. At zero field, x = E D + I - 4, where E D is the desorption energy of the atom. However, the field not only shifts the

1METAL >

ATOM

-

I I

I

I

I

I ,

I

I

I I I t , ,

I , ,

,

I

METAL

I ,

I

,

,

, , , I ,

ATOM

I

, , I ,

I

I

I

I

-.. ,

Fro. 56. Effective potential and energy levels for electron in the proximity of an atom core a t a surface (75). E,.-Fermi level; Ee-electron level in free atom. Coordinate axes based on core show probability density (+') for electron a t the core as a function of energy. Occupied states shaded.

image plane, as above, but also lowers the energy of the ground state relative to its zero field value by Fex,. The condition for field evaporation a t T = 0 [Eq. (47)] therefore becomes

Fe3 = ( E D + I

-

4

+ FexJ2.

(48) The important difference between this and the previous examples lies in the presence of the ion at the very start of the process. Ionization

359

MODERN METHODS I N SURFACE KINETICS

in the field is not necessary. Inasmuch as the potential energy of the ion is known, an explicit relation for the desorption field is possible. This is, however, an entirely hypothetical situation. Even if an ion core were deposited on a surface, it would be screened by the conduction electrons.” A more appropriate way of describing real systems is therefore the following: I n bringing up an atom to the surface, as in Fig. 56, the electron level E , in the atom splits and becomes broader as the separation from the surface diminishes, provided it lines up with the conduction band of the metal. Indeed, the electron is no longer localized

2r-

1

F-

h O

i 5

-2

ZERO FIELD 0

1

2

3

4

5

6

1

1

0

1

2

3

4

5

6

x,A FIG. 57. Field desorption of electropositive adatom (Z - 4 0).Ionization occurs at z,.For evaporation, Schottky saddle a t 2, must still be surmounted ( 5 8 ) .At T = O”K, field evaporation occurs when e(Fe)’’? = E, + Z + F z (a, - u+). N

+

a t the atom, but becomes part of the compound system metal-plusadatom in which the electronic equilibrium is defined by the Fermi energy E,. In the adsorbed state, or close to the surface, the levels of the atom up to the Fermi energy are filled. If the maximum in the atom “band” is above the Fermi level, then the positive core will not be completely compensated and we have a positive surface layer. When the well representing the atom core is displaced slowly to the right, away from the surface, the atomic levels sharpen again into a single state. As long as the barrier between atom core and metal is small, however, electronic equilibrium can be maintained. The probability of finding an electron a t the level E , is therefore given by the Fermi distribution function fP =

(1

+ exp

E , -’ [-E, kT-]) -

* This is based on the ideas of Gurney ( 7 4 ) ;see also Dobrezow ( 7 5 ) .

360

GERT EHRLICH

The chance of finding an ion or an atom can now be altered by applying an electric field. If this raises the level E , sufficiently above the Fermi energy, the probability of occupying this level with an electron will be negligible. An ion may appear, even though on the surface itself a bare ion did not exist as such. Provided this occurs a t a distance x, < x8, the ion must still pass over a Schottky saddle, as shown in Fig. 57. With one exception, the situation is analogous to that for the desorption of a true ion. The polarization energy did not enter in Eq. (48); i t has the same value a t the Schottky saddle as a t the equilibrium position x, and therefore cancels. Now the ground state is atomic, and the effect of the field must therefore be represented as a polarization term t F 2 (a,, - a,). Evaporation a t T = 0 will occur when the field satisfies the condition

+I

Fe3 = [ED

-

+ + 4 F2(a, - .+)I2

(49)

and the desorption energy is given by

EiD

=

ED + I -

+ + iF2 (ao- a,)

- F"2es/2.

(50)

B. HIGH FIELDEFFECTS IN PRACTICE 1. Formation and Stability of the Emitter

Since evaporation can be induced at high fields without thermal activation, a sample can be prepared without thermal disorder of the surface. The fields required for this are estimated most conveniently by assuming that evaporation occurs by depression of the Schottky saddle and resorting to Eq. (49). The magnitude of the field required for evaporation a t T = 0°K for the more interesting metals are listed in Table V. Inasmuch as the field for a helium image is on the order of 4.5 volts/A such materials as Cu, Ni ,and Fe should prove difficult to examine in this way. This expectation has been experimentally verified by Miiller (69). A possible remedy here is to employ an imaging gas with a low ionization potential, requiring lower fields and sacrificing some resolution [as demanded by Eq. (45)]. Neon appears particularly suitable but suffers from the following disadvantages: (1) Its higher polarizability (0.398A3against O.2lA3 for He) may result in considerable adsorption on the surface, when high fields are imposed. (2) The collision cross section for Ne+-Ne is greater than for He+-He (21.7A2compared to 16.7A2). Lower pressures must therefore be maintained yielding smaller ion currents. (3) The efficiency of Ne+ in exciting the phosphor is smaller than for He+. A complete resolution of these difficulties has not yet been

MODERN METHODS IN SURFACE KINETICS

361

accomplished but will probably depend upon the use of image intensifiers rather than high aperture lenses for photography. Field evaporation yields a surface distinct from that attainable by thermal methods. This is evident from the field emission and field ion patterns of tungsten emitters shaped by the two methods (76). TABLE V Field Evaporation at T

W

Ta Ir Nb Mo Pt Au Fe Si Ni Ge

cu U

8.69 8.11 8.88 7.73 6.86 6.86 3.67 4.36 4.60 4.42 3.93 3.63 6.08

7.98 7.7 9.2 6.88 7.13 9.0 9.22 7.90 8.16 7.63 7.88 7.72 4.7

=

0°K

4.60 4.19 6.3 4.01 4.22 6.32 4.70 4.17 4.82 4.74 4.80 4.62 3.47

10.2 9.4 8.1 7.8 6.6 6.3 4.7 4.6 4.4 3.8 3.4 3.2 2.8

The electron field emission photograph of a tungsten point formed by field evaporation a t low temperatures, shown in Pig. 58, is dominated by the extensive nonemitting (110) plane. The emission comes primarily from the higher index planes around the (100)s with a much smaller contribution from the triangular (111)s. The (100) poles, as well as the {211)s, are only ill defined. This is in contrast to the usual thermally shaped emitters, in which the (loo), as well as (211) poles appear as definite and prominent areas of low emission. The ion image of the same point, also displayed in Fig. 58 is typical of a substantially clean surface. Over the regions around the (110) and (100) the regular ordered array of bright dots corresponds to the more protruding tungsten atoms constituting the surface. Only in the region around the (100) poles, and on the periphery of the point, is there apparent a more random disordered structure, probably indicative of impurities. There are some variations in intensity over the ion image, most pronounced in moving from the completely nonemitting (110) toward the outer regions of the point.

362

GERT EHRLICH

This has been rationalized by Muller ( 7 7 ) ,who pointed out that the rate of evaporation is exponentially dependent upon the desorption energy for the ion, as given by Eq. (50). The desorption energy for an atom E , is independent of orientation provided this occurs from a kink site, Only the work function and the field F vary with orientation. The lower the work function, the higher Ei7, and therefore the slower the evaporation. This leaves relatively untouched regions of lower a t the expense of regions having a higher work function, and results in the formation of sharp edges.

+

+

a

-

FIG. 58. Comparison of electron and ion emission photograph of a clean tungsten surface, formed by field evaporation a t T 20°K. (a) Electron image at 5 x lo-' amp and 1.28 kv. (b) He ion image of identical surface a t 14.9 kv. Enlargement 16% greater than in (a). (c) Orientation diagram.

The bright band around the {loo}, formed from planes with low work functions, and therefore with high desorption energy for the ion, indicates regions having a very small radius of curvature. This shape tends to equalize the rates of field desorption a t different atom sites, since the variation of the local curvature and thus of the field, counteracts differences in the energy required for removing metal ions. These variations in curvature are also responsible for the appearance of the electron image in Fig. 58. The (110) and surrounding regions are

MODERN METHODS IN SURFACE KINETICS

FIGS.58b and 58c

363

364

GERT EHRLICH

flattened by field desorption. The nonemitting center of the field emission pattern is larger than usual in the thermally shaped emitter. The latter tends toward a configuration in which the low index planes are flat minimizing the total surface free energy. I n the field formed point, however, the (211)s form part of the sharp ridges. Electron emission is therefore enhanced and these areas become less distinct than in the thermal pattern, in which the (21 1) planes are flat. These characteristics serve as a convenient indication of the shape and also the cleanliness of the surface achieved by field desorption, without involving observation of the ion image. 2. Surface Cleaning

I n order to carry out ion microscopy conveniently, the sample must have a radius of approximately l000A or less. Reasonable voltages, on the order of 10-20 kv, then suffice to establish the fields required for image formation without breakdown problems. This requirement eliminates the possibility of removing impurities by heating, inasmuch as the desorption energies of oxygen, hydrogen, nitrogen and carbon monoxide exceed the activation energy for surface migration even of tungsten. Thermal desorption alone would quickly blunt the emitter into uselessness. The field itself can, however, be used to clean the surface. Field desorption of electronegative impurities follows Eq. (46). When the ionic level is depressed down to the atomic, field ionization occurs and the ion is swept away. The depth of the well binding the atom therefore enters, as well as the ionization energy, and these conspire in many systems to raise the desorption energy above that necessary for image formation. The fields necessary for desorption of impurities are in many instances sufficient to bring about evaporation of the clean emitter surface itself, and this imposes an upper limit. The electric field, however, drops off from the apex of the tip toward the shank. It is clear from Fig. 58 that only a small cap centered on the apex is really cleaned by field desorption. Alternate heating and field desorption must therefore be used in order to establish a surface which does not suffer through contamination by diffusion from the shank. Blunting as a result of field desorption or heating can be counteracted by ion bombardment (44 a ) as carried out in Section 11, B, 2. By operating the tip as a field emitter in an inert gas, such as argon (at p 10-4 mm) very fine clean tips can be regained. Subsequent field evaporation smooths this badly damaged surface, and may restore the emitter for further observations by ion microscopy. N

MODERN METHODS I N SURFACE KINETICS

365

3. Quantitative Determination of Potentials

I n principle the formal theory of field desorption (outlined in Section 111,A, 3) makes it possible to predict the stability of adatoms in a high field, or to deduce from experiment the nature of the potential curve of the adatom. I n order t o carry this promise into practice it is necessary to examine more closely the additional information on high fields and their effects which is required for quantitative work. The field itself is of course the primary parameter that must be known. Experimentally it is the voltage between tip and screen that is determined. However, a fairly accurate value ( & 15%) of the field a few atom spacings above the surface can be derived from the over-all shape of the emitter, by procedures already worked out for field electron microscopy (44 b ) . At a given atom site, a rnicrofield value can then be estimated from the geometry. Thus for a hemispherical knob on a plane, the field is enhanced by a factor of 3, a t a kink site on a crystal surface, the enhancement has been estimated a t 1.8. This procedure yields microfields in fair agreement with those estimated from Eq. (49) without any polarization correction (78).Refinements are still in order, however, when dealing with gas adatoms. As was pointed out in Section 111, A, 1, this field shifts the image plane; because of the high density of conduction electrons in a metal, this shift should be small even a t the fields necessary for ion microscopy. Using classical electrostatics Muller (79)estimated the displacement as

A,

-

Fd/4rn8e;

d denotes the diameter of the outermost layer of the crystal, containing n, atoms per unit area. At a field of 5 volts/A, A , amounts to 0.5A for a (110) surface of tungsten. Following Mott and Watts-Tobin (80),the decay of the field inside the metal (that is for negative values of the displacement z) can be more adequately represented as

F

=

F,exp

b].

For a free electron metal, A , the penetration depth as deduced using the Fermi-Thomas approximation, is

where A is the number of electrons per unit cell, a, the Bohr radius, and d, the lattice spacing. A is again on the order of 0.5A, substantiating the

366

GERT EHRLICH

original estimates. The effect of the field on the electron distribution in the metal can therefore be readily accounted for. The applied field also interacts with the electrons in the adsorbed atoms. All these effects were in Eq. (46) formally lumped into a single polarization term 9 (ao - a+)F2.For the adatom this term represents the lowering of the energy of the adsorbed entity in the field, both through a redistribution of electrons, as well as through the interaction of the field with partial net charges on the atom, With its image in the metal a net charge on an adatom establishes a dipole; the field interacts directly only with the charge on the adatom, which can be approximated knowing the effective dipole moment M , and the bond length. The image charge, however, is affected both by the shift in the electron distribution in the metal as well as by the change in the equilibrium position of the adatom (with its charge) in the field. Even when it is not necessary to consider the net charge on the adatom (either because it is small or in dealing with field desorption of the metal itself) the correction for polarization is difficult. Polarizabilities are known for only a few atomic species existing in the gas phase; experimental values for ions are sparser still. However, both are likely to contribute significantly t o the energy changes in field desorption, as appears from the polarizability data (81)for three typical systems: 0 - a0 = 0.89A3,0 4 - - a+ = 0.49A3; Ne - 0.40, Nef - 0.21; H, - 0.79, H i - 0.38. I n a field of 5 volts/A the energy of an oxygen atom is lowered 0.7 ev; the polarization effect on the ion amounts to 55% of this. Even for free atoms in the gas phase a single polarizability term of the form $aoF2 fails to account for the energy changes that occur in high fields. Thus Sewell (82) has estimated that for atomic hydrogen this change is given (in ev) by -

AE

=

2.315 x

F2

+ 4.38 x

lo-* F 4

+ 2.16

x

F6.

I n a field of 2 voltsiA, the higher terms contribute 4% even for so small an atom. At a surface, representation of the energy changes through a low field polarizability must become poorer still. The redistribution of electrons within the metal confronts an adatom with an array of essentially bare cores, and this should profoundly change the nature of the chemical bond to the adatom. There is one further experimental precaution that must be observed in evaluating experiments-it is important to establish the nature of the entities desorbed in the high field. This is necessary for two reasons: Although throughout we have written all relations for singly charged ions only, multiple ionization is likely in field desorption of materials

MODERN METHODS I N SURFACE KINETICS

367

with a low ionization potential (I 5 ev). For more electronegative adatoms, especially for those strongly bound t o the surface, field desorption may involve removal not only of the adatom, but of a complex consisting of adatom plus metal. This will be touched upon in Section 111) c, 1, a. It is clear that in the present state of knowledge of the behavior of materials in very high fields, the quantitative interpretation of field desorption must be made with care. However, the theory of field desorption is already far enough advanced to give a good qualitative account of the processes likely t o occur in the ion microscope, and this is most important in using the microscope for the direct observation of surface events. N

C. DIRECTOBSERVATION OF ADATOMS The ability of the ion microscope to resolve the details of the atomic arrangement of a surface is well established; this by itself is exceedingly important for the study of adsorption phenomena. However, the full power of the microscope can be brought to bear upon surface studies only if adsorbed gas atoms can be resolved as well. This question has been examined in detail for nitrogen atoms on tungsten ( 7 6 ) . Because of its importance, the evidence showing that direct observation of adatoms is indeed possible will be outlined in some detail. 1. Observations on Nitrogen

a. Adsorption. The changes in a helium ion image of a tungsten surface after adsorption of a small dose of nitrogen at room temperature are shown in Fig. 50. I n the regions which for the clean emitter (Fig. 59a) were well ordered a few additional emitting spots, somewhat more prominent than the atoms of a tungsten substrate, can be detected. Beyond these additions, the lattice itself is hardly disturbed. Only at the rim of the (l2T) is there any indication of atom removal. After a few more doses of nitrogen almost the entire surface is covered by this irregular array of dots and the sense of order is lost, as in Fig. 59c. The appearance of additional spots in the ion image must be assigned to the action of the added nitrogen. Without introducing nitrogen to the system, but otherwise following the same sequence, changes in the tungsten surface are negligible. It is thus established that the effects of adsorption can be successfully viewed in the field ion microscope. This still leaves unspecified the nature of the additional emission centers. These bright spots, a t which field ionization of helium occurs readily, need not necessarily define the location of nitrogen on the surface.

368

GERT EHRLICH

Enhanced ion emission could conceivably arise a t protruding tungsten atoms. Such a change in the atomic arrangement of the substrate might occur in two ways: (1) Nitrogen may bond so strongly to the lattice that the tungsten structure is weakened. I n the field necessary for an ion image a tungsten atom already perturbed by the act of adsorption

-

FIG. 59. Interaction of N, with W at T = 300"K,aa observed in the ion microscope. (a) Clem surface, formed by field evaporation at 20°K. Radius 400A. (b) Tip after first room-temperature addition of nitrogen. Arrow points to one of the new emission centers formed by interaction with nitrogen. (0) Same surface after heavy dosing. All photographs at 20°K.

-

may be pulled out of place. The local field around such an atomic protrusion would be enhanced and a bright spot could therefore become visible a t the site of adsorption. By this mechanism the location of adatoms would be revealed in the microscope, even though only indirectly. (2) The energy released during chemisorption of a molecule of nitrogen (3.7 ev) may be expended in displacing a tungsten atom from its lattice position and onto the crystal surface some distance from the actual adsorption site. The loci of enhanced emission on this model would correspond to the final resting places of displaced tungsten atoms which leave behind them holes in the surface.

MODERN METHODS IN SURFACE KINETICS

FIGS.59b and 59c

369

370

OERT EHRLICH

A decision between these different possibilities can be reached from more detailed examination of the properties of the emission centers. b. Surface Diffusion. It has been established in the field emission microscope ( 4 1 ) that nitrogen adatoms migrate over the regions sur400°K. If the emission rounding the (111) planes of tungsten at T centers in Fig. 59 indicate the presence of nitrogen atoms, they should redistribute a t T 2 400°K; tungsten adatoms on the other hand are known to be completely immobile a t these low temperatures (83) and should therefore remain fixed in place. Such an experiment is shown in Fig. 60 in which the area around the (111)plane bounded by the [1111 and [ 1113 zone lines has been isolated.

-

FIG. 60. Surface diffusion of nitrogen emission centers. (a)View of (111)and surroundings on a clean tungsten surface. (b) After nitrogen adsorption at T = 300°K. Arrow indicates location of one of the new adatoms. (c) Identical tip after heating to T 420°K for 1 min. Emission centers have migrated away from the sites occupied in (b) without disturbing the lattice.

-

On a clean tungsten surface nitrogen is allowed t o adsorb in small quantities. New emission centers appear in the region between the zone 420°K Fig. 60c reveals lines as shown in Fig. 6Ob. After heating to T a drastically altered distribution of these emitters. Every one of the emission centers has changed its position. The centers are therefore

-

MODERN METHODS I N SURFACE KINETICS

FIGS.60b and 60c

371

372

GERT EHRLICH

mobile a t these temperatures and in this respect resemble nitrogen, not tungsten atoms. More than that, the sites a t which changes have taken place after adsorption (Fig. 60b) are vacated in Fig. 60c, again revealing the underlying structure of the lattice. At the original adsorption sites this remains unchanged from the atomic arrangement prevailing prior to adsorption. Indeed the only changes of the lattice in the area under study occur a t

-

FIG. 61. Tungsten point at ! I ' 20°K shadowed by a beam of N,. (a) Clean surface, (b) Same point after exposure to nitrogen stream; beam direction indicated by arrow. Changes in the ion image are, with a few exceptions, concentrated on the side of the tip directly in the beam.

the ( 1 1 1 ) and (334), where no adsorption was apparent. The edges of these planes are slightly rearranged by the movement of one or two atoms. The observations on diffusion therefore indicate that there is no permanent displacement of lattice atoms around the emission centers. It is also dear that the emission centers are not due to tungsten atoms pulled partly out of the lattice. Even if displaced atoms were able to migrate a t low temperatures, they would on moving either leave behind them holes in the surface, or else jump back into the lattice and thus

MODERN METHODS I N SURFACE KINETICS

373

disappear. The latter did not happen. The emitting centers are still visible after diffusion. The former did not take place either. There were no holes left. If, on the other hand, the emitting centers were tungsten adatoms displaced over long distances during the adsorption of nitrogen, there should appear an equivalent number of holes. No such disappearance of lattice atoms can be detected. This mechanism for image changes therefore does not appear appropriate.

FIG.61b

It could still be argued that holes are created in the areas not accessible to direct observations such as the (110) and (211), and other low index planes for which the individual lattice atoms are not clearly defined. We can invoke thermochemical data on the energetics of nitrogen adsorption to eliminate this eventuality. The energy change xu, for the transfer of tungsten atoms out of the (110) to an evaporation site, estimated on the assumption of pairwise additive lattice forces, amounts to 5.5 ev. It is still 3.16 ev for an atom abstracted from the (211). If the adsorption of a molecule of nitrogen as atoms resulted in the displacement of a single tungsten atom the heat of adsorption will be

QERT EHRLICH

374

roughly equal to 2% - Do - xu,.Here x is the binding energy of the nitrogen atom, Do the dissociation energy for the gas phase molecule. I n desorption a t high temperatures, however, the rate of evolution of molecules is limited by the energy barrier 2x - Do. Tungsten atoms displaced during the adsorption may, a t the higher temperatures, return

Fro. 62. Effect of prolonged exposure to N, a t T = 300°K on tungsten tip. (a) Clean, field desorbed surface. (b) Surface after 160 min at nitrogen pressure of 6 X 10-@mm. Photographs a t T 20°K. After 2 X lo4 impingements per atom site, the {110} as well as {130} and {IOO} still appear empty. N

to their equilibrium positions, but this would not affect the desorption of molecular nitrogen. On this model we must expect a difference of xw between the heats of adsorption and desorption, with ,yw as large as 3.16 ev. The desorption energy measured for nitrogen is 3.52 ev. The heat of adsorption should therefore amount to only 0.35 ev. Experi-

MODERN METHODS I N SURFACE KINETICS

376

mentally determined heats of adsorption and desorption (84) are, however, in good agreement (differing by less than 0.2 ev) and this type of process therefore does not appear important. That there is no long range displacement of lattice atoms over the surface can also be demonstrated directly, by shadowing the emitter

FIG. 62b

with a beam of nitrogen (61).On such a surface, shown in Fig. 61, the changes are with a few isolated exceptions limited to the side of the emitter in the direct beam. These various considerations appear to confirm that nitrogen is directly visible. This is further substantiated by focusing upon the adsorption on a single plane-the (110). c. Adsorption on the (110). I n the adsorption sequence of Fig. 58 nitrogen appears only on the highly stepped surface regions. Low index

376

OERT EHRLICH

planes [{l lo}, (211) and {130}] are left entirely bare. If a tungsten point, mm for 150 min, correFig. 62, is exposed to nitrogen a t p == 6 x sponding to 2.1 x 1010 impacts per unit area, there still does not appear any adsorption on the (110) plane. Only the edges of the lattice plane have changed as nitrogen congregates there. I n the electron field emission microscope, direct observations on the (1 lo} and (211) planes were impossible; the high work function shrouds in darkness adsorption processes that might be occurring there. This, however, should not prove a limitation in the ion microscope, but the act of forming the helium ion image may conceivably denude this particular region of adsorbed material, without exerting such a strong effect elsewhere.

FIQ.63. Adsorption of nitrogen on the ( 1 10). (a) Helium ion image of clean W surface at 10.3 kv. (b) After nitrogen adsorption at 20°K. Image at 10 kv. (c) Surface after warming to T = 300°K for -5 min without field applied. Population of ( 1 10) has disappeared.

I n Fig. 63 the area around the central ( 1 10) is now isolated for study. If nitrogen is allowed to interact with this a t 20°K adsorbed material is apparent on the ( 1 10). However, this can again be eliminated by raising the sample to 145" < T < 190°K. I n their thermal behavior emission centers differ sharply from W atoms. W atoms deposited on the (110) from a filament at 3000°K

MODERN METHODS IN SURFACE KINETICS

FIGS.63b and 630

377

378

GERT EHRLICH

remain on this plane even when it is warmed to 300°K for minutes (as in Fig. 64). Also, the (110) can be filled with tungsten both a t 300 and a t 20°K. Field desorption of the nitrogen emission centers also occurs differently from that for tungsten atoms. At 20"K, the former 4.6 volts/A; the latter require a field F 2 5.0 are eliminated a t F volts/A, both in the absence of He. N

-

FIU. 64. Behavior of W adatoms on the (110). (a) View of clean (110) and environs. (b) W atoms deposited onto surface a t T 20°K from direction indicated by arrow. (c) (110) plane after warming to T = 300°K for -10 min. W adatoms on (110) have realigned, but remain on this plane.

It is therefore clear that the emission centers appearing after nitrogen adsorption cannot be displaced W atoms, and that adsorbed nitrogen is directly detectable. For proof that these are atoms rather than molecules we must still rely on the macroscopic flash desorption results, outlined in Section I, B, 3, a. I n any event, these observations establish the direct visibility of adsorbed gases in the ion microscope. It should be noted parenthetically that these studies on the (110) also reveal an unusually pronounced structural specificity in adsorption.

MODERN METHODS IN SURFACE KINETICS

FIGS.64b and 64c

379

OERT EHRLICH

380

d. High-Field Removal of Surface Complexes. Having established that nitrogen adatoms can be made visible in the ion microscope, we can now return to a point raised in Section 111, B, 3-removal of lattice atoms in field desorption of adsorbed material. An atom A firmly enough bonded to the surface to withstand the high fields necessary for a helium ion image may affect the cohesion of the metal surface itself. At fields above those needed for the image, a molecular ion MA+ may therefore desorb rather than just A . +

a

b

C

-

-

FIQ.66. Desorption of tungsten-nitrogen complex at high fields. (a)Clean tungsten. He image at T 20°K, 16.3 kv. (b) After adsorption of nitrogen at T 20°K. (c) Same surface after applied potential has been raised to 19 kv (in the absence of He). W atom in left corner of the plane is removed together with adatom.

With nitrogen, both events occur. A single atom plane of tungsten is shown in Fig. 65 both before and after exposure to small amounts of N, at T 20°K. When the field is raised 17% above that required for the image, the atom adsorbed in Fig. 65b again disappears, and with it disappears the corner W atom over which it rested (61).This removal is limited to certain types of surface sites. The observations of surface diffusion (Fig. 60) proved that a t low fields the tungsten lattice is not N

a

b

C

-

FIQ.66. Field desorption of nitrogen leaving tungsten lattice undisturbed. (a)Clean tungsten; ion image at 16.3 kv. (b) Nitrogen adsorption at T 20°K. (c) Surface after potential hae been raised to 18 kv (inthe absence of He).

permanently disturbed. It is also clear from Fig. 66, for example, that nitrogen can be field deaorbed without removing the underlying metal. The extent to which the substrate is affected in field desorption varies with the nature of the adsorbed gas and each individual system must be explored in using the ion microscope to estimate bond energies.

MODERN METHODS I N SURFACE KINETICS

381

2. Observations on Adsorbed Hydrogen, Oxygen, and Carbon Monoxide

The ability to detect with the ion microscope individual atoms or molecules held on a surface depends upon several parameters that are difficultto predict and that can only be estimated empirically. A high binding energy helps to insure the stability of an adatom, even when its ionization potential is only 10-15 ev. What is of interest of course is the binding energy for the adsorbed entity and not necessarily the heat of adsorption of the gas phase molecule. Thus for nitrogen on tungsten, the heat of adsorption per mole is 85 kcal; for nitrogen atoms, however, the binding energy is 155 kcal/mole. It is the latter value that dictates the stability of the layer at low temperatures, under conditions such that diffusion cannot occur. Stability in the applied field is only one factor, however. The ion image gives a picture of the rate of field ionization of the image gas. To become visible, an adatom must be. able to affect this rate of ionization, either by changing the local field or the shape of the barrier through which the electron tunnels. In order to establish the general usefulness of the ion microscope for adsorption studies and to obtain a better qualitative feeling for these parameters, several systems besides nitrogen have been explored (85). The most interesting of these is hydrogen, whose simple electronic structure should in principle permit a better insight into the process of image formation. Indeed it has been suggested that because of their small size hydrogen atoms may be invisible to helium ions. A photograph of a tungsten surface exposed to hydrogen is shown in Fig.67 .The ion image has remained essentially unchanged after hydrogen adsorption. This does not necessarily indicate that hydrogen atoms are invisible, however. A picture of the surface by electron emission, both before and after taking the ion image, reveals that the hydrogen has been completely removed during the act of observation in the ion microscope. The fact that no change in the image can be detected must therefore be related to the removal of the adsorbed layer, which is sufficient, but perhaps not necessary to establish the absence of an image change. This observation does, however, bring into focus another distinct effect influencing the operation of the microscope (61).The action of the applied field alone on the adlayer can be followed in the electron emission microscope. This is accomplished by applying a suitable positive potential t o the screen for testing the electron emission and then reversing the polarity to impose a higher field in order to bring about desorption. I n measurements of this type, it becomes clear that the hydrogen

GERT EHRLICH

3 82

-

layer is only eliminated a t a field F > 5.4 volts/A compared to F 4.5 volts/A used in taking an ion image. The removal of the hydrogen layer in the ion microscope therefore is not a simple field desorption, but involves the presence of helium. The helium may have three effects: (1) The point is continuously bombarded with neutral particles having a translational kinetic energy of 0.15 ev.

-

-

FIG. 67. He ion image of W surface before (a) and after ( b ) hydrogen adsorption a t T 20°K. Surface appears essentially unchanged, and there is no indication of any additional atoms.

(2) Electrons are showered onto the tip when ionization takes place. These come directly from the helium atoms ionizing close to the edge of the ionization region, from occasional impacts between ions and neutral gas atoms, and from secondary emission a t the screen. (3) He+ formed close t o the surface may promote the removal of H adatoms. With helium present, the hydrogen layer is attacked even when the field is greatly reduced. Preliminary indications are that this effect cuts

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out only when field ionization becomes negligibly small, a t fields F 3 volts/A. Bombardment by neutrals is unlikely to play an important role. The translational energy is only halfed a t the cutoff. Even a t picture voltage the kinetic energy of the helium atoms should be too low to bring about evaporation. Bombardment by high energy secondaries can be eliminated by N

FIG. 67b

interposing a mesh grid negatively biased with respect to the screen. Experiments prove that this suppresses the emission of secondary electrons; despite this the hydrogen layer is removed when a He+ image is formed. Electrons freed in ion-atom collisions can simiIarly be eliminated by operating a t low He pressures (< 10-3 mm) without changing the instability of the adsorbed hydrogen. The effect of He+ close to the surface is difficult to establish directly, However, it has been possible to demonstrate that hydrogen is also stripped when Ne and Ar are used for image formation. With Ar, complete ionization can be brought about some distance from the surface in a field of 4.5 volts/A. Under these conditions the ion should not N

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affect the adsorbed layer directly, yet i t is found that even then the hydrogen is completely eliminated from the surface. This promotion of field desorption must therefore be ascribed to bombardment of the adsorbed hydrogen by electrons (- 106 electrons per surface atom per minute) tunneling out of the image gas. This, a t the moment appears as one of the most important limitations on the

-

FIQ. 68. Observation of the interaction of carbon monoxide with tungsten in the ion 20°K, 16.2 kv. (b)Same surfee after adsorption microscope.( a )He ion image of W at T at T = 300°K. Image at 14.6 kv, to reduce perturbation of adsorbed material.

use of the ion microscope in adsorption studies, and again emphasizes the great care required both with execution and interpretation of the ion microscopy image. The effects of this bombardment are not as severe with gases other than hydrogen, however. It is possible to observe changes in the ion image of a tungsten surface after exposure to either 0, or CO as in Fig. 68. Oxygen is seriously

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perturbed by the act of observation, and images must be obtained at much reduced fields. Here the early onset of field ionization a t protuberances, that is a t the adatoms, is of great help; still most of the adsorbed layer appears to be lost. Adsorbed carbon monoxide is much less disturbed by ion microscopy. I n contrast to nitrogen, CO does appear to populate the (110) planes,

FIG. 68b

even at room temperature. This not only suggests a marked difference in the way surface structure affects different molecules; it also establishes that the appearance of individual emission centers is not a good indication of their chemical constitution. Figure 69 shows an ion photograph of CO on the (110).Although carbon monoxide should be held as a molecule, the appearance of the emission centers does not differ markedly from that of N, or W in Fig. 63 and 64. A distinction between different chemical entities can, however, be made through their stability to

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changes in temperature or applied field. At 20"K, CO for example is eliminated from the (110) a t F 2 5.3 volts/A, W a t P > 5.0 volts/A, and nitrogen a t F 2 4.6 volts/A. Although adsorption is affected in observing an ion image, and experiments must be carried out and interpreted with great care, it appears that in the field ion microscope significant information can be

FIQ.69. CO adsorption on the (110). (a) (110) and surroundings on clean W. He image

-

a t 16 kv. (b) Ion image at 14.6

emitter at T

kv, after 60 min exposure to CO at 6 X lo-' mm, and 300°K. Arrow points to material on the (110).

attained for a variety of adsorbed materials. Furthermore preliminary experiments carried out in this Laboratory indicate that by reducing the image currents by a factor of 103, the side reactions encountered even with hydrogen can be suppressed. Image intensification should therefore in the near future extend the usefulness of this microscope.

D. PRACTICAL IMPLEMENTATION 1. Vacuunz Requirements for Adsorption Studies I n the examination of surface structures, only minimal demands are usually imposed on the vacuum conditions in the ion microscope. The

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field of approximately 4.5 volts/A required for good helium ion images is generally considered sufficient to ionize any reactive gas molecules that may be present, sweeping them toward the screen before they can approach the surface. The tip, itself shaped and cleaned by field desorption, should therefore remain clean, provided it is cooled with liquid hydrogen to prevent migration of impurities along the field free shank.

FIG.69b

At higher temperatures, diffusion may deposit contaminant on the surface. In contrast, the study of adsorption processes requires the utmost cleanliness of the ion microscope. An unequivocal identification of surface changes can be attempted only if all extraneous adsorbable materials are excluded from the tube. Unusual precautions are necessary because : (1) The high field imposed for image formation can no longer be used to maintain surface cleanliness. The rate of autoionization in a field F (in volts/A) is proportional to exp [ - 0.68 13’’/F].For the more common reactive gases, the ionization potentials I are comparable (CH, - 13.1 ev, CO - 14.0, H, - 15.4, N, - 15.6, 0,- 12.2). Effective admission of only one of these gases without contamination by the

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other, solely through the regulation of the field, is therefore not feasible. (2) Information on adsorption processes occurring under the usual thermal conditions can be attained only in the absence of fields. A field as low as 1.3 volts/A (in which the rate of autoionization is negligible) imparts to a molecule with an ionization potential of 15 ev and a polarizability of 1.8A3a kinetic energy of 0.105 ev. This is equivalent to an effective translational temperature perpendicular t o the surface of 1200°K. More than that, examination of the dependence of surface

-G

2

1

V

.

I

S

N2 He

Trap and Pump

I

Trap a..d 'ump

FIG. 70. Ultrahigh vacuum system for field ion microscopy of adsorption phenomena. 1-emitter assembly developed into Dewar; 2-trapping Dewar; 3-shielding Dewar; A-electrode terminals; B-gas inlet; C c o l d finger for liquid hydrogen; G-selective mm); N-liquid nitrogen trap; Pgetter; I-inverted ion gauge (10-1O < p < magnetically operated ground glass port; R-grounding ring; S-Schulz-Phelps gauge (p> mm); T-thermocouple gauge; V,-Granville-Phillips valve (type C); Wwillemite screen.

properties on the temperature of the tip is possible only if the emitter temperature can be changed over an appreciable range without contamination from the shank. Cleanliness can be assured only if the supports are not exposed to a continuous supply of contaminant. (3) The ion image is formed in the presence of helium at approximately l p pressure. The purity of the image forming gas, and therefore of the gas handling system, must be such as to prevent contaniination of the sample, or a t least the sample shank.

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2. Vacuum Techniques These severe requirements have been met in this Laboratory by an ultrahigh vacuum system of the type sketched in Fig. 70. The system consists of two distinct parts: the gas handling train and the microscope proper. These two can be pumped out independently with mercury diffusion pumps, and are separated by an ultrahigh vacuum valve V p ,which controls the admission of gas to the sample. I n the gas train, the pressure either of helium or of the gas under study, can be regulated by adjustment of the valve V, or V, leading to the spectroscopically pure gas supply (Airco) and by valve V, to the pumps. Prior to admission of any gas the microscope and gas train are subjected to the standard ultrahigh vacuum treatment (Section IV, D) reducing the background pressure to p < 2 x 10-lo mm. As an additional precaution, both the gas supply as well as the microscope itself could be cleaned by selective gettering. 3. Sample Preparation, Microscope Operation, and Image Evaluation

The preliminary treatment of the sample for ion microscopy is precisely that customary in work with the electron emission microscope (69). With tungsten, for example, three mil 218 wire is spot welded onto ten mil tungsten supports and briefly heated to 2500°K in a hydrogen atmosphere. The sample is then pointed by electrolytic etching in 3.5N KOH, a t 2 volts ac. Final shaping is carried out in the microscope after the pressure has fallen below mm. Some of the adsorbed material is removed by heating the sample t o 1150°K. The tip is then further cleaned and smoothed by slow evaporation at room temperature in fields up to 4.5 volts/A as described in Section 111,B, 1. The progress of this operation can be followed by periodic reversal of the potential between tip and the screen and observation of the electron emission pattern. If the tip remains intact, i t is cooled by introducing liquid hydrogen into Dewar {jl, with liquid nitrogen in shielding Dewars {j2 and 3. Shaping is completed by field evaporation at 5.6 volts/A. That the vacuum techniques outlined above suffice to maintain a clean surface is easily established by direct observation. A tungsten 50 min without tip has been maintained in the helium image gas for any field applied, and without showing any change after this exposure. Similarly, after appropriate preparation, the point can be warmed to room temperature with less than a dozen atoms being added. I n observing an adsorption sequence, it has been our practice to start N

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with a photograph of the clean emitter cooled with liquid hydrogen. The image gas is then pumped out and a small amount of the material to be studied is introduced. After adsorption has taken place t o the extent desired (as monitored by field emission), helium is readmitted for ion microscopy. When studying adsorption a t temperatures T > 20°K the contamination rate of the tip on warming without gas addition must first be established. Gas is then adsorbed a t the predetermined temperature without coolant in Dewar #/2 (immediately surrounding the sample). With adsorption over, Dewar #/2 is cooled with liquid hydrogen and only thereafter is the tip itself brought to 20°K. Any residual gas not pumped out should settle on Dewar #/2 rather than upon the sample, thereby preventing sperious temperature effects. For photography, Polaroid film is particularly useful. Type 47 film has a speed of 3000, and even with an ordinary fll.9 lens exposure times may be as short as two minutes. Lenses with focal ratios of unity or less have, however, been used in fluoroscopy and television and are available from such companies as Wray (Optical Works) Ltd., Farrand Optical and Old Delft. 35 mm cameras withf/0.95 optics are also on the market (Canon Camera Company), but these use standard film, which is both slower and less convenient. For an immediate indication of changes in an ion image, such as after adsorption, i t is convenient to resort t o color superposition. The ion image after adsorption is projected in red on top of the clean image, in green; additions then appear in red, atoms removed from the surface in green, and the areas that remain unchanged in yellow. However, the fine details of structural changes are more readily apparent by comparing the black and white originals. Considerable care is called for in interpreting such photographs. As pointed out in Section 111, A, 2, the formation of the image of an adatom is a cooperative effect, in which the local environment plays an important role. Thus in Fig. 59 the “size” of the atoms in the vicinity of the nitrogen indicated by the arrow appears to have diminished. Indeed, under some conditions neighboring atoms seem to disappear. This is well illustrated in Fig. 65. After adsorption, the lattice atoms in the vicinity of the adsorbed atom vanish from the ion image. As soon as the protrusion is removed, however, they reappear quite clearly. These image changes do not reflect a real change in surface structure; they indicate rather a depletion of the helium supply by rapid ionization a t a protruberance, which throws the immediate surroundings into a shadow. Such difficulties are of course to be expected in any new technique. However, considering the unique information that can be attained in the ion microscope, its successful operation is really not difficult.

