Advances in Catalysis and Related Subjects, Volume 4

ADVANCES IN CATALYSIS AND RELATED SUBJECTS

VOLUME I V

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ADVANCES IN CATALYSIS AND RELATED SUBJECTS VOLUME I V EDITED BY

V. I. KOMAREWSKY

W. G. FRANKENBURG

Chicago, I l l .

Lancaster, Pa.

E. K. RIDEAL London, England

EDITORIAL BOARD

H. S. TAYLOR

P. H. EMMETT

Princeton, N . J .

Pittsburgh, Pa.

1952 ACADEMIC PRESS INC., PUBLISHERS NEW YORK, N. Y.

Copyright, 1952, by ACADEMIC PRESS INC. 125 East 23rd Street New York 10, N. Y.

All Rights Reserved NO PART OF T H I S BOOK MAY B E REPRODUCED I N A N Y FORM, BY PHOTOSTAT, MICROFILM, OR A N Y OTHER MEANS, WITHOUT WRITTEN PERMISSION FROM THE PUBLISHERS.

Library of Congress Catalog Card No. 49-7755

PRINTED I N THE U N I T E D STATES O F AMERICA

CONTRIBUTORS TO VOLUME I v

J. H. BAXENDALE, Chemistry Department, University of Manchester, England

R. P. BELL,Balliol College, Oxford, England PHILIPGEORGE, Department of Colloid Science, University of Cambridge, England GEORGED. HALSEY,Mallinckrodt Chemical Laboratory, Harvard University, Cambridge, Massachusetts*

R. C. HANSFORD, t Socony-Vacuum Laboratories, Paulsboro, New Jersey TERRELL L. HILL,Naval Medical Research Instgtute, Bethesda, Maryland HELMUT PICHLER, Hydrocarbon Research, Inc., Trenton, New Jersey HERMAN E. RIES,JR., The Research and Development Department, Sinclair ReJining Company, Harvey, Illinois JOSEPHWEISS, University of Durham, King’s College, Newcastle-uponTyne , England

* Present address: Department of Chemistry, University of Washington, Seattle, Washington. t Present address: Research Center, Union Oil Company of California, Brea, California. V

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PREFACE I n preparing this fourth volume of the Advances in Catalysis, the editors have continued their efforts t o present a picture of important new contributions by various workers t o our scientific and technical knowledge of catalysts, catalytic reactions, and related fields. To some it may appear desirable that all the papers of this and preceding volumes be in complete harmony, being variations of one and the same “leitmotif” of an universal theory and of a resulting technical mastery of catalytic phenomena. However, the present state of catalytic research has not yet led t o such an universal theory, and any enforced agreement among the papers appearing in these series would, in the editors’ opinion, hamper the free expression of opinion by the individual authors. The diversity of the material presented in the Advances is a natural result of the great number of variables that enter into the over-all effect of catalytic action, particularly of catalytic action on solid catalysts. The chemical composition of solid catalysts and of their surfaces, their fine structure and electronic configuration are equally important for a full understanding of catalytic reactions as the specific effects exerted by the catalysts on the reacting molecules, such as their physical and chemical adsorption, as well as their distortion, polarization or ionization under the influence of the catalyst surface. There are further the chemical and physical processes that follow these primary acts of activation on the catalyst surface and which include the formation of intermediates, side reactions, and desorption of the various products, often combined with changes of the catalyst surface. There are the effects of catalyst poisons and of catalyst aging. There are the transport processes t h a t control the flow of all the species of reactants t o and from the specific sites of the catalyst surface, diffusion processes through pores of the catalyst and many other factors. Each of them influences the one or other step in the complex sequence of elementary processes, and changes, often profoundly, the over-all results that we call a LLcatalytic reaction.” The papers in the present volume include contributions dealing with catalytic effects in homogeneous media which, in some respects, are less complicated than heterogeneous catalytic reactions. We trust that this and the preceding volumes of the Advances will be of substantial help in keeping both the specialist and the novice in the field of catalysis informed on its various aspects, and on the methods used in its scientific exploration and technical exploitation. W. G. FRANKENBURG V. I. KOMAREWSKY April, 1952 E. K. RIDEAL vii

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CONTENTS CONTRIBUTORS TO VOLUME IV . . . . . . . . . . . . . . . . . . . . . .

v

EDITORS’PREFACE . . . . . . . . . . . . . . . . . . . . . . . . . . .

vii

Chemical Concepts of Catalytic Cracking

BY R . C . HANSFORD. Socony-Vacuum Laboratories. Paulsboro. New Jersey

.

I Introduction . . . . . . . . . . I1. Commercial Cracking Catalysts . I11. Catalyst Evaluation . . . . . . IV Chemistry of Catalytic Cracking . References . . . . . . . . . . .