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E. PROSPECTS The use of the ion microscope in adsorption studies is still in its infancy. It is therefore difficultto compare in detail the achievements of this technique with flash desorption or with the standard field emission microscope. At the moment the range of systems that can be fruitfully studied in the ion microscope is circumscribed by the high fields required for image formation, by field desorption of the adsorbed layer, as well as by electron bombardment. However, the tremendous power of the ion microscope for studies of adsorption is clearly apparent from the work already done, and further technical improvements should widen the range of applicability. By ion microscopy, the processes occurring on low index planes can be readily detected, whereas these are usually obscured by the high work function of these planes when observations are attempted by field emission of electrons, The ability of the ion microscope to reveal individual atoms, both in the adsorbed layer as well as in the substrate, opens up entirely new possibilities in the study of surfaces. This increase in resolution should permit a most detailed assessment of the relation between structure and surface processes. The ion microscope can also be expected to prove particularly important in probing atomic collisions a t an interface. When, for example, a diatomic molecule dissociates on a surface, the distribution of adatoms as recorded in the microscope will yield information on the efficiency with which the dissociation fragments give up their kinetic energy to the lattice. Similarly, observation of adsorption occurring on small clusters of metal atoms should establish the dependence of energy transfer upon the size of these agglomerations, as well as possible lattice rearrangements. I n short, provided that in each instance i t is established that the act of observation has not seriously influenced the event, it will be possible to observe atomic distributions arising from a variety of chemical processes directly, instead of having to deduce these by indirect means. Further in the future, but still in the realm of the possible in view of the unexpectedly high stability of a variety of adatoms to high fields, lies the direct examination of the atomic arrangement of complex molecules held on a surface,

IV. Ultrahigh Vacuum Techniques Flash desorption, as well as electron and ion emission microscopy are based on a mastery of ultrahigh vacuum techniques, that is upon the mm. I n ability to establish and maintain pressures well below

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the next few pages we will therefore touch upon some practical considerations" in achieving vacuum conditions suitable for adsorption studies. These are based primarily upon experiences in this Laboratory. Three points should be noted a t the very start: (1) Achievement of ultrahigh vacuum conditions is simple, involving only slight extensions of normal vacuum practice. It can be done routinely, with little training, and depending entirely upon commercially available components. Indeed, ultrahigh vacuum work is in many ways easier and more certain than the manipulation of standard vacuum lines for gas analysis and surface area measurements. (2) Pressures of 10-lo mm are not a t all new. It is clear that work under these conditions (91-93) was already done in the 1930's. New are the gauges to measure these pressures directly, and the general availability of commercial equipment to facilitate the achievement of these low pressures. (3) Surface phenomena of basic importance have been successfully studied without resorting to ultrahigh vacuum techniques, that is, to the routine and equipment described below. Conversely, following this routine has not always provided automatic insurance against errors. VACUUMTECHNIQUESNECESSARY? A. ARE ULTRAHIGH I n studying the interaction of a gas with a solid, two requirements must be met: (1) Both the surface and the gas must be in a known, reproducible state. (2) The environment must be adjusted to prevent the intrusion of impurities. I n a Maxwellian gas of molecular weight m, maintained a t temperature T and pressure p (in mm), 3.5 x loz2p / d % molecules cross unit area of surface every second. For nitrogen a t 298°K and 1 mm, this amounts to 3.84 x lozomolecules/cm2 sec. If we can tolerate 1% contamination a t the end of an experiment lasting 7 seconds, and based on the data in Fig. 3 assume that 1 out of every 4 molecules striking the surface is successful in contaminating it, then the experiment must be completed in a period T

-

lo-' x p-l sec.

Significant results can therefore be insured by working very rapidly: At mm, 100 msec are available for experimentation. If the pressure can be reduced, the experiment can be performed more leisurely.

* The general principles underlying ultrahigh vacuum practice are adequately covered in standard references (87-89), and the latest developments in technique and equipment are recorded in the annual Transactions of the National Vacuum Symposium. Access to the older literature has bcen greatly simplified by NBS Monograph No. 36 (90).

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The speed of contamination of clean metal (94)surfaces was already realized in the early 1930's. The limit set by Eq. (51)is therefore generally accepted and beyond dispute. What has not always been clear are the conditions under which the pressure has been sufficiently lowered. A system at a temperature T having a working volume of 1 liter contains 9.66 x loz1p / T molecules; a t 298°K this means 3.24 x 1013molecules a t a pressure of mm. If as little as 3 cm2 of clean surface can be formed (by evaporation, for example) once this pressure has been established, the adsorption of all these molecules from the gas phase would still only contaminate 1% of the surface. I n such a system recourse to ultrahigh vacuum procedure appears unnecessary. This approach will only be successful, however, if the molecules in the free volume account for all the gas, and there are no further additional supplies of contaminant on the walls of the system in a steady state with the gas phase. If this assumption is not valid, and this steady state is disturbed, contaminant will evaporate and maintain a fairly constant pressure. This may suffice to cover the entire sample immediately. That the glass parts, grease, and mercury of the usual high vacuum line do act as significant sources of gas can be readily established by observing the pressure rise on raising the wall temperature slightly, or when adjusting a stopcock. There is one sure way of carrying out the experiment without danger of contamination. The pressure even before formation of the surface must be reduced to the same Iow value needed for the execution of the entire experiment. If we expect to operate successfully for a time interval measured in minutes, this requires a pressure of 10-lo mm.

B. THE GOLDENRULE The pressure in any vacuum line is dictated by the competition between two processes. (1) The rate of removal of gas by the pumps; (2) The rate a t which gas is introduced into the system, either by evaporation from the interior walls or by permeation through them. To go from the pressure regime designated as high vacuum ( > 10-6mm) to the ultrahigh vacuum range (< 10-smm) it is only necessary to follow the prescription: Minimize gas evolution, maximize pumping.

C.

VACUUMSYSTEMS

The implementation of this golden rule dictates the form assumed by an ultrahigh vacuum system. To maximize pumping speed, duct work is kept a t a minimum and the diameter of the connections is maximized

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consistent with the available space. Gas evolution can be suppressed in two ways: (1) By heating the system above the normal operating temperature while under vacuum, thereby reducing the concentration of adsorbed material; (2) By cooling the walls to lower the rate of evaporation. In either event the system must be designed to sustain large temperature changes and must be compact enough to fit into ovens or refrigerators. The components necessary for a laboratory sized ultrahigh vacuum system can be conveniently considered under three headings: 1. Pumps

a. True Pumps. I n these the dilute gas is subjected to multiple collisions, directing the gas molecules to a high concentration chamber, which is evacuated by an ordinary rotary pump. Directed collisions can be attained by allowing the gas to interact with a vapor jet, as in diffusion pumps, or with a rapidly moving surface, as in the turbomolecular pump. (1) D{fusion pumps. Both oil and mercury can be successfully employed as a working medium in the ultrahigh vacuum range. Both have to be trapped to prevent the escape of vapor into the high-vacuum part. With mercury, ordinary liquid nitrogen cooled traps are sufficient. With oil, trapping can also be accomplished by interposing high surface area solids (95) to adsorb the vapors of the working medium and its cracking products. Oil diffusion pumps have in the past been favored because of the high pumping speeds. This, however, has been achieved in the absence of any traps. Once oil diffusion pumps are trapped their performance is comparable to mercury pumps. They have the advantage of not requiring a continuous supply of refrigerant, since zeolite traps can operate a t room temperature. However, any oil vapor that penetrates into the ultrahigh vacuum chamber causes extremely serious contamination problems. Penetration of mercury into the system is completely circumvented by interposing a nitrogen cooled surface and automatic filling devices eliminate the need to attend traps. Even if these fail, the damage is confined to those materials that amalgamate readily-gold and silver. For general cleanliness mercury pumps appear preferable and have been used exclusively in this Laboratory. Their particular design appears of little importance, except insofar as it affects pumping speed. We have attained pressures of the order of 10-10 mm with a single dual or triple stage pump, of the type in Fig. 71, as well as with two single stage pumps in series. Diffusion pumps, in general, are ideally suited for adsorption studies. They can operate up to high pressures p (- lo-' mm)

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FIG.71. Triple stage mercury diffusion pump for ultrahigh vacuum system with 42 mm diameter pumping conduits. To prevent excessive mercury losses, cooling of both high and low vacuum conduits is desirable.

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yet do not become saturated with gas. At the worst the trap may become clogged with condensate; this can be remedied by warming up. (2) The turbo-molecular pump. This device has a high, constant pumping speed over a wide pressure range ( 10-2-10-9 mm) and appears very suitable for evacuation of large systems. A model for small scale installation is not yet available. b. Getter Devices. Instead of expelling the gas from the system, it can be taken up inside it through ad- or absorption on a suitable getter (96). One of the most widely used of these is titanium. Although its capacity is lower, molybdenum is easier to clean and as generally useful. New getter surface may be formed by thermal evaporation (preferably under pulse heating) or else by sputtering; there the chemical effects are further assisted by excitation and dissociation in a gas discharge and by enveloping physically adsorbed gases by sputtered metal. Such getter- or sputter-ion pumps (97) are available in a variety of models; they do not require trapping, are compact, and appear well suited for the continuous maintenance of a system at low pressures. When gases are to be handled, these devices have to be used with circumspection. Getter-ion pumps do saturate and the possibility of reeniitting small amounts of unwanted gases on changing the nature of the gas introduced by a chemical reaction must be kept in mind. As an alternative to chemical interaction with the surface, or burial in a stream of material, gases can also be condensed by bringing them in contact with a cold finger. Inasmuch as nitrogen, oxygen, carbon monoxide, carbon dioxide, and water must all be removed, this requires at least liquid hydrogen and preferably liquid helium. This mode of achieving low pressures is therefore not yet suitable for continuous operation. As a convenient auxiliary pump on an ultrahigh vacuum line, hydrogen and helium cooled surfaces are extremely convenient. Moreover, if clean conditions must be sustained for short periods only, complete immersion in liquid helium can quickly transform a routine vacuum of mm into an excellent ultrahigh vacuum of 10-lo mm and better. I n such a system, interaction of an electron beam with the walls may of course liberate contaminants. 2. Vacuum Envelopes

a. Materials and Assembly. The construction of ultrahigh vacuum systems depends upon the same materials (98) and follows the same considerations applied to standard vacuum work. The only exceptions to standard practice are imposed by the need to change the temperature of the system as a whole. These are essentially twofold: (1) The expansion coefficient of different construction materials must be matched

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and the dimensional changes of the system as a whole must be taken into account; (2) Materials of construction must be inert and have a negligible vapor pressure at the highest bakeout temperatures. This eliminates use of organic materials such as rubber, grease and plastics. Systems are usually built of either boro-silicate glass (Corning 7740) or stainless steel. Glass parts are joined by the usual methods. Where permeation by helium must be avoided, lead or alumino-silicate glass (Corning 0120, 1720) may be used. Standard graded seals make it possible to introduce metal parts into the glass system. Metal systems are assembled by heliarc welding, or more easily by sticking parts together with commercial shear seals. The advantage of glass as an envelope material lies in its transparency to radiation, its ease of manipulation, dielectric strength, relative chemical inertness and simplicity of leak checking. Metal systems, on the other hand, are less fragile and, with the advent of shear seals, more readily demountable. For laboratory systems of five liters volume and less, glass is usually preferable, and only such systems will be considered. b. Practical Design Coneiderations. The conductance of cylindrical tubing in Knudsen flow is directly proportional to the cube of the diameter d , and inversely proportional to its length L. The rate at which gas is evolved is proportional to the surface area, that is, to d x L. To minimize the pressure drop from the pump to the end of the system, it is therefore desirable to keep the ratio d / L as large as compatible with the space available. A reasonable compromise in our systems has been to use tubing at least 14 in. 0.d. for main conduits. To withstand the rigors of baking, the system is mounted on a base of insulating material, such as Transite or Marinite. Connections to the base are kept at a minimum (preferably 1) and the tubing is allowed to remain essentially self-supporting, Supporting rods and clamps, just like the vacuum envelope itself, must be able to withstand bakeout temperatures. Cast-alloy ware must therefore be rejected in favor of brass, iron, and stainless steel. The latter is particularly desirable since it deteriorates least through oxidation. Iron, brass and other metals readily attacked may, however, be protected by painting with high temperature aluminum paint (G.E. Glyptal q'i86018). Where a rigid connection to the base is made, bellows are useful for relieving stress without increasing the length of connecting tubing. A typical ultrahigh vacuum system incorporating these features is shown in Fig. 72. The system is evacuated by a three stage mercury diffusion pump, backed up by a rotary pump. This is attached to the glass tubing through flexible stainless steel bellows. To minimize the

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Fro. 72. High conductance ultrahigh vacuum system. (a) Working area (front): P-magnetically operated ground glass port; I-inverted ionization gauge; F-field emission microscope; V-ultrahigh vacuum valve with hydraulic driver (G-P 2-12-11); S-selective getter; G-glass splinter trap; M-Marinite mounting board.

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FIG. 72. (b) Pumping and service area (back): B-gas bottles; H-heating tape for outgassing; N-liquid nitrogen trap (topped by silicone rubber ring) D-triple stage mercury diffusion pump; T-thermocouple gauge (for fore-pressure); C-conduit t o oil trap and rotary pump.

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possibility of contamination by oil, the fore pump is isolated from the diffusion pump by an intervening trap, cooled with liquid nitrogen as long as the trap is in the system. Main ducting is of 45 mm tubing. Equipment is mounted on an ordinary vacuum rack; the working parts are attached to a vertical Marinite board, to which a bakeout oven can

FIG.72. (c) General view of vacuum system and ancillary equipment.

be brought up. The working space, which is limited by the oven dimensions, is 2 ft x 2 ft, and equipment can be mounted up to 14 ft away from the board. This vertical arrangement, which closely follows vacuum practice

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standard among chemists, has several advantages over the more usual vacuum table. On a rack, the pumps, gas reservoirs and ducts can be mounted on the back, where they are out of the way yet easily serviced and repaired. The working parts of the system can also be more readily thermostated when this is necessary, and photographic equipment, particularly for ion microscopy, is more easily accommodated. 3. Bakeout Ovens. Size and shape of ovens varies with the individual system. We have found a vertically mounted oven on coasters very convenient in combination with vertical racks (shown in Fig. 72c). As many as five systems have been serviced by one such oven, which does not encroach on the space around the system. With Fiberfrax bulk fiber insulation between aluminum walls (to minimize weight), 44 kw suffice for reliable operation a t 460°C. Better insulation may unnecessarily prolong cooling times. Heat is delivered by high temperature calrod units, G.E. 48419-G2, running on 208 volts and regulated by a programmed temperature controller. For intermittent use, ovens can also be quickly made up from Fiberfrax paper sandwiched between aluminum foil (99). Trap ovens are constructed of anodized aluminum, or even quartz, wound with nichrome wire or commerical heating tapes. With mercury pumps it is important to avoid cool spots anywhere except at the pump, as otherwise mercury will collect there. Glass tubing outside the main oven is therefore baked with tapes. When working with pyrolyzable gases, such as CO or hydrocarbons, the gas container itself must remain cool. Accumulation of mercury can be prevented by sealing two break-offs in series, and evacuating the space between them (9). The glassware can then be safely baked to the outer seal, without difficulty either from mercury or decomposition.

D. SYSTEM OPERATION 1. Leak Checking

Just as in ordinary vacuum technology the primary instrument for leak detection in a glass ultrahigh vacuum system is the Tesla coil. Care is necessary in sparking graded seals. I n the vicinity of field emission sources, the induction coil should not be employed since the emitter may be destroyed by arcing. mm and below, Small holes, affecting the pressure in the range of can be picked up by observing the pressure change on applying organic liquids such as acetone, carbon tetrachloride, or else by spraying with helium or hydrogen. I n leak checking, an ion gauge with unstabilized

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electron current is useful since the emission current, as well as the ion current, are changed when the test vapor enters the system. Hydrogen is very effective in pin-pointing leaks and does not contaminate the interior nor plug leaks temporarily, as will the hydrocarbons. In the pressure range 10-7-10-8 mm, leak checking with He or H, is as effective and simple as a t higher pressures. Detection of any mercury vapor is particularly easy after bakeout. Cooling with liquid nitrogen reduces the mercury pressure permanently ( p < mm); with contaminant gases, the pressure drops only temporarily, and rises as soon as the cooled surface is saturated. Very small pinholes, which do not prevent the achievement of pressures less than mm are, however, difficult to locate. Isolating parts of the system from the pumps t o determine the pressure rise, and patience are the only remedies. 2. Bakeout

mm and established as Once a system has been brought to p < tight, it must be processed to suppress gas evolution. Baking is really only a n extension of the torching procedure customary in ordinary vacuum work. We have favored the following procedure: While the traps are warming to room temperature an oven is placed over the system, and heating is started. Parts outside the oven are at the same time heated with tapes. Ovens are placed around the traps and the temperature of the entire system, starting a t the throat of the diffusion pumps, is raised slowly. The oven temperature is then increased to 450°C; the 300°C. A temperature rise of 3"/min is suitable, outer parts to at least even for new, complex systems. I n a well-constructed system, properly mounted, baking actually relieves residual strains in the glass parts. This is, however, the most dangerous period in ultrahigh vacuum work, since any exposure to the atmosphere while a t 450°C will result in severe oxidation and possibly destruction of metal parts in the system. When a system does crack during baking, it is usually because of overly rigid mounting or a failure in temperature control. Even then it is possible t o avoid extensive damage by immediately flushing the interior of the oven with an inert gas (N, or Ar) and lowering the temperature. This danger must not be overemphasized. Lines used for field emission in this Laboratory receive a t least 60 bakeouts a year, yet we have never suffered a single breakdown of this type. After a t least six hours a t temperature, the trap ovens are turned down, and cooling with liquid nitrogen is started one hour later. At that 4 hr. point the oven is allowed to cool slowly, reaching 100°C after The oven is then removed, and heating tapes are turned off. If the presN

N

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sure is below 5 x mm outgassing of metal parts is begun immediately with the glass envelope still hot. Gauges are heated either by electron bombardment or with an induction coil. Electron bombardment is convenient, since it does not require placing heavy equipment around the system. With a heavily contaminated system, electron bombardment is not too effective, however;

500W, 500 kc OSCILLATOR

I

1

FIQ. 73. Radio frequency supply for outgassing ionization gauges. Capacitance in microfarads ( p f ) when not otherwise indicated.

rf heating may therefore be preferable when first processing a system, and when the possible formation of metal films on the gauge walls must be suppressed. A circuit for a low-power rf oscillator suitable for outgassing both the inverted ionization gauge and the high pressure Schulz gauge is shown in Fig. 73. Outgassing is usuaIIy continued for at least one hour, preferably in several stages. The baking cycle is then repeated. With a tight system it

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is possible to achieve the 10-lo mm range after a single bakeout, by continuous ion pumping, Despite this, the walls of the system are still contaminated. I n placing a new system into operation, we have found at least three cycles of heating and outgassing desirable. The total time for bakeout can be reduced, however, if metal parts are outgaased (by hydrogen or vacuum firing) prior to incorporation into the system. A line that has been operating under vacuum will become cleaner with time. Once a system has been baked thoroughly it is easier to achieve a good vacuum, even after it has been let down to atmospheric pressure. When this becomes necessary, it is helpful to fill the line with a relatively inert gas, such as nitrogen, and to maintain the cooling of all traps. Ultrahigh vacuum lines are somewhat like wine: they improve with age (upto a point).

E. PRESSURE MEASUREMENTS 1. Ionization Gauges

a. Operation. The most important single event contributing to the rise of ulti-ahigh vacuum technology was the development of the inverted ionization gauge by Bayard and Alpert (100).With this device pressures can now be measured over the range from 10-10-10-5 mm. The inverted ionization gauge shown in Fig. 74 consists of four parts: (1) A thermionic emitter, the electrons from which are accelerated through a potential of 150 volts toward (2) a cylindrical anode. Ions created within the volume of the anode structure by collision with energetic electrons are collected by (3) a fine wire probe. The whole is enclosed in (4) a glass envelope. This gauge differs from the form of the ionization gauge familiar for half a century only in inverting the arrangement of ion probe and electron source, and in changing the ion collector from a cylinder enclosing all the electrodes to a fine wire. This configuration lowers the ejection of electrons from the ion collector by soft x-rays emitted by the anode under electron bombardment; in the standard lo-@amps, equivalent gauge this effect simulated an ion current of to a pressure limit of 10-8 mm. The ionization gauge is not an absolute instrument and does not measure pressure, but rather gas density. It is, however, customary to relate the ratio of the ion current i p over the electron current i, to the pressure through

-

The sensitivity g p is determined empirically, by calibration against an absolute manometer such MI a McLeod. Sensitivities for the ionization

MODERN METHODS IN SURFACE KINETICS

406

gauge WL5966* are listed in Table VI. These values are appropriate only to the temperature of calibration (300°K). The sensitivity is temperature independent if expressed in terms of gas density. The

FIG. 74. Bayard-Alpertinverted ionization gauge 6966, wodi5ed for 1+ in, OD etrdght through pumping leads (0).1-Thermionic emitter; 2-Hlectron eolleotor; 3-Ion collector.

pressure sensitivity should vary as 1/T, but this has so far not been established experimentally. For careful adsorption studies, the gauge should in any event be maintained in an oil thermostat.

* Values for the Veeco gauge RG76 have been given by McGowan (IOla).However, them were taken with the gauge in a bistable condition, and refer only to the lower sensitivity mode.

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Ion gauge sensitivity varies with the exact geometry of the electrodes. The calibration for the two filaments of the same gauge may differ by 10%. On the other hand, the relative sensitivities for different gases appear independent of gauge structure; the calibration need therefore be established for one gas only, Standardization is not a t all simple (21, 104). The ratio ip/ie can be taken as linear in pressure only up to p 5 x mm; a high sensitivity McLeod gauge is therefore necessary for successful calibration.

-

TABLE VI Ionization Gauge Senaitivitiea (9,)

Gas

9, (mm1-l Schulz (101b) G E ( I 4 , 2 1 , 6 1 ) 12

b

(102)

1

1

1

4.4 11

0.42

0.43 1.1

0.47

0.21 0.33 1.6 28

= g,

GE

10

2.6 4 18

r

r, other

r , 6966"

Schulz

(103) 1 0.81 0.63 1.1 1.4

0.16 0.24

2.6

1.2 1.9 2.7 3.4

(gae)/g,(nitrogen).

The operation of the Bayard-Alpert gauge becomes nonlinear at high pressures for three reasons: (1) The free path of the ionizing electrons and therefore the sensitivity becomes pressure dependent; (2) At high pressures the efficiency with which positive ions are collected by the probe is altered; (3) Secondary electrons formed during the act of ionization contribute significantly to the electron current. These difficulties can be avoided by the design shown in Fig. 75, in which the electron path is kept small and efficiency of collection high. This gauge, developed by Schulz and Phelps (105), is operable over a range from 1 0 - e l mm. The collector plates in this version are difficult to outgas, but have the advantage of shielding the cathode filament from the glass walls. The potentials required for the operation of these gauges (BayardAlpert gauge: anode +170 volts, cathode $25 volts, collector -0.5 volts; Schulz-Phelps gauge: anode + l 2 0 volts, cathode +60 volts,

MODERN METHODS IN SURFACE KINETICS

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FIG.76. High pressure (10-6 < p < 1 mm) ionization gauge (105). E-Electron collector; I-Ion collector; F-Tungsten filament cathode. Sensitivities gP (mm)-l: N,, 0.6; CO, 0.62; H,, 0.32; He, 0.09.

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collector -0 volts) can be supplied by standard electronic circuits, many of which are commercially available, or simply by batteries. Battery operation actually has an advantage: It eliminates the 120 cycle ripple that appears in the electron current when the cathode is heated with ac (106). The only expensive instrument that is really required for proper operation is the ion current meter. Observation of pressure changes occurring over a time interval of 1 sec or less, at or below p mm demands a separate detector such ES, for example, a Keithley 603 or 415. It is the ion current detector and the associated leads, rather than the gauge itself, that limit the response time. Millisecond transients can be readily followed on an oscilloscope. b. Limitations of the Bayard-Alpert Gauge. Its simplicity, ease of operation, sensitivity and speed of response make the inverted ionization gauge a truly powerful instrument. It is not without its difficulties, however. Electrical disturbances, such as Barkhausen oscillations and variable wall potentials, do occasionally vitiate measurements. Once recognized, these are readily eliminated.* More serious are the limitations inherent in the mode of operation of the ion gauge. (1) Gauge pumping. The ionization gauge indicates pressure by ionizing a fraction bf the gas in the vacuum system; these ions are trapped and, at least temporarily, lost. The pressure can therefore be seriously affected during such a measurement, More than that, gases can be driven out of the envelope in continuous operation and impurities may therefore be introduced. The ionizing electrons are drawn from a thermionic emitter. With easily dissociable gases, and in practice this means all except nitrogen, carbon monoxide, and the rare gases, a chemical reaction such as dissociation will occur a t the hot filament even when it is not emitting. The dissociation fragments are free radicals and are particularly reactive. Ionic pumping has been recognized as a serious problem, even in conventional ion gauge operation (108).It is easily suppressed, however. The amount of gas removed during any interval of time is proportional to the ion current (provided there is no re-evolution from the walls) and this can be minimized either by lowering the electron current or else by operating the gauge intermittently. The ionic pumping speed of a clean gauge is on the order of 0.1 liters/sec a t 10 ma emission. Operation a t 10 pa emission cuts this by a factor of 1000, yet the ion current at a pressure of 10-lo mm will still be -10-14 amps. This has the additional benefit of eliminating Barkhausen oscillations. The limitation, in either continuous operation a t low electron current or on intermittent work at higher currents, lies in the response time of the ion N

* For discussions of these difficulties and their remedies, see refs. 21 and 107.

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current detector, At low current levels (10-13-10-16amps), the detector may not be able to keep up either with transients in pressure or in the electron current. One possible solution to this problem has been worked out by Lafferty (109))who focussed the ions from a hot-cathode magnetron gauge onto an ion multiplier system. Even for an emission current of only lo-' amp, amp at 4 x 10-l6 mm. the multiplier current was estimated a t 2 x Chemical interactions at the hot electron source present a more severe limitation. It has long been appreciated that dissociable gases are rapidly removed by interacting with the hot cathode (108).What was not clearly realized, until recently, is the possibility of chemical changes and the creation of new species (20). For example, in an operating ionization gauge, both oxygen (110)and hydrogen (111)have been found to form carbon monoxide and other carbon compounds by interacting with carbon containing metal parts and the glass walls. These reactions are thermally initiated and can therefore be suppressed only by lowering the emitter temperature. Taking the dissociation of a diatomic gas as an example (36) the temperature dependence of the rate of atom formation will vary as exp [- Do/2kT] where Do is the dissociation energy. The rate of electron emission, however, is proportional to exp [ - +/kT]. The values of Do for common gases range around 5 ev, comparable to the work function of most clean refractory metals. Inasmuch as D0/2 enters into the rate expression, dissociation can only be suppressed by resorting to a thermionic emitter with a work function of -2.5 ev or less. LaB, supported on rhenium has been successfully used for measuring hydrogen, causing only little dissociation. However, materials such as oxygen, that can be strongly bound on the emitter, may raise the work function of the emitter, necessitating high temperatures and therefore again leading to dissociation. More than that, LaB, is difficult to apply uniformly to the cathode, it is gassy and does evaporate. Thoriated tungsten and iridium are superior in these respects, even though emission is lower, An all-purpose low temperature cathode is still to be realized, however. Another approach to this problem is to allow dissociation to occur, but to provide a mechanism for recombination. With hydrogen this can be accomplished by surrounding the emitter with a clean metal surface. A gauge of the type built by Klopfer (112) should avert such difficulties. This remedy is likely to be less successful with oxygen and other gases for which recombination is not as efficient as for hydrogen. In some circumstances the chemical effects of the thermionic emitter can be minimized by operating in a flow system. A rapid flow of gas N

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through the gauge can then be used to compensate for a high pumping speed out of the system which serves to remove the reaction products. So far, no general solution to the chemical interaction at the hot gauge filament has been found, and this problem must be examined on its own merits for every system. The metal parts of the gauge, as well as the glass walls will, in operation, acquire a gas layer. Just after cleaning, the gauge is an effective getter; some of this gas may be slowly re-evolved in continuous operation. When the filament temperature is raised under electron bombardment, or when the nature of the gas under study is changed, such evolution may become troublesome. I n special systems adsorption on the electrode structures could be eliminated entirely by making or coating these with an inactive material. Thus, for hydrogen, a gold coating would be inert, and for nitrogen, the platinum metals. At the very least it is desirable to establish a constant environment by saturating the gauge with the gas under study prior to any serious work. (2) Gauge mounting. The side reactions just outlined make it appear desirable to operate the gauge a t the lowest possible electron current compatible with reasonable detection of the ion signal. There is one further effect influenced by the pumping action of the gauge: The validity of the pressure reading as an indication of the gas density at the point of interest. I n its usual commercial form the inverted ionization gauge (WL5966) is supplied with a one-half inch diameter connection to the vacuum system. This constriction isolates the gauge, and the gas density in the system may not correspond to that measured. Two situations can be profitably distinguished. ( a ) Pumping by the gauge. The probability s, that a Maxwellian gas will pass through a tube of cross-section A,, length L and diameter d is just 4 d / 3 L . The pumping speed of the gauge, S,, can be similarly expressed in terms of an effective capture cross section AGs,, such that*

N,, the gas density in the gauge, is then related to that in the system, N,, through

* is equal to an effective sticking coefficient only if there is no re-emission, and a separate symbol is therefore used.

MODERN METHODS IN SURFACE KINETICS

41 1

The cross section A,s, in cm2is roughly equal to S,jl2 when the pumping speed is given in liters/sec. Operating at 10 ma emission, a BayardAlpert gauge connected by a 10 cm length of 1 cm i.d. tubing will therefore indicate a pressure 8% lower than in the system. As already suggested this discrepancy can be reduced by operating at reduced emission. ( b ) Pumping by the system. Operation a t reduced emission intensity will not circumvent the difficulties introduced by narrow connections when monitoring the pressure in a system in which gas is being rapidly pumped, either by a clean surface or a wide aperture. The density in the gauge and in the system are then given by

-

2 cm2 a t which I n following the pressure a t a metal surface of area A p the probability of adsorption sF is 0.5, the gauge will read ten times the actual value if connected by a 10 cm length of 1 cm diameter tubing. The only recourse here is to enlarge the pumping leads as much as possible. If, for example, the tubing is made both shorter and wider, so that Assc 8 cm2 (corresponding to L = 13, d = 4 cm) the density difference diminishes to less than 15%. Inverted ionization gauges without glass bodies are now also available as standard items for insertion directly a t the point of interest. With these, it must be realized that the calibration is different from the gauge in the glass body and may depend on its immediate environment. I n operation the glass walls of the ion gauge normally acquire a potential close to that of the cathode, which strongly affects the electron free path. Indeed, by making the walls conducting, or introducing a separate grid surrounding the normal electrode structure, it is possible to vary the gauge sensitivity appreciably and to control the various electrical disturbances to normal gauge operation (21,113). N

2. Miscellaneous Gauges Almost all pressure gauges customary in vacuum work can be incorporated in an ultrahigh vacuum system, provided they have been adapted to withstand the rigors of baking. Although they may lack the versatility and simplicity of the ionization gauge they are important for specialized measurements. Commercially available types are therefore briefly listed. a. Cold Cathode Gauge. The Penning discharge gauge has been adapted for measurements in the range 10-la mm by Redhead (114). I n this device ionization is initiated by a discharge, which is sustained by extending the electron path through spiraling by a magnetic field.

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The need for a separate hot electron emitter is therefore eliminated. The electronic pumping speed of this gauge is high (1 liter/sec for N,) and cannot be adjusted. Its use in gas-filled systems is therefore limited, especially since this gauge cannot be readily keyed. The high rate of clean-up does make it quite suitable as a pump, however. b. Thermal Conductivity Gauges. Both Pirani and thermocouple gauges can be advantageously incorporated in an ultrahigh vacuum system. Their primary advantage is inertness. Operation does not depend upon a hot electron source, and therefore does not disturb the system by ionization and high temperature fragmentation. They are, in turn, but little affected by accidental operation at high pressures. For measuring reactive gases these devices are incomparably better than the ionization gauge. Some chemical reactions may still be catalyzed by the warm filaments-this must be established for each system. The useful pressure range of the standard designs is small (10-3-10-1 mm), however, and the response slow, on the order of tens of seconds. By special attention to thermostating, the Pirani gauge has been used successfully at lo-' mm (115). Similarly the response time of thermocouple gauges can be reduced below 0.1 sec by miniaturization of parts. Even in their standard form these gauges are invaluable for establishing that the pressure in a vacuum system is sufficiently low to allow safe operation of the ionization gauge. c. Knud5en Manometer. With all the previous gauges the response is sensitively dependent upon the nature of the gas. With the radiometer gauge this dependence is markedly reduced and pressures in the range 10-7-10-2 mm are readily measured. Designs suitable for incorporation in a bakeable system have been described by Onchi (116); a model in a glass envelope is also available commercially. Although this is not an absolute instrument, the response depending upon the thermal accommodation coefficient of the particular gas, this gauge operates essentially without disturbing the system under study. It is, however, much more sensitive to vibration and shock than other devices and requires considerably more time to put' into operation. d. Diaphragm Gauges. The usual liquid manometers for pressure measurements in the range of -1 mm cannot be directly incorporated into an ultrahigh vacuum line. These can, however, be coupled into a system through thin metal diaphragms, the position of which is monitored through capacitance changes, and is sensitive to pressure differmm. Such diaphragms serve not only as nullences as low as 5 x indicators, but after calibration, they can also be used to indicate pressure differences directly. e. Ma55 Spectrometers. The ion current measured in the ionization

MODERN METHODS IN SURFACE KINETICS

413

gauge is proportional to the pressure only if the gas composition is unvarying. In practice, even when only a single gas is introduced into a vacuum, residual contaminants are always present. Since the current is proportional to C N p i , where Ni is the concentration of gas i and K~ its ionization cross-section, even small amounts of impurities of high cross-section may contribute significantly. Ions of different mass can, however, be resolved by introducing an analyzer section. There is a considerable literature dealing with designs useful in ultrahigh vacuum investigations (90).Two general types of spectrometers are available commercially. (1) Magnetic dejection. As in standard mass analyzers, a beam of ions uniformly accelerated to a potential V is forced to describe a circular trajectory of radius p by imposing a magnetic field H (gauss) at right angles to the beam. The radius varies with the mass M (in atomic mass units) and charge q of the ion as

- 1H 43.6Jy

cm;

only ions with the proper radius pass through defining slits to be counted. Instruments are now available with good resolution (117)up to mass 150 and operable down to pressures of 10-13 mm. (2) Resonance spectrometers ( a ) Omegatron. Separation is achieved on the cyclotron principle (89,90, 118),by spiraling ions of a given q/M in a magnetic field H crossed by an rf field. Ions are collected only when the impressed rf frequency just equals the cyclotron frequency, v, = 1.54 x lo3 Hq/M sec -1. Pressure range: 10-5-10-10 mm; resolution drops as mass increases, but peaks up to mass 50 can usually be distinguished. ( b ) Topatron. This radio frequency mass spectrometer (119) is similar to a linear accelerator. Ions in resonance with the radio frequency signal are accelerated through a series of parallel grids and can penetrate into a retarding potential which rejects nonresonant ions. Pressure range: 10-~-10 -7 mm. (c) Farvitron. Here ions are brought into oscillation in a region of parabolic potential. A superimposed, high frequency signal imparts sufficient energy to ions in resonance to allow them to reach a collector (120). Designed for total pressures from 10-*-lo-* mm, but the sensitivity can be increased for work at lower pressures (122). The ideal spectrometer combines high resolution, high sensitivity, low response time, minimum bulk, and minimal interactions with the system together with simplicity of construction and circuitry. The

414

OERT EHRLICH

different spectrometer designs represent different compromises of these ideals. A discussion of the relative merits of these models, as well as the designs which are not as yet commercially available, is out of place in this sketch. I n general the additional power of the spectrometer is attained only a t the cost of a stronger perturbation of the system, primarily through the very considerable increase in the amount of metal required in the system. I n studying the interaction of gases with solids it is important t o take careful account of this, as well as of the ability of the spectrometer to sample the gases in the system impartially.

F. GAS HANDLING 1. Introduction of Gases

Greased stopcocks have no place in an ultrahigh vacuum system. This is not too great a loss since even in ordinary vacuum practice they are far from trouble-free. There are several simple and reliable ways of introducing and controlling gases. a. Dijfusion and Evaporation Sources. A selected few gases can be introduced into an ultrahigh vacuum system by diffusion ( M a ) through a membrane able to withstand baking. These are: hydrogen and deuterium through palladium, palladium-25 % silver alloys, or nickel; oxygen through silver; helium through quartz. Diffusion through molybdenum may conceivably prove advantageous for introducing small amounts of nitrogen, provided hydrogen and oxygen are excluded from the hot metal. Hydrogen diffuses rapidly a t much lower temperatures than nitrogen; oxidation of molybdenum becomes catastrophic a t high temperatures (T > 1000°K). This general technique has been widely used for many years; it is unfortunately limited by the selection of membrane materials available. The rate of gas flow in all these devices is controlled by changing the temperature of an external heater. This is satisfactory for constant flow systems, but too slow for accurate dosing of small quantities. Vapors of high melting metals and inorganic materials can be produced in situ by evaporation, either from resistively heated filaments of the material itself or else from bulk crucibles. These are all standard procedures (122). I n a variant of this approach, permanent gases have been evolved by high temperature decomposition of a suitable solid, as for example: H, from ZrH,, 0, from CuO ( 4 4 4 , or from oxygen dissolved in silver (123), and CO from Mo(CO), (124). I n all these it is

MODERN METHODS IN SURFACE KINETICS

415

important t o remove contaminants from the evaporator by thorough heat treating prior to evaporation. For filamentary samples this is easily done by pulse heating, while the evaporator is surrounded by a shroud that prevents deposition of contaminating material on the target. When evaporation of a 10v temperature material is carried out from a refractory heater, it may be necessary to separately outgas the former by heating to incandescence prior to loading. I n systems which are not baked, but in which gas evolution from the wells is suppressed by cooling, evaporation can also be employed for liberating permanent gases. Thus Gomer ( 4 4 4 has successfully condensed carbon monoxide as well as the rare gases on a liquid helium cooled surface, and then evolved these a t will by controlled heating. To insure purity of the material evolved by this and other evaporation techniques, it is necessary to start with a highly refined product and, where possible, to subject this to further purification in the vacuum system itself by chemical treatment (such as hydrogen reduction or decarburization with oxygen) or fractional distillation. b. Gas CapsuEes. Predetermined amounts of gas can be injected into an ultrahigh vacuum line from ampules provided with break-off seals (125). This method is generally applicable to any gas. It is limited, however, to situations in which the amounts desired are known before the experiment begins. The minimum pressure in these capsules is limited by the necessity of minimizing impurities. Contaminants may enter from the walls, as well as when the capsule is sealed after filling, and the pressure should therefore be kept above 1 p. c. Variable Leaks (1) Metal valves. The most versatile instrument for introducing gases is the all-metal valve. A host of designs (90) in a variety of sizes are now available commercially and are described in detail by a number of vacuum instrument manufacturers. Valves suitable for general small-scale laboratory use have 8 in. pumping openings. These valves are operated by forcing a hard nose piece, mounted on a flexible diaphragm, against a soft metal backing to establish a seal with a leak rate of 10-14 liters/sec. Hydraulic drivers now available give smooth fingertip control over several decades of pressure. When used in conjunction with gas containers at 1 atm or above, there is still sufficient gas flow through these valves (> 2 x 108 molecules/sec) even when they are completely closed to contaminate the system itself. For dosing, a series arrangement of valves, with differential pumping between them, may therefore be necessary. I n mounting, particuIarly of more than one valve, care must be exercised to retain flexibility of pumping leads. The valves themselves

416

UERT EHRLICH

must be rigidly attached to a supporting frame so that the closing torque will not shift their position. When a T-connection is made, as in Fig. 70, the tubing must be kept small (- Q in. diameter) and bellows must be inserted in every arm. Only in this fashion is the system able to withstand bakeout. Considerable gas is squeezed out when a valve is first closed. Repeated closures reduce this source of contamination until no pressure change is mm. This cleaning can also be observed with the system at p accelerated by flushing with the gas to be admitted. I n the other extreme of very high flow rates, when the nose piece is on the point of losing contact with the backing, the control of the valve may become erratic. Torque wrenches do not give a fine enough adjustment, and the rate of flow may suddenly rise by an order of magnitude. This difficulty is eliminated by hydraulic drivers that permit accurate manual positioning of the nose piece. The useful life of a valve is limited by the deterioration of the sealing surface, as well as of the metal body. Kovar connections are severely affected by oxidation and may develop leaks, as will the flexible diaphragm. This process can be delayed by painting with aluminum bronze. The sealing surface, on the other hand, is scoured by glass fragments formed in opening the break-off to the glass container. These particles may be effectively confined by interposing a plug of glass wool between the gas bottle and the valve. This, however, lowers the pumping speed significantly. An alternative arrangement, shown in Fig. 72, is to insert a small-scale version of the familiar oil catcher. This does not restrict the flow of gas through the system. Properly used, all-metal valves are both simple and reliable. They permit accurate, fast control of gas inflow for a large variety of gases and vapors. (2) Glass ports. Magnetically operated ground glass ports (126) serve t o separate parts of an ultrahigh vacuum system without complete interruption of flow. These devices do not occupy much space and when completely opened do not appreciably affect the pumping speed. They are helpful in leak checking a system, as well as in defining a fixed volume for adsorption studies. There are two general types-the flat plate, resting on a ground shoulder and the hemispherical ball and socket joint. The former can be made more compact; it does require careful polishing. The ball and socket joints occupy a considerable amount of space; however, to attain the same efficient closure as in the flat plate 10-4 liters/sec) requires only a trivial amount of grinding. (leak rate (3) Liquid seals. Materials which have a negligible vapor pressure, even a t bakeout temperatures, but which are liquid below the softening

-

N

MODERN METHODS I N SURFACE KINETICS

417

point of glass, can be used to seal parts of the system. The principle is, of course, precisely that employed in mercury cutoffs in a standard vacuum line. Only the material must be altered. Indium and gallium, which at 700°K have vapor pressures of 1.32 x 10-0 and 8.63 x 10-12 mm, respectively, are suitable for this purpose in glass systems. Care must be exercised with these to insure adequate outgassing and to prevent alloying with metal parts. Tin is not as readily handled, but at 700°K has a vapor pressure of 1.66 x 10-14 mm, and is more inert toward metals. It is clear that these methods of confining gases are not as versatile and convenient as the use of valves. (4) Capillary leaks. Gases can be introduced through a capillary leading to a reservoir. The rate of admission can then be controlled by varying the pressure or else the effective length of the capillary. The former requires a separate pump for the gas handling system. A capillary of variable length has been achieved by Hagstrum and Weinhart (127). They admitted gas through a rod of porous ceramic, the length of which was controlled by changing the level of a mercury bath surrounding the rod. Thorough outgassing is a problem with both methods, which again cannot compete in simplicity with the use of valves. 2. Gas Purijication

A large selection of gases with an impurity content of only 1 in l o s or less is available in glass containers, equipped with break-off seals. These assayed reagents are satisfactory for routine work; however, for adsorption studies at high pressure'or long times, even minute amounts of contamination may obscure the validity of the experiment. Further purification in the system itself may therefore be necessary. Each gas and each specific application will demand special procedures. I n the following are listed a number of techniques with illustrative examples drawn from experience in this Laboratory. a. Diflusion. Although selective membranes suffer by comparison with valves for use in admitting gases, they are ideal for gas purification. After passing through palladium or palladium-silver alloys, tank hydrogen has an impurity concentration (128) of less than 1 in 1O'O; with a nickel membrane the lowest contaminant concentration is of the order of 1 in 108. The pressure diminution as a result of diffusion through such a thimble can actually serve as a test for hydrogen, as shown long ago by de Boer and Lehr (129),and most successfully used by Young (128). For the other gases which can be introduced by diffusion, the following impurity concentrations (130) have been reported: He through quartz: 10 ppm H,, less than 1 of A, 0,, N,, CO, or CO,. Purification

418

GERT EHRLICH

has also been accomplished by allowing the super fluid a t T < 2.18"K to penetrate into a very fine opening in the vacuum envelope (131). 0, through Ag: Impurities on the same order of magnitude as for reagent bottled gas (dictated primarily by decarburization reactions in the mass spectrometer). Diffusion through a selective membrane yields the purest gas so far achieved; it is unfortunate that this technique is not applicable t o a larger variety of gases. b . Selective Gettering. Reagent grade gases introduced through a metal valve may be further purified by passage over a surface that selectively removes contaminants (41). The principle of operation here is just the opposite of that in diffusion. In gettering the impurities are removed because of a preferential interaction with the surface. I n diffusion it is the primary component that undergoes a unique interaction and is therefore selected. This distinction is important, for it suggests that complete removal of all impurities by selective gettering should not be expected. Several getter systems employed in investigations in this Laboratory are listed in Table VII. TABLE VII Selective Getters

Gas (to be purified)

Rare gases

Getter

W, Mo

NZ CH,

co

Ni Ni Ge

1

Impurities removed

Reference

All chemically active gaces

45, 132

CO, COz,

co,,

0,

0 8 ,

Hz,C J %

132 133

The choice of getter can usually be based on a qualitative tabulation of selectivity in chemisorption, such as the periodic table in Fig. 76. There are marginal systems, however, which demand quantitative data. Thus for the removal of CO from hydrogen it would appear from Fig. 76 that gold might prove useful. However, Trapnell's (134) data indicate that a t room temperature the partial pressure of CO usually found in reagent hydrogen ( -0.5 mm) would be lowered less than 1yo over gold. This purification can, however, be readily accomplished by other techniques. A single getter, properly selected, may sometimes serve t o clean up two gases; thus in the ion microscopy of adsorption, both the image gas,

MODERN METHODS IN SURFACE KINETICS

419

helium, as well as the gas under examination should be carefully purified. I n studying nitrogen, both can be accomplished with nickel (76). Mass spectrometric tests of the efficiency of purification are not yet available. Indirect evidence, such as work function measurements, suggest that with N,, the impurity content is lowered to less than 5 ppm. These getters are employed in the form of thin films, prepared in the system itself by direct evaporation from a resistivity heated filament of the parent metal, or with germanium from a wire spiral.

CHEMISORPTION ON THE ELEMENTS

GASES ADSORBED

ma NZlHZlCO FIG. 76. Selectivity in chemisorption of common diatomic gases.

Although the selectivity of getters is not as great as of diffusion membranes, the latter have a wider scope and are effective even at room temperature. c. Condensation and Evaporation. A cold finger, such as shown in Fig. 70, makes it possible to purify through condensation gases already in the system. Reagent grade materials have already been fractionated by the manufacturer, and liquid nitrogen cooling of the ultrahigh vacuum parts is therefore generally ineffective. At lower temperatures the following separations become possible. (1) Triple-poi-nt nitrogen. A t the nitrogen triple point, 63.1°K, the vapor pressure of carbon dioxide is reduced to 10-12 mm; CO, in equilibrium with the solid (135), still has a pressure of 40 mm. By pumping on liquid nitrogen in a cold finger it is therefore possible to remove this troublesome impurity of carbon monoxide. The onIy precaution required is to bubble nitrogen into the cold finger through a capillary leak while pumping in order to suppress bumping.

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GERT EHRLICH

CO on the high pressure side of a valve can be cleaned up even a t the normal boiling point of nitrogen (77.35OK). There the vapor pressure of CO, is 1.05 x mm, that of CO 453 mm; the former can therefore be limited to 2.3 parts in lo9. (2) Liquid hydrogen. In equilibrium with its own vapor at 760 mm, liquid hydrogen has a temperature of 20.4"K. Only hydrogen helium, and neon exert an appreciable vapor pressure a t this temperature. To remove the reactive impurities from hydrogen gas it is therefore only necessary to pass it over a liquid hydrogen-cooled surface. Liquid

U

GAS TRANSFER SYSTEM

FIG.77. Ultrahigh vmuum system for filling low pressure gas containers ( p N 1 mm). T-Thermocouple gauge; I-Ionization gauge; N-liquid nitrogen trap; P-Ground glass port. Break-off seal B is opened only after bakeout and after sealing off at C. Gas from cylinder G ( p 800 psi) admitted through copper tubing L to GranvillePhillips valve (2101-S-HP)V. For pyrolyzable gases, bottles should be equipped with a doubIe break-off seal.

-

hydrogen is very convenient to use; it can be handled and transferred precisely like liquid nitrogen, provided ventilation is adequate to keep the hydrogen concentration in the atmosphere below the explosive limit (4.1%). Because of its ease of transfer, even in small quantities,

MODERN METHODS IN SURFACE KINETICS

42 1

liquid hydrogen is also ideally suited for use in cold fingers, as an auxiliary pump. (3) Liquid helium. At 4.2"K, the normal boiling point of helium, only helium and hydrogen have a vapor pressure in excess of lo-' mm. Purification of helium can therefore be accomplished a t this temperature, provided the helium is collected at one atmosphere. This, however, is not as convenient a technique as gettering or diffusion; liquid helium is more useful for unselective cryopumping. 3. Gas Transfer Many interesting gases, such as deuterium, ammonia, and all the hydrocarbons are available only in high pressure cylinders. These are not suitable for direct use in adsorption and catalytic studies a t low pressures; for this type of work glass bottles, with break-off seals, holding the gas a t one atmosphere or less, are most suitable. I n transferring gas to these flasks it is important to remember that at low pressures the residual gases in the flask, as well as on the walls, make a proportionally greater contribution than in the high pressure cylinder. Baking of the glass containers under vacuum is therefore desirable. A system for bottling gases is shown in Fig. 77. Some leakage from the high pressure cylinder is inevitable. To prevent the build-up of a high pressure of gas that may pyrolizt during baking, the high pressure part of the system behind the break-off seal is evacuated by a separate pump. After baking and establishing a good vacuum in the main part of the system this pump is sealed off. The valve is a special model that can be used a t pressures up to 800 Ib/sq in. Since each sealing operation introduces impurities the system should be flushed with fresh gas after filling and removing every bottle.

G . SOURCES OF EQUIPMENT The availability of commercial equipment, more than any other factor, has changed ultrahigh vacuum work. Rather than constituting a test of experimental skill and patience, the achievement of good vacua now demands nothing more strenuous than a choice between alternatives in the design and construction of ready-made equipment. A listing of manufacturers supplying the special instruments and equipment discussed in this section is therefore appended. Under the stimulus of the space age (which has transformed the lowly vacuum line into a space simulator) the number of manufacturers specializing in equipment suitable for adsorption studies is rapidly increasing. *The following list is therefore fragmentary, and is based mainly on personal experience; *For example, a complete line of ultra-high vacuum equipment has recently been announced by General Electric.

422

OERT EHRLICH

a more detailed appreciation of the variety of instruments available can be attained at the bi-annual equipment exhibit of the Vacuum Society.

1. System Components (a)

Pumps Mercury Diffusion

Glass Metal

Turbo-Molecular Getter-ion Sorption (b) Shear Seals, Fittings

Kontes Glass Company H. S. Martin Edwards High Vacuum The Welch Scientific Co. Ultek Varian Associates Granville-Phillips Ultek Varian Associates

(c) Heaters Bakeout Ovens Heating Tapes (d) Temperature Controllers Programmed Ordinary (e) Dewars Glass Metal-for liquid hydrogen (f) Liquid Level Controllers

Gruenberg Electric Wilt Briskeat Glas-Col West Instrument General Electric Honeywell

H. S. Martin Superior Air Products Fisher Scientific (Levelator) Johns and Frame, Livermore

MODERN METHODS IN SURFACE KINETICS

423

2. Pressure Detectors (a)

Ionization Gauges Bayard-Alpert, 5966

Modified Inverted

Westinghouse Sylvania Tung-Sol Thoriated I r Emitter Nottingham Miscellaneous

Veeco NRC

cvc

W.R. & W., Port Washington, N.Y. Schulz-Phelps (High Pressure) (b) Miscellaneous Gauges Redhead Magnetron Thermocouple

Westinghouse NRC General Electric Hastings-Raydist

Pirani

cvc

Knudsen Diaphragm

Sy lvania Edwards High Vacuum C-ranville-Phillips Atlas Werke, Bremen

McLeod (c) Gauge Power Supplies With Adjustable Electron Current (below 10-4 amp)

For Magnetron (d) Ion-Current Detectors Vibrating Reed Electrometer Electrometers, etc. Oscilloscopes Operational Amplifiers (e) High-speed Recorders

cvc

cvc Edwards High Vacuum Nuclide Analysis Associates Verian Associates NRC Cary Instruments Keithley Tektronix

Brush Electronics Honeywell (Visicorder) Sanford ( f ) Mass Spectrometers

Magnetic Deflection

Omegatron Radio Frequency (Topatron) Farvitron

General Electric Nuclide Analysis Associates CEC Sy lvania Leybold Leybold Vactron Engineering Leybold

424

UERT EHRLICH 3. Gas Train

(a) Variable Leaks

Metal-Valves

Gas-admiasion Large Bore

Glass Ports (b) Controlled Diffusion Sources

Granville-Phillips Granville-Phillips Varian Associates Westinghouse Leybold

K-B Glass Apparatus, Schenectady (0) Bottled Gases Matheson

ACRNOWLEDQMENTB Many people at this Laboratory have contributed to the experience on which this review is besed. I am especially indebted t o the Glass Development Shop, and to those members of the Metallurgy and Ceramics Research Department with whom I have at one time or another been associated in the surface program: T. W. Hickmott, E. G. Brock, F. G. Hudda, T. A. Vanderslice, C. A. Neugebauer, E. J. Covington, and B. McCarroll. The burden of preparing the manuscript itself was eased through the friendly help of Messrs. Hudda and McCarroll.

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MODERN METHODS IN SURFACE RINETICS

425

20. a. Ehrlich, G., and Hickmott, T. W., Nature 177,1046 (1966);b. Hickmott, T. W., 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38.

39. 40. 41. 42. 43. 44.

45. 46. 47. 48. 49.

50. 51. 52.

53. 54.

55. 56.

57. 58. 59. 60.

and Ehrlich, G., J. Chem. Phye. 24, 1263 (1966). Ehrlich, G.,J . AppZ. Phya. 32, 4 (1961). Smith, A. W., and Aranoff, S., J. Phya. Chem. 62,684 (1968). Redhead, P.A., Tram. FaradaySoc. 67, 641 (1961). Randall, J. T., and Wilkins, M. H. F., Proc. Roy.Sot. A184,366 (1946). Booth, A. H., Can. J. Chem. 82, 214 (1964). Clausing, P.,Ann. Phyaik [6] 6,489 (1930). Hagstrum, H. D., Rev. Sci. Znstr. 24, 1122 (1963). Busch, H., Ann. Phyeik [a] 64,401 (1921). Ehrlich, G., J . Chem. Phya. 36, 1499 (1962). Kington, G. L.,and Aston, J. G., J . Am. Chem. SOC.73,1929 (1961). Beeck, O.,Advancea in Catal. 2, 161 (1960). Brennan, D.. Hayward, D. O., and Trapnell, B. M. W., Proc. Roy.SOC.A266, 81 (1960). McCarroll, B., and Ehrlich, G., J. Chem. Phya. 88, 623 (1963). Hill, T. L., Advances i n Catal. 4, 212 (1962). Hill, T. L.,J. Chem. Phya. 14,441 (1946). Ehrlich, G.,J. Chem. Phya. 31, 1111 (1969). Langmuir, I., and Kingdon, K. H., Science 67,68 (1923). Davisson, C., and Germer, L. H., Phys. Rev. 30,706(1927). Frisch, R., and Stern, O., i n “Handbuch der Physik” (H. Geiger and K. Scheel, eds.), Vol. 22, p. 313. Springer, Berlin, 1933. Mignolet, J. C. P., Mdm. uoc. roy. aci. W g e [6] 1, faac. 3 (lD68). Ehrlich, G.,and Hudda, F. G., J . Chem. Phye. 86, 1421 (1961). Muller, E. W., Phyeik. 2.87, 838 (1936). Gomer, R.,Advancea in Cutal. 7,93 (1966). a. Muller, E. W., Ergeb. ezakt. Naturzu. 27, 290 (1963);b. Dyke, W. P., and D o h , W. W., Advances in Electronics and Electron Phye. 8,89 (1966);c. Good, R. H., and Muller, E. W., i n “Handbuch der Physik” (S. Flugge, ed.), Vol. 21,p. 176.Springer, Berlin, 1966; d. Gomer, R.,“Field Emission and Field Ionization,” Harvard Univ. Press, Cambridge, Massachusetts, 1961. Ehrlich, G., and Hudda, F. G., J. Chem. Phye. 30,493 (1969). Fowler, R. H., and Nordheim, L., PTOC. Roy.SOC.A119, 173 (1928). Stern, T. E., Gossling, B. S., and Fowler, R. H., PTOC. Roy.SOC.Al24,699 (1929). Alpher, R. A., and White, D. R., Phya. Fluids 2, 163 (1969). Landolt-Bornstein, “Zahlenwerte und Funktionen,” 6th edition, Vol. I, Part 1, p. 401.Springer, Berlin, 1960. Liepack, H., and Drechsler, M., Naturwiuaenschajten 43,62 (1966). Hess, E., and Drechsler, M., Proc. 4th Intern. Conj. on Electron Microscopy, Berlin, 1958, I, p. 811. Springer, Berlin, 1960. Topping, J., PYOC. Roy. Sot. A114, 67 (1927). Herring, C., and Nichols, M.H., Reva. Modern Phys. 21, 186 (1949). Mignolet, J. C. P., Rec. trav. chim. 74, 686 (1966). Kisliuk, P.,Phya. Rev. 122, 406 (1961). Gomer, R., and Hulm, J. K., J. Chem. Phya. 27, 1363 (1967). Gomer, R., Wortman, R., and Lundy, R . , J . Chem. Phya. 26,1147 (1967). Utsugi, H., and Gomer, R., J. Chem. Phys. 37, 1706 (1962). Moore, G. E., and Allison, H. W., J . Chem. Phys. 23, 1609 (1966). This waa derived independently by Klein. R., J . Chem. Phya. 81, 1306 (1969).

QERT EHRLICH

426

61. Ehrlich, G., and Hudda, F. G., Unpublished work. 62. Swanson, L. W., and Comer, R., Proc. 8th Field Emiesion Symposium, Williamstown, 1961. 63. Langmuir, I., J. Am. Chem. SOC.54, 2798 (1932). 64. Langmuir, I., J. Am. Chem. SOC.40, 1361 (1918). 65. Morrison, J. L., andRoberts, J. K. Proc. Roy. SOC.A173,13 (1939). 6 6 . Ehrlich, G., “Surfaces: Structure, Energetics and Kinetics,” ASM-AIME Symposium, New York, 1962, in press. 67. Burton, W. K., Cabrera, N., and Frank, F. C., Phil. Trans. Roy. SOC.A243, 299 (1951). 68. Muller, E. W., 2. Phyeik 131, 136 (1951). 69. Muller, E. W., Advances in Electronics and Electron Phys. 13, 83 (1960). 70. Lanczos, C., 2.Phyeik 68, 204 (1931). 71. Rice, M. H., and Good, R. H., Jr., J. Opt. SOC.Am. 52, 239 (1962). 72. Hiskes, J. R., Phya. Rev. 122, 1207 (1961). 73. McCarroll, B., Private communication (1962). 74. Gurney, R. W., Phys. Rev. 47, 479 (1935). 75. Dobrezow, L. N., “Elektronen- und Ionenemission,” p. 110. VEB Verlag Technik, Berlin, 1954. 76. Ehrlich, G., and Hudda, F. G., J.Chem. Phya. 36,3233 (1962). 77. Muller, E. W., Proc. 4th Intern. Conf. on Electron Microscopy, Berlin, 1958, I, p. 820. Springer, Berlin, 1960. 78. Muller, E. W., and Young, R. D., J. Appl. Phys. 82, 2425 (1961). 79. Muller, E. W., Phys. Rev. 102, 618 (1956). 80. Mott, N. F., and Watts-Tobin, R. J., Electrochim. Acta 4, 79 (1961). 81. Parkinson, D., Proc. Phys. SOC.(London)75, 169 (1960). 82. Sewell, G. L., Proc. Cambridge Phil. SOC.45, 678 (1949). 83. Muller, E. W., 2.Phyaik 126, 642 (1949). 84. Ehrlich, G.,J. Chem. Phys. 36, 1171 (1962). 85. Ehrlich, G., and Hudda, F. G., Proc. 9th Field Emission Symposium, Notre Dame, 1962. 86. McGee, J. D., Repta. Progr. i n Phys. 24, 167 (1961). 87. a. Dushman, S., “Scientific Foundations of Vacuum Technique,” Wiley, New York, 1949; b. Pirani, M., and Yarwood, J., “Principles of Vacuum Engineering.” Reinhold, New York, 1961. 88. Alpert, D., J. Appl. Phys. 24, 860 (1953). 89. Alpert, D., i n “Handbuch der Physik” (S. Fliigge, ed.), Vol. 12, p. 609. Springer, Berlin, 1958. 90. Brombacher, W. G., “Bibliography and Index on Vacuum and Low Pressure Measurements,” N.B.S. Monograph No. 35. Natl. Bureau of Standards, Washington, D.C., 1961. 91. Anderson, P. A., Phya. Rev. 47, 958 (1935); see also later publications of the same

author. 92. Nottingham, W. B., J. Appl. Phye. 8, 762 (1937). 93. Coploy, M. J., and Phipps, T. E., Phys. Rev. 48,960 (1935). 94. a. Roberts, J. K., “Some Problemsin Adsorption,” Cambridge Univ. Press, London and New York, 1939; b. Wheeler, A., in “Structure and Properties of Solid Surfaces,” p. 439. Univ. of Chicago Press, Chicago, Illinois, 1953. 95. Biondi, M. A., Rev. Sci. Instr. 30, 831 (1959). 96. Wagener, J. S., Proc. 4th Natl. Conf. on Tube Techniques, 1958, p. 1 (1959).

MODERN METHODS IN SURFACE KINETICS

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97. a. Hall, L. D.,Science 128,279 (1958);b. Klopfer, A., Vakuum-Tech. 10,113 (1961). 98. a. von Ardenne, M., “Tabellen der Elektronenphysik, Ionenphysik und Ubermikroskopie.” Deut. Verlag der Wissenschaften, Berlin, 1956. b. Espe, W., “Werkstoffkunde der Hochvakuum-technik.” Deut. Verlag der Wissenschaften, Berlin, 1959; c . Knoll, M., “Materials and Processes of Electron Devices.” Springer, Berlin, 1959; d. Kohl, W. H., “Materials and Techniques for Electron Tubes.” Reinhold, New York, 1960. 99. Prescott, L.E., Rev. Sci. Imtr. 33, 485 (1962). 100. Bayard, R. T., and Alpert, D., Rev. Sci. I W ~ T21, . 571 (1950). 101. a. McGowan, W., and Kerwin, L., Can. J. Phys. 38, 567 (1960);b. Schulz, G.J., J. Appl. Phya. 28, 1149 (1957). 102. Dushman, S.,and Young, A. H., Phys. Rev. 68,278(1945). 103. Wagener, S., and Johnson, C. B.,J.Sci. Imtr. 28, 278 (1951). 104. Nottingham, W. B., and Torney, F. L., Jr., Natl. Symposiumon Vacuum Technol., Tram. 7 , 117 (1960). 105. Schulz, G. J., and Phelps, A. V., Rev. Sci. Imtr. 28, 1051 (1957). 106. Jones, H. A,, and Langmuir, I., Uen. Elec. Rev. 30,408(1927). 107. Redhead, P. A., Natl. Sympoeium on Vacuum Technol., Trans. 7,108 (1960). 108. Schwarz, H.,2.Physik 122, 437 (1944). 109. Lafferty, J. M., Natl. Symposium on Vacuum Technol., Tram. 8,460 (1961). 110. Wagener, J. S., and Marth, P. T., J. Appl. Phya. 28, 1027 (1957). 111. Hickmott, T. W . , J . Appl. Phys. 31, 128 (1960). 112. Klopfer, A., Natl.Symposium on Vacuum Technol., Trans. 8,439 (1961). 113. Nottingham, W. B., Natl. Sympoaium on Vacuum Technol., Trans. 1, 76 (1954). 114, Redhead, P. A., Nail. Sympoaium on Vacuum Technol., Trans. 5, 148 (1958). 115. Ellett, A., and Zabel, R. M., Phya. Rev. 37, 1102 (1931). 116. Onchi, M., Natl. Sympoaium on Vacuum Technol., Trans. 8,544 (1961). 117. Davis, W. D., and Vanderslice, T. A., Natl. Symposium on Vacuum Technol., Trans. 7. 417 (1960). 118. Diimmler, S.,Vakuum-Tech. 10, 131, 184 (1961). 119. Varadi, P. F., Sebestyen, L. G., andRieger, E., Vakuum-Tech. 7 , 13,48(1958). 120. Tretner, W.,2. angew. Phyeik. 11, 395 (1959). 121. Behrisch, R., Blauth, E. W., Melzner, F., and Meyer, E. H., Physik. Verhandl. 12, 224 (1961). 122. Holland, L.,“Vacuum Deposition of Thin Films.” Wiley, New York, 1956. 123. Miiller, E.W., Private communication. 124. Klein, R., Bull. Am. Phys. SOC.[ 2 ] 8, 47 (1958). 125. Mitchell, J. W., Private communication. 126. Decker, R. W., J. Appl. Phys. 25, 1441 (1954). 127. Hagstrum, H.D., and Weinhart, H. W., Rev. Sci. Inatr. 21,394 (1950). 128. Young, J. R., To be published. 129. de Boer, J. H., and Lehr, J. J., 2.phyaik. Chem. B22,423 (1933). 130. Young, J. R., and Whetten, N. R., Natl. Symposium on Vacuum Technol., Tram. 8, 625 (1961). 131. Biondi, M. A., Rev.Sci. Imtr. 22,536 (1951). 132. Bond, G. C., “Catdysis by Metals,” Chapter 5. Academic Press, New York, 1962. 133. Bennett, M. J., and Tompkins, F. C., Tram. Faraday SOC.58, 816 (1962). 134. Trapnell, B. M. W., Proc. Roy.SOC.A218, 566 (1953). 135. Johnson, V, J , , “A Compendium of the Properties of Materials a t Low Temperatures,” Part 1. WADD Tech. Report 60-56(1961).

This Page Intentionally Left Blank

Catalytic Oxidation of Hydrocarbons L . YA . MARGOLIS Ilzetitute of Chemical Phy@cs Academy of Sciences, MOSCOW. U.S.S.R.

. 429 437 111. Reaction Mechanism ................................................. 439 A . Oxygen Adsorption on Oxidation Catalysts ........................... 440 B. Isotopic Oxygen Exchange ......................................... 442 C. Chemisorption of Hydrocarbons on Oxidation Catalysts ................ 444 D. Charging of the Surface in Adsorption ................................ 446 E . Oxidation Schemes................................................ 448 F. Heterogeneous-Homogeneous Reaction Steps......................... 467 IV. Reaction Kinetics .................................................... 468 A. The Effect of Macroscopic Factors ................................... 468 B. Kinetics of Propene and Ethylene Oxidation over Vanadium............ 469 C. Kinetics of Naphthalene Oxidation over Vanadium Oxide .............. 470 D. Kinetics of Benzene Oxidation to Maleic Anhydride over V,O, ........... 472 E . Kinetics of Propene Oxidation to Acrolein over Cu,O ................... 473 F. Kinetics of Ethylene Oxidation to Ethylene Oxide over Silver........... 474 C-. Kinetics of High Conversion of Hydrocarbons ......................... 476 479 V. Modified Catalysts ................................................... VI . Mixed Catalysts ...................................................... 493 References .......................................................... 496 I Introduction .........................................................

I1. Catalysts ............................................................

I. Introduction Organic compounds involving oxygen are the main stock for synthesis of various plastics. lacquers. resins. and other material . Oxidative processing of hydrocarbons haa long ago attracted the attention of chemists as one of the main trends in organic synthesis. The results of investigations on the oxidation of certain hydrocarbons over various catalysts are summarized in Tables I to 111. 429

430

L. YA. MAROOLIS TABLE I Hydrocarbon Oxidationa on Metals -

~~~~

Catalyst

Platinum

Platinum

~

Support

-

Asbestos

~

~

Reaction Compound tempera- undergoing ture ("C) oxidation 800

Methane

250-600 Methane

Ethylene Acetylene Propene Pentane Benzene 150-300 Toluene

Platinum

-

Platinum

-

Platinum

Silica gel

Toluene Methane 100-300 Ethylene

Platinum

Asbestos

250-500 Methane

Platinum

Asbestos

Platinum

Asbestos

Pentane Hexane 172-700 Methane

Platinum

Silica gel

200-400 Methane

Platinum

-

150-350 Methane

Platinum

-

Platinum

Barium sulfate

Platinum

Asbestos

Platinum Palladium Palladium Palladium

800

800

700

20-180 Propene 250-500 Pentane

Heptane Octane Methane

Asbestos Asbestos Asbestos

Ethylene

Methane 450-700 Methane 150-600 Methane Ethylene Acetylene Pentme Benzene 200

Main reaction products

CO and H,O at times traces of HCOOH COPand H,O HCOOH

...

CO, and H,O . . HCOOH CO, and H,O HCOOH CO, and H,O HCOOH CO, and H,O . . HCOOH CO, and H,O . HCOOH CO, and H,O . . . HCOOH CO, and H,O HCOOH CO, and H,O . . . HCOOH CO, and H,O HCOOH CO, and H,O . . HCOOH CO, and H,O HCOOH

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

.. .

. ...

CO, and H,O HCOOH CO, and H,O C O , and H,O CO, and H,O

...

References

CATALYTIC OXIDATION OF HYDROCARBONS

43 1

TABLE I (cont.)

Catalyst

Support

Reaction temperature "C

Compound undergoing oxidation

Palladium Palladium Palladium Silver

Silica gel Silica gel Asbestos Silica gel

100-300 200-400 172-700 100-300

Silver

Aluminum oxide Asbestos

260-290 Ethylene

Silver

Ethylene Methane Methane Ethylene

263-293 Ethylene

Silver

-

225-325 Ethylene

Silver

-

263-293 Ethylene

Silver

Porous

220-225 Ethylene

Silver

Skeletal

220-250 Ethylene

Silver

Silica gel

200-400 Methane

Silver

Pumice

172-700 Methane

Silver

-

Silver

-

475 Methane (pressure 200 atm) 500 Methane

Pumice

Copper Copper

Silica gel

Copper

Pumice

Copper

Silica gel

200-400

Copper Copper

Pumice

172-700 400

-

-

-

100-300 150-300

300

CO, and H,O CO, and H,O CO I and H ,O Ethylene oxide GO, and H,O Ethylene oxide CO, and H,O Ethylene oxide CO, and H,O Ethylene oxide CO, and H,O Ethylene oxide CO, and H,O Ethylene oxide CO, and H,O Ethylene oxide CO, and H,O CO,, H,O, and traces of aldehydes CO,, HzO, . . . aldehydes CO,, HzO, . . aldehydes

.

Cozy

. . . aldehydes Maleic acid CO,, H,O, and traces of HCOH CO, and H,O Ethylene CO,, H,O, and Toluene traces of benzaldeh yde Toluene Cow HzO, . . .benzaldehyde CO,, H,O, and Methane traces of CH,COOH CO, and H,O Methane 3.3-dimethyl- CO,, H,O, and pentane traces of aldehydes

263-295 Benzene 500 Methane

Silver Copper

Main reaction products

Referencea

432

L. YA. MARUOLIS

TABLE I (cont.)

Catalyst

Gold Gold Gold

Support

Asbestos Silica gel Asbestos

Reaction temperature "C

Compound undergoing oxidation

250-500 Methane 100-300 Ethylene 150-500 Methane

Ethylene Propene Pentane Benzene Methane

Nickel

-

Nickel

-

160-350 Methane

Nickel

__

850-900 Methane

Nickel

-

150-300 Toluene

500

Main reaction products

C O , and H,O CO, and H,O GO, and H,O

References

(6) (5)

(3)

CO,, H,O, and (24) traces of formaldehyde COB, HsO, (9) . . formaldehyde COs, H.0, (27) , formaldehyde Traces of (2) benzaldehyde

. ..

CATALYTIC OXIDATION OF HYDROCARBONS

433

TABLE I1 Hydrocarbon Oxidatwne an Metal Oxidea

Catalyst

Vanadium pentoxide Vanadium pentoxide Vanadium pentoxide Vanadium pentoxide

Support

Reaction temperature "C

-

400

Asbestos

Asbestos

Vanadium pentoxide

Vanadium pentoxide

Vanadium pentoxide

Compound undergoing oxidation Benzene

Main reaction products Maleic acid

Naphthalene Phthalic anhydride 400 Ethylene Formaldehyde Acetaldehyde Acids 400 Ethylene Formaldehyde Acetaldehyde Mdeic and fumaric acids Benzaldehyde 400-600 1.3-butaMaleic acid dime Formaldehyde Glyoxal Acetic acid Methylacrolein 400-600 Isobutene 400-600 1- and 2Maleic and butene acetic acids Acetaldehyde Formaldehyde Methylvinylketone 400 Butane Maleic and metic acids Formaldehyde GIyoxal CO, CO,, and 600

H,O Vanadium pentoxide Vanadium pentoxide Vanadium pentoxide Vanadium pentoxide

Asbestos

400

Benzene

Asbestos

300

Toluene

Asbestos

460

Asbestos

Maleic anhydride Benzaldehyde

Naphthalene Phthalic anhydride 172-700 Methane CO, and H,O

References

434

L. YA. MAROOLIS

TABLE I1 (cont.)

Catalyst

Support

Reaction Compound tempera- undergoing ture ("C) oxidation

Asbestos

36&600 2-pentane

Asbestos Asbestos

Pumice Pumice

Trimethylethylene Amy1ene Heptane 360-400 Cyclohexane 400-450 Benzene

Vanadium pentoxide

Pumice

290-420 Naphthalene

Vanadium pentoxide Vanadium pentoxide Vanadium pentoxide Vanadium pentoxide

Pumice

Vanadium pentoxide Vanadium pentoxide Vanadium pentoxide Vanadium pentoxide Copper oxide (CUO) Copper oxide (CUO) Copper oxide (CUO) Copper oxide (CUO) Copper oxide (CUO) Cuprous oxide G O )

Silica gel Fused asbestos Alundum Silica gel

Asbestos Asbestos Pumice Pumice

Asbestos Asbestos Asbestos Pumice Pumice Silicon carbide

600

Methane

Main reaction products Maleic acid Aldehydes Acids

Maleic acid Maleic anhydride Phthalic anhydride Meleic acid CO, and H,O

380-480 Toluene

Benzaldehyde Maleic acid 426 Phenanthrene Phenanthrenequinone 400-600 Naphthalene Naphthaquinone Phthalic anhydride, CO,, and HSO 400-460 Propene Formaldehyde CO, and H,O 400 Toluene Benzaldeh yde Benzoic acid 400 PhenanPhenanthrenethrene quinone 400 Naphthalene Phthalic Toluene anhydride Benzoic acid 160-200 Acetylene CO, and H,O 260-360 Methane

100

Methane

360-1000 Methane

400

CO,, H,O, and formaldehyde CO, and H,O CO, and H,O

Phenanthrene Phthalic acid

360-400 Propene

Acrolein CO, and H,O

References

CATALYTIC OXIDATION O F HYDROCARBONS

435

TABLE I1 (cont.)

Catalyst

Support

Tungsten Pumice oxide (WO,) and molybdenum oxide (MOO,) WO,andMoO, Pumice W0,and MOO, Asbestos

Main reaction products

400-460 Toluene

Benzaldehyde Benzoic acid CO, and H,O

400-460 Benzene 300-426 Toluene

160-200 Acetylene

Benzaldehyde Benzaldehyde Benzoic acid CO,, CO, H,O CO, and H,O

400-460 Benzene

Benzaldehyde

400-460 Benzene

Maleic acid

160-200 Acetylene

CO, and H,O

Asbestos

250-600 Methane

CO, and H,O

Asbestos

100-350 Methane

Asbestos

360-600 Methane

Formaldehyde CO, and H,O CO, and H,O

Asbestos

360-1000 Methane

Asbestos

360-1000 Methane

Manganese oxide (MnO,) Manganese Pumice oxide (MnO,) Nickel oxide (NiO) Nickel oxide (NiO) Nickel oxide (NiO) Nickel oxide (NiO) Lead and ceric oxides Zinc and lead oxides Zinc and lead oxides Chromium oxide

Reaction Compound tempera- undergoing ture ("C) oxidation

Pumice

-

Asbestos

600

Ethylene

Formaldehyde CO, and H,O CO, and H,O Acetaldehyde Formaldehyde CO, and H,O

References

436

L. YA. MAROOLIS

TABLE I11 Hydrocarbon Oxadationa on Mixturea of Ildekzb with Certain Salts

Catalyst

Carrier

Reaction temperature ("C)

Compound undergoing oxidation

Main reaction products

References

-

Zinc Asbestos vanadate Bismuth and lead vanadates Vanadates of various metals Lead vanadate -

Chromium, tungsten, vanadium and molybdenum plumbites Uranium and berryllium oxides Bismuth and Pumice lead vanadates Manganese Pumice and lead vanadates Copper, iron, Pumice and magnesium ehromites

Zinc and Pumice manganese oxides with chromium oxide Cadmium and Pumice zinc vanadates

300-600 Toluene 200-380 Toluene

Benzaldehyde Benzoic acid Benzoic acid

(34)

153)

naphthalene

460

Naphthalene Phthalic anhydride

(35)

-

Naphthalene Phthalic acid Toluene Benzoic acid 300-400 Toluene Benzaldehyde Benzoic acid

600-000 Methane

Formaldehyde

(55)

400-460 Benzene

Benzaldehyde

(38)

290-420 Naphthalene Acetaldehyde

Cyclohexane Decalin 800-600 Pentane Heptane Ethylene Benzene Cyclohexane Tsooctane - Acetylene

-

Acetylene

Phthalic anhydride CO, and H,O

(56, 1 3 )

Acetone

(57)

Acetone

(58)

CATALYTIC OXIDATION OF HYDROCARBONS

437

Oxygen enters into diverse reactions with hydrocarbons, such as addition, substitution, and destructive and polymerizing oxidation. There are a t present three fundamental directions of pursuit in the investigation of oxidative processing of hydrocarbons : (1) heterogeneous oxidation by molecular oxygen in the gas or liquid phase; (2) homogeneous oxidation by molecular oxygen in the gas or liquid phase; (3) radiation-induced oxidation by molecular oxygen, oxidation under electric discharge, electrochemical oxidation in solutions, etc. Certain chemical features, such as the formation of peroxide compounds, of atomic oxygen, etc., are common for all the three groups of reactions. A theory of heterogeneous catalytic oxidation of hydrocarbons would be impossible without knowledge of the elementary mechanism of oxidation, of the fundamental laws governing this process, and of its rate-determining steps. Insufficient theoretical treatment of the wide amount of experimental information on the proper choice of catalysts for hydrocarbon oxidations also hampers advances in this field.

II. Catalysts Metals (platinum, silver, etc.), metal oxides-semiconductors (vanadium pentoxide, molybdenum, tungsten, and copper oxides, etc.), and complex semiconductors-spinels, are the catalysts most widely used for oxidation of hydrocarbons. Oxidation catalysts almost ihvariahly contain transient metals with open inner electronic shells (Tables I V and V). The catalytic process comprises a number of consecutive elementary acts, such as the breaking and formation of bonds in reactant molecules, and electron transfers between the latter and the solid catalyst. Consequently, the electronic properties of catalyst surfaces influence the catalytic and adsorption processes. The catalytic activities of metals and semiconductors would be expected to differ due to their different electronic properties. However, under conditions of oxidation catalysis many metals become coated with a more or less thin semiconducting film of the given metal oxide, and this might be the reason why the mechanism of hydrocarbon oxidation on metals and semiconductors has much in common (59). Kalish and Burstein (60)found that the amount of oxygenadsorbed by a platinum layer adjacent to the surface was about one hundred times that required for a monolayer. Temkin and Kul’kova (61)have noted a similar phenomenon for oxygen adsorption on silver. The amount of

L. YA. MARGOLIS

438

oxygen adsorbed by the adjacent-to-surface layer of silver was equal to that for five monolayers. Hiroto and Kobajashi (62)report that even with prolonged reduction of silver in hydrogen a t 275" oxygen atoms are not removed from the metal. When the oxide film is fairly thick (about several tens of atomic layers), the chemical and electronic characteristics of the catalyst surface will be determined by the oxide film properties and the metal will exert no considerable effect on catalysis, whereas with a thin layer (of about several atomic layers) the catalytic properties will depend on the nature of the metal support. TABLE IV Characteristic Catalysts for Low Conversion of Hydrocarbons

Initial hydrocarbon

Catalyst Silver Cuprous oxide Vanadium pcntoxide

Molybdenum oxide

Ethylene Propene Propene Benzene Toluene o-Xylene Naphthalene Anthracene Butane

Main reaction products Ethylene oxide Acrolein Aldehydes Maleic anhydride Benzaldehyde Phthalic anhydride Phthalic anhydride Anthraquinone Maleic acid

TABLE V Characteristic Catalyats for High Conversion of Hydrocarbone

Catalyst Platinum, palladium Copper oxide Manganese dioxide Spinels (Co,Mn,Cr)

Initial hydrocarbon Methane, pentane, isooctane, ethylene, propene, acetylene, benzene, cyclohexane. Methane Acetylene Ethylene, propene, isooctane, cyclohexene, benzene

Recent developments in the theory of electronic structure of solids and more extensive data on the mechanism of physical phenomena in the solid phase revealed a relationship between coloration in the visible and the electronic structure of the solid. Elovich et al. (63)investigated the catalytic activities of various oxides with respect to high con-

CATALYTIC OXIDATION OF HYDROCARBONS

439

version of hydrocarbons, such as isooctane, cyclohexane, etc., as a function of the coloration of oxides and the valences of cations entering into the catalyst composition. Intensively colorated compounds display strong catalytic activity. From the standpoint of electronic concepts this dependence is indication that there is a certain class of catalytic processes occurring only with a oatalyst possessing free electrons or electrons readily passing to an excited or free state. The specificity of elementary steps of catalysts is determined above all by chemical bonds in the reactant molecule and by the structure of intermediate species appearing in the course of the reaction. There are yet no considerable achievements in this field and the choice of active oxidation catalysts is still based on extensive empirical information and on rather coarse approximations. The reactant mixture containing hydrocarbons of various structure and oxygen induces changes in the phase composition of the catalyst during the reaction. Simard et al. (64) suggested that the active surface of a vanadium catalyst represented a dynamic system involving Vs+, V4+and 0; ions. Roiter et al. (65) investigated a vanadium catalyst that was under operation in an industrial reactor for 2& years. They found that half the amount of V,O, of the upper layer was converted into V,O,. Thus the composition of a catalyst is controlled by the mixture composition. Similar results were obtained by various scientists (66, 67) in studying the behaviour of a cuprous oxide-copper oxide catalyst in the oxidation of propene into acrolein. The appearance of copper by further reduction of the catalyst results in a decreased selectivity of the process. According to Roginskii (68) such variations in the catalyst in the course of the reaction almost exclude possible modifications in the technique of preparation of active catalysts having to ensure the formation of structures with increased energy. The effect of reactants may be so great as to destroy these active structures. This is probably the reason for a certain primitivism in the choice of techniques for the preparation of oxidation catalysts, chief attention being paid to the surface area, pore sizes, etc.

111. Reaction Mechanisms The working out of an explicit scheme for hydrocarbon oxidation over various catalysts was hindered until lately due to insufficient development of the theory and to lack of extensive experimental data. Any catalytic mechanism implies that adsorption represents the

440

L. YA. MAROOLIS

primary step of catalysis and controls the transition of a reactant molecule to the active state. Molecules of oxygen or of a hydrocarbon are adsorbed on the catalyst surface during hydrocarbon oxidation. The state of these adsorbed molecules, their interaction, and their reactions with the gas phase molecules would account for different routes of the process.

A.

OXYQEN

ADSORPTION ON OXIDATIONCATALYSTS

Extensive investigations carried out by scientists of various countries (69-72) have shown that chemisorption of oxygen on metals may occur even a t relatively low temperatures ( - 80', 0'). According to Trapnell (73) the chemisorption of 'oxygen on various metals over a range of different temperatures is so fast as to make kinetic measurements impossible; this is indication that the activation energy for chemisorption is very low. Fast chemisorption is followed by slow uptake of oxygen by the metals. The dissolution of oxygen in nickel, copper and certain other metals results in the formation of oxides of these metals. With noble metals, such as platinum and silver, oxygen will be dissolved in layers adjacent to the surface, bringing about changes in the electronic properties of the latter. TABLE V I Activation Energies for the Heats of Adsorption and Kinetic Equation for Oxygen Chemieorption on Semiconductor Catalyets

Semiconductor

Heat of adsorption (kcclljmole)

Activation energy (kcal/mole)

Equation for adsorption kinetics

References

___.

-

NiO

36-62

2-8 7

43-70 25-64

16-18 32-35

Banghem RoginskiiZeldovich Bangham

(74)

~ 5 ~ 7 7 ) (76, 73)

The activation energy for adsorption of oxygen on silver and platinum varies with surface coverage, which is indication that the metal surfaces are heterogeneous and the oxygen-metal bonds have different energies. Oxygen bonds on the typical metal catalysts, namely platinum and silver, are of different strength, and two forms-molecular and atomiccoexist on the surface.

CATALYTIC OXIDATION O F HYDROCARBONS

44 1

The adsorption of oxygen on certain semiconductors, such as NiO and Cu,O was studied in fair detail, The activation energies, heats of adsorption and kinetic laws for oxygen sorption on simple semiconducting catalysts are summarized in Table VI. I n studying the chemisorption of oxygen on cuprous and copper oxides Garner et al. (77) found that this process is accompanied with slow incorporation of oxygen into the crystal lattice. Jennings and Stone (75) reported that complete coverage of Cu,O with oxygen took place at 20". The heat of oxygen chemisorption was 55-60 kcal/mole; oxygen was strongly fixed on the surface and represented a negatively charged monolayer. Adsorption of oxygen on V,O, was studied less explicitly. However, all scientists concerned report that it is insignificant. Clark and Berets (78) suggest that chemisorption of 0, on V,O, leads to a decrease in the concentration of surface defects. Margolis and Plyshevskaya (79) studied oxygen sorption on vanadium over a temperature range of 250 - 400°C. The kinetic curves obtained are described by Bangham's equation q = atl/"

(1)

Here q is the adsorbed amount, t is time, and a and n are constants. Roginskii (80) has shown that this is a typical kinetic equation for activated sorption on a heterogeneous surface characterized by the exponential shape of the distribution function with respect to activation energies for adsorption. The adsorption of oxygen on NiO occurs a t 200-360". According to Roginskii and Tzellinskaya (81)the kinetics of this process obeys the equation q =a

dT+ c

(where a and C are constants). This fact was associated with dissolution of oxygen in the NiO lattice. Keier and Kutzeva (82) have found by detailed investigation that the dissolution of oxygen in the NiO lattice may be disregarded and chemisorption may be considered as a process typical for a heterogeneous surface. With a coverage not exceeding 10% of the monolayer activation energies increase from 10 to 15 kcal/ mole. Thus oxygen is chemisorbed on all simple semiconducting catalystsmetal oxides-and in a number of cases it is partly dissolved in the lattice. Very little work on adsorption properties of complex semiconductors --spinels-was reported in literature. Magnesium and copper chro-

442

L. YA. MAROOLIS

mites, manganese cobaltite and cobalt manganite, as well as certain ferrites, were used as catalysts for high conversion of hydrocarbons. Margolis (83) studied the kinetics of oxygen adsorption at elevated temperatures close to or identical to the temperature of oxidation and found that the Bangham's law was followed in this case. An increase in coverage with oxygen from 12 to 20% resulted in a rise of activation energy from 8 to 25 kcal/mole. Linde (84)investigated oxygen adsorption on CoMn,O, and MnCo,O, spinels. The bilogarithmic law held for MnCo,O, (q = ~ t l ' ~and ) , the semilogarithmic law (q = u'e-aq) for CoMn,O,. This was another indication that the surfaces of these catalysts were heterogeneous. The activation energy for chemisorption of oxygen on MnCo,O, was found to vary exponentially with coverage, while uniform distribution of activation energies was observed for CoMn ,O,. A heterogeneous surface is characteristic of these catalysts. Thus chemisorption of oxygen on simple semiconductors and spinels involves changes in heats of adsorption and, consequently, in energies of oxygen-solid surface bonds.

B. ISOTOPIC OXYGENEXCHANGE The mobility of atoms or molecules of oxygen on the surface and in the lattice of the solids may be determined by using the isotopic exchange method. Isotopic oxygen exchange was studied in the last years by many investigators, of whom Winter (85))Vainstein and Turovskii (86, 87), and Boreskov and Kassatkina (88) have shown that oxygen of the metal oxide lattice is of a low mobility. Significant isotopic oxygen exchange on oxides occurs a t temperatures by 100-200° degrees higher than the temperature of catalysis. It may be seen from information available on isotopic oxygen exchange on silver and platinum that the values for exchange activation energies and for oxygen adsorption are close. Appreciable isotopic oxygen exchange on platinum begins only a t a temperature exceeding 350") whereas the oxidation of hydrocarbons on platinum is fast even at 20". The strongly bound oxygen does not seem to participate in catalytic oxidation. Oxygen is more mobile on silver than on platinum, since oxygen exchange on platinum sets in at 220") while catalytic oxidation may readily occur at 200". Comparison of data on adsorption and exchange on silver shows that under conditions of catalysis only a part of the oxygen adsorbed is mobile. The oxygen dissolved in the adjacent-to-surface layer does not seem to participate in catalysis.

CATALYTIC OXIDATION OF HYDROCARBONS

443

The nature of intermediates formed in adsorption may be established from investigation of oxygen exchange. The formation and breaking of chemical bonds during isotopic exchange may occur with and without unpairing of electrons (homolytic and heterolytic exchange, respectively). Investigation of homolytic oxygen exchange on various solids would show whether oxygen dissociation occurs on the surface. Margolis (74)studied homolytic oxygen exchange on platinum and silver, and on vanadium pentoxide and manganese dioxide. No homolytic exchange was observed at low temperatures for platinum, silver and semiconducting oxides (MnO, and V,O,), as well as for the spineltype compounds. Certain information on the mechanism of exchange and the nature of active intermediates may be obtained from comparison of rate constants for heterolytic and homolytic exchange involving adsorbed oxygen with the desorption constant, at the same temperature and coverage. The rate constants for homolytic exchange on silver at 232" are lower than those for heterolytic exchange by a factor of 2.5. The rate constant of desorption a t this temperature is equal t o that of oxygen exchange. This is indication that the dissociation of oxygen molecules on silver a t 232" occurs only in part. As the temperature is raised, the rate constants for homolytic isotopic exchange become close to those for ordinary isotopic exchange and desorption, and a t 290" these are almost equal. It may be inferred that under such conditions the oxygen adsorbed on silver dissociates into atoms. For platinum this equality of rate constants is observed only a t temperatures higher than 400". Phenomena similar to those observed for silver take place at lower temperatures. With semiconducting oxides such as MnO,, V,O,, etc., oxygen exchange may be observed only at elevated temperatures. Consequently, it is yet impossible to find out whether oxygen dissociates a t the temperature of catalysis. Dzisyak et al, (89) investigated homomolecular oxygen exchange on V,O, proceeding a t a temperatwe higher than 450". The initial rate of oxygen adsorption was shown to be by a power of ten higher than the rate of exchange. Correlation in rates, activation energies, and reaction orders with respect to oxygen indicates that the dissociation of oxygen into atoms is the rate-limiting step common to homomolecular and to isotopic exchange, The above mentioned scientists reported data on exchange at 500-550", when the occurrence of oxygen molecules a t the surface is scarcely probable and, consequently, only the dissociative mechanism of chemisorption and exchange would be valid under these conditions. A t lower temperatures (350-450"), such as those for

444

L. YA. MAROOLIS

catalysis on vanadium pentoxide, both the molecular and atomio oxygen forms may coexist. It may be seen from investigation of homolytic oxygen exchange on various oxidation catalysts that oxygen may be present on the surface as molecules without dissociating into atoms. The ratio of molecular to atomic oxygen content of the surface is a function of temperature. OF HYDROCARBONS ON OXIDATION CATALYSTS C. CHEMISORPTION Chemisorption of hydrocarbons on various metals, such as nickel, platinum, copper, etc., was investigated in great detail (9, 90, 91, 92). Information on chemisorption of ethylene, acetylene and methane on various metals may be found in Trapnell's review (93). However, direct application of the relations obtained to metal oxide catalysts would scarcely be justifiable, As a rule, oxygen covers the whole surface of the metal, and chemisorption of hydrocarbons occurs either on a thin layer of the given metal oxide formed as an individual phase, or on oxygen that was sorbed on the surface and has filled the adjacent-to-surface layers. Thus data on chemisorption of hydrocarbons on oxides of these metals may be of use in the above cases. Margolis (94) studied ethylene chemisorption on silver proper and on silver covered with oxygen. When the silver sudace is completely oxygen-free, ethylene sorbs in equilibrium and reversibly, and the coverage degree is low. Things are quite different when hydrocarbons are adsorbed on silver covered with oxygen, Adsorption occurs in time; adsorption kinetics follows the Roginskii-Zel'dovich equation characteristic for heterogeneous surfaces. Only a certain amount of the ethylene sorbed may be removed by outgassing, a considerable part of it becomes strongly bound with oxygen of the surface. The activation energy for the over-all process is very low, 2-3 kcal/mole. Butyagin and Elovich (12) studied propene adsorption on platinum and found that the oxygen-covered platinum surface takes up propene, and the rate of this process follows an exponential law. The effect of the metalsurface heterogeneity is very marked for oxygen-covered metals, probably due to regions with different electronic potentials. Chemisorption of hydrocarbons on semiconducting oxides was not investigated extensively. Turkevich, Howard, and Taylor (95) studied ethylene, ethane, and propane adsorption on MOO,, Cr,O, over a temperature range of 0-400" and established the boundaries between physical and chemical adsorption. They emphasized the difficulties arising in the investigation of hydrocarbon adsorption on metal oxides, due to side processes occurring at high temperatures, for instance pyrolysis and oxidation.

CATALYTIC OXIDATION OF HYDROCARBONS

445

The chemisorption of unsaturated and saturated hydrocarbons on various simple semiconducting oxides, namely vanadium pentoxide and cuprous oxide, and on spinels (CuCr,O, and MgCr,O,) was investigated by Margolis (94). The rates of hydrocarbon adsorption on all oxidation catalysts (V,O,, Cu,O, etc.) are so high as to make kinetic studies impossible. Characteristic “equilibrium” isotherms for ethylene sorption on magnesium chromite are shown in Fig. 1. This is a spurious equilibrium,

I00

200

FIQ.1. Isotherms for ethylene adsorption on magnesium chromite at 110”;

1-

primary, 2-secondary.

as at the given temperature the process is only partly reversible. The hydrocarbon will not desorb from the catalyst even with repeated intermediate high-temperature training. The catalyst surface may be freed from the strongly sorbed hydrocarbon only by treating it with oxygen. Isotherms for propene adsorption on NiO, vanadium pentoxide, and cuprous oxide at 100” are shown in Fig. 2. All isotherms are described by Freundlich’s equation and correspond to a heterogeneous surface with exponential distribution as to heats of adsorption. The structures of intermediates appearing in adsorption of hydrocarbons are yet unknown. Voevodskii, Vol’kenstein, and Semenov (96) consider that the generation of free radicals in chemisorption is accounted for by free valences present on the solid surface. According to Syrkin (97)the chemisorption mechanism involves the formation of a three-center bond by sorption of unsaturated hydro-

446

L. Y A . MARGOLIS

carbon molecules on metals (Fe, Co, Ni, Pt, etc.) and semiconductors (chromium oxide). The formation of donor-acceptor bonds involves electrons of the solid and of the adsorbate, as well as the electrondonating bond formed between the antibonding orbit of the adsorbed molecule and the metal atom orbit occupied by an electron pair,

FIG.2. Isotherms for propene adsorption on NiO ( l ) ,vanadium pentoxide (2), and ouprous oxide (3) at 100".

I). CHARGINGOF

THE

SURFACE IN ADSORPTION

A number of works are concerned with charging of the surface in chemisorption (98-101).Chemisorption on metals and semiconductors is accompanied by electron transfer between the adsorbed molecule and the catalyst. The direction of electron transfer depends on the Fermi level position in the crystal, and on the energy level of the chemisorbed molecule. The Fermi level position on a semiconducting surface may be determined from the change in the electron work function. Adsorption of acceptor molecules results in an increase in the electron work function, while a fall-off is observed for donor molecules. Measurement of the electron work function during adaorption on solids permits determination of the charge sign for the adsorbed molecule (see Table VII). Thus various hydrocarbon molecules are donors when adsorbed on simple and complex semi-conductors, such as NiO, MnO,, Cr203,Cu,O, MnCo,O,, and CoMn,O,. Margolis (202) found from determination of the electron work function for silver during oxygen and ethylene adsorption that the oxygen molecule obtains a negative charge and is an electron acceptor, while the positively charged ethylene molecule is a donor. The electron

TABLE VII Charging of the Surfaces of Metal Semiconductors During Adsorptiun of Hydrocarbons and Oxygen (103) Catalyst ~

NiO

Hydrocarbon Electroconductivity Charge with respect to the work function

Propene Propane Decrease

+

p-type semiconductors Cr*O, MnO

,

cu,o

Butane

Ethane

Propene

Decrease

Decrease

Increase

+

+

+

Spinels MnC0,Ol C0Mn,Od Propene Not

measured

+

n-type semiconductors

Metals

Ag

Propene Propane Increase

+

Pt

Propene

Ethylene

Ethylene

Increase

Increase

Not measured

not determined

+

+

448

L. YA. MARQOLIS

work function values for oxygen and propene adsorption on platinum are close to those for silver. Thus for the majority of catalysts used the adsorption of oxygen molecules results in a negative charge of the surface, while a positive charge is observed for hydrocarbons irrespective of their structure.

E. OXIDATIONSCHEMES The working out of schemes for catalytic oxidation of hydrocarbons on metals and semiconductors is difficult due to the complexity of reactions, the diversity of products formed and the lack of information on the nature of elementary acts. The genetic relationships between oxidation products are much easier to determine. The high conversion to CO, CO,, and H,O hinders the formation of oxygen-containing products in direct oxidation of hydrocarbons with molecular oxygen. It is considered as firmly established that the reason for low selectivity of catalytic hydrocarbon oxidations lies in the ready subsequent oxidation and decomposition of oxygen-containing products formed by synthesis. The accepted scheme for olefine oxidation is olefine + aldehyde -+ acid -+ CO -+ CO f hefine oxide ------+ aldehyde

Pongratz (103)) Sosin (104), and Dolgov (105) suggest that the oxidation of aromatics to phthalic and benzoic aldehydes proceeds in steps. For example, naphthalene is converted into naphthol, then t o naphthahydroquinone, to phthalic anhydride, and, eventually, to CO 2. Thus optimal conditions for fast elimination of oxygen-containing compounds from the reaction zone are usually sought for in the attempts to obtain products of low conversion. Under certain conditions low selectivity may be due to unfavorable formation to conversion ratios for stable oxygen-containing intermediates, such as aldehydes, olefine oxides, acids, which seem to prevent the accumulation of valuable products. However, this reason cannot be the sole and general one. For instance, with typical catalysts for high conversion of hydrocarbons, such as magnesium and copper chromites, aldehydes undergo oxidation a t 200-400" mainly to acids and only in part to CO,. Under the same conditions hydrocarbons oxidize to water and carbon dioxide, minor amounts of aldehydes and acids being found in reaction products. Charlot (106) investigated the activity of a great number of various metal oxides in the reaction of toluene with oxygen, and stated that the high and low conversion processes are independent of each other.

CATALYTIC OXIDATION OF HYDROCARBONS

449

Marek ( 4 )believes that a catalyst accelerating the first step will accelerate all the subsequent steps as well, since it is difficult to conceive a catalyst that would accelerate the oxidation of hydrocarbons and have no effect on the oxidation of formaldehyde to water and carbon dioxide. The ratio of parallel to consecutive reaction rates is a function of the hydrocarbon structure, the catalyst type, and the reaction temperature. I n studying the kinetics of catalytic benzene oxidation to maleic anhydride Hammar (107) came to the conclusion that benzene oxidizes by two independent routes: the formation of maleic anhydride, and the complete combustion to carbon dioxide and water via unidentified intermediates. He suggested the following scheme for benzene oxidation over a vanadium catalyst:

+ 0, + X I -+ C4H,0, + CO + CO, + H,O C,H,O, + 0, -+ X, -+ CO + CO, + H,O C,H,

+

COHO 0, -+ X, -+ CO

+ CO, + H,O

where X,, X,, and X, are hypothetic intermediates. Ioffe (108) and later on d’alessandro and Parkas (42) have found from kinetic studies that heterogeneous catalytic oxidation of aromatic hydrocarbons involves a number of parallel and parallel-consecutive reactions, and that high conversion does not necessarily include the formation of low conversion products. For example

*

Anapht haquinone

naphthalene

-

C\O,

phthalic anhydride H20

+

-

maleic anhydride

Naphthalene oxidizes by three independent routes to naphthaquinone, phthalic anhydride and carbon dioxide. The latter may also be formed by oxidation of naphthaquinone and phthalic and maleic anhydrides. Ushakova, Korneichuk, and Roiter (109) suggest that various reactions of naphthalene oxidation proceed independently (Fig. 3). Unpublished results obtained by Simard et al. on the oxidation of o-xylene on V20, by a supposedly parallel-consecutive mechanism are mentioned in a review by Longfield and Dixon (110). Suvorov, Rafikov, and Anuchina (111) came to the conclusion that the first step of toluene oxidation over vanadium oxides yields hydroperoxide which may decompose by various routes. The formation of high-conversion products does not occur in parallel. Bretton, Wan, and Dodge (112) investigated the composition of products formed in the oxidation of a hydrocarbon containing four

450

L. YA. MAROOLIS

carbon atoms over V,O, as a function of the hydrocarbon structure. They suggested that the diversity of products formed is due to different structures of radicals yielded by oxidation. Thus a number of reactions proceeding a t different rates occur simultaneously on semiconductors, such as metal oxides.

+ H20

FIG. 3. Scheme of naphthalene oxidation over vanadium pentoxide, according to Roiter.

The most widely used metal oxidation catalysts are platinum, palladium, copper, and silver. With all these, except silver, hydrocarbons undergo oxidation with complete destruction of the molecular skeleton into CO, and H,O. Silver is the sole catalyst for obtaining ethylene

4

FIG.4. Scheme of ethylene oxidation to ethylene oxide over silver, according to Twigg.

CATALYTIC OXIDATION O F HYDROCARBONS

45 1

oxide from ethylene. This reaction was the object of many investigations and several oxidation schemes were proposed. Twigg (18) suggested on the basis of extensive kinetic studies that oxygen adsorbed on a silver catalyst dissociates into atoms; ethylene does not sorb on silver (Fig. 4). When ethylene comes into contact with two oxygen atoms, a number of reactions occur (Fig. 4). The two formaldehyde molecules formed undergo subsequent conversions CH,O -+ CO (ads)

+ 2 H (ads)

CO (ads)

+ 0 (ads) -+

2 H (ads)

+ 0 (ads) + H,O

CO,

According to Twigg, when an ethylene molecule comes into contact with one oxygen atom, the interaction H,C=CH,

0

/ \ Ag

Ag

-

H ,C-CH

v 0

I

ethylene oxide

Ag - Ag

takes place. Ethylene oxide may isomerize to acetaldehyde which is readily oxidized on silver to CO and H ,O. The oxidation of ethylene oxide on silver yields carbon dioxide and water, but the amounts of these are not equivalent to C,H,O consumption. Twigg believes that this is accounted for by a n adsorbed organic residue formed on the catalyst surface. Ethylene also detected in oxidation products was thought to be formed by ethylene oxide decomposition to ethylene and adsorbed oxygen. Todes arid Adrianova (113)studied the kinetics of ethylene oxidation over pumice-supported silver and suggested that ethylene was converted to ethylene oxide, which yielded C O , and H,O. I n other words two consecutive reactions took place

,

+ 0 -+ C,H,O C,H,O + 2 0 , -+ 2 C 0 , + 2H,O

C,H,

and

Orzechowski and McCormak (21) proposed the following kinetic mechanism of ethylene oxidation (Fig. 5 ) . Two parallel reactions were believed to occur: the formation of ethylene oxide, and the formation of carbon dioxide and water from ethylene. These reactions would be accompanied by slow decomposition of

L. YA. MARGOLIS

462

ethylene oxide which might be sorbed on silver: C*H,Oo,

I::

+ 0 - Ag + 0

\

. . . 0 . . . Ag -+ X -+COa + Ha0 + Ag

(X is an intermediate, possibly formaldehyde). As the temperature is raised, the ratio of formation to oxidation rates of oxygen-containing products (aldehydes, ethylene oxide) suffers a change and consecutive oxidation takes place under certain conditions.

C2H40

+ Ag

y+Ag-COz

+H20

FIG.6. Scheme of ethylene oxidation over silver, according to Orzechowski and McCormak.

The contribution of consecutive reactions is especially important when the oxygen-containing compound formed is very reactive. The kinetic reaction scheme shown in Fig. 6 was proposed by Isaev, Margolis, and Roginskii (114).The main route of the reaction was supposed to be propene -+ acrolein -+ COa

i.e., a consecutive conversion of the hydrocarbon to aldehyde and carbon dioxide. Information on propene oxidation to acrolein over a copper catalyst was published recently by Belousov et aE. (115).They consider that at a low temperature (320") acrolein and CO, are formed mainly by two parallel reaction routes acrolein propene

1 carbon dioxide

A parallel-consecutive scheme holds a t higher temperatures of 350 to 400". The relationship between different oxidation routes is a function not only of temperature, but of the catalyst surface condition as well. Isaev, Golovina, and Sakharov (116)have shown that as the surface becomes treated with the reactants, the contribution of parallel propene

CATALYTIC OXIDATION O F HYDROCARBONS

453

oxidation to the formation of CO, becomes less important. Assuming that the oxygen-containing products formed (aldehydes, ethylene oxide, etc.) readily undergo subsequent reactions with oxygen, these substances may be considered as intermediates in the formation of carbon dioxide, carbon monoxide and water.

FIG.6. Scheme of propene oxidation to acrolein over cuprous oxide.

The nature and importance of intermediates may be established by kinetic studies through comparison of formation, decomposition and oxidation rates. At present the mechanism of complex processes may be investigated by the more precise radioactive tracer technique. The genetic relationship between individual reaction products may be determined by using this technique together with kinetic studies. The origin of reaction products and the rates of individual reactions involved in the overall process may be established by C144abeling of the hydrocarbon molecules, or of the products suggested as intermediates. Neiman (117) worked out an interesting kinetic tracer method for determining the reaction kinetics from variations in radioactivity and concentration of intermediates with time. When a substance A is converted into X, and then X is converted into B (A + X + B), then the concentration and specific radioactivity of X may be determined a t any time by labeling A or X and adding the labeled substance to the reactant mixture. The rates of formation and consumption kx follow the equation

_ -aaat

(/3 -

-

a) w

z

specific activity of X; /Ithe specific activity of the initial product A; x the concentration of X.

a is the

454

L. YA. MARGOUS

WhenB = 0, then

ax - _ = W - k x

(4) at When a~ = 0, then -dx{dt = 0 and w = 0. Consequently, there is only consumption of X without its formation. When 01 changes with time, w is found before k x . 1. Ethylene Oxidation over Silver

When ethylene is oxidized over Ag, carbon dioxide may be formed either directly from ethylene, bypassing ethylene oxide, or by stepwise oxidation of ethylene oxide. To obtain confirmation for the occurrence of these reactions Margolis and Roginskii (118) carried out investigations with a mixture of C’*-labeled ethylene, ethylene oxide, and oxygen. Variations in concentrations of initial compounds and reaction products as a function of the contact time were determined, as well as the activity distribution. Direct quantitative and systematic information on the carbon dioxide formed via ethylene oxide and ethylene was obtained for the first time by means of the kinetic tracer method. Further discussion thus became unnecessary. Adrianova and Todes (113) and also Orzechowski and McCormak (21)allowed for one process of carbon dioxide generation only. This was not correct as carbon dioxide is always formed by two independent routes. However, having accepted this fact, there is no need in considering, as Twigg does (18),the predominance of one of these processes over the other, or to estimate the contribution of each process to losses of ethylene as carbon dioxide. The part played by stepwise oxidation via ethylene oxide, or by “direct” oxidation bypassing ethylene oxide, in the process of CO formation is different for various silver catalysts. For certain catalysts stepwise oxidation will be predominant a t low temperatures, and for others ‘(direct” oxidation will be more important. But this is also conditional, as the contribution of carbon dioxide formed by any route is dependent to a considerable extent on external factors, first of all on temperature. The CO Bstcp to CO edir ratio increases exponentially with temperature. Many workers consider that ethylene oxide isomerizes to acetaldehyde (18). Twigg suggests that formaldehyde and acetaldehyde are intermediates both in the stepwise and in the “direct” formation of carbon dioxide. The contribution of aldehydes to oxidation processes may be found out by oxidizing ethylene-acetaldehyde and ethylene-formalde-

CATALYTIC OXIDATION OB HYDROCARBONS

455

hyde mixtures under static and dynamic conditions and tracing the distribution of activity among various reaction products (118). The oxidation and adsorption of ethylene oxide on a silver catalyst is stronger in the presence of acetaldehyde, and the former hinders the formation of carbon dioxide from acetaldehyde. The latter fact is probably due to interaction between acetaldehyde and ethylene oxide a t the surface. A conjugated oxidation of this kind is often encountered with homogeneous reactions. The oxidation of mixtures of ethylene with ethylene oxide, formaldehyde and acetaldehyde on a stationary surface of a silver catalyst that was under operation for several hundred hours is of a different nature. The reaction rates for mixtures of radioactive ethylene with ethylene oxide, as well as with formaldehyde and acetaldehyde, are shown in Fig. 7. Oxidation proceeded under dynamic conditions a t 220".

':I

3

10

FIQ. 7. Diagram of formation rates ( w ) for ethylene oxide and carbon dioxide in the oxidation of ethylene and its mixtures with aldehyde over silver at 220'; w is given in arbitrary units.

I n the presence of ethylene oxide, formaldehyde, and acetaldehyde, the rate of ethylene consumption remains unchanged. There is also no inhibition of ethylene oxide formation a t its gas-phase concentration of about 1yo.This is probably due to the blocking off of the most active surface sites. At a prolonged use of the catalyst with an ethylene-formaldehyde mixture the catalyst activity gradually decreases and complete poisoning of the catalyst takes place. The acetaldehyde effect is similar to that of formaldehyde. All oxygen-containing compounds such as formaldehyde, acetalde-

466

L. YA. MARQOLIS

hyde, and ethylene oxide are taken up by the catalyst surface to form an adsorbed residue (Twigg calls it the "organic residue"). The removal of carbon dioxide by heating the oxygen-containing compounds in the absence of gas-phase oxygen is indication that the residue contains both oxygen and carbon atoms. The silver activity varies until the complete formation of the residue, after which it becomes constant. It will be noted that inhibition of ethylene oxidation with the ethylene oxide formed is associated with the blocking off effect of the residue. No inhibiting effect of ethylene oxide is observed for stationary catalyst surfaces. Various organic compounds display a different capacity for forming the surface residue. The compounds most readily adsorbed on the surface are CH,O, CH,CHO, and C,H,O. The contribution of aldehydes to catalytic oxidation is of great interest. Aldehydes are usually considered as primary oxidation products yielding the majority of other products. Isotopic analysis of ethylene oxidation over V,O, shows that neither acetaldehyde nor ethylene oxide can be the main intermediates in the formation of carbon dioxide from mixtures of ethylene with various oxygen-containing compounds (Fig. 8).

co

-

-

7

0,

FIQ. 8. Diagram of formation rates ( w ) for aldehydes, carbon monoxide, and carbon dioxide in the oxidation of ethylene and its mixtures with ethylene oxide and acetaldehyde over V,O, at 400";w is given in arbitrary units.

When ethylene is oxidized on vanadium oxides, a considerable amount of carbon dioxide and carbon monoxide is yielded directly by the hydrocarbon, bypassing acetaldehyde and ethylene oxide. It was found by investigating the oxidation of an acetaldehyde-ethylene oxide

CATALYTIC OXIDATION O F HYDROCARBONS

467

mixture that neither of these components was formed as the main intermediate in the formation of carbon dioxide and carbon monoxide. When ethylene is oxidized over Ag or V,O, the consecutive formation of several stable oxygen-containing compounds either does not occur at all, or is of secondary importance. It may be seen from comparison of results on ethylene oxidation over silver and vanadium pentoxide that with both catalysts the oxidation of unsaturated hydrocarbons will proceed by the same mechanism. CO, generation is not accelerated in the presence of aldehydes and these cannot be intermediates in ethylene combustion. When aldehydes are introduced into the reactant mixture, the ratio of ethylene oxide to CO, formation rates undergoes a change, due to strong adsorption of aldehydes on the catalyst surface. Ethylene oxide will form on silver and is in fact absent on vanadium oxides. It was shown experimentally that the absence of acetaldehyde and formaldehyde in the products of oxidation over silver, and the low absolute content of these substances for vanadium oxides is due to the fact that they are not formed a t all, or formed at a low rate, and not to their oxidation or decomposition. 2 . Oxidation of Propene Gorokhovatskii and Rubanik (119) came to the conclusion that the catalytic combustion of propene over silver does not proceed via the propene oxide formation step, since the rate of propene oxide conversion a t 240" is considerably lower than that for propene. Zimakov (120) suggested earlier that the impossibility of obtaining propene oxide from propene over silver was due to peculiarities of the propene oxide structure and the readiness of its further oxidation. However, Gorokhovatskii and Rubanik have shown that this is not so. Adsorption of propene on the silver surface seems to be different from ethylene adsorption. De Boer, Eischens, and Pliskin (121) suggest that ethylene sorbs on a silver surface covered with oxygen to form complexes

which may readily convert into ethylene oxide (the strong reactivity of hydrogen atoms contained in the CH, group, the so-called 6 - rr

458

L. YA. MAROOLIS

conjugation, was established for propene). Thus the probability is great that propene is adsorbed on the silver surface with abstraction of a hydrogen atom and the formation of OH O-CH+2H=CH,

I I

Ag - Ag

Under the action of oxygen this compound readily oxidizes to carbon dioxide and water. This might be the main reason for absence of propene oxide in the oxidation of propene over silver. This question might have been solved by spectroscopic investigation of the adsorbed propene structure. Acrolein and carbon dioxide are formed in propene oxidation over cuprous oxide. Minor amounts of acetaldehyde and formaldehyde were found in reaction products, and this is indication that the reaction proceeds by breaking of the propene double bond. The oxidation of unsaturated hydrocarbons over other metal oxides, such as V,O,, WO,, etc., is of the same nature. Let us consider what are the intermediates in the formation of carbon dioxide. Isaev, Margolis, and Sazonova (116) attempted to elucidate the mechanism of propene oxidation to acrolein using the kinetic tracer method. The specific radioactivity of carbon dioxide was found t o be higher than that of acrolein, but lower than the propene radioactivity. If carbon dioxide and acrolein were yielded by propene separately, the specific radioactivity of CO, would be identical to that of propene. It follows from radiometric data that, within experimental error, there is no parallel route of carbon dioxide generation. It will be noted that with a fresh catalyst having a non-stationary surface up to 30% of carbon dioxide are yielded directly by propene, bypassing acrolein. Thus the parallel-consecutive scheme of propene oxidation to acrolein over cuprous oxide was revealed by the kinetic tracer method. 3. Mechanism of Catalytic Oxidation

Isotopic investigations of various reaction mechanisms yielded the scheme for catalytic oxidation of hydrocarbons shown in Fig. 9. This scheme is of great help in the proper choice of catalysts. The first task in improving the results of hydrocarbon oxidation is to find catalysts ensuring the necessary reaction route. Parallel reaction routes are to be suppressed by adjusting the chemical composition of the catalyst through addition of donor and acceptor impurities. Due to lack of experimental information on the nature of elemen-

CATALYTIC OXIDATION OF HYDROCARBONS

45 9

tary steps in the catalytic oxidation of hydrocarbons, it was difficult to establish the true mechanism of catalytic heterogeneous reactions. However, it was suggested by a number of scientists, by analogy with reactions occurring in the homogeneous phase, that catalytic oxidation of hydrocarbons proceeded via the generation and interaction of radicals. Schultze and Theile (122) pointed to the possibility of interaction between ethylene and oxygen molecules to form peroxides. Pokrovskii (123) reported a scheme of catalytic ethylene oxidation to ethylene oxide involving radicals.

Fro. 9. Scheme of Catalytic hydrocarbon oxidation; H-hydrocarbon, R, to R,-labile intermediates,probably of the peroxide type.

C-catalyst,

Lyubarskii (124) believed that upon interaction with molecular oxygen ethylene was converted to ethylene oxide. Sosin (104)suggested that in the oxidation over V,O, aromatic hydrocarbons lose hydrogen and convert into radicals to form a peroxide radical by interaction with oxygen. Bretton et al. (112) consider that the catalyst effect is one of abstracting the hydrogen atom from the hydrocarbon molecule to form a radical converting into a peroxide radical, and then into a hydroperoxide. The high rates and low activation energies of heterogeneous catalytic processes made many scientists believe long ago that active intermediates play an important part in these processes. Semenov and Voevodskii (125, 126) reported that radicals were yielded by reactions between gaseous molecules and the surface. The electronic theory of catalysis also implies that reactive radicals are generated on the catalyst surface by adsorption of molecules. Let us summarize the essential points of the hydrocarbon oxidation scheme, as derived from published data and from the electronic theory of catalysis. (1) A molecule with a double bond adsorbed on a semiconducting catalyst surface converts into a radical bound with the lattice and having a free valence. A molecule with a single bond emerging from the gas phase may react with the free valence of such a radical and dissociate. (2) An adsorbed saturated molecule with a single bond may dissociate into two radicals, one saturated with the surface valence, and

460

L. YA. MAROOLJS

the other having a free valence; free radicals will be generated by desorption of the latter radical into the gas phase. (3) It may be considered from isotopic data and electron work function measurement that negatively charged ions of molecular and atomic oxygen are present on semiconducting surfaces. The ratio of these is a function of temperature and the chemical properties of the solid. (4) Hydrocarbons are sorbed on semiconducting catalysts either weakly-reversibly, or strongly-irreversibly . The ratio of weak to strong adsorption is a function of temperature and the chemical composition of the catalyst. ( 5 ) Various types of ion-radicals are formed in adsorption of reactant molecules on the semiconducting surface; the formation of these is a function of electronic properties of the solid, and the structure and kind of bonds. (6) The catalyst surface is markedly heterogeneous both with respect to oxygen and to hydrocarbon adsorption. ( 7 ) The heterogeneous-homogeneous step occurs only for certain catalysts, such as platinum and spinels, and is not observed with oxide catalysts over the temperature range up to 400'. (8) Reaction products, such as aldehydes, olefine oxides, etc., are strongly sorbed on the catalyst surface, contributing to formation of the organic residue and representing additional sources of carbon dioxide generation. (9) It may be considered on the basis of data obtained by means of the radioactive tracer technique that the various stable oxygen-containing products on semiconducting oxides are generated by different routes, through active intermediates. (10) The oxygen in metal oxide lattices, as well as that sorbed on lattice surfaces is of a low mobility. Under certain conditions the hydrocarbon will react with oxygen of the catalyst lattice. A t low temperatures this side reaction is of small importance for oxidation. Scientists working in catalysis were long ago concerned with the theory of reduction-oxidation mechanisms of catalytic oxidation reactions. The nature of oxygen contribution to the oxidation process may be established by using the radioactive tracer technique. The reductionoxidation mechanism of reactions is a function of the lattice oxygen mobility. Vainstein and Turovskii (86) investigated the distribution of 0 1 8 in the oxidation of carbon monoxide over MnO, and CuO and found that there will be no transfer of the catalyst oxygen to the reaction producte, when water is carefully removed from the solid oxides.

CATALYTIC OXIDATION OF HYDROCARBONS

46 1

Vassiliev et al. (127)found that CO, exchanges oxygen with the MnO, surface even a t low pressures. The rate constants for CO, exchange and reduction of the surface are almost equal, whereas those for CO oxidation are by a power of ten higher. Kinetic and isotopic studies show that oxidation of carbon monoxide over V,O, cannot be considered solely as alternate reduction-oxidation of the surface. Vanadium pentoxide exchanges its oxygen with oxygen of the gas phase at temperatures above 45OoC, and catalytic oxidation proceeds within the same temperature range. Roiter et al. (128) compared the rates of isotopic oxygen exchange and catalytic oxidation of sulfur dioxide. They found that a t 500" the exchange rate is 10 times lower than that of oxidative-reductive catalysis. The rate of CO, oxidation by alternate reduction-oxidation over V,O, will be many times lower than that of oxygen exchange, and the latter is, in its turn, 10 times lower than catalysis with the given catalyst. Consequently, it may be seen from comparison of rates that the process is not a reduction-oxidation one in this case as well. Even more convincing is the comparison of exchange and catalytic oxidation rates for naphthalene. Naphthalene oxidation proceeds a t 400"; exchange over V,O, at this temperature will be considerably slower than the oxidation reaction, and the difference in rates of these processes cannot be explained on the basis of the reduction-oxidation scheme. Kinetic isotopic analysis of different reactions, such as the oxidation of CO, SO,, and hydrocarbons of various structures over MnO,, V,O,, etc., shows that the reduction-oxidation mechanism cannot be considered as the main route of these processes. Certain reaction schemes making allowance for surface charging will be discussed. Peroxidt radicals are chain-carrying centres in the homogeneous oxidation of hydrocarbons. The leading part in heterogeneous oxidation is played by peroxide ion-radicals (Margolis, 129). Let us consider three types of interaction between: (1) adsorbed hydrocarbon ion-radicals and oxygen of t h e gas phase; (2) adsorbed oxygen and hydrocarbon of the gas phase (oxygen adsorbed both with and without dissociation); (3) adsorbed oxygen and adsorbed hydrocarbon. Saturated aldehydes and acids containing less carbon atoms than in the molecule of initial hydrocarbon, as well as carbon dioxide and water, are formed in the first case. The second type of interaction yields unsaturated aldehydes, olefine oxides, carbon monoxide, carbon dioxide, and water for the oxidation of unsaturated hydrocarbons; and saturated aldehydes, carbon dioxide, carbon monoxide, and water for the oxidation of saturated hydrocarbons. The third type of reaction gives

L. YA. MAROOLIS

462

unsaturated aldehydes, carbon monoxide, carbon dioxide and water (olefine oxidation). The latter type of interaction would be impossible in the oxidation of saturated hydrocarbons. Radical oxidation schemes for unsaturated and saturated hydrocarbons may be derived on the basis of investigation of various reaction types. For instance:

+ O , + L - 0, 2L + 0 , + 2 L - it

1. L

2.

+ C,H, -+ L - H + isH, 4. L + 6 , -+ ~L ~ C,H, 6. L - 6, + C , H , - + L - 0 - O H + &H, 6. L - 0 , + CSH, -+ L - C,H7 3. L

Op

7.

L - 0, - CsH,

=

(a) CH,CHO (b) C H I 0 ( 0 ) COs, H , O

- 0 , + L - C,H7 -+ L - O O H + L - C,H, 9. L - CSH, + 0 , -+ L - C,H, - 0 , 8. L

10. L

- C,H, - 0, -+ (a) CH,CHO (b) CH*O CH,COOH (d) H C O O H (8) CO, CO,, Ha0

(c)

11. L

- C,H, - 0 , + C3H, -+

polymerization

All types of reactions between oxygen and hydrocarbons yield oxygen-containing compounds, such as aldehydes, acids, etc., present together with the products of complete combustion, i.e., with carbon monoxide and water. The reaction selectivity seems to be determined by the strength of bonding between the surface and the ion-radicals formed, and may be increased solely by changing the chemical composition of the catalyst. Let us consider schemes for certain oxidation reactions over metals and semiconducting catalysts. 4. Metals a. Ethylene Oxidation to Ethylene Oxide over Silver. The scheme for ethylene oxidation over silver suggested by Margolis in 1957 (118, 129) is based on the following considerations. It was found by investigating homolytic exchange on silver a t temperatures of catalytic ethylene oxidation that the surface-adsorbed oxygen may be present as mole-

CATALYTIC OXIDATION OF HYDROCARBONS

463

cules and atoms. Measurement of electrical conductivity and the electron work function of silver in oxygen adsorption shows that the surface oxygen is of a negative charge. Ethylene will not sorb chemically on a silver surface free of oxygen and its electrical conductivity will remain unchanged. Oxygen-containing surface ethylene will be sorbed obtaining a positive charge, and sorption kinetics will follow an equation characteristic of a heterogeneous surface. It was found by isotopic studies that ethylene and carbon dioxide were formed independently and by parallel reactions, unstable oxygencontaining products being intermediates. Aldehydes cannot be the main intermediates in the generation of CO, and H,O. Due to strong adsorption of ethylene oxide, an organic film will be formed on the silver surface, blocking off the active centers. I n accord to Enikeev et al. (102) variations in the electron work function ( A 4 ) in adsorption of the above compounds on silver are indication that ethylene and ethylene oxide are electron donors, and oxygen and CO, electron acceptors. Water would slightly decrease the +value. Following reactions occur a t the Ag surface : C,H,

-e

CaH40

-e

+e

CO,

-+ (CaH4)+ -+ (C,H40Jf -+ (C0,)-[possibly (CO,)-]

Carbon dioxide may be sorbed reversibly only on an oxygen-covered surface, probably to form a (C0,)- complex. Twigg has shown that a t temperatures about 200 to 250°C ethylene oxide will decompose in part to ethylene and adsorbed oxygen. Variations in the contact potential difference in adsorption of ethylene oxide at the above temperatures are indication that the electron work function will continuously increase with surface coverage. The following scheme of ethylene oxidation over silver may be proposed on the basis of experimental information available: Ag

+ 0 , -+ Ag - ( 0 a ) -

2Ag Ag

+ 0,-+ 2Ag - (0)-

- Oa + CaH4 -+ Ag - (0, - CaH4) Charged complex 1.

+o

- 0 , - CaH4 -d(COa) + (HaO)++ Ag Ag - 0 , - CZH4 --+ (CaHaO)' + Ag - (0)Ag - 0 , - CaH4 + CaH, + Ag - 0 , - (CZH,) Ag - 0 + CSH,+=Ag - (0 - CaH4) Ag

*

L. YA. MAROOLIS

464

Charged complex 2.

- 0 - CaH, + 0,+ (COJ + 2 (HnO)' Ag - (C,H,O)+ + e + A g + C,H,O Ag - (COJ - e - + A g + CO, Ag - (HpO)+ + e + A g + H,O Ag

Hayes (130) proposed in 1959 a scheme for ethylene oxidation over silver. He believed that ethylene oxide might be yielded only by interaction between ethylene and atomic oxygen, while carbon dioxide and water would be formed by a reaction of ethylene with molecular oxygen. If this were consistent with reality, the marked selectivity of silver with respect to ethylene oxidation to ethylene oxide would be inexplicable. Atomic oxygen is present on the surface of other metals as well, for instance on platinum and palladium, but no ethylene oxide is formed by reaction of these with ethylene. b. Olejine Oxidation over Platinum. The oxidation of hydrocarbons over platinum is of a markedly different nature compared with that observed for silver. Solely carbon dioxide and water are always present in reaction products in the oxidation over platinum, within a very wide range of experimental conditions, such as temperature, concentration of components, pressure. It was shown by Elovich and Butyagin (12) that the propene sorbed becomes strongly bound with platinum and may be eliminated only by treatment of the surface with oxygen, after which propene is taken up in greater amounts. I n oxygen adsorption on platinum oxygen may be dissolved in the adjacent-to-surface layers in an amount equal to that required for several tens of monolayers. I n accord to Frumkin (70) the binding of oxygen with platinated platinum becomes stronger with longer contact times and its reduction is then hindered. It was shown by investigating the electron work function in the adsorption of oxygen on platinum than on the platinum surface oxygen is an electron acceptor. It may be seen from results on isotopic exchange than 0; ions of molecular oxygen might be present a t the surface. Oxygen was found to be strongly sorbed on the platinum surface. Isotopic oxygen exchange becomes important at temperatures by 200-250" higher than the temperature of catalysis. Thus the binding of oxygen and hydrocarbons on the platinum surface is different from that for silver. Due to the great strength of oxygenhydrocarbon bonds, the formation of ethylene oxide on platinum seems to be scarcely probable. It was found, moreover, that the propene oxidation reaction passed into the gas phase even a t 70". With other metals, such as Ni and Cu, a new phase, i.e., a film of NiO

CATALYTIC OXIDATION O F HYDROCARBONS

465

and Cu,O oxides will be formed during catalytic oxidation and, consequently, the latter should be considered as a reaction proceeding on semiconductors. 5 . Semiconductors Various reactions, such as high conversion, low conversion with partial destruction of hydrocarbon molecules, and without essential distortion of the molecular structure may occur in the oxidation over semiconducting catalysts. Let us consider several reactions of propene oxidation over various semiconductors. a . Propene Oxidation over Cuprow Oxide. The cuprous oxide surface becomes charged by adsorption of oxygen, propene, and acrolein. The charge of molecules adsorbed may be determined from the electron work function. At the surface of CuO propene and acrolein are electron donors, like the majority of organic compounds. Adsorbed water slightly decreases the electron work function. Consequently, water also is an electron donor. The following reactions occur in the adsorption of various gases and vapors on cuprous oxide: 0,

+ e-+(O&

0

+e-+(O)-

C,H, - e -+ (C,H,)+ CHO H,O

- CH = CH, - e -+ (C,H,O)+ (acrolein) - e -+ (HnO)+

These data make possible the determination of a number of steps and elucidation of the nature of electron transfers in adsorption. Propene oxidation to acrolein over cuprous oxide seems to follow the scheme: 0, 0

+ 20

+ e + (0)-

0,

+

CsH6

e

-+

(0,)-

+ (0s)-

-+ (CaHsOOH)

charged complex 1 (hydroperoxide)

(CaHsOOH)+3 (CJH40)++ (H,O)+

+e (CsH@) + 0 , (C8H40)+

--+

-+

(CsH40)gas (CsH40

- 0,)

charged complex 2.

466

L. YA. MARGOLIS

+ H,O + (RH)’ (RH)’ + 0, -+ CO, + H,O + (R‘H)’ - 0,)-+

(C,H,O

- e -+

C,H,

(CsH,)+

CO,

(C,H,)+

+ 2(0)-

-+

(CsH, 00) a

charged complex 3. (C,H,

- 00)-+ C,H,O + HCHO

(C,H,

*

(RH)’

00) -+ ROO

+ 0, -+

+ (RH)’

R’OO

+ (R”H)+

etc.

b. Propene Oxidation over Spinels (MgCr,O,,CoMn,O,,MnCo,O,, etc.). Hydrocarbons are oxidized over these catalysts to carbon dioxide and water, i.e., hydrocarbon molecules become completely destructed. No acids are found in reaction products. It was suggested above that a chain reaction occurred a t the surface resulting in complete oxidation of hydrocarbon molecules. The charge signs of the components adsorbed measured from the electron work functions are similar to those for simple oxides. The reaction scheme may be written as follows: 20

0,

$

0

+ e -+ (0)-

Oa

+e

C,H,

-+

(0,)-

- e + (CsH,)+

+

(0,)- C,H,

--f

(CSH600) Charged complex I.

(C,H,)+

+ 0,

-+

(C,H,OO) Charged complex 11.

(C,H,

- 00),-+ (R’H) + RO,

(C3HeOO)11+ 0 , -+ (HOa)

+ ROa

+ 0, -+ RO, -t(R”H) etc. (RO,) + O,+R‘O, + (R”H) eto. (R’H)

The charged complex I (C,H,O,) is bound with the spinel lattice through oxygen. The oxygen-surface bonding is known to be very strong for these compounds; oxygen desorption is not observed even a t temperatures above 5OO0C. Thus complex I may suffer destruction

CATALYTIC OXIDATION O F HYDROCARBONS

467

only by further interaction with oxygen. Cationic sites of the Mn3+ and other types seem to be the active centers for oxygen adsorption on spinels. These sites are located at a considerable distance from each other and, consequently, a reaction of the complex with the oxygen adsorbed will be much less probable. It may be suggested that complex I involving an unpaired electron may enter into reaction with oxygen of the gas phase. The charged complex I1 is bound with the spinel lattice through carbon. The bonding is weak; hydrocarbon adsorption is reversible at 100200’. Thus the decomposition of such a complex or its desorption and subsequent oxidation in the gas phase are probable.

F. HETEROGENEOUS-HOMOGENEOUS REACTION STEPS The various hydrocarbon oxidation schemes discussed above were believed to proceed a t the catalyst surface only. The present concepts accept the occurrence of complex heterogeneous-homogeneous reactions proceeding in part at the solid surface and in part in the gas or liquid phase. Many catalytic oxidation processes considered recently as purely heterogeneous appeared to proceed by the heterogeneoushomogeneous mechanism. Such are the oxidations of hydrogen, methane, ethane, ethylene, propene, and ammonia over platinum a t elevated temperatures, as studied by Polyakov et al. (131-136). When hydrocarbons are oxidized over platinum the reaction sets in on the catalyst surface and terminates in the gas phase. However, until recently, the homogeneous continuation of the reaction could be established only by indirect means, using kinetic and other methods. In 1946 Koval’skii and Bogoyavlenskaya (137)proposed a method of differential calorimetry permitting a more accurate and unambiguous solution of this problem. With catalysts for high conversion the process of oxidation will proceed by the heterogeneous-homogeneous mechanism, while surface reactions only will occur with low conversion catalysts, at ordinary temperature of catalysis. However, the classification of catalysts as belonging to two groupsfor low and high conversion is not always justifiable with respect to surface and heterogeneous-homogeneous .processes. Thus, in using cobalt-manganesespinels catalyzing propene oxidation to carbon dioxide and water only, Linde (138)did not observe the passing of the reaction to the gas phase. With these catalysts the reaction will proceed on the surface only. Popova et al. (139) have shown that the conversion of carbonyl compounds in the oxidation of propene to acrolein over a copper catalyst involves a heterogeneous-homogeneous step.

468

L. YA. MARGOLIS

IV. Reaction Kinetics It is very difficult to establish kinetic laws for hydrocarbon oxidation first of all due to the high endothermicity of this reaction resulting in sintering of the catalyst, in surface changes, and in the intensification of side processes. This is probably the reason why the kinetics of a number of hydrocarbon oxidation reactions is insufficiently known, and the data reported in literature are scarce. A. THE EFFECTOF MACROSCOPICFACTORS

A number of physical side processes, such as the diffusion of initial compounds and reaction products, the liberation and distribution of heat, the dynamics of gases and liquids exert an influence on hydrocarbon oxidation under working conditions. All these factors are of prime importance for the design of catalytic apparatus, and moreover, may bring a change in the main oxidation characteristic, i.e., in the selectivity. The diffusion coefficient for catalyst pores is usually calculated approximately. Roiter et al. (140, 141) worked out a method for experimental determination of the diffusion coefficient and calculation of reaction rates without errors induced by diffusion. A catalyst possesses a considerable inner surface the access to which is difficult for gases, due to windings and small diameters of pores. Oxidation proceeds mainly in the internal diffusion region and as the temperature is raised to 400", it passes into the external diffusion region, with sharp heating up of the catalyst to 100-1 10" (at a reaction temperature 420"). The phthalic anhydride yielded by naphthalene inside the pores oxidizes to the end-products, C,O and H,O, due to hindered diffusion and the increased contact time (Fig. 10); this results in lower selectivity. Boreskov (142) found that the contribution of the inner surface is different for various industrial catalysts. Inside the catalyst grains the reaction is several times slower than at the external surface. Kholyavenko and Rubanik (143) investigated the effect of internal diffusion on the ethylene oxidation rate, using the diaphragm method, and calculated the effective diffusion coefficients for ethylene and carbon dioxide diffusing through a silver diaphragm. Ethylene oxidation on a porous silver catalyst proceeds over a wide temperature range in the transient internal region (190-260"). The selectivity of ethylene oxidation a t 190-250" is the same inside the grain and a t its external surface. The nature and structure of the support exert a considerable effect

CATALYTIC OXIDATION O F HYDROCARBONS

469

on the efficiency and selectivity of reactions. For example, Gorokhovatskii et al. (144)studied the oxidation of propene to acrolein over cuprous-copper oxide catalysts on various supports, namely on glass, aluminum oxide, and silicon carbide. The supports differed in specific activities. The selectivity of these catalysts was a function of the support structure. For aluminum oxide-supported catalysts the Selectivity decreased with increasing of the inner surface of the siipport.

Fro. 10. The efficiency ( U )of catalytic naphthalene oxidation over V,O, as a function of temperature; 1-in naphthalene; 2-in phthalic anhydride. U in molelmin g X lo6.

Frank-Kamenetzkii (145) suggested that theoretical analysis of exothermic regimes would make possible the proper choice of conditions ensuring the reaction development with weak heating up in the kinetic and strong heating up in the diffusion region.

B. KINETICSOF PROPENE AND ETHYLENE OXIDATIONOVER VANADIUM Roginskii et al. (146) found that the rate of olefine oxidation is proportional to oxygen concentration, and almost independent of hydrocarbon concentration. The activation energy values for olefine oxidations are summarized in Table VIII. The activation energy for carbon dioxide generation, Eco,, is higher than those for carbon monoxide and aldehyde (Eco and E,). E,, increases with an increasing number of carbon atoms in the hydrocarbon molecule. The activation energies for the formation of aldehydes from propene and isobutene are low and close to each other.

470

L. YA. MAROOLIS

The low activation energies for aldehydes and carbon dioxide formation might be due t o the fact that these reactions occur in the diffusion or transient regions. The activation energies for the formation of aldehydes and acids from propene over V,O, were measured in the kinetic region by Kutzeva and Margolis (147) and were found t o be 13 and 14 kcal/ mole, respectively. T A B L E VIII Activation Energies for the Formation of Aldehydes, Carbon Monoxide, and Carbon Dioxide from Various Hydrocarbona Reactions

Hydrocarbon ~

_

_

~

Activation energies References kcal/mole ____ _ _ __

Formation of aldehydes Formation of aldehydes Formation of acids Formation of CO, Formation of CO

Propene Propene Propene Propene Propene

22 2

Formation of CO, Formation of CO

Ethylene Ethylene

21 5

Formation of CO Formation of formaldehyde

Isobutene Isobutene

7

-

4

13 14

5

(146)

(30)

(314

It was noted before by the same workers that acrolein was formed along with saturated aldehydes in propene oxidation over metal oxides of the Vth and VIth groups of the Periodic System of Elements ( 5 1 ) . The kinetics of acrolein generation from propene over V,O, was also studied recently, and the activation energy for this reaction was found to be 12 kcal/mole. C. KINETICSOF NAPHTHALENE OXIDATION OVER VANADIUM OXIDE Ushakova et al. (148) studied the kinetics of naphthalene oxidation over V,O, under conditions preventing the effect of macroscopic factors on the reaction rate. Experiments were made over a temperature range of 380 to 410" with large crystals of nonporous V,O, in a flow-circulating system. As the gas circulation rate was 3.5 liters/min/cycle, there were no variations in naphthalene concentration in the catalyst layer. Calculation of experimental information gave following equations for

CATALYTIC OXIDATION OF HYDROCARBONS

47 1

individual reaction rates in the oxidation of naphthalene over V,O,: formation of phthalic anhydride Wpht.an.

= kph. Cnaph.

formation of maleic anhydride 0.5

w m a ~ a n .= kmal. Cnaph.

(6)

formation of naphthaquinone 2

Wnaphthaquin.

= knq. Cnaph.

(7)

high conversion wCO,

= kCO, *Cnaph.

(8)

The activation energies for these reactions were found to be for phthalic anhydride 37.4 kcal/mole for maleic anhydride

31.6 kcal/mole

for naphthaquinone

32.7 kcal/mole

for carbon dioxide

37.2 kcal/mole

Ioffe and Sherman (149)studied the kinetics of naphthalene oxidation to phthalic anhydride on a more complex vanadium-potassium-sulfate catalyst over a wide range of conversions and temperatures. The naphthalene oxidation was found to be independent of naphthalene concentration, This reaction is first order with respect to oxygen concentration and is inhibited with reaction products.

Here [O,] is the oxygen concentration in the reactant mixture; [A,] is the concentration of reaction products expressed as the amount of oxygen consumed for their formation; k , and k , are constants. The activation energy for this process is 27 kcal/mole; the preexponential factor 5 x 109. The zero reaction order obtained earlier by Calderbank (150) is jmitated for low conversion. The kinetic laws for naphthalene oxidation over a mixed vanadium catalyst and pure vanadium pentoxide are different. The adsorption capacities of naphthalene and oxygen seem to suffer a change when potassium sulfate is added to V,O,. According to Boreskov and Kassatkina (88) the rate of isotopic oxygen exchange on V,O, increases in

47 2

L. YA. MARQOLIS

the presence of potassium sulfate, while the activation energy for exchange decreases. The rate-limiting step for this catalyst probably differs from that for V,O,. Ioffe and Sherman suggest that the rate of naphthalene oxidation is limited by that of desorption of reaction products. Roiter et al. (148)studied the kinetics of naphthalene oxidation to 1.4-naphthaquinone over a mixed vanadium-potassium-sulfatesilica gel catalyst over a wide range of naphthalene and oxygen concentrations at 330-360". Certain information on the kinetics of naphthalene oxidation over various vanadium catalysts is given in the review by Dixon and Longfield (110).This reaction was found to vary from zero to first order with respect to naphthalene, and be close to first order for oxygen.

D. KINETICSOF BENZENE OXIDATIONTO MALEIC ANHYDRIDE OVER V20, Hammar (107) studied benzene oxidation over catalysts consisting of V,O, on supports, and found that the rates of maleic anhydride, CO, and CO formation were proportional to benzene concentration (the reaction order was close to first, as calculated from the degree of conversion aa a function of the contact time). The activation energies were equal and amounted to 2 4 f kcal/mole. The reaction order with respect to benzene, as determined from initial rates, varied from zero to second order. I n studying the kinetics of benzene oxidation to maleic anhydride and in order to eliminate diffusion hindrance, Ioffe and Lyubarskii (151) used the flow-circulating method. The rate of this reaction was found to be proportional to benzene concentration to the power of 0.78, and that of high conversion to the power of 0.71. The rate of maleic anhydride oxidation followed a first order equation. Kinetic equations were derived from experimental results : mole/liter: For oxygen concentrations lower than 4 x The rate of maleic anhydride formation (MA)

(C6HB)0-7fJ w 1 = k,(O2)Z - -____ (MA)OS~~ The rate of maleic anhydride oxidation w2 = k, (MA) The rate of high conversion of benzene wg

=

(C H )O.'l k3(0,)2S _ _ B - _ (MA)O*'O

CATALYTIC OXIDATION OF HYDROCARBONS

473

The over-all rate of benzene conversion (13)

For oxygen concentrations higher than 4 x lo-, molejiter: w1

=

(C8H6)0'78 k1 ( m ) 0 . 7 4

w 2 = k,(MA) (C6H6)O*'11 W, = k, __(MA)o*74

(15)

1

w,, = k 1(C6H6)0*78 + k3(C8H6)0*71 (MA)O.74 Activation energies: For the formation of maleic anhydride El = 22.6 kcal/mole For the oxidation of maleic anhydride E , = 12.6 kcal/mole For high conversion of benzene E , = 37.0 kcal/mole Ioffe and Lyubarskii believe that the main amount of benzene is decomposed by the oxygen adsorbed on V,O,. The lattice oxygen is responsible for the decomposition of a small amount of C,H,, but the rate of this reaction is considerably lower. Maleic anhydride is readily sorbed on V,O,. No allowance is usually made in kinetic studies for the side reaction of decomposition by the lattice oxygen. Ioffe and Lyubarskii (151) derived an equation for the benzene oxidation rate, allowing for phase transitions in the V,O, lattice.

E. KINETICS OF PROPENE OXIDATIONTO ACROLEIN OVER Cu,O Almost no papers were published on the kinetics of acrolein formation from propene. Those available report that cuprous oxide on silicon carbide or on pumice was used as catalyst, Isaev and Margolis (152) studied the kinetics of acrolein synthesis from propene under dynamic conditions a t atmospheric pressure. The rates of acrolein and carbon dioxide formation were found t o be proportional to oxygen concentration in the gas phase and independent of propene concentration. It will be of interest to note that the two reactions are of the same order. A similar result was obtained earlier for the formation of aldehydes, CO and CO, in the oxidation of propene over vanadium pentoxide.

L. YA. MARGOLIS

474

The activation energy for C 0 2 formation is 28-30 kcal/mole, and for acrolein it amounts to 12-14 kcal/mole. With acrolein transition to the external diffusion region occurs a t a lower temperature than that for CO,. The selectivity of propene oxidation to acrolein is determined by the ratio of the reaction rate constants for acrolein and carbon dioxide. Over the temperature range lower than 320' the reaction rate constant for acrolein formation may be higher than that for carbon dioxide, while the reverse is observed a t temperatures higher than 350". Belousov et al. (115,153) studied the kinetics of propene oxidation t o acrolein on a cuprous-copper catalyst, using the flow-circulating method. The reaction products were shown to exert a considerable effect on the reaction rate, and this permitted the derivation of a more precise kinetic equation for this reaction. Over the temperature range of 305-365" the rates of propene oxidation to acrolein (wacr,) and carbon dioxide (wco,)are fairly well described by equations

k , [O 21'

wco, = [C,H40]o*7[CsHal0*2 Here 0,, C3H40,and C3H, are the oxygen, acrolein, and propene concentrations per cycle in volume percentage; A, is the concentration of products expressed as the amount of oxygen consumed for their formation; k,, k , are constants that may be obtained from kinetic data. The activation energies obtained are:

E,,,

=

30 & 2 kcal/mole,

Eco, = 36 & 2 kcal/mole

By calculating the heats of adsorption of products it is possible to obtain the activation energies for these reactions under conditions ruling out inhibition by products. Then E,,, = 20 & 1 kcal/mole, and Eco, = 26 & 1 kcal/mole, and this is in agreement with results obtained by other workers. The kinetic equations derived by using the flow and flow-circulating method show good agreement.

F. KINETICSOF ETHYLENE OXIDATIONTO ETHYLENE OXIDE OVER SILVER The data on kinetics of ethylene oxidation to ethylene oxide are summarized in Table IX. I n all papers published the rate of ethylene oxidation is but slightly

CATALYTIC OXIDATION OF HYDROCARBONS

475

dependent on ethylene concentration (the reaction order within zero to 0.45) and proportional to oxygen concentration (first order in oxygen). Temkin et al. (159) studied the kinetics of ethylene oxidation over a stationary silver surface. It was shown by means of the flow-circulating method that the rate of ethylene oxide and carbon dioxide formation was proportional to ethylene concentration in the gas phase, and that there was inhibition with reaction products. TABLE IX Kineties of Ethylene Oxidation to Ethylene Oxide over Silver Catalyata

Reaction equation

Reaction

Oxidation of.eth~lene d[C&I to ethylene oxide -__ dt

=

Activation energy (kcal/mole)

k l [ C p H 4 1 0 . 4 S [ O ~ 1 0 . ~ ~ 16

Formation of CO, from d[CO ] 2=k n[C,H4]o*S [0n]'" ethylene dt

16

Ethylene oxide oxidation

20

- [o,x& dCC,H,OI -dt

Oxidation of ethylene d [ C 2 H 4 1= k , [ C p H 4 ] [ C , H 4 ] o [ 0 , ] 1 ~ U13 7 to C , H 4 0 and C O , Ethylene oxide oxidation

d[C,H,Ol dt

=

k,[C8H4]1[Oz]1

17-19

Oxidation of ethylene d [ C d h I = k[C,H410.8[0,10., to ethylene oxide - dt-

19

-d[C,H,I ~--__ dt

- [o,ll

-

C0,l'

-

Ethylene oxidation Ethylene oxide oxidation Ethylene oxidation

--

d[C&O] dt

H ' kL [C,H,I'[O*I0 -d[C 2 =---dt [C$&Ol k[COJ

+

15(C8H40)

19 (CO,)

The regularities observed experimentally may be classified as characteristic of three cases, when the rates of ethylene and carbon dioxide formation are proportional to: (1) oxygen concentration (being independent of ethylene concentration) ; (2) ethylene concentration (being

476

L. YA. MAROOLIS

independent of oxygen concentration); and (3) ethylene and oxygen concentrations to fractional powers. These dependences seem t o be determined by the surface content in molecular and atomic oxygen, and by the rates of oxidation and decomposition of peroxide radicals.

G. KINETICSOF HIGH CONVERSION OF HYDROCARBONS Almost no data were reported in literature on the kinetics of the high conversion of hydrocarbons t o carbon dioxide and water. This is probably due to the fact that the strong exothermicity of high conversion processes makes difficult the obtaining of reliable kinetic characteristics. Margolis and Todes (13)have studied the kinetics of catalytic oxidation of various hydrocarbons a t an excess of oxygen, using a number of catalysts, such as magnesium chromite, copper, and platinum under isothermal conditions. For hydrocarbons of aliphatic isostructure (2.2.4-dimethylpentane),and for naphthene hydrocarbons (cyclohexane and methylcyclohexane) the reaction kinetics is second order. For other classes of normal paraffins, such as n-pentane and n-heptane, and for unsaturated hydrocarbons (C,H,, C,H,) the reaction rates are proportional to the first power of hydrocarbon concentrations. For unsaturated and normal aliphatic hydrocarbons the kinetics of oxidation is proportional to the hydrocarbon concentration to the first power. At 200’ the specific catalytic activities will be in the order: platinum > copper chromite > magnesium chromite. As the temperature is raised to 400°,this order changes. Platinum remains the most active catalyst, while the catalytic activity of magnesium chromite becomes almost by 3 powers of ten higher than that of copper chromite. It will be of interest to note that over certain tem’perature ranges catalysts of higher activation energies are less active than those of lower activation energies. Thecurveslog k v s 1/Tfor two different catalysts may intersect (Fig. 11) due to the fact that E and K Ovalues change in the same sense. This was found to be characteristic of all catalysts used in the oxidation of various hydrocarbons. Simultaneous changes in E and K O were studied for a large number of organic reactions, such as hydrogenation and dehydrogenation (160). The activation energies and preexponential factors for the oxidation of saturated and unsaturated hydrocarbons over various catalysts are summarized in Table X. I n the oxidation of hydrocarbons of various structure over CuCr,O, and MgCr aO,, the activation energies and preexponential factors

TABLE X

The E and K O Valuesfor Catalytk Combvatbn of Hydrocarbons over Various Catalysts E (cal/mole) Hydrocarbon Platinum n-pentane n-heptane n-octane Benzene Ethylene

7800 8000 7300 -

Ka

Manganesechromium

Copper chromium

19,300 2700 46,000 4700 8100

18,300 3900 33,000 6400 25,000 20,000 5300 17,500

-

5500 28,000

Platinum

lo5..‘ 1054 105

Manganesechromium

Copper chromium

108.4 102.2 1019.5 102.7 104 108.1 1011.6

107 102 10’2 108 109 109.8

1074

Temperature range

280-350 350-550 300-330 330-500 300-500 300-400 400-550 280-350

0

ki

t P

l

4

L. YA. MAROOLIS

478

increase with the number of carbon atoms in the hydrocarbon suffering oxidation. This dependence is not observed for the oxidation of normal saturated hydrocarbons, such as pentane, heptane, and octane over platinum. Linde et al. (138) have studied propene oxidation under static conditions at low pressures over manganese-cobalt spinels. The reaction rate was found to be proportional to coverage of the surface with oxygen and independent of the hydrocarbon concentration. bk

-I

.o

'

20040-8

I

150.10-5

I

1o040-8

FIG. 11. The rate constant logarithm as a function of the temperature reciprocal in the oxidation of iso-octane over various catalysts: 1-copper chromium, 2-copper chromite, 3-platinum, 4-copper aluminum and iron-aluminum, 5-silver manganate, 6-iron-chromium, 7-magnesium chromite, B-copper chromium with addition of PbO.

The rates of high conversion of hydrocarbons over simple metal and oxide catalysts were studied by few workers. Pigulevskii ( 4 3 ) has shown that the rate of propene combustion over V,O, was proportional to oxygen concentration. Reyerson and Swearingen ( 9 ) found that the rate of ethylene oxidation over platinum was directly proportional to oxygen concentration and inversely proportional to ethylene concentration. Elovich and Butyagin (12) investigated the high conversion of

CATALYTIC OXIDATION OF HYDROCARBONS

479

propene on platinum and found that at a moderate surface coverage with propene the rate of carbon dioxide formation was independent of oxygen concentration, while with extensive coverage it was independent both of propene and oxygen concentrations. Branson, Hanlon, and Smythe (161)have studied recently the oxidation of methane, ethane, propane, butane, and isobutane on copper oxide over a temperature range of 330-650" and found that a pseudoinduction period was observed almost for all hydrocarbons. The rate of the first oxidation step follows the equation

w=

kb

(U - X )

1 - b ( a -x)

where lc and b are constants. The rate of the second step obeys the Zeldovioh-Roginskii equation. Thus the formal kinetics of high conversion of hydrocarbons is primarily a function of molecular structure and is but slightly affected by the nature of catalysts (chromites and platinum). The greater the number of carbon atoms in a molecule the higher are the preexponential factor and the activation energy for high conversion. This regularity holds both for saturated and unsaturated, as well as for simple cyclic hydrocarbons. Change in the order of the kinetic equation as a function of the molecular structure of a hydrocarbon provides evidence for a rate-determining step that seems to be related to the nature of hydrocarbon radicals formed in adsorption, I n certain cases the rate-determining step is the chemisorption of oxygen. The above considerations on reaction kinetics seem to show that surface charging must exert an effect on the kinetics of catalytic and adsorption processes. Inhibition of the reaction rate with reaction products is one of the characteristic features of hydrocarbon oxidation. Oxygen-containing compounds, such as aldehydes, olefine oxides, etc., are strongly sorbed on various catalyst surfaces. Measurement of the electron work function in the adsorption of these compounds showed that all these were electron donors, like hydrocarbons, and were probably sorbed at the same surface sites, thus inhibiting the reaction rate.

V. Modified Catalysts The phenomenon called promotion was discovered long ago. This is the increase in catalytic activities of various semiconductors and metals by addition of certain impurities. Margolis and Todes (162) studied the effect of impurities on the rate

480

L. YA. MARGOLIS

of high conversion. They found that the main kinetic characteristics, i.e., the activation energy and the preexponential factor, change in the same direction depending upon concentration of impurities in the catalyst. This and the different effect of impurities on the catalytic activity induced Roginskii to discard the old definitions (“promotion,” “poisoning”) and to introduce the word ‘(modification” meaning a dual change in the cataIytic activity. Impurities added to a solid may either fill the lattice vacancies, or substitute certain ions in the lattice, or else become diluted in the solid. Adsorption of impurities on the solid surface or substitution of sodium

: :-

1 2 3 4 5 6 7 8 9

15%

r

0.7 Li,O

FIG.12. Variations in the electron work function (A+ in ev) various impurities (I’in atomic yo)in ZnO.

VB

the concentration of

or potassium ions for hydrogen ions (the so-called ion-exchange adsorption) are possible. Of especial interest are semiconducting systems with a part of atoms (ions) replaced by atoms (ions) of about the same dimensions, but of a different valence or charge. According to de Boer and Verwey (163, 164) considerable modifications of the electron properties of such systems may be obtained in this way. The electron theory of catalysis established the relationship between chemisorptive and catalytic activities of solids and the Fermi level position. The Fermi level of a solid will be shifted in different directions with respect to the conduction band, depending upon the nature of impurities added, i.e., upon their being electron donors or acceptors.

CATALYTIC OXIDATION OF HYDROCARBONS

48 1

The electrical conductivity of the solid would also vary as a function. of the impurity nature and the chemical effect of these must be different. However, experimental results show that the relationship between conductivity and catalytic activity is much more complex. This is probably due t o the fact that the conductivity o! polycrystalline semiconductors often is not affected by changes in the Fermi level of the surface. Thus there must be another connection between changes in electron work functions of modified catalysts and their adsorptive and catalytic activities. The effects exerted by surface impurities, and by impurities located in the adjacent-to-surface layer at a depth exceeding the Debye length, on electrical conductivity and the electron work function will be different.

201 --

-A

E

FIO. 13. Variations in the activation energies for chemisorption calculated in kcel/ mole from initial rates vs the electron work function in kcal; ZnO containing various impurities.

Zhabrova et al. (165) report that with a modified zinc oxide catalyst a lithium cation impurity, located at a depth having no effect on the electron work function, is an electron acceptor and raises the activation energy for conduction. The same impurity on the grain surface is an electron donor and decreases the electron work function. Thus the Fermi level shift on the surface of modified catalysts cannot be determined unambiguously from change in the electrical conductivity. According to Enikeev, Margolis, and Roginskii (166), lithium and thorium cations added to zinc, copper, and nickel oxides change the

482

L. Y A . MAROOLIS

electron work function, and this change is a function of their concentration (Fig. 12). I n other words, modification of semiconductors makes possible a wide range of variations in the electron work function (d4). On the other hand, the same workers (167) found that A+ was a linear function of the activation energy for chemisorption of oxygen (Fig. 13). The kinetics of the over-all process of catalytic hydrocarbon oxidation will be considerably affected by impurities. An important part in the catalytic oxidation of various compounds is played by the mobility of the surface-adsorbed oxygen; it may be determined from oxygen exchange. Margolis and Kisselev (168) have studied isotopic oxygen exchange on typical oxidation catalysts such as silver (catalyzing the oxidation of ethylene to ethylene oxide) in the presence of halides, and copper oxide (a catalyst for propene oxidation to acrolein) in the presence of lithium, chromium, and bismuth oxides, and of copper sulfate. The logarithmic dependence of variations in isotopic exchange with electron work functions are shown in Fig, 14. The exchange rate in.o

-0.5L

FIG.14. Variations in the rate logarithm (w,) for isotopic oxygen exchange VE varistions in the electron work function ( A $ in ev). Exchange on CuO; additions made at 412'.

creases with A$. This is indication that the surface concentration of oxygen is controlled by the electron work function. The addition of donor or acceptor impurities to CuO, V,O,, or Ag changes the activities of the latter substances, as well as the selectivity of reactions. To find out the nature of the effect of impurities on the selectivity of hydrocarbon oxidation reactions, Enikeev, Isaev and Margolis (102) attempted to find out the relationship between the electron work function of a modified catalyst and the reaction rates and activation

CATALYTIC OXIDATION O F HYDROCARBONS

483

energies for two reactions: the formation of acrolein from propene, and of ethylene oxide from ethylene and carbon dioxide. Three major steps are observed for these reactions WI

W1

C,H,

-+

acrolein -+

L

W8

CO,

CJ

To a first approximation, without allowing for variations in the true activation energy with surface charging, the rate of acrolein formation may be expressed as w1 = KolCo,exp

( - E o l +RT

-

A+

Here K O ,is the preexponential factor, C,, is the surface concentration of oxygen; Eol the true activation energy for acrolein formation; Q the heat of oxygen adsorption on CuaO; A+ the change in the electron work function with addition of impurities: ( + A + ) for acceptors, and ( - A + ) for donors. Variations in the rates of CO, formation from acrolein and propene with surface charging may be represented as:

w3 = KO3*CO, exp

It was necessary to determine experimentally the nature of E and K O variations with A + . The relationship between variations in the electron work function of a modified copper oxide and the activation energies for CO, and acrolein formation relative to and E values for pure cuprous oxide are shown in Fig. 15. An increase in the electron work function brings about a decrease in the activation energy for carbon dioxide formation. The variations in activation energies for acrolein and CO, formation are linearly dependent upon the electron work function

A compensation effect is observed for propene oxidation on modified catalysts, as well as for other catalytic reactions, The dependence of the preexponential factor K O on is shown in Fig. 15. For acrolein forma-

+

L. YA. MARGOLIS

484

tion E,, increases with decreasing 4; the increase in w is slight. This seems to be due to the compensation effect of K O . The dependence of two rates (wzand w3)on the surface potential is essential for CO, formation. The rate of its formation from acrolein is a function of acrolein concentration and must decrease with increasing 4. The rate ( w J of

FIQ. 15. Variations in activation energies (LIE kcal/mole) end in the preexponential factor logarithm (logKO) for the formation of acrolein and carbon dioxide in the oxidation of propene over copper catalysts with admixtures vs variations in the electron work function (44 in ev). 1-CuO; 2-CuO Fe,O,; 3-CuO Cr,O,; 4-CuO+ Li,O; 6-CuO Ci-; 6-CuO SO:-.

+

+

+

+

parallel CO formation from propene, bypassing acrolein, is proportional to C,, and must increase with 4. It was shown experimentally (Fig. 15) that the activation energy Eco, builds up with 4 and, consequently, the rate of acrolein oxidation controls the rate of CO, formation, The may be expressed as the wllw,ratio reaction selectivity (8)

The rate of acrolein formation (wl)increases with 4, and that for CO, (wz)falls down. Variations in the selectivity of acrolein synthesis as a function of those in the electron work function ( A $ ) are shown in Fig. 16. Acceptor impurities added to CuO raise the selectivity by 15 to 20%, and the electron work function by 0.2 to 0.3 ev. To obtain a still greater increase in selectivity it would be necessary to investigate

CATALYTIC OXIDATION OF HYDROCARBONS

486

impurities capable of inducing a considerable increase in the electron work function. Extensive investigations on the effect of diverse impurities added to CuO were carried out by Margolis et al. (169). The effect of these substances was found to proceed by an electronic mechanism. The results obtained are summarized in Table 11. The elements are given in the order of increasing electronegativity . Those involving atoms with

-AS LO.I FIG. 16. Variations in the selectivity of propene oxidation ( A S )over CuO vs variations 3-CuO+ Cr,O,; P C u O in the electron work function. d-$/ev. 1-CuO; 2-Cu0+Fe,O3; Li,O; 5-CuO C1-; 6-CuO SO:-.

+

+

+

an electronegativity lower than that of CuO decrease the electron work function, and those with E exceeding that of CuO increase it (see Table XI). The qualitative relationship between electronegativity values for impurity atoms, and the increase or decrease in the electron work function of a semiconductor reveals that the solid solutions or . microheterogeneous systems formed exert a prevailing effect on L I ~It was established by electron diffraction of a number of systems that the impurities added form a separate phase on the CuO surface. This is indication that the effect of microheterogeneous inclusions on the electronic properties of the surface is predominant. Modifying additions increasing the A $ of CuO raise, and those decreasing it lower the selectivity of oxidation reactions. The electronic mechanism of semiconductor modifications associating variations in activation energy with A $ was investigated by Hauffe ($70)’ Vol’kenstein (274,and Roginskii (272). The oxidation of ethylene to ethylene oxide is a typical process occurring by a parallel scheme. Kummer (173) reports that A 4 is increased by adsorption of oxygen on silver. A similar effect is exerted on C$ by acceptor impurities C1, I , S, Se, etc. Similar variations in $ for silver catalysts were observed by Wilson

486

L. YA. MAROOLIB

et al. (174) upon addition of metalloids, such as chlorine, sulfur, and phosphorus. with adsorption of the components of ethylene The changes in oxidation are indication that on the catalyst surface these substances possess a charge. As the electron work function is related to the heat of adsorption, it may be supposed that with modified silver it will increase for acceptor gases (02,CO,) and decrease for donors (ethylene). It may be seen from results on isotopic exchange on silver samples modified with ion-chlorine, that with increased chlorine concentration in Ag (increased A + ) the amount of surface oxygen will fall off. Kinetic studies of oxidation over silver revealed discrepances in the data obtained by various scientists (Table IX). The rates of C,H40 and CO formation are proportional to oxygen and ethylene concentrations, but the reaction order with respect to these components varies from zero to one. The dependence of reaction rates on C,H4 and 0, concentrations is expressed more often than not in fractional powers. Temkin et al. (158) consider that this diversity in kinetic laws is due to varying oxygen content in the adjacent-to-surface layer. The dependence of partial surface concentrations of 0, and C2H4on variations in 4 of the catalyst seems to explain the unsteady nature of reaction kinetics. It will be expected, moreover, that addition of acceptor impurities will raise the reaction order for oxygen and decrease it for ethylene. The zero reaction order with respect to oxygen in the formation of C,H40 and CO,, as found by Temkin et al., may apparently be explained as follows. The electron work function is related to oxygen pressure A+ = ylogC,, Substituting this value into the equation for the rate of ethylene oxide formation, we obtain

+

,

w ~ , =~KolCo, , ~ exp

RT

Then

r

Y

=

Y

RT

When y' = 1 the reaction rate will be independent of oxygen concentration. It is also necessary to take into account the effect of reaction products on kinetics of the reaction. The rate of ethylene oxide formation is known to be inhibited both by ethylene oxide and CO,. The measured shifts of C$ with adsorption of these substances on silver have shown

TABLE XI Changes in the Electron Work Fu7aetwn of Mali$& Coppw O d d 8 and in the Selectivity of Propene Ozidotion to Acrolein 0

Ba

Elements

wu 5

Li

Cr

Pb

Fe

Mo

CuO

Bi

P

S

1.0

1.6

1.8

1.8

1.9

2.0

2.0

2.1

2.5

3.0

-400

-150

+400

+300

+lo0

4

-3

-3

+5

+20

+24

3U TI

H

Electronegstivity 0.85 Changes in the electron work function: (mv) in vacuum in 6 mixture of C,H, 0, ( 1 : l ) - 120 Changes in the selectivity of propene oxidation to acrolein -10

+

- 400 -12

- 150 -5

- 70 41

0

0

+300 +3

Modified CuO samples were reduced to cuprous oxide during the catalytic process. Pure CuO and Cu,O samples differed in the electron work function but slightly. Thus Ad brought about by modification should be comparable. The E-value for metal copper was taken as the CuO electronegativity.

d

C

0

2

L. YA. MARBOLIS

488

that at low temperatures C,H,O is a donor, and at high temperatures an acceptor of electrons. The shift of d, in the adsorption of C,H40 is greater than with carbon dioxide. Inhibition of the reaction by its products cannot be explained solely by blocking off of the surface, the change in with adsorption should also be taken into account. Water does not considerably decrease at the temperature of reaction and probably only blocks off certain surface sites. The effect of C,H40 and C 0 8 on the electron work function is similar to that of oxygen and metalloid impurities. Thus in the presence of these substances the surface activity would be expected t o fall off and the reaction selectivity to raise. However, the effect of these products on is less marked than that of metalloids and, consequently, the selectivity would scarcely be considerably increased by admission of C O , into the reactant gas. The presence of sulfur impurities in silver results in an increased ethylene oxide yield; further increase in sulfur concentration would lead to poisoning of the catalyst. Roginskii et al. (167)found that the same was observed for silver samples modified with chlorine. The work function of sulfur-containing samples, as well as of those containing chlorine, will increase. The activity maximum is controlled by the ratio of ethylene to oxygen surface concentrations. The reaction order with respect to ethylene will decrease with increasing sulfur concentration, i.e., with +, while that for oxygen will buiId up. At a certain value the rate to1 = K,,C$C&, will be maximum, when n = m. This is apparently the reaaon why silver is modified with sulfur. When chlorine is added to silver, no increase in the ethylene oxide yield is observed, as “pure” silver involves ion-chlorine8 in excess over the optimum amount. It was shown above that the activation energy is related to +. There is almost no information available on as a function of activation energies and rates of catalytic reactions over silver. According to Hayes (130)the activation energy for N,O decomposition on alloys of silver with various +decreasing metals will be low. Sosnovsky (175)has investigated the catalytic activities ( E and K O )for different planes of silver crystals, with respect to the decomposition of formic acid. E and K O were found to increase with plane indexes. The relation between and the rate of ethylene oxidation to ethylene oxide waa not established. With modified silver the activation energy for ethylene oxidation to ethylene oxide will not suffer considerable changes. This may be explained by increased concentration of donor molecules compensating the change in under the action of metalloids. The differences in activation energy values for ethylene oxidation, as reported by various

+

+

+

+

+

CATALYTIC OXIDATION OF HYDROCARBONS

489

scientists (Table IX) did not exceed 4 kcal/mole. Different reaction rates observed upon modification are probably due to changes in surface concentrations of the reaction components. The reaction selectivity (8) is determined by the ratio of C,H,O (wl) to CO, (wz) formation rates. The relationship between the electron work function and the over-all activity and selectivity of oxidation may be seen from Fig. 17. The A S 100

r

1

I

I

I

0

0.I

0.2

0.3

tA4

FIG. 17. Variations in the activity ( A ) and selectivity (8)of ethylene oxidation to ethylene oxide vs variations in the electron work function (A+ in ev).

activity of silver decreases and the reaction selectivity increases with increasing electron work function. Changes in reaction rates with the electron work function are different for low and high conversions. is proportional to oxygen concenThe rate of C,H,O formation (wl) tration to the power of 0.4-0.7, and that of CO, formation t o the power of 1.1. Thus the changes in w,and w,with increasing work function will be different. The higher is the reaction order with respect to oxygen, the greater will be the decrease in reaction rate with 4. Higher 4 results in greater selectivity of the reaction. Analysis of data available seems to reveal that the modification of silver by addition of acceptor impurities proceeds by an electronic mechanism. As no experimental data on the effect of alkali and earth alkali impurities on electrical and catalytical properties of silver were available, corresponding studies were undertaken with a silver catalyst. Elements of the 1st and IInd groups of the Periodical System of Elements decrease, and those of the VIth and VIIth groups usually increase the electron work function for metals. The mechanism of 4 variations with adsorption of atoms and molecules of various substances on metal surfaces is due to formation of a double electrical layer

490

L. YA. MAROOLIS

accounted for by dipoles of adsorbed particles (121).The relationship values for metals, the ionization potential, and electron between affinity of adsorbed atoms was reported by Mikhailov et al. (176).With a great number of elements added the electronegativity ( E ) would be a more convenient value than the ionization potential or the electron affinity for investigating in metals. Experimental data on variations, and also on the selectivity of silver in the presence of various impurities are summarized in Table XII. Alkali and earth-alkali metal impurities decrease, while metalloids increase the electron work function for silver. The elements studied are given in the order of increasing electronegativity. I n the presence of elements with a n electronegativity lower than that of silver the value falls off, while high-electronegativity elements raise it. This law holds for measured in vacuum and in a hydrocarbon-oxygen mixture. Only the sign of change in the electron work function dependent on c should be taken into account. I n many cases the decrease in heat of adsorption with greater coverage may be accounted for by interaction of dipoles in the adsorbed layer. Consequently, the effect of impurities on the reaction rate may be explained as follows. The impurities localized on the silver surface and displaying an electronegativity higher than that of silver, thus increasing the work function, will form dipoles with the negative charge outwards. Measurement of in the adsorption of oxygen and ethylene has shown that such is the case for oxygen, while for ethylene outwards is the positive charge. As a result of an electrostatic reaction of 0, and C,H, dipoles with a metalloid, the heats of adsorption become lower and, consequently, there is a decrease in the surface coverage with oxygen; the reverse is observed for coverage with ethylene. The probability of complete combustion of an ethylene molecule will be considerably lower in this case. Modification of V,O, was investigated for propene oxidation t o saturated aldehydes, acrolein, etc. The impurities added to V,O, may be classified as falling into two groups: acid (metalloid) anions of SO,, P,O,, and other alkali cations such as Na, K, etc. The additions of SO:- results in a sharp rise in the activation energy for the generation of reaction products, while in the presence of K and Na the activation energy for the formation of acrolein, aldehydes, and acids falls off. Sodium was found to decrease and acids to increase the electron work function. Activation energies for the formation of acrolein, acetaldehyde, acid, and carbon dioxide change with A + . A characteristic feature of hydrocarbon oxidation over Cu ,O, V,O,,

+

+ +

+

+

+

d

Elements

'

Electronegativity Changes in the electron work function (mv) in vacuum (20") in a mixture of C,H, 0,(1:l)

+

K

BE

Na

Ca

Be

Ag

Mo

Bi

S

J

C1

0,

0.8

0.85

0.9

1.0

1.5

1.8

1.9

2.0

2.5

2.6

3.0

3.5

-50

-80

-300

-130

-70

0

+200

+lo0

+200

+300

+600

+300

-100

-60

-40

0

-3

0

( 1800)

+300

Changes in the selectivity of

ethylene oxidation 60 ethylene oxide

14

-3

-2

-11

+3

Poisoning

+15

+18

+23

0 M

492

L. YA. MARQOLIS

and Ag is the possibility of suppressing complete combustion by increasing the electron work function. Thus choice can be made of substances changing the reaction selectivity. It is apparent from qualitative data that alkali cations, such as Na, K, Ba that lower the work function, increase the rate of high conversion to COa and HaO and decrease the selectivity. All acid impurities, such as C1, I, Br, SO,, PO,, etc., cause an increase in the electron work function, a decrease in the rate of CO, formation, and a rise in selectivity, but also exert an effect on the formation of oxygen-containing products, such as ethylene oxide, acrolein, acetaldehyde. Thus high activity and selectivity would be ensured only at a strict optimum concentration of the compound added. Complex semiconducting spinels of a defect structure are catalysts for high-degree oxidation of hydrocarbons. Transfer of charge in the spinel lattice seems to occur by a relay mechanism and, consequently, of importance is the spacing of cations that are responsible for charge transfer. Anyway the electrical and catalytic properties of these compounds would be changed only by addition of considerable amounts of other substances. This is indication that the collective effects observed for simple semiconductorswill be negligible with the above systems. Notwithstanding the great amount of work concerned with the effect of modifying impurities, there is no information on their behavior in catalysis. Stepanov, Margolis, and Roginskii (177) used the radioactive tracer technique in studying the mobility of impurities in metals. Labeled chlorine ( C P ) , iodine (P), and sulfur (SsS)were added to silver, which was then heated in a flow of various gases. It was shown that in a reducing medium the concentration of modifying impurities rapidly falls off. Reduction of silver halide in ethylene and hydrogen yielding hydrogen halide detected in exit gases occurs at the catalyst surface. As a result of similar experiments with a single crystal of silver containing C P as KCP, it was established that chlorine does not diffuse over the silver lattice at 300'. The same conclusion was drawn by Mikulski and Werber (178) in studying the diffusion of sulfur in silver at 400'. A great amount of work was reported in literature on the effect of various organic halides on the selectivity of ethylene oxidation to ethylene oxide (123), but the mechanism of this effect was not studied. Organic impurities may sorb on the catalyst surface and react with oxygen in the gas phase, or else decompose. Stepanov et at?. (177)found that the rate of oxidation of metalloid-

CATALYTIC OXIDATION OF HYDROCARBONS

493

containing organic impurities over silver is a function of the molecular structure of these compounds. The sequence of processes occurring on a catalyst containing a metalloid impurity follows the scheme RHsl

4

reduction Ag +------

Ag

+ AgHal+

CO,

+ H,O

adsorption

!CHd

oxidation

4% The difference in activation energies for the oxidation of organic halides (Box= 4.6 kcal/mole) and the reduction of AgCl (Bred= 15 kcal/mole) is indication that the rates of oxidation and reduction may vary for different silver samples, depending upon the ion-chlorine wntent of the surface. Changes in the ratio of these processes accounted for by different temperature and the structure of organic volatile compounds result in different distribution of impurities over the catalyst layer, and this may exert a marked effect on the selectivity.

VI. Mixed Catalysts Various solid mixtures that are either converted into solid solutions or remain polyphase are used, along with simple and complex semiconducting oxides, for conducting different reactions. Information was reported by Rienacker (l79),and in patent literature, on the catalytic activity and the selectivity of mixed catalysts with respect to dehydrogenation of alcohols, oxidation of hydrocarbons, etc. More recent work on the activity of such catalysts (180,181) revealed that the change in specific surfaces with increasing content of one metal oxide in the other is not additive. Mixtures involving oxides of molybdenum and vanadium, molybdenum and cobalt, iron and chromium, etc., were used as catalysts for hydrocarbon oxidation. The concentration of defects may increase a t the interphase boundary during the preparation of mixed catalysts, and these would then display higher catalytic activity. Frequently the changes in catalytic activity of mixed contacts are due to formation of spinels. The effect of a polyphase system of this kind, consisting of metal oxide spinels (often as solid solutions), on the rates of various reactions is a complex problem. Various mixed catalysts were used for the catalytic oxidation of hydrocarbons. Variations in activity and selectivity of benzene oxida-

494

L. YA. MARQOLIS

tion to maleic anhydride as a function of the composition of vanadiummolybdenum catalysts, as established by Ioffe et al. (180)are shown in Fig. 18. The activity maximum coincides with the limit on the dissolution of MOO, in V,O,. The effect of MOO, dissolved in V,O, is suggested to result from the appearance of defects in the V,O, lattice. Voevodskii et al. (182) have shown by ESR studies that for samples with low MOO, content (up to 30%) V4+ and Mae+ ions form solid solutions in the V,O, lattice. Compounds containing four-valent vanadium ions are formed a t a considerable MOO, content.

MO0 3

"2 '5

FIG. 18. Selectivity of benzene oxidation to maleic anhydride as a function of the composition in etomic yo of vanadium-molybdenum catalysts. I-over-all conversion of C,H,; 2-conversion in ?/, of C,H, to C4H8O4.

Mixed catalysts are complex polyphase systems, the electron work function of which may vary either a t the interphase boundaries or due to formation of solid solutions by substitution and incorporation. The effect of these catalysts on high conversion will differ with preparation conditions. The activation energies for the formation of saturated aldehydes, acids, and CO, decrease with propene oxidation on a catalyst MOO, solution. consisting of a solid V,O, With a mixed molybdenum vanadium catalyst the selectivity of acrolein formation remains the same, while decreasing on a solid solution of these oxides. The formation of oxygen-containing products (acrolein plus saturated hydrocarbons and oxides) is more selective in the oxidation of propene over a mixed catalyst.

+

CATALYTIC OXIDATION OF HYDROCARBONS

495

at. % Bi at. 'lo Mo;W

A4

0.5

0.3

Mo

IOOat. X Bi

0.20.3 0.4 0.1

\

0.4

(b) Fro. 19. (a)The selectivity of propene oxidation to acrolein as a function of the composition of bismuth-molybdenumand bismuth-tungstencatalysts. (b) Variations in the electron work function (A+ in ev) vs the composition of mixed bismuth-molybdenumand biemuth-tungsten Catalysts.

496

L. YA. MARGOLIS

It was shown before by Margolis et al. (51) that propene will suffer oxidation over molybdenum oxides at a high temperature (560') yielding acrolein, carbon dioxide, and water. Variations in the selectivity of propene oxidation as a function of the catalyst composition are shown in Fig. 19a and b. If the suggested electronicmechanism of the action of mixed catalysts is true, the electron work function (4) of mixtures should be higher than that of pure molybdenum and bismuth oxides. The dependence of A# on the composition of a molybdenum-bismuth catalyst is shown in Fig. 19b. The maximum change in the electron work function corresponds to highest selectivity. Such a proportional change in catalytic and electronic properties seems to provide evidence for the electronic mechanism of the effect of these mixed catalysts. Tungsten yields compounds displaying chemical properties approaching those of molybdenum. Like molybdenum, WO, is a mild catalyst for low conversion of hydrocarbons. Variations in electron work functions of mixed tungsten-bismuth catalysts of various composition are shown in Fig. 19, also. A relationship was observed between selectivity and the electronic properties of these catalysts. * According to Rienacker (179) either the electron mechanism of the effect of mixed catalysts, or a change in the strength of oxygen bonds in metal oxides are possible. Under the assumption that the oxygen bond strength is important for catalysis, the oxidation-reduction mechanism will have to be accepted. However, data were given above (Section 11, E, 3) showing the invalidity of this scheme. Consequently, the changes in reactiop rates in the presence of mixed catalysts are probably due to changes in the amount and nature of the defeots at the phase boundary, Further experimental inveatigatiov of the relationship between wtalytio aptivity and electronjp propsrties of mixed catqlyets would Qoptrfbutato the duoidation ~f the e#e& of bhse sptems on oatal$rt;te

reactlone.

REF~EREYCEO 1: pavy, H., Phil. Tram. Roy. Soc. (1817). 2. F o b , S . A., Zhur. R u e . F k-Khim . Obahchlva 40,216 (1908). 3. Philips. O.,J . Am. Ohm. Soc. 16, 104,266 (1894). *New reauh were obtained reoently by the author of this paper. Theee will be published in the near future.

CATALYTIU OXIDATION OF HYDROCARBONS

497

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CATALYTIC OXIDATION OF HYDROCARBONS

499

88. Kassatkina, L. A., and Boreskov, G. K., Zhur. Fiz. Khim. 29,455 (1969). 89. Dzisyak, A. P., Boreskov, G. K., Kaesatkina, L. A., and Kochurikhin, V. E., Kinetika i Kataliz 2, 386 (1961). 90. Stewie, E. W. R., and Stovel, H. V., J.Chem. Phys. 2,581 (1934). 91. Elovich, S.Yu., end Zhabrova, G. M., Zhur. Fiz. Khim. 13, 1761 (1939). 92, Klar, R., 2 . physik. Chem. A168, 215 (1934). 93. Trapnell, B. W. M., Quart. Reva. (London) 8, 404 (1954);Kavtaradze, N.N., Zhur. Fiz. Khim. 36, 628 (1962). 94. Margolis, L. Ya., Izueat. Akad. Nauk S.S.S.R., Otdel. Khim. Nauk No. 10, 1175 (1968). 95. Turkevich, J., and Taylor, H., J. Am. Chem. SOC.56, 2254 (1934);Howard, J.,and Taylor, H., ibid. 2259. 96. Voevodskii, V. V., Vol’kenstein, F. F., and Semenov, N. N., i n “Voprosy Khimicheskoi Kinetiki, Kataliza i Reektzionnoi Sposobnosti” (V. N. Kondratiev and N. M. Emanuel, eds.), p. 423.Akad. Nauk S.S.S.R., Moskva, 1955. 97. Syrkin, Ya. K., Zhur. Strukt. Khim. 1, 189 (1960). 98. Aigrain, P., and Dugas, C., 2 . Elektrochem. 56, 363 (1952). 99. Germain, J. E., Compt. rend. acad. aci. 238,345 (1954). 100. Engell, H. I., and Hauffe, K., 2 . Elektrochem. 57, 762 (1953). 101. Vol’kenstein, F.F., Vmtnik Moskov. Univ. No. 4, 79 (1957). 102. Enikeev, E. Kh., Isaev, O.V., andMargolis, L. Ya., Kinetikai Kataliz 1,431 (1960). 103. Pongratz, A,, and Scholtis, K., Ber. 76, 1179 (1943). 104. Sosin, S. L., and Sladkov, A. M., Uspekhi Khim. 23, 377 (1954). 105. Dolgov, B. N., “Kataliz v Organicheskoi Khimii.” Goskhimizdat, Moskva, 1959. 106. Charlot, V.,Ann. chim. [ll]2, 415 (1934). 107. Hammer, C. G. B.,Svemk Kem. Tidskr. 64, 165 (1952). 108. Ioffe, I. I., i n “Voprosy Khimicheskoi Kinetiki, Kataliza i Reaktzionnoi Sposobnosti” (V. N. Kondratiev and N. M. Emenuel, eds.), p. 232. Akad. Nauk S.S.S.R., Moskva, 1955. 109. Ushakova, V. P., Korneichuk, G. P., and Roiter, V. A., Ukrain. Khim. Zhur. 23, 310 (1957). 110. Dixon, J. K., and Longfield, J. E., i n “Catalysis” (P. Emmet, ed.), Vol. VII, p. 183. Reinhold, New York, 1960;Parks, W. G., and Allerd, C. E., Ind. Eng. Chem. 31, 1162 (1939). 111. Suvorov, B.V., Rafikov, S. R., and Anuchina, I. G., Doklady Akad. NaukS.S.S.R. 88, 79 (1953);Solomin, A. V., Suvorov, B. V., and Rafikov, S. R., Trudy I m t . Khim. Nauk Akad. Nauk Kazakh. S.S.R. 2, 192 (1958). 112. Bretton, R. H., Wan, S. W., and Dodge, B. F., Ind. Eng. Chem. 44,594 (1952). 113. Todes, 0.M., and Andrianova, T. I., Zhur. Fiz. Khim. 27,1485 (1953). 114. Isaev, 0.V., Margolis, L. Ya., and Roginskii, S. Z., Chur. Obshchei Khim. 29, 1522 (1959). 115. Belousov, V. M., Gorokhovatskii, Ya. B., Rubanik, M. Ye., and Gershingorn, A. V., Doklady Akad. NaukS.S.S.R. 132, 1125 (1960). 116. Isaev, 0. V., Margolis, L. Ye., and Sazonova, I. S., Doklady Akad Nauk S.S.S.R. 129, 141 (1959);Golovina, 0.A., Isaev, 0. V., and Sakharov, M. M., ibid. 142, 619 (1962). 117. Neiman, M. B., Intern. J. Appl. Radiation and Iaotopes 3,20 (1958). 118. Margolis, L. Ya., and Roginskii, S. Z., i n “Problemy Kinetiki i Kataliza” (S. Z. Roginskii el al., eds.), Vol. IX, p. 107.Akad. Nauk S.S.S.R., Moskva, 1957. 119. Gorokhovatskii, Ye. B., and Rubanik, M. Ya., Ukrain. Khim. Zhur. 24, 63 (1958). 120. Zimakov, P. V., Khim. Nauka i Prom. 2,24 (1957).

500

L. YA. MAROOLIS

121. de Boer, L. N., Advances i n Catalysis 8, 151 (1956); Eischens, R. P., and Pliskin, N. A,, ibid. 10, 1 (1958). 122. Schultze, G. R., and Theile, H., Erdol u. Kohle 5 , 552 (1952). 123. Pokrovskii, V. A., Uspekhi Khim. 25, 1446 (1956). 124. Lyubarskii, G. D., Doklady Akad. NaukSi3.S.R. 110, 112 (1956). 125. Voevodskii, V. V., i n “Problemy Kinetiki i Kataliza” (S. Z. Roginskii et al., eds.), Vol. VIII, p. 97. Akad. Nauk S.S.S.R., Moskva, 1955. 126. Semenov, N. N., “Some Problems of Chemical Kinetics and Reactivity.” Princeton Univ. Press, Princeton, New Jersey, 1958. 127. Vassiliev, V. N., Elovich, S. Yu., andMargolis, L. Ya., Doklady Akad. NaukS.S.S.R. 101, 703 (1956). 128. Roiter, V. A., Stukanovskaya, N. A., and Volkovskaya, N. S . , Ukrain. Khim. Zhur. 24, 37 (1968); Margolis, L. Yu., “Geterogennoye kataliticheskoye okislenye uglevodorodov.” Gostoptekhizdat, Moskva, 1962. 129. Margolis, L. Ya., Uspekhi Khim. 28, 615 (1959). 130. Hayes, K. E., Can. J . Chem. 37,583 (1959). 131. Polyakov, M. V., and Stadnik, P. M., Zhur. Fiz. Khim. 4,449 (1933). 132. Urizko, V. I., and Polyakov, M. V., Ukrain. Khim. Zhur. 22, 713 (1956). 133. Vysotzkii, Z. Z., and Polyakov, M. V., Ukrain. Khim. Zhur. 22, 180 (1956). 134. Polyakov,M.V., andShoikhet,P. A., Doklady Akad. NaukS.S.S.R. 89,1057 (1953). 135. Trotzenko, M. A., and Polyakov, M. V., Doklady Akad. Nauk S.S.S.R. 96, 115 (1954). 136. Shoikhet, P. A., Trotzenko, M. A., and Polyakov, M. V., Doklady Akad. Nauk S.S.S.R. 89. 619 (1953). 137. Bogoyavlenskaya, M. L., and Koval’skii, A. A., Zhur. Piz. Khim. 20, 1325 (1946). 138. Linde, V. R., Margolis, L. Ya., and Roginskii, S. Z., Doklady Akad. Nauk S.S.S.R. 136, 860 (1961). 139. Popova, N. I., and Vermel, E. E., Izue.at.Sibir. Otdel. Akad. NaukS.S.S.R., No. 11, 89 (1960). 140. Roiter, V. A., Korneichuk, G. P., Leperson, M. G., Stukanovskaya, N. A., and Tolchina, B. I., Zhur. Fiz. Khim. 24, 459 (1950). 141. Roiter, V. A,, and Gaukhman, S. S., in “Kataliz” (A. I. Brodskii, ed.), p. 161. Akad. Nauk Ukrain. S.S.R., Kiev, 1950. 142. Boreskov, G. K., in “Problemy Kinetiki i Kataliza” (S.Z. Roginskii et al., eds.), Vol. VI, p. 404. Akad. Nauk S.S.S.R., Moskva, 1946. 143. Kholyavenko, K. M., and Rubanik, M. Ya., Ukrain. Khim. Zhur. 24, 55 (1958). 144. Gorokhovatskii, Ya. B., Rubanik, M. Ya., and Popova, N. I., Kinetkka i Kataliz 3, 230 (1962). 14.5. Frank-Kamenetzkii, D. A., Zhur. Fiz. Khim. 18, 738 (1939). 146. Margolis, L.Ya.,Malyarova, E . P., andRoginskii, S. Z . , Izvest. Akad. NaukS.S.S.R.. Otdel. Khim. Nauk, 6 p. 958 (1954). 147. Kutzeva, L. N., and Margolis, L. Ye., Zhur. ObsheheZ Khim. 32, 102 (1962). 148. Roiter, V. A., Ushakova, V. P., Korneichuk, G. P., and Skorbilina, T. G., Kinetika i Kataliz 2, 94 (1961). 149. Ioffe, I. I., and Sherman, Yu. G., Zhur. Fiz. Khim. 28,2095 (1954). 150. Calderbank, P. H., Ind. Chemist 28, 291 (1952). 151. Ioffe, I. I., and Lyubarskii, A. G., Kinetika iKalaliz 8, 261 (1962). 152. Isaev, 0. V., and Margolis, L. Ya., Kinetika i Kataliz 1,237 (1960). 153. Belousov, V. M., Gorokhovatskii, Ya. B., Rubanik, M. Ya., and Gershingorn, A. V., Kinetika i Kataliz 8, 221 (1962).

CATALYTIC OXIDATION OF HYDROCARBONS 154. 155. 156. 157. 158. 159. 160. 161. 162. 163. 164. 165. 166. 167. 168. 169. 170. 171. 172. 173. 174. 175. 176. 177.

178. 179. 180. 181. 182.

60 1

McKim, F. L., and Cambron, A., Can. J . Reaearch B27, 813 (1949). Murray, K. E., Australian J . Sci. Reaearch A3, 433 (1950). Wan, S.W., f n d . Eng. Chem. 45, 234 (1945). Fognanini, F.,and Montarnal, R., Rev. inst.franc. pitrole 14, 191 (1959). Kurilenko, A. I., Kul’kova, N. V., Rybakova, N. A., and Temkin, M. I., Zhur. Fiz. K h i m . 32, 1043 (1958);Kinetika i Kataliz 3, 208 (1962). Ostrovskii, V. E., Kul’kova, N. V., Lopatin, V. L., and Temkin, M. I., Kinetika i Kataliz 3, 189 (1962). Cremer, E.,Advances in Catalysis 7, 75 (1955). Branson, S.,Hanlon, L., and Smythe, B., Trans. Faraday SOC. 52,672 (1956). Margolis, L. Ya., andTodes, 0. M., Doklady Akad. NaukS.S.5.R. 58,421 (1947). de Boer, J. H., and Verwey, I. W., Rec. trav. chim. 55,631 (1936);. de Boer, J. H., and Verwey, I. W., Proc. Phys. SOC. (London). Extra Part 49, 59 (1937). Vladimirova, V. I., Enikeev, E. Kh., Zhabrova, G. M., and Margolis, L. Ya., Doklady Akad. Nauk S.S.S.R. 131, 342 (1960). Enikeev, E . Kh., Margolis, L. Ye., and Roginskii, S. Z., Doklady Akad. Nauk S.S.S.R. 130, 807 (1960). Enikeev, E. Kh., Margolis, L. Ye., and Roginskii, S. Z., Doklady Akad. Nauk S.S.S.R. 129, 372 (1959). Margolia, L. Ya., and Kisselev, V. A., Doklady Akad. Nauk S.S.S.R. 130, 1071 (1960). Margolis, L. Ye., Enikeev, E. Kh., Isaev, 0. V., Krylova, A. V., and Kushnerov, M. Ya., Kinetika i Kataliz 3, 181 (1962). Hauffe, K., Advancer, in Catalyaia 7,213 (1955). Vol’kenstein, F.F., “Electronnaya Teorya Kataliza.” Fizmatizdat, Moskva, 1960. Roginskii, S.Z., Kinetika iKataliz 1, 15 (1960). Kummer, J. T., J . Phys. Chem. 63,460 (1969). Wilson, J. N., Voge, H. H., Stevenson, D. P., Smith, A. E., and Atkins, L. T., J . Phya. Chem. 68, 463 (1959). Sosnovsky, H.M., Phya. and Chem. solids 10,304(1959). Mikhailov, G.S., Kutovskaya, L. A., and Pospelov, L. A., Radiotekh. i Elektron. 5, 658 (1960). Stepanov, Yu. N., Margolis, L. Ye., and Roginskii, S. Z., Doklady Akad. Nauk S.S.S.R. 135,369 (1960);Stepanov, Yu.N.,and Margolis, L. Ya.,Kinetika iKataliz 2, 684 (1961). Mikulski, J. M., Mrowec, S., and Werber, T., Bull. acad. polon. aci., aer. sci., chim., geol. geograph. 7, 737 (1959). Rieniicker, G.,2. anorg. Chem. 258, 280 (1949);Acta Chim. Acad. Sci. Hung. 14, 173 (1958). Ioffe, I. I., Ezhkova, Z. I., and Lyubarskii, A. G., Zhur. Piz. Khim. 35, 2348 (1960). Kernos, Yu. D., and Moldavskii, B. L., Zhur. Priklad. K h i m . 33,2593 (1960). Kazanskii, B. B., Ezhkova, Z. I., Lyubarskii, A. G., Voevodskii, V. V., and Ioffe. I. I., Kinetika i Kataliz 2, 862 (1961).

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Author Index Numbers in parentheses are reference numbers and are included to assist in locating references when the authors’ names are not mentioned in the text. Numbers in italics indicate the page on which the reference is listed. A Aston, J. G., 425 Adbersberg, W. I., Jr., 18311131, 186(113), Atkins, L. T., 486(174), 501 Augstein, J., 227( 11l), 251 200 Aalsberg, W. Ij., 18(33), 33 Austen, D. E. G., 183(116), 200 Acker, D. S . , 22(37), 33 Ackerman, J. F., 219(86), 220(86), 250 B Adams, R., 204, 209(10), 213(46), 213(47), 224(10), 236(10), 238, 247(180), 248, Bade, H., 73, 78, 79, 112 Bader, H., 237, 237(143), 252 249, 253 Adkins, H., 83, 112, 204(6), 206(11), 206, Baeder, D. L., 189(127), 200 210, 211, 216(3), 217(66), 236(3), Baker, B. R., 227(112), 250, 251. Balaceanu, J. C., 97, 11.3 238(11), 238, 245(11), 247, 248, 249 Balandin, A. A., 81, 97, 112, 113 Agnello, E. J., 233(131), 252 Aigrain, P., 446(98), 499 Ballou, E. V., 140(68), 199 Alexander, W. A., 431(22), 497 Eerham, H. N., 90, 91, I12 Allard, C. E., 472(110), 499(110), 499 Barter, C., 186, 200 Allen, M. B., 30(73), 34 Bartlett, P. D., 171(101), 200 Berth, R. T., 140(68), I99 Allison, H. W., 320(69), 425 Alpert, D., 392(88), (89), 404, 413(89), Basila, M. R., 130, 198 Bassham, J. A., 2(2-8), 32 426,427 Alpher, R. A,, 319(48), 425 Batta, I., 83, 84, 112 Altehuller, A. P., 180(133), 200 Baughan, E. C., 193(141), (142), 201 Amberg, C. H., 119(4), 198 Baur, J. P., 272(17), 424 Ames, D. E., 205(12), 213(12), 248 Bayard, R. T., 404, 427 Amesz, J., 28(66), 30(66), 34 Becker, J. A., 262( lo), 424 Anderson, L. C . , 170(97), 200 Beeck, O., 306(3I), 425 Anderson, P. A., 392(91), 426 Beets, M. G . J., 233(129), 251 Andrianove, T. I., 101, 113, 461, 464, 476 Behrisch, R., 413(121), 427 Bellamy, L. J., 192(140), 201 (113),499 Androes, G. M., 33 Belousov, V. M., 462, 474, 499, 500 Anuchina, I. G., 449, 499 Benesi, H. A., 10, 33, 123(19), 126(19), Apker, L., 268, 424 127(19), 140, 140(67), 141, 198, 199 Aranoff, S., 274(22), 425 Bennett, M . J., 418(133), 427 Armbruster, M., 440(72), 498 Benson, A. A., 2(1), 32, 213(41), 249 Armit, J. W., 219(82), 250 Benson, R. E., 22(37), 33 Arnold, W., 23(39, 40), 28(39), 29, 30(39), Benton, A., 440(71), 498 33 Berets, D. J., 441, 498 Arnon, D. I., 3(8, 9), 4(16), 6(9), 32 Bergstrom, F. W., 213(41), 249 Arnott, R. Y . , 439(64), 498 Beyerman, H. C . , 223(101), 251 Arnwine, B., 230(120), 251 Bickelhtlupt, F., 220(864), 250 Asai, S., 69, 60, 62, 63, 99, 111 Biegeleisen, J., 91, 113 603

504

AUTHOR INDEX

Biel, J. H., 217(68), 218, 250 Bierstedt, P. E., 22(38), 33 Bigeleisen, J., 141(60), 199 Billica, H. R., 206(11), 238(11), 246(11),

248 Biondi, M. A., 394(95), 418(131), 426,427 Bircumshaw, L. L., 93, 104, 113 Birkhofer, L., 221, 251 Bishop, N. I., 30(66), 34 Blahe, K., 228(116), 251 Blake, P. G., 83, 90, I12 Blauth, E. W., 413(121), 427 Blessing, G., 236(134), 252 Block, J., 41, 43, 66, 57, 62, 95, 109, I10 Blue, R. W., 88, 112 Boehringer, A., 228(113), 251 Boekelheide, V., 220(88), 233(131),(132), 234(132), (133), 250, 252 Bogoyavlenskaya, M. L., 467, 500 Bolton, F. H., 141(61), 199 Bond, G. C., 108, 113, 418(132), 427 Bone, W. A., 430(14), 497 Bontekoe. J. S., 223(101), 251 Boomer, E. H., 431(23), (24), 432(24),

497 Booth, A. H., 282, 425 Boreskov, G. K., 442(89), 468, 471, 499, 500 Borisov, P., 206, 249 Bowden, K., 219(83), 250 Bowen, A. R., 434(45), 435(46), 497 Bowman, R. E., 206(12), 213(12), 248 Bradley, W. F., 127(38), 198 Brady, L. E., 210(32), 249 Branch, G., 168(93), 200 Branson, S., 479, 501 Braun, W., 124(22), 198 Bremer, H., 48, 111 Brennan, D., 306(32), 425 Bretton, R. H., 433(31a), 449, 469, 470 ( 3 1 4 , 497, 499 Bridges, D. W., 272(17), 424 Brihta, I., 78, 79, 112 Brindley, G. W., 119(5), 198 Britton, E. C., 223(98), 251 Brivati, J. A., 166, 199 Brombacher, W. G . , 392(90), 413(90), 416(90), 426 Bronevskii, A. I., 211(36), 249 Brown, F. W., 272(18), 424

Brown, H. C., 213(42), 220, 249, 250 Brown, O . , 434(48), 436(48), 497 Brunner, W., 204(9), 210(9), 212(9), 241 (9), 243(9), 248 Bruns, B. P., 101, 113 Bryant, B., 230(120), 251 Buiten, J., 92, 113 Burns, R. M., 433(32), 497 Burstein, R. Kh., 437, 498 Burton, W. K., 346(67), 426 Busch, H., 300(28), 425 Buttram, J. R., 230(119), 251 Butyagin, P. Yu., 430( 12), 444, 464, 478,

497 C

Cabrera, N., 346(67), 426 Calderbank, P. H., 471, 500 Calvin, M., 2(1 , 2, 3, 4, 6 , Be), 3(11, 12, 13), 4(14), 6(20), 9(14, 27, 28), lO(28). l l ( 3 0 , 31, 32), 18(34), 23(43, 44), 26 (45, 46), 27(47, 62, 53, 54), 32(28), 32,

33 Cambron, A., 431(22), 497, 501 Campbell, B. K., 219(86), 220(85), 242 (167), 243(167), 246(167), 246(167), 247(167), 250, 252 Campbell, J. R., 430(8). 431(8), 433(8), 434(8), 497 Campbell, K. N., 219(86), 220(85), 242 (167), 243, 246(167), 246, 247(167).

250, 252, 253 Cannon, W. N., 204(7), 248 Carney, J. J., 218(71), 250 Carney, T. P., 213(43), 249 Carter, L. A., 27(47), 33 Caujolle, F., 104, 113 Cavellito, C. J . , 217, 218, 219, 244, 250 Cederstrand, C., 27(48), 30(48), 33 Chalmers, R., 27(48), 30(48), 33 Chance, B., 28(68), (62), (63), 30(63), 34 Charlot, V., 448, 499 Chaussidon, J., 127(34), 198 Chauvette, R. R., 204(7), 248 Chernikova, E. A., 83, 86, 87, 88, 102, 104, 106, 112 Chesnut, D. B., 22(36), 33 Chessick, J. J., 160(64), 299 Chiang, Y., 272(19), 424

AUTEOR INDEX

606

Chiaverelli, S., 243(171), 243, 246(171), Day, A. R., 100, 113 252 de Boer, E., 182(111), 183(111), 200 Chmeil, J., 190, 201 de Boer, J. H., 100, 113, 119(9), 121(10), Clark, H., 441, 498 123(21), 134, 134(47), 136(47), 136(9), Clark, L. W., 90, 91, 112 136(9), 137(48), 198, 199, 417, 427, Clarke, J. K. A., 64, 100, Ill, 164, 199 480, 501 Clausing, P., 297(26), 425 de Boer, L. N., 467, 490(121), 500 Clayton, R. K., 23(39), 28(39), 29(39), Decker, R. W., 416(126), 427 30(39), 33 de Diesbach, H., 247, 248(183), 253 Clemo, G. R., 230( 124), 251 Degueli, Y., 226(110), 252 Cockborn, W. F., 217(69), 250 Dendurent, M., 434(36), 497 Cohn, G., 206(16), 249 Deno, N. C., 168(96), 200 Commoner, B., 23, 28(42), 33 Dessibourg, J., 247(182), 253 Condon, F. E., 171(101), 189(130), 197, Deyrup, A. J., 140(63), 168(87), 199 200 Dienes, G. J., 272(18), 424 Connor, R., 204(4), (6), 248 Dietz, W., 39, 43, 110 Conover, C., 433(36), 436(36), 497 Dixon, J. K., 449, 472, 499 Conredi, J., 182(111), 183(111), 200 Dobrezow, L. N., 368(76), 369, 426 Constable, F. H.,63, 111 Dodge, B. F., 433(318), 449, 469(112), Cook, M. A., 172, 200 470(31a), 497,499 Copley, M. J., 892(93), 426 Doering, W. E., 240, 242, 252 Corwin, A. H., 216(59), 250 Dolan, W. W., 317(44b), 365(44b), 425 Courtois, M., 110, 113 Dolgov, B. N., 448, 499 Covert, L. W., 204(4), 248 Dorgelo, 0. J. H., 61, 111 -Covington, E. J., 269(7), 280(7), 282(7), Doub, L., 132(46), 198 301(7), 319(7), 334(7), 424 Dowden, D. A., 62, 78, 79, 111 Craig, R. A., 189(136), 190(136), 201 Downs, C. R., 433(32), (33), 434(44), 497 Cramer, H. I., 211(36), 238, 249 Drake, L., 440(71), 498 Crane, F. L., 30(67), (68), (69), 34 Drechsler, M., 319(60), (51), 425 Crawford, V., 118(2), 198 Duel], M. J., 64, 66, 66, 111 Cremer, E., 63, 81, 83, 84, 111, 112, 476 Dummler, S., 413(118), 427 (l60), 501 Dugas, C., 446(98), 499 Crocker, A. J., 49, 61, 96, 111 Dushman, S., 392(87a), 406(102), 426, 427 Crook, K. E., 230(121), 252 Duyaens, L. N. M., 28(66), (60), (66), 30 Cryder, D. S., 83, 112 (66), 33, 34 Cunningham, R. E., 49, 111 Dyke, W. P., 317(44b), 366(44b), 425 Custera, J. F. H., 119(9), 134(47), 136(47), Dzisyak, A. P., 443, 499 138(9), 198, 199 E Cutler, I. B., 172, 200 Cvetanovic, R. J., 189(128), 200 Eastman, J. W., 19(36), 33 Eckstrom, H. C., 122, 161, 198 D Edwards, J., 93, 104, 113 d’Alessandro, A. F., 434(42), 449, 497 Ehrlich, B., 30(68), 34 Dalmai, G., 37, 84, 96, 97, 100, 110 Ehrlich, G., 266(1), 269(6), (7). 260(7), Danes, V., 104, 113 262(7), 270(13), 271(1), (9), 272(6) Darzene, G., 238, 252 (g), (20)9 273(21), 279(6), (20), 281(1), Davies, A. G., 189(138), 193(142), 201 282(6), 286(6), 289(8), 298(6), 299(6), Davis, W. D., 413(117), 427 301(7). 304(29), 306(33) 308(36), 312 Davisson, C., 312(38), 425 (41),319(7), (8), 320(41), (4% 321(41), Davy, H., 430( l), 496 327(61), 330(41), 332(41), 334(7), (46),

606

AUTIEOR INDEX

338(66), 345(46), 361(76), 367(76), 370(41), 376(84), 380(61), 381(61), (85), 401(9), 406(21), 408(21), 409(20a, b), P a ) , 411(21), 418(41), (45), 419(76), 424, 425, 426 Eischens, R. P., 62, 64, 66, 67, 70, 71, 88, 95, 111, 112, 113, 116(1), 123, 124, 160, 161, 152, 155, 198, 199, 467, 500 Eldred, N. R., 213(42), 220, 249, 250 Eley, D. D., 9, 33, 43, 79, 110, 112 Eliel, E. L., 195(143), 201 Ellett, A., 412(115), 427 Elovich, S. Yu., 430(12), 438, 444(91), 444, 461(127), 464, 478, 497, 498, 499, 500 Emerson, R., 27(48), 30(48), 33 Emmert, B., 219(81), 250 Emmett, P. H., 171(100), 200 Engell, H. I., 446(100), 499 Enikeev, E. Kh., 446(102), 463, 481(166), 481, 482(167), 486(169), 488(167),

499, 501 Erhard, F., 95, 112 Ernest, I., 233, 251 Espe, W., 396(98b), 427 Eucken, A,, 87, 89, 112 Evans, A. G., 168(94), 177, 178, 200 Evans, P. L., 83, 104, 112 Eyring, H., 44, 60, 63, 111 Ezhkova, Z . I., 493(180), 494(180), 494 (182),501

F Fahrenfort, J., 39, 42, 43, 44, 56, 66, 60, 64, 65, 66, 93, 94, 100, 104, 105, 107, 109, 110, I l l , 123, 164, 156(69), 198,

Fogo, J. K., 183, 200 Fokin, S. A., 430(2), 431(2), 432(2), 496 Folman, M., 119(8), 124(24), 127(8), 127 (24), 128, 129(8), 133, 198 Fornasir, V., 218(75), 219(76), 250 Forti, G., 5(16), 32 Foster, H., 22(36), 33 Foster, N. F., 189(128),200 Frachon de Pradel, A., 37, 83, 84, 89, 96, 97, 100,110, 112 Fowler, R. H., 317(46), (47), 425 Francis, S. A., 67, 70, 111, 112 Franpois-Rossetti, J., 99, 113 Frank, F. C . , 426 Frank, R. L., 224(102), 251 Frank-Kamenetzkii, D. A., 469, 500 Freedman, L., 222(96), 251 Freidlin, L. Kh., 83, 84, 112, 113 Freifelder, M., 204(8), 206( l6),(19), 208(8), (la), 210(19),(32),(33),211(37),212(16), 213(16), (214(8),(16),(49), 216(8),(49), (61),(62),(63), 216(64), 217(8),( l6), 218 ( W , 220(8),(87), 221(8), 223(8),(16). 227( l6), 237(8),(18), 240(161), 243(16), (49), 248, 249, 250, 252 French, C. S., 30(72), 34 Fricano, R., 251 Friedman, H. L., 218(73), 250 Fripiat, J. J., 127(34), 198 Frisch, R., 312(39), 425 Frohnsdorff, G. J. C., 130, I98 Frumkin, A. N., 440(70), 464, 498 Fuderer-LuetiO, P., 78, 79, I12 Fueki, K., 97, 99, 113 Fujimori, E., 27(53, 64), 33

199 Faith; W., 434(36), 497 Faith, W. L., 431(26), 497 Farkas, A., 434(42), 449, 497 Farkas, E., 240(160). 252 Farlow, M., 204, 210(3), 216(3), 236(3), 247(3), 248 Faslod, K., 229, 251 Fassel, V. M., 272(17), 424 Fastovskii, V . G., 434(49), 435(49), 497 Ferles, M., 230(126), 251 Fischer, H., 6(18), 32 Fischer, J. B., 83, 97, 112 Flinn, R. A., 140(68), 199 Fognanini, F., 475(157), 501

G Gaffron, H., 7(21), 33 Galinovsky, F., 216(66), 219, 250 Garland, C . W., 123(12), 123, 161, 152(66), 153(66), 198, 199 Garner, W. E., 110, 113, 440(76), (77), 441,

498 Garrett, A. B., 189(136), 190(136), 201 Gaukhman, S. S., 468(141), 500 Gautier, J., 219(77) (78). 250 Gel’bshtein, A. I., 90, 91, 112 Georgian, V., 245 (176a). 253 Germain, J. E., 446(99), 499 Germer, L. H., 312(38), 425

AUTHOR INDEX

507

Gerehingorn, A. V., 452(115), 474(115), Hall, L. D., 396(97a), 427 Hall, W. K., 124, 167, 168(27), (82), (163), 500 Gibbs, H. D., 433(35), 436(35), 497 171(99), 172(82), 178, 178(27), 179, Gillam, E., III 183(114), 184, 198, 199, 200 Oilman, G., 205(15), 249 Hamilton, T. S., 204(10), 209(10), 224(10), Glasstone, S., 44, 50, 53, 111 236(10), 238, 248 Gmelin, W., .241(163), (164), 242(163), 252 Hammer, C. G. B., 449, 472, 499 Gold, V., 168(88), 178(88), 199 Hammel, E. F., 171(98), 200 Golding, D. R. V., 228, 251 Hammett, L. P., 90, 112,140, 168(87),199 Goldschmidt, St., 237, 252 Hanlon, L., 479, 5-01 Golovina, 0. A,, 452, 458(116), 499 Hannan, R. B., 124, 125(28), 127(28), Gomberg, M., 168(90), 200 128(28), 198 Gomer, R., 312(43), 320(44d), (56), (573, Hansen, N., 36, 46, 61, 62, 110 ( 5 8 ) , 326(57), 329(62), 330(44d), 350 Hantke, E., 97,113 (44d), 366(44d), 358(68), 359(58), 414 Hantzsch, A., 168(92), 200 (44d), 415(44d), 425, 426 Harder, R. J., 22(37), 33 Gonzalez-Vidal, J., 189(137), 201 Hare, D. G . , 189(138), 201 Good, G. M., 167(85), 172(85), I99 Harrison, R. J., 245(176a), 253 Good, R. H., Jr., 349(71), 426 Hartman, C. D., 262(10), 267, 424 Goode, W. E., 230(120), 246(177), X I , 253 Hasegawa, H., 123(14), 165, I98 Gorokhovatskii, Ya. B., 452(115), 457, Hashemeyer, R. H., 91, I I3 469, 474(115, 153), 499, 500 Haskell, T. H., 217(66), 218(66), 219(66), Gossling, B. S., 317(47), 425 244, 250 Govindjee, 27(49, 51), 30(49, 51), 33 Ham, H. B,, 431(19), 497 Grace. J. A., 178, 200 Hauffe, K., 446(100), 485, 499, 501 Gracheva, T. A., 434(51), 470(51), 496(51), Hawes, B. W. V., 168(88), 178(88), 199 498 Hawk, C. O . , 430(10), 497 Graeber, E. G., 83, I12 Haxo, F. T., 30(71), 34 Grave, T. B., 218, 219(74), 221(74), 222, Hayes, K. E., 464, 488,500 Hayward, D. O., 305(32), 425 250 Gray, A. P., 240( 158), 252 Hazebroek, H. F., 64, 65, 100, I I l , 123, Gray, T. J., 440(76), 498 154, 198 Green, P. N., 219(83), 250 Hedemann, B., 240(157), 252 Greensfelder, B. G., 167(85), 172(85), I99 Heise, J. J., 23, 26(42), 33 Grob, C. A., 216(60), 224(106), 225, 226, Heitmeier, D. E., 240(158), 252 226( l06), 232( 1043, 250, 251 Helbing, C. H., 246(179), 253 Gurevich, D. A., 434(41), 497 Helfer, L., 239, 252 Gurney, R. W., 359, 426 Hempel, W., 430(15), 497 Gwathmey, A. T., 49, 111 Herd, R., 230(119), 251 Herring, C., 319(53), 358(53), 425 Hertler, W. R., 22(37), 33 H Hess, E., 319(51), 425 Haag, W. O., 189(129), 196(145), 200, 201 Hem, K., 213(44), 249 Haring, W., 236(139), 252 Hettche, O., 243(170), 244(170), 252 Hagstrum, H. D., 298, 417, 425, 427 Heuer, K., 87, I12 Hahn, D. A., 103, JJ3, 430(4), 431(4), Hickmott, T. W., 259(6), 272(6), (14), (20), 449(4), 497 279(6), (20), 282(6), 285(6), 298(6), Hddeman, R. G., 171, 200 299(6), 319(14), 409(20a, b), ( l l l ) , Hall, H. K., 250 424, 425, 427 Hell, H. K., Jr., 216, 217(70), 250 Hildebrand, J. H., 10, 33

508

AUTHOR INDEX

Hildebrandt, H., 79, 112 Hill, G. R., 172, 200 Hill, T. L., 306(34), 307, 425 Hindin, S. G., 88, 112 Hinehelwood, C., 39, 83, 90, 96, 110, I12 Hirota, K., 59, 60, 62, 63, 69, 71, 84, 97, 99, Ill, 112, 113, 154, 188(125), 199,

200,438,498 Hirota, N., 188(124), 200 Hirschler, A. E., 140(59), 199 Hiskes, J. R., 349(72), 426 Hobson, M. C., 178, 183(105), 185, 200 Hoffmann, R. W., 119(6), 198 Hofman, V., 216(55), 250 Hofmann, K., 62, 103, 104, I l l Hohenlohe-Oehringen, K., 230(125), 251 Hoijtink, G. J., 18(33),33,183(112), (113), (119), 184(119), 185(113), 200 Hollabaugh, C. M., 150(64), 199 Holland, L., 414(122), 427 Holm, N. C. F., 88, 112 Horsley, L.H., 223(98), 251 Hoshina, T., 243(170), 244(170), 252 Howard, J., 444, 499 Howton, D. R., 228, 251 Hiiclrel, W., 239, 252 Hudda, F. G . , 312(41), 320(41), (45), 321 (41),327(61), 330(41),332(41), 334(45), (7), 345(45), 361(76), 367(76), 370(41), 380(61), 381(85), (61), 418(41), (45), 419(46), 425, 426 Huitema, R., 434(48), 436(48), 497 Hulm, J. K., 320(56), 425

I Ichimura, S., 27(50), 30(50), 33 Ikuma, S., 238(145), 252 Imalik, B., 37, 83, 84, 89, 96, 97, 99, 100, 110, 112, 113 Ingrarn, D. J. E., 183(116), 200 Inoue, K., 208(26), (28a), 249 Ioffe, I. I., 449, 471, 472, 473, 493(180), 494(182), 499, 500, 501 Ipatief, V., 238, 252 Isaev, 0. V., 439(66), 446(102), 452, 458, 463(102), 473, 482, 485(169),

498, 499, 500, 501 Ishii, T., 219(84), 250 Isogai, K., 242, 252 Ivashchenko, Y. N., 221, 251

J Jagendorf, A. T., 4(15), 5(16), 32 Janes, A., 435(52), 498 Janknich, P. E., 91, 113 Jsruzelski, J. J., 168(96), 200 J w u d a , B., 436(55), 498 Jedlioke, J., 104, 113 Jennings, T. J., 440(75). 441, 498 Johns, I. B., 241, 252 Johnson, C. A., 431(23), 497 Johnson, C. B., 406( 103), 427 Johnson, O., 140(55), 199 Johnson, V. J., 419(136), 427 Jones, A. C., 123(19), 125(19), 127(19), 198 Jones, H., 74, 112 Jones, H. A., 408(106), 427 Jones, J. I . , 208(25), 249 Jones, N., 178, 200 Jones, P. M. S., 178. 200 Jost, M. R., 26(48), 33 Jungers, J. C . , 97, 113 Jura, G., 123(12), 198 Jurinak, J. J., 187, 200

K Kalish, T. V., 437, 498 Kallischnigg, R., 236(137), 252 Kanien, M. D., 28(57), 30(57), (70), 34 Kamp, B. M., 28(65), 30(85), 34 Kaneko, C., 242(168), 243(168), 252 Kao, C. H., 218(72), 250 Karagounis, G., 119(3), (7), 198 Kariyone, K., 219(80), 250 Karyakin, A. V., 124(25), 127(25), 128,198 Kassatkina, L. A., 442(89), 471, 499 Kataoka, A,, 244(174), (175), 253 Katz, J., 436(54), 498 Kavtaradze, N. N., 444(93), 499 Kawasaki, C., 242( l68), 243( 168), 252 Kawazoe, Y., 245, 253 Kazanskii, B. B., 494(182), 501 Kearns, D. R., 11(30), (31), (32), 18(34).33 Keen, N., 155, 199 Kegel, L . P., 30(68), 34 Keier, N. P., 441, 498 Keller, A., 104, 113 Kepler, R. E., 22(38), 33 Kernos, Yu. D., 493(181), 501 Kerwin, J. F., 242(167), 243(167), 245( 167), 246(167), (179), 247(167), 252, 253

AUTHOR INDEX

509

Kerwin, L., 405(101a), 427 Kiihnis, H., 248(183), 253 Kesarev, V. V., 434(48), (47), 497 Kuffner, F., 238, 252 Keyes, D. B., 431(26), 497 Kuhn, R., 236, 252 Kholyavenko, K. M., 468, 500 Kuick, L. F., 204, 210(3), 216(3), 236(3), Kindler, K., 240, 252 247(3), 248 King, J. A., 216(55), 250 Kul'kove, N. V., 437, 475(159), 475(158), Kingdon, K. H., 312(37), 425 486(158), 498, 501 Kington, G. L., 130, 198, 425 Kullich, E., El, 83, 84, 112 Kirsanov, A. V., 221, 251 Kummer, J. T., 485 Kirschenlohr, W., 236, 252 Kupryanov, G. N., 434(38), 435(38), 436 Kiselev, A. V., 127(35), 132(43), 198 (381, 497 Kisliuk, P., 320(55), 425 Kurilenko, A. I., 475(158), 486(158), 501 Kisselev, V. A., 482, 501 Kursanov, D. N., 189(135),201 Kitamura, S., 208(26), 249 Kushnarev, M. Ya., 439(66), 498 Kim, R., 444(92), 499 Kushnerov, M. Ya., 485(169), 501 Klauser, H. E., 119(4), 198 Kutovskaya, L. A., 490(176), 501 Kleiman, M., 216(57), 250 Kutzeva, L. N., 441, 470, 498, 500 Klein, R., 414(124), 425, 427 Kuwata, K., 59, 60, 82, 63, 84, 99, 111, Klopfer, A., 396(97b), 409, 427 154, 188(125), 199, 200. Kloppenburg, C.C., 233( 129), (130), 251, Kuznetsov, S. G., 225, 251 Kuznetzov, M. I . , 433(34), 435(34), 436 252 Knoll, M., 396(98c), 427 (341, 497 Kobajashi, Y., 438, 498 Koelsch, C. F., 218(71), 250 L Kohl, W. H., 396(98d), 427 Kohn, E., 189(137), 201 Lafferty, J. M., 409, 427 Kohn, E. J., 208(29), 249 LsForge, R. A., 242(167), 243(167), 245 Komarov, 0 . N., 83, 86, 87, 88, 102, 104, (167), 246(167), 247(167), 252 105, 112 Laidler, K. J., 44, 50, 53, 111 Komarov, N. A., 83, 86, 87, 88, 102, 104, Lamborg, M., 234(133), 252 105, 112 Lanczos, C., 349(70), 426 Koeneichuk, G. P., 439(65), 449, 168(140), Langmuir, I., 257(3), 258, 312(37), 329, 470(148), 472(148), 498, 499, 500 336, 408(106), 424, 425, 426, 427 Kornfeld, E. C., 204(7), 247, 248, 253 Larkworthy, L. F., 189(138), 201 Kornienko, V. P., 93, 103, 104, I13 Lasslo, A., 217(61), 225(81). 250 Korosy, F., 104, 113 Lavignino, E. R., 204, 240(160), 248, 252 Kortiim, G., 124(22), 158, 159(75), 180 Lavrushin, V. F., 168(95), 178, 189(131), (75), 161(75), 162(22), 198, 199 (135), 200, 201 Kosuye, T., 242(188), 243(168), 252 Lazier, W. A., 204(5), 248 Kotov, E. I., 163, 164, 185(78), 184(78), Leftin, H. P., 124, 125(29), 125(30), 141 199 (62), 143(62), 144(62), 145(82), 148 Koval'skii, A. A., 467, 500 (62). 147(62), 148(62), 149(62), 167, Kovar, J., 228(llS), 251 168(27), (29), (82), 168(91), 170(29), Krel, H., 43, 56, 57, 62, 109, 110 171(99), 172(82), 177(103), 178, 178 Krasnovskii, A. A., 7(23), 33 (27), (106), 179, 183(105), 184, 186, Krastinat, W., 236(137), 252 187(122), 188(122), 190, 192(122), 193 Kmshkov, A. P., 434(39), 497 (103), 198, 199, 200,201 Krogmenn, K., 104, 113 Lehr, J. J., 417, 427 Krohnke, F., 229, 251 Leigh, J. S., 178, 183(105), 185, 200 KI'ylove, A. V., 485,501 Leiser, H. A., 218(73), 250

510

AUTHOR INDEX

Lemke, G., 243(173), 253 Leonchik, Z. I., 83, 86, 87, 88, 102, 104, 106, 112 Leone, S. A., 230(l22), 251 Leperson, M. G., 468( 140), 500 Lester, R. L., 30(69), 34 Leugering, H. J., 102, 113 Levit, A.M., 83, 84, 112, 113 Levy, J., 206(13), 248 Lewis, G. N., 141(60), 166, 199 Libman, N. M., 226, 251 Lichtin, N. N., l68(91), 200 Liebrandt, F., 213(46), 249 Liepack, H., 319(60), 425 Linde, V. R., 442, 467, 478, 498, 500 Linn, W. J., 234(133), 252 Lipkin, D., 166, 184(120), 199, 200 Lippincott, B. B., 23(42), 26(42), 33 Little, L. H., 119(4), (6), 124(6), 166, 167, 198, 199 Livingston, R, C., 7(24), 33 Long, F. A., 90, 112, 140(64), 199 Longfield, J. E., 449, 472, 499 Longwell, J. P., 189(127), 200 Lopatin, V. L., 476(169), 505 Low, M. J. D., 88(97), 112 Lucchesi, P. J., 167, 189(127), 199, 200 LuetiO, P., 43, 79, 110, 112 Luke;, R., 228(116), 230,251 Lundegardh, H., 28(66), 33 Lundy, R., 320(67), 326(67), 425 Lygin, V. I., 127(36), 198 Lyle, G. G., 224(104), (106), 226(104), (106), 232(106), 251 Lyle, R. E., 224( 104), 224, 225( 104), (106), 226, 230, 232(106), 251 Lyubarakii, A. G., 472, 473, 493(180), 494(180), (182). 500, 501 Lyubarski, G. D., 469, 500 '

M Mackor, E. L., 183(113), 186(113), 200 Macks, E. B., 240(166), 252 Maday, H. K., 23(40), 29(40), 33 Mahler, W., 22(37), 33 Mailhe, A,, 36, 37, 83, 110 Makor, E. L., 18(33), 33 Malyarova, E. P., 469(146), 470(146), 500 Malyusov, V. A,, 434(49), 436(49), 497 Marec, L. F., 103, 113

Marek, L. F., 430(4), 431(4), 449, 497 Margolis, L. Ye., 430(13), 434(61), 436(13), (66), 438(63), 439(66), 440(74), 441, 442, 443, 444, 446, 446, 462(116), 464, 456(118), 468, 461(127), (128), 462, 463(102), 467(138), 469(146), 470(46, 61), 473, 476, 478(138), 479, 481(166), 481,482(167),486,488(167), 492(177), 496, 497, 498, 499, 500, 501 Marine, W. M., 217(61), 225(61), 250 Marini-Bettal, G. B., 243(171), 246(171), 252 Marion, L., 217(69), 250 Mars, P., 37, 81, 83, 87, 96, 97, 104, 110, I12 Marth, P. T., 409(110), 427 Mathieu, M. V., 131, 198 Matsen, F. A., 189(137), 201 Matui, T., 436(66), 498 Maxted, E. B., 208, 208(31), 209, 209(30), 216(30), 226(30), 228(30), 229(30), 246(30), 246(30),249, 436(53), 498 McBee, E. T., 431(19), 497 McCarroll, B., 306(33), 364, 425, 426 McCarthy, F. J., 237, 252 McCormak, B. E., 431(21), 461, 464, 497 McDonald, R. S., 127(31), 127(32), 129, 198 McElvain, S. M., 213(43), 230(121), 249, 251 McElvain, S. S., 213(47), 249 McEvoy, F. J., 227(112), 250, 261 McEwen, W. E., 246(178), 253 McGee. J. D., 426 McGowan, W., 406, 427 McKim, F . L., 476(164), 501 McMillen, F. H., 216(66), 250 Medvedev, S. S., 430(6), 432(6), 434(6), 436(6), 497 Melby, L. R., 22(37), 33 Melzner, F., 413(121), 427 Menon, P. G., 83, 96, 99, 112 Merrifield, R. E., 22(38), 33 Metayer, M., 240, 252 Meyer, E. H., 413(121), 427 Meyer, W. A., 238, 239,252 Mignolet, J. C. P., 312(40), 320(64), 338, 425 Mikhailov, A. K., 206(22), 207, 249 Mikhailov, G. S., 490,501

61I

AUT3OR INDEX Mikulski, J. M., 492, 501 Miles, N., 434(37), 497 Miller, J. G., 127(36), 198 Milliken, T. H., Jr., 140(60), 167(60), 199 Mills, G. A., 140(60), 167(60), 199 Mitchell, J. W., 416(126), 427 Miyamoto, T., 244(174), 253 Miyaki, K., 244(176), 253 Miyaahita, I., 182(110), 200 Mizuguchi, J., 207(24), 239, 249 Mochel, W. E., 22(37), 33 Moldavskii, B. L., 493(181), 501 Molnar, J. P., 267, 424 Montarnal, R., 476(167), 501 Moore, 0 . E., 320(69), 425 Moore, M. A., 220(90), 250 Moore, M. L., 240(166), 252 Morard, F., 247(182), 253 Moraw, R., 28(61), 34 Morigaki, K., 188(126), 200 Morrison, J. L., 337(66), 426 Morton, A. A., 141(61), 199 Mott, N. F., 74, 112, 366, 426 Mrowec, S., 492(178), 501 Miiller, E. W., 312(42), 312(44), 320(44a), 347, 348, 360, 362, 364(44a), 366(78), 366, 370, 389(69), 414(123), 425, 426,

427 Muller, A., 28(61), (64), 34 Murat, 204(1), 206(1), 248 Murata, Y.,242(168), 243(168), 252 Murchio, J. C., 30(73), 34 Murray, K. E., 476(166), 501 N Nakai, Y., 69, 62, 84, 97, 99,111, 113, 164.

199 Naldrett, S. N., 431(24), 432(24), 497 Nash, A. W., 434(46), 436(46), 497 Neiman, M. B., 463, 499 Nelson, D. A., 230(123), 251 Neumffinn, K., 434(60), 436(60), 498 Neumark, I. E., 96, 113 Neville, 0.. 91, 112 Newman, M. S., 189(136), 190(136), 201 Nichols, M. H., 319(63), 368(63), 425 Nienburg, H., 220, 250 Nikitin, V. A., 124(26), 127(26), 128, 198 Nishimura, M., 28(63), 30(63), 34 Nolan, T. J., 219 (82), 250

Noller, H., 97, I13 Norbert, R. E., 23(42), 26(42), 33 Nordheim, L., 317(46), 425 Norton, T. R., 213(41), 249 Nottingham, W. B., 392(92), 406(104), 411(113), 426, 427 Nowaczinaki, W., 247(182), 253

0 Oblad, A. G., 140(60), 167(60), 172, 199,

200 Ochiai, E., 238(146), 242(168), 243(168),

246,252,253

Ogilvie, G. J., 63, 111 O’Grady, P., 234(133), 252 Okuda, M., 123(16), 166, 184(16), 198 Oldham, W., 241, 252 Olmer, F., 272(16), 424 Onchi, M., 412, 427 O’Neill, C. E., 164, 199 Oppenheimer, H., 431(17), 497 O’Reilly, D. E., 171(99), 178(106), 200 Orlov, E. I., 431(26), 497 Oroshnik, W., 237, 237(143), 252 Orthner, L., 221, 250 Orzeohowski, A., 431(21), 461, 464, 497 Ostermayer, F., 224(106), 226, 226, 226 (106), 232(106), 251 Ostrovskii, V. E., 476( 169), 501 Otaki, T., 68, 69, 60, 62, 63, 99, 111 Overhoff, J., 212(38), 236, 237, 238, 239,

249 P Papa, D., 251 Park, R. B., 7(26), 8(26), 33 Parkins, W. E., 272(18), 424 Parkinson, D., 366(81), 426 Parks, W. G., 434(40), 436(64), 449(110), 472(110), 497, 498, 499 Parrffivano,G.. 171(98), 200 Passoneau, J. V., 23(42), 26(42), 33 Paul, D. E., 184(120), 200 Paul, M. A., 90, 112, 140(64), 199 Paul, R., 220(86), 250 Pavlova, E. N., 132(44), 198 Pelletier, F., 224(102), 251 Peri, J. B., 124, 126(28), 127(28), 128(28),

198

612

AUTHOR INDEX

Perlowski, E. F., 224(106), 226(106), 232(106), 251 Peter, O., 119(3), (7), 198 Petroutsos, G., 83, 86, 112 Petzold, A., 241(162), (164), 242(162), 243(162), 252 Pewny, N. I., 110, 113 Phelps, A. V., 406, 407(106), 427 Philips, G., 430(3), 431(3), 432(3), 496 Phillips, W. D., 22(36), 33 Phipps, T. E., 392(93), 426 Pichler, H., 103, 113 Pickering, H. L., 122, 161, 198 Piette, L. H., 30(73), 34, 183(117), 200 Pigulevskii, V. V., 434(43), 478, 497 Pimentel, G. C., 123(12), 198 Pines, H., 189(129), 196(146), 200,201 Pink, R. C., 183(116), 184(116), 200 Pirani, M., 392(87b), 426 Pitha, J., 233, 251 Plakidina, V. A., 436(68), 498 Platonov, M. S., 92, 113 Pliskin, N. A., 467, 500 Pliskin, W. A., 62,84,66,67, 70, 71,88(97), 96, 111, 112, 113, 116(l), 123, 124, 160, 161, 162, 166,198, 199 Plotnikov, M. S., 436(68), 498 Plyshevskaya, E. G., 441, 498 Pokrovskii, V. A., 431(20), 469, 492(123), 497,500 Polyakov, M. V., 430(11), 467, 497, 500 Pon, N. O., 7(26), 8(26), 23(44), 33 Ponec, V., 104, 113 Pongratx, A., 447( 103), 448, 499 Popova, N. I., 439(67), 467, 469( 144), 498, 500 Porter, G. B., 193(141), 201 Pospelov, L. A., 490(176), 501 Prescott, L. E., 401(99), 427 Profft, E., 223(97), 251 Promyslov, M. Sh., 212(40), 249 Pugin, A., 247(182), 253 Pullin, A. D. E., 64, 99, 111, 164, 199

R Raasch, M. S., 216(60), 249 Rabinowitch, E., 27(49, 60, 61), 30(49, 60, 61), 33 Rafikov, S. R., 449(111), 499 Randall, J. T., 281, 425

Raney, M., 204, 248 Rapala. R. T., 240, 252 Raper, R., 230(124), 251 Rastrenenko, A. I., 96, 113 Ratusky, J., 217(62), (83), 250 Redhead, P. A., 278, 408(107), 411, 425, 427 Reischauer, H., 440(69), 498 Reiser, A., 217(63), 250 Renk, E., 216(60), 250 Reyerson, L. H., 430(6, Q ) , 431(6, 9, 17), 432(6, 9), 444(9), 478, 497 Reynolds, P. W., 62, 78, 79, 111 Rice, M. H., 349(71), 426 Rice, R. G., 208(29), 249 Richard, H., 430(18), 434(16), 497 Riegel, B., 230(118), 251 Rieger, E., 413(119), 427 Rieniicker, G., 36, 39, 42, 43, 46, 47, 48, 49, 61, 62, 73, 74, 78, 79, 80, 104, 110, 111, 112, 113, 493, 496, 501 Roberts, J. K., 337(86), 393(94a), 426 Roberts, R. M., 186, 200 Robertson, A. J. B., 49, 61, 64, 66, 68, 96, 111 Robin, M., 123(13), 132(13), 133(13), 198 Robinson,R. M., 206(16), 208(16), 210(33), 212(16), 213(18), 214(16), 217(16), 218 (la), 223(16), 227(16),237(16),243(16), 249 Roev, L. M., 162, 199 Roginskii, S. Z., 434(61), 437(69), 438(63), 439, 441,462,464, 466(118), 462(118), 487(138), 469, 470(46, 61), 478(138), 481, 482(167), 486, 488, 492(177). 496(61), 498, 499, 500, 501 Roiter, V. A., 439, 449, 481, 467, 470(148), 472, 498, 499, 500 Rooney, J. J., 183(116), 184(116), 200 Ropp, G. A., 91, I12 Rosenbaum, J., 189(132), (134), 189,200. 20 I Rosenblatt, E. F., 206(14), 249 Rothchild, S., 220(88), 233(132), 234(132). 250, 252 Royen, P., 84, 112 Rozengart, M. I., 438(67), 498 Rubanik, M. Ye., 462(116), 467, 488, 489 (144), 474(116), (163), 499, 500 Rubinshtein, A. M.,102,113

AUTHOR INDEX Rubstov, M. V., 221(93), 251 Rumberg, B., 28(61), 34 Ruzicke, L., 218, 218(76), 219(76), 250 Ruzicke, V., 113 Rybakova, N. A., 476(168), 486(168), 501 Rylander, P. N., 206(17), 249

613

Serratose, J. M., 127(38), 198 Setkine, V. N., 189(136), 201 Sewell, G. L., 366, 426 Shabrowa, G. M., 84, 112 Shapiro, S. L., 222(96), 251 Sherman, S. H., 196(144), 201 Shcheglova, G. G., 90, 91, 112 S Shepard, A. R., 240( l60), 252 Sabatier, P., 36, 37, 83, 110, 204, 206, 248 Sheppard, N., 124(26), 130, 131, 166, 198, Sachtler, W. M. H., 39, 42, 43, 44, 66, 66, 199 69, 61, 64, 66,66, 69, 93, 94, 104, 106, Sherlock, M., 251 107, 108, 109, 110, 111, 113, 164, Sherman, Yv. G., 471, 500 166(69), 199 Shimeda, I., 242(168), 243(168), 252 Sadikov, V. S., 206(22), 207, 249 Shindo, K., 97, 99, 113 Sakharov, M. M., 462(116), 468(116), 499 Shoikhet, P. A., 467(134), (136), 500 Saksena, B. D., 127(37), 198 Shostak, F. I., 434(38), 435(38), 436(38), San Pietro, A., 4(16), 32 497 Shreve, R. N., 433(29), 436(29), 497 Sawa, Y.,208, 208(27e), (28a), 249 Shuba, R. J., 146(63), 199 Sazonova, I. S., 462(116), 468, 499 Sidorov, A. N., 124(26), 127(26), 128, 198 Schiirfe, E., 240(167), 252 Sidorova, A. I., 164, 199 Schenck, G. O., 7(22), 33 Sigel, L. A., 439(64), 498 Scheuing, G., 228(114), 251 Sohibsted, H., 62, 103, 104, I l l Sigli, P., 99, 113 Simard, G. L., 439, 498 Schierz, E., 91, 113 Schlier, R. E., 269(6), 424 Skaletzky, L. L., 245(176a), 253 Skits, A., 204(9), 210(9), 212, 238, 239, Schmialek, P., 236(139), 252 Schmidt, F., 139(49), 199 241(9), (166), 2 W 9 ) , (166), 248, 252 Schneider, A., 171(101), 200 Skorbilina, T. G., 470(148), 472(148), 500 Scholtis, K., 447(103), 448(103), 499 Sladkov, A. M., 448(104), 459(104), 499 Schriesheim, A., 168(96), 200 Sleva, S. F., 189(133), 200 Schuit, G. C. A., 108, 113 Smith, A., 140(59), 199 Schultheiss, A., 241(163), 242(163), 252 Smith, A. E., 486(174), 501 Schultze, G. R., 469, 500 Smith, A. W., 274(22), 425 Schulz, G. J., 406(lOlb), 406,407(106), 427 Smith, C. R., 223, 251 Schwab, G.-M., 36, 39, 43, 46, 47, 72, 73, Smith, H. A., 211, 212, 249 74, 76, 76, 77, 78, 79, 80, 83, 86, 97, Smith, L., 28(68, 69), 34 110, 111, 112, 113 Smothers, W. J., 272(19), 424 Schwab-Agallidis, E., 43, 72, 73, 78, 79, Smythe, B., 479, 501 83, 110, 112 Sogo, P. B., 23(43, 44), 26(46), 27(47), 33 Schwerz, H., 408(108), 409(108), 427 Solomin, A. V., 449( 11l), 499 Schwarz, R., 433(31), 436(31), 497 Solomon, W., 243(172), 252 Schwenk, E., -? Soloway, H., 222(96), 251 Sebba, F., 83, 97, 112 Solymosi, F., 83, 84, 96, 112 Sebestven, L. G., 413(119), 427 Sorm, F., 214, 216(48), 217(62), (63), 249, Seely, G. R., 6(20), 33 250 Seeman, J., 241(162), 242(162), 243(162), Sosin, S. L., 448, 469, 499 252 Sosnovsky, H. M., 488, 501 Seibert, R. A., 213(41), 249 Sosnovsky, H. M. C., 43, 61, 62, 110, 111 Semenov, N. N., 446, 469, 499, 500 Spiith, E., 238, 252 Serafinovne, M., 224(103), 226(103), 251 Sperber, N., 223(99), 251

614

AUTHOR INDEX

Sprengler, E. P., 218(73), 250 Stach, K., 220(88a), 250 Stadnik, P. M., 487(131), 500 Stmier, R. Y.,3(10), 5(10), 32 Starks, F. W., 224( 102), 251 Steacie, E. W. R., 444(90), 499 Steele, D. R., 205(17), 249 Steger, J. F., 439(84), 498 Stepanenko, M. A., 433(34), 435(34), 438 (341, 497 Stepanov, Yu. N., 492, 501 Stepf, F., 239, 252 Stepski, M., 430(7), 497 Stern, A., 5(18), 32 Stern, E., 219(79), 250 Stern, O., 312(39), 425 Stern, T. E., 317(47), 425 Stevenson, D. P., 488(174), 501 Stoll, A., 5(17), 32 Stone,F. S., 83, 104, 110,112,113,440(76), (76), (77), 441(77), 441, 498 Stone, 0. R., 204(8), 205(18), 208(8), (18), 210(32), 212(18), 213(18), 214(8), (l8), .215(8), 217(8), (l8), 218(18), 220(8), 221(8), 223(8), (la), 227(18), 237(8), (l8), 240(161), 243(18), 248, 249, 252 Stone, H., 185, 200 Strouf, O., 230(128), 251 Stovel, H. V., 444(90), 499 Stukanovskaye, N. A., 481(128), 488(140), 500 Sudzuki, J., 238, 239, 252 Sugasawa, S., 225(110), 251 Sugino, K., 207(24), 239, 249 Suhrmann, R., 81,111 Supniewski, J. V., 224(103), 225(103), 251 Suvorov, B. V., 449(111), 499 Swam, S., 431(28), 497 Swanson, L. W., 329(82), 426 Swearingen, K., 430(6), 431(5), 432(5), 497 Sweeringen, L. E., 430(9), 431(9), 432(9), 444(9), 478, 497 Symons,M.C.R., 155,178,189(132), (134), 189, 199,200, 201 Syrkin, Ya. K., 445, 499 Szeb6, Z. G., 83, 84, 98, 112 T Tachibane, T., 123(15), 166, 184(15), 198 Taft, R. W., Jr., 225(108), 251

Takahashi, T., 219(80), 250 Tameru, K., 41, 42, 82, 84, 110 Tamele, M.W., 140(51), 187(51), 199 Taylor, H., 444, 499 Taylor, H. S., 171(98), 200 Taylor, J. B., 257(2), (3), 258, 424 Tohelitoheff, S., 220(86), 250 Teichner, S. J., 110, 113 Temkin, M. I., 90, 91, 112, 437, 475(158), 486,498,501 Terenin, A. N., 62, I l l , 124(23), 152, 184, 198, 199 Theile, H., 459, 500 Thiel, M.,220(88a), 250 Thomas, C. L., 167(84), 199 Thomas, J. B., 27(49), 30(49), 33 Thomas, J. H., 172(84). 178, 200 Thomas, V., 431(23), 497 Tikhomirova, M.,433(30), 470(30), 497 Tiley,P. F., 110,113,440(77), 441(77), 498 Tilford, C.H., 237(141), 252 Todes, 0. M., 430(13), 438(13), (56), 451, 454, 475(113), 478, 479, 497, 498, 499, 501 Tolchina, B. I., 468(140), 500 Tollin, G., ll(31, 32), 27(52, 53, 54), 33 Tomilov, V. I., 92, 113 Tomita, K., 225(109), 251 Tompkins, F. C., 418(133), 427 Topley, B., 39, 96, 110 Topping, J., 319, 425 Torney, F. L., Jr., 406(104), 427 Touillaux, R., 127(34), 198 Toursel, W., 104, 113 Townsend, J., 23, 26(42), 33 Trapnell, B. M. W., 305(32), 418, 425, 427, 440, 444, 498, 499 Trautmann, G., 73, 74, 112 Tretner, W., 413(120), 427 Trevalien, P. A., 155, 199 Troscianiec, H. J., 224(105), 225(105), 230 (122), 232(105), 251 Trotzenko, M. A., 487(135), (138), 500 Trueblood, K. N., 123(13), 132(13), 133 (13), 198 Truitt, P., 230(119), (120), 251 Tsuda, K., 238(145), 252 Tsushme, S., 238, 239, 252 Tuley, W. F., 253(180), 253 Turkevich, J., 444, 499

AUTHOR INDEX

515

Turovskii, G. Ye., 442, 460, 498 Twigg, G. H . , 431(18), 451, 454, 475(18),

Vol’kenstein, F. F., 445, 446(101), 485,

497 Tye, F. L.,168(88), 178(88), 199 Tzellinskaya, T. F., 441, 498

Volkovskaya, N. S., 461(128), 500 Volmsn, D. H., 187, 200 Voltz, S. E., 140(59), 199 von Ardenne, M., 396(98a), 414(98a), 427 von Braun, J., 236(134), 241(162), (163), (1641, 242, 243(162), (173), 252, 253 Vorhees, V., 213(46), 249 Vorozhtzov, N. H., 434(41), 497 Vysotzkii, Z . Z . , 467(133), 500

499, 501

U Umio, S., 236(136), (138), 252 Unger, S., 48, 111 Urbach, F., 272, 277, 424 Urizko, V. I., 467(132), 500 Ushakov,M. I., 211(36), 212(39), (40),249, W 436(57), 498 Wadsworth, M. E., 172, 200 Ushakova, V. P., 439(65), 449, 470, 472 Wagener, J. S . , 267, 268, 396(96), 406 (ME), 498,499,500 (103), 409(110), 424, 426, 427 Utsugi, H., 320(58), 358(58), 359(58), 425 Walba, H., 168(93), 200 Walden, P., 168(89),200 Waldschmidt-Leitz, E., 2 17(67), 244(67),

V

Vainstein, F. M., 430(11), 442, 460, 497,

498 Van Campen, M. G., Jr., 237(141), 252 Vendenbelt, J. M., 132(45), 198 van der Meij, P. H., 183(119), 184(119),

200 Vanderslice, T. A., 413(117), 427 van Niel, C. B., 3(7), 32 van Reijen, L. L., 39, 43, 64, 65, 66, 93, 94, 104, 105, 107, 108, 109, 110, 111, 113, 154,156(69), 199 Varadi, P. F., 413(119), 427 Vassiliev, V. N., 461, 500 Veer, W. L. C., 237, 252 Vedeneeva, N. E., 165, 199 Veltistov, K. K., 436(58), 498 Verhoek, F. H., 46, 47, 92, 93, 111 Vermel, E. E., 439(67), 467(139), 498, 500 Verwey, I. W., 480, 501 Vipond, H. J., 230(124), 251 Vladimirova, V. I., 481(165), 501 Volter, J., 42, 45, 46, 48, 110, 111 Voevodskii, V. V., 445, 459, 494,499, 500, 501 Voge, H. H., 167(85), 167(86), 172(85), 199, 486(174), 501 Vogel, J., 124(22), 158, 159(75), 160(75), 161(75), 162(22), 199 Vogl, J., 41, 95, 110, I98 Vogl, O., 216(56), 250

250 Walker, A. GI., 208(31), 209, 249 Walker, G. N., 205(18), 210, 213(18), 220 (go), 237, 249, 250 Waller, P. D., 217(61), 225(61), 250 Walling, C., 140, 199 Walsh, W., 434(37), 497 Walton, D. K., 46, 47, 92, 93, 111 Wan, S. W., 433(31a), 449, 459(112), 470 (31a), 475(156), 497, 499, 501 Wang, H., 434(50), 435(50), 498 Ward, H. F., 91, 113 Ward, R. L., 188(126), 200 Waring, W., 38, I10 Warner, G. H., 230(122, 123), 251 Wartburg, A. F., 189(133),200 Watts-Tobin, R. J., 365, 426 Webb, A. N., 123(20), 167, 178, 178(20), 179(20), 185(20), 186, 191, 198 Webb, J. L.,216(59), 250 Wedler, G., 61, 111 Weijland, W. P., 18(33), 33, 183(112), 183(113), 185(113), 200 Weinberger, A. J., 91, 112 Weinhart, H. W., 417, 427 Weinhouse, S., 216(57), 250 Weiser, R., 216(56), 250 Weiss, J. M., 433(32), (33), 497 Weissman, S. I., 182(111), 183(111), 184 (120), 188(124), 200 Weitz, E., 139(49), 199

616

AUTHOR INDEX

Welborn, R. W., 433(29), 436(29), 497 Weller, S. W., 88, 112 Werber, T., 492, 501 Wessing, G., 73, 74, 112 Wheeler, A., 393(94b),426 Wheeler, R. E., 430(14), 497 Wheland, G. W., 217, 218(64), 221, 250 Whetten, N. R., 417(130), 427 White, D. R., 319(48), 425 Whitmore, F. C., 100, 113 Wibaut, J. P., 212(38), 233(129, 130), 236, 237, 238, 239, 249, 251, 252 Wicke, E., 87, 89, 112 Wieland, H., 243(170), 244(170), 252 Wilkins, M. H. F., 281, 425 Wilson, J. N., 486, 501 Wilstatter, H., 5(17), 32 Willstatter, R., 217, 244(67), 250 Winstein, S., 195(144), 201 Winter, E. R. S., 442, 498 Winterfeld, K., 227(111), 236, 251, 252 Winterhalder, L., 228(114), 251 Wiseman, P. A., 431(19), 497 Wissing, F., 213(44), 249 Witkop, B., 239, 240, 252 Witt, H. T., 28(61), (64), 34 Wittcoff, H., 230(118), 251 Wojcik, B., 204, 210(3), 216(3), 236(3), 247(3), 248 Wolfsberg, M., 91, 113 Woodward, R. B., 5(19), 32, 240, 242, 247,

252, 253 Wortman, R., 320(57), 326(57), 425 Wright, P. E., 246(178), 253 Wulff, C., 241(166), 243(166), 252 Y

Yakhontov, L. N., 221(93), 251 Yakovleva, E. V., 212(39), 249

Yang, A. C., 123, 198 Yant, W. P., 430(10), 497 Yaroslavskii, N. G., 62, I I I , 124(23), 198 Yarwood, J., 392(87b), 426 Yarzhemskaya, E. Ya., 434(43), 478(43),

49 7 Yates, D. J. C., 119(8), 123(24), 124(26), 127(8), 127(24), 127(33), 128(33), 129(8), 130, 131, 133, 154, 156(33), 156, 157, 198, 199 Yatcs, J. T., Jr., 151, 152(66), 153(66), 199 Yatsenko, S. V., 221(93), 251 Yokozawa, Y., 182(110), 200 Yoschikawa, K., 432(27), 497 Young, A. H., 406(102), 427 Young, G. J., 123(18), 127(18), 198 Young, J. R., 417(128), (130). 427 Young, R. D., 365(78), 426 Young, W. G., 195(144), 201 Yula, R. W., 434(40), 497

Z Zabel, R. M., 412(115), 427 Zalkind, Yu. S., 433(28), 434(46), (47), 497 Zapletal, V., 104, 113 Zelinsky, N. D., 206, 249 Zenchelsky, S. T., 145(63), 199 Zenitz, B. L., 240, 252 Zettlemoyer, A. C . , 150, 199 Zhabrova, G. M., 438(63), 444(91), 481,

498,499,501 Zhigailo, Ya. V., 439(65), 498 Zieger, G., 28(61), 34 Zimakov, P. V., 457, 499 Zobel, F., 236(134), 252 Zolotarev, S., 433(28). 497

Subject Index A Acetylene, adsorption on nickel, 166 Activation energy, for formic acid decomposition, 41, 43, 47, 49, 61, 64, 78, 83,

Alkanes, spectra of adsorbed, 130 Alumina, 82, 96, 97, 99, 127, 161 spectra for formic acid adsorbed on, 99 Aluminium formate decomposition, 104 oxide, see Alumina 2-Aminopyridine, hydrogenation of, 22 1 Ammonia, spectra of adsorbed, 133 Aniline, spectra of adsorbed, 132, 133, 163 Anthracene. spectra adsorbed, 186 Aromatics, oxidation of, 448 spectra of adsorbed, 131 Arylalkanes, spectra of hydrocarbons chemisorbed on, 167 Arylalkenes, spectra of hydrocarbons chemisorbed on, 167

84, 91, 97, 101, 166

for oxidation, 470-476, 483, 484 for oxygen adsorption, 440, 442 Adatoms, ion microscope observation of, 367

Adenosine triphosphate, role in photosynthesis, 4 Adsorption, calorimetric measurements in, 304

cell design, 124 charging of surface in, 446 effects on spectra, 118 extinction coefficients in, 119, 133 flow systems in study of, 267 heat of oxygen, 440 high temperature, of HCOOH on metals, 69

iodine, 136 isotherms, 306 kinetics, meaaurements in, 266, 307, 310 in closed systems, 269 in flow systems, 263 light, of chromatophores, 29 measurements in macroscopic rate studies, 266 of carbon monoxide, 161, 381 of hydrogen, 381 of isobutane, 187 of p-nitrophenol, 134 of oxygen, 381, 440 of rare gases on metals, 334 of xenon, 334 spectra of acetylenes, 167 spectral changes induced by illumination, 28 vacuum techniques for study of, 389, 394

Aliphatic hydrocarbons, spectra of chemisorbed, 186 Alizarin, adsorption on alkaline earth halides, 137

B Bayard-Alpert inverted ionization gauge, 408

Benzene, spectra of adsorbed, 132, 133 Benzene oxidation, 448 over V806, kinetics of, 449, 472 Benzidine, spectra of-chemisorbed on bentonite, 166 Bronsted acids, in cracking catalyst, 163 Butene-2, adsorption on silica-alumina, 186

C Carbon monoxide, chemisorption on supported metals, 161 Carbonium ion formation, 168, 170, 172, 174-178, 186, 187, 191, 194

Catalytic activity and Fermi level, 480 Catalytic decomposition of formic acid, 36-110

Catalytic oxidation of hydrocarbons, 429496

Catalytic systems, study by spectrophotometry, 116-197 surface structure study in, 126 Catalytic ‘superactivity’, on nickel and copper, 63

617

518

SUBJECT INDEX

Catalysts, chemisorption of hydrocarbon on oxidation, 444 for pyridine and quinoline hydrogenation, 204 mixed, for oxidation, 494 modification in hydrocarbon oxidation, 479 oxygen adsorption on, 440 poisoning, in pyridine hydrogenation, 208 semiconducting, 466 Cell, conductivity, for measurement of charge separation, 10 design of adsorption, 124 Charge separation, photoinduced, 10 processes, 16 Chemisorption, electron transfer in, 440 and Fermi level, in oxidation, 480 mechanism of ptwaffin, 174 of butene-2, 186 of carbon monoxide, 161 of hydrocarbons on oxidation catalysts, 444 of oxygen, on metals, 440 on semiconductors, 441 of propene, 446 studies by spectrophotometry, 150 Chlorobenzene, spectra of adsorbed, 133 Chlorophyll, photochemistry of, 6 Chloroplast luminescence, 27 Chloroplast materials, ESR in, 23 Chloroplasts, quantum conversion in, 1-32 Chromatophores, light adsorption of, 29 Chromium formate decomposition, 104 Chromium-manganese spinel, aee Spinels Cold cathode gauge, 412 Conductivity, electrical, in adsorption, 463 dark, in violanthrene systems, 15 in model for chlorophyll energy transfer, 11 Complexes, surface, removal of in field ion microscopy, 380 Copper, 42, 49, 431, 440 absorption of CO on, 161 Copper-Gold alloy, in formic acid dehydrogenation, 72 Copper oxide, as catalyst, 434, 440, 466, 473

Copper-Palladium alloy, in formic acid dehydrogenation, 73 Cracking catalysts, Bronsted and Lewis acids in, 163 spectra of adsorbates on, 139, 163 aee also Silica-alumina Crystallographic surface orientation, effect on kinetics, 47 Cupric oxide, kinetics of propene oxidation over, 473, 476 Cuprous oxide, chemisorption, on 445 mechanism of propene oxidation over, 473. 476

D Dehydration, of formic acid on metals, 92 on oxides, 96 Dehydrogenation, of formic acid on oxides, 81 Desorption, and field evaporation, 366 flash meaaurement of, 271 heat of, 281 kinetics, 273, 276 Dimethylacetylene, spectra of adsorbed, 167 2, 6-Dimethylhexane, from isobutane, 197 Diffusion, by field emission microscopy, 326, 330 by field ion microscopy, 370 effect on kinetics of hydrocarbon oxidation, 468 surface, determination of, 330 2-Diethylaminopyridine, hydrogenation of, 222 N, N-Dimethylaniline, spectra of adsorbed, 133 1,l-Diphenylethylene, spectra of chemisorbed on silica-alumina, 178 E Electron spin resonance (ESR), in chloroplast materials, 23 measurement of adsorbates on silicaalumina, 183 spectra of chemisorbed hydrocarbons, 167 transfer in chemisorption, 446 Emerson effect, 27 Emission intensity, in field electron micro. scopy, 312

61 9

SUBJECT INDEX

Energetics, of surface processes, in adsorption, 304 Energy transfer in quantum conversion, 7 1-Ethyl-2,3-dimethylpyridineby N-alkylation, 207 Ethylene oxidation, over silver, 454, 458, 474 over vanadium, 469 Ethylene oxide, kinetics of formation over silver, 474 Extinction coefficients, in adsorption, 119, 133

F Fermi level, and catalytic activity, 480 Field electron microscopy, emission in, 312 experimental consideration in, 330 Field evaporation and description, 356 Field ion microscopy, 347 high field effects in, 360 surface cleaning in, 380 surface complexes in, 364 Flash desorption measurements, 27 1 temperature control in, 298, 300 of gaa density, 296 Formic acid decomposition, 38, 154 activation energy for, 41, 47, 49, 51, 64, 91, 97 catalytic, 35-1 10 Hammett acid function in, 90 isotope effects in, 56 over Pt-SiO, and Ni-SiO,, 95 dehydrogenation, 39, 81 dehydration, 90, 92, 96 high temperature adsorption of, 69 Fowler-Nordheim equation, 314, 316

G Gas purification techniques, 417 Gold catalysts, 45, 60, 431

H Hammett acid function, 139 in formic acid decomposition, 90 Hammett dyes, spectra of adsorbed, 141, 142 Heat of adsorption, of oxygen on semiconductors, 446 Helium, field ionization, of 349

Hume-Rothery alloys in formic acid decomposition, 74 Hydration-dehydration, effect on hydroxyl surface groups, 127 Hydrocarbons, chemisorption of, 155 oxidation of, on metals, 430 on metal oxides, 433 Hydrogen, chemisorption, on nickel, 155 on Pt, 150 on rhodium films, 151 Hydrogenation, of aminoquinolines, 245 of aminopyridines, 220, 221, 222 of dipyridyls, 223 of esters and amides in the pyridine series, 216 of hydroxypyridines, 217 of hydroxyquinolines, 244 of isoquinoline, 238, 241 of nicotinic acid, 214, 223 of picolinic acid, 214 of pyridine carboxylic acids, 213 of pyridines and quinolines, 203-248 of pyridyldkanoic acids, 216 of pyridylalkanols, 220 of pyridylalkylamines, 222 of pyridinium salts, 224 of quinoline, 238, 241 of quinoline and isoquinoline carboxylic acids, 242 of quinoline esters and amides, 243 of quinolyl ketones, 245

I Infrared study of formic acid dehydration, 97 of formic acid decomposition, 61, 84, 97 Infrared spectra of deuterated formic acid, 71 Iodine, adsorption on calcium floride, 136 Ion exchange resin, for formic acid decomposition, 101 Ionic salts, spectra of adsorbates on, 134 Ion image, formation of in field ion microscopy, 353 Ionization, field, of helium, 349 gauges, 404, 408, 412 Iron catalysts, adsorption of CO on, 161 formate decomposition, 104

520

SUBJECT INDEX

Isobutane, adsorption of, on silica-alumina, 187 conversion to 2, 6 dimethylhexane over AICl,, 197 Isoquinoline, hydrogenation of, 239 Isotopic oxygen exchange, 442

K Kinetics, adsorption, 269, 2 6 5 266 in formic acid decomposition, 40 pretreatment effect on, 47 product formation effect on, 46 methods in surface, 266-424 of aromatic oxidation, 469 of hydrocarbon oxidation, 468-470,472, 476

of olefin oxidation, 469 radioactive tracer method in determination of, 463, 468 Knudsen manometer, 412

L Lewis acids, in cracking catalysts, 163 Luminescence, chloroplast, 27 M Magnesium chromite, aee Spinels Magnesium formate decomposition, 104 Magnesium oxide, 82, 161 Maleic anhydride, kinetics of formation, 472

Manganese dioxide, 436, 443 Manganese-cobalt spinel, aee Spinels Mechanism, of catalytic oxidation, 439, 461, 468

of formic acid decomposition, 106 Metal alloys, in formic acid dehydrogenation, 72 Metal oxides, formic acid decomposition on, 81 Methane, spectra of adsorbed, 130 Methylacetylene, spectra of adsorbed, 167 Microscopy, field electron, 311 field ion, 347 “Mixed” catalysts in hydrocarbon oxidation, 493 Mobility, of atoms or molecules, determination of, 442

N Naphthalene, oxidation kinetics of, 470 N-alkylation of pyridine, 208 Nitrogen. adsorption observed by field ion microscopy, 367, 369 desorption for tungsten, 284, 380 Nickel catalysts, 46, 61, 64, 92, 161, 166, 166, 431

Nickel formete decomposition, 104 Nickel oxide, 42, 436, 440 Nickel-silica catalyst, formic acid decomposition with, 66, 164 in hydrocarbon adsorption, 166 Nicotinic acid, hydrogenation of, 214, 223 p-Nitrophenol, adsorption on calcium fluoride, 134 4-Nitroaniline, spectra of adsorbed, 133 0

Olefh oxidation, 464 Oxidation of aromatics, 448 of benzene, 472 of hydrocarbons, 429-496 of olefins, 469 of naphthalene, 470 over semiconducting catalysts, 466 reaction steps in, 467 Oxygen, activation energy for adsorption, of, 440 P Paliadium catalysts, 161, 430 Paraffin chemisorption, mechanism of, 174 p-Phenylenediamine, adsorption spectra of, 166 Photochemistry of chlorophyll in solution, 6 Photoinduced charge separation, 10 Photophysicd effects, 7 Photosynthesis, energy sources for, 3 organic semiconductors and, 22 role of adenosine triphosphate in, 4 aee also Quantum conversion Phthalocyenine, &B a model for chlorophyll energy trensfer, 9 semiconductor properties of, 9 Physical adsorption, spectrophotometric study of, 130 Picolinic acid, hydrogenation of, 214

521

SUBJECT INDEX

Piperidine, as pyridine hydrogenation poison, 209 Platinum, 160, 161, 430, 442 Potential measurement, in field ion microscopy, 366 Pressure measurements in ultrahigh vacuum systems, 404 Propene adsorption isotherms, 446 Propene oxidation, 462, 467, 466, 469, 473 Pyridine, hydrogenation, side reactions in, 206 N-alkylation of, 208 spectra of adsorbed, 133 Pyridine nucleotide (TPNH), 4 Pyrolysis, of bulk nickel formate, 93 Pyrophosphate linkage in ATP, 4

Q Quantum conversion, in biological material, 29 in chloroplasts, 1-32 photoinduced, 10 Quinolines, effect of substituents on reduction of, 241

R Rate studies, macroscopic, 266 Reduction, of pyridines, i n the presence of a benzenoid ring, 236 of other heterocycles, 237 of ketone function, 227 of nitrile function, 234 af pyridine salts, 209 Resistance, change of electric, during formic acid decomposition, 61 msonance ppectrqmeters in ywuum

weaasremeqt, 413 Rhodium cetdysts, 161

0 Welegtivity, in pyridipe recluction, 826 Semiconductom, chemisorption on, 441, 446

in formic acid decompoaition, 64 organic, in photosynthesis, 22 oxidation over, 466 Silica, 82, 96, 127, 130, 132, 142, 147, 149 Silica-alumina, 118, 142, 147, 149, 166, 173, 178, 180, 186, 193

carbonium ion formation in, 170, 174178, 186, 187, 191, 194 Silver, see also silver catalysts in ethylene oxide formation, 462 in formate decomposition, 104 in propylene oxidation, 467 Silver catalysts, 42, 431, 442, 464, 462, 474 ethylene oxidation over, 460, 464, 474 Spectra, effect of adsorption on, 118 of adsorbates on cracking catalysts, 139, 163, 178 on ionic salts, 134 on silica, 142, 147 on silica-alumina, 118, 142, 147, 149, 163, 166, 173,186 of adsorbed alkanes, 130 amines, 166 ammonia, 133 aniline, 163 anthracene, 186 aromatics, 131 arylalkanes, 167 arylalkenes, 177 benzidine, 166 1, 1-diphenylethylene, 178 triphenylchlorometha, 161 1,3,6-trinitrobenzene, 162 triphenylmethane, 161, 173 of Hammett dyes, 141, 142 Bpectral efficienoy, appsrent, 26 E(peotrophotometry, for Ptudy at' cstalyet systems, ii6-1e7 Bpectroscapy, optical, 116 Spectrum, ESR, of o-chlqranil, 18 Bpinels, propene oxidation over. 466, 478 aterip effeots in pyridine hydrogenation, 810 Burface, kinetic meassremento by ultrahigh vacuum techniques, 391 structure etudiee, 126 @urfaoechsrge formation, in adsorption, 446 BwfBace dlfisiop, observation by Add electrqn mioroscopy, 330 T Thiophenol, spectra of adsorbed, 133 Tin formate decomposition, 104 Titanium dioxide, 127 Trigonelline, hydrogenation of, 216

622

SUBJECT INDEX

1,3,5,-Trinitrobenzene,adsorption spectra of, 162 Triphenylchloromethane, adsorption spectra of, 161 Triphenylmethane, adsorption on silicaalumina, 161, 173 Toluene, oxidation of, 448, 449 Tungsten, adsorption of xenon on, 334

U Ultrahigh vacuum techniques in surface kinetic measurements, 391

V Vaouum requirements for adsorption study. 386, 392 Vacuum techniques for adsorption study, 389, 391, 394

Vanadium catalysts, 469 Vanadium pentoxide, aa catalyst, 433,440,449, 470, 472 chemisorption on, 445 oxidation over, 449, 470, 472 Violanthrene systems, dark conductivity measurements on, 15

W Work function, in adsorption, 463 in field electron microscopy, 316, 332 in oxidation studies, 482, 489

Z Zinc formate decomposition, 103, 104 oxide, chemisorption on, 481