.

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

1 4 8 . . . . . . . . . . . . . . . . . 14 . . . . . . . . . . . . . . . . . 29 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Decomposition of Hydrogen Peroxide by Catalysts in Homogeneous Aqueous Solution

BY J . H . BAXENDALE. Chemistry Department. Un.iversity of Manchester. England

I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . I1. Halides and Halogens . . . . . . . . . . . . . . . . . . . . . . . I11. Iodate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

IV . V VI . VII . VIII . I X.

Iron Salts . . . . . . . . . . . . . . . . . . . . . . . . . . . .

. Ferrocyanide and Ferricyanide Copper Compounds . . . . Permanganate . . . . . . Chromate . . . . . . . . . Molybdate and Tungstate .

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

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

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31 35 43 46 67 71 73 75 80

Structure and Sintering Properties of Cracking Catalysts and Related Materials

E. RIES. JR.,The Research and Development Department. BY HERMAN Sinclair Refining Company. Harvey. Xllinois I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . 88 I1. Experimental Procedures and Interpretation . . . . . . . . . . . . . 90 I11. Structure and Sintering Properties of Representative Cracking Catalysts 99

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IV Structure and Sintering Properties of Various Forms of Materials . . . . . . . . . . . . . . . . . . . V. Summary and Conclusion . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . ix

Silica and Related . . . . . . . . 124 . . . . . . . . 146 . . . . . . . . 147

CONTENTS

X

Acid-Base Catalysis and Molecular Structure

.

BY It P . BELL.Balliol College. Oxford. England

I . Introduction . . . . . . . . . . . . . . . . . . I1. The Empirical Laws of Acid-Base Catalysis . . . I11. The Molecular Mechanism of Acid-Base Catalysis IV . The Velocity of Acid-Base Reactions . . . . . . .

. . . . . . . . . 151 . . . . . . . . . . 153 . . . . . . . . . . 164 . . . . . . . . . . 192 V The Importance of Molecular Structure . . . . . . . . . . . . . . . 201 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . 207

.

Theory of Physical Adsorption BY TERRELL L . HILL.Naval Medical Research Institute. Bethesda. Maryland

I. I1. I11. IV .

Introduction . . . . . . . . . Monolayer Adsorption . . . . Multilayer Adsorption . . . . . Thermodynamics of Adsorption References . . . . . . . . . . Glossary . . . . . . . . . . .

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

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. 212 . 212 . 225

. 242 . 255 . 258

The Role of Surface Heterogeneity in Adsorption

.

BY GEORGE D HALSEY.Mallinckrodt Chemical Laboratory. Ha.rvard University. Cambridge. Massachusetts

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . I1. Chemisorption . . . . . . . . . . . . . . . . . . . . . . . . I11. Physical Adsorption . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . .

. 259

.

259

. 263 . 269

Twenty-Five Years of Synthesis of Gasoline by Catalytic Conversion of Carbon Monoxide and Hydrogen BYHELMUTPICHLER.Hydrocarbon Research. Inc., Treiaton. N . J .

I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . 272 11. Historic Review of Research Work Concerning the Catal.ytic Conversion of Carbon Monoxide and Hydrogen to Higher Hydrocarbons . . . . . 273 111. Solved and Unsolved Problems of Hydrocarbon Synthesis . . . . . . . 319 References

. . . . . . . . . . . . . . . . . . . . . . . . . . . .337

The Free Radical Mechanism in the Reactions of Hydrogen Peroxide BY JOSEPHWEISS. University of Durham. King's C'ollege. Newcastle.upon.Tyne. England

I . The Reaction between Hydrogen Peroxide and Ferrous Ions . . . . . . 343 I1. The Reaction between Hydrogen Peroxide and Ferric Ions . . . . . . 347 I11. The Reaction between Hydrogen Peroxide and Cuprie Ions . . . . . . 351

CONTENTS

xi

IV . The Decomposition of Hydrogen Peroxide a t Different Metal Surfaces . 352 V. The Photochemical Decomposition of Hydrogen Peroxide . . . . . . . 354 VI The Decomposition of Hydrogen Peroxide by Ionizing Radiations . . . . 357 VII . The Reaction between Ozone and Hydrogen Peroxide and the Decomposition of Ozone in Aqueous Solution . . . . . . . . . . . . . . . . 358 VIII Detection of Free OH Radicals . . . . . . . . . . . . . . . . . . . 361 I X Some Thermodynamics Data Concerning the Radicals OH and H 0 2 . . 361 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . 364

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. .

The Specific Reactions of Iron in Some Hemoproteins BY PHILIP GEORGE.Department of Colloid Science. University of Cambridge. England

.

I Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . I1. General Chemistry of Hemoglobin. Myoglobin. Peroxidase. and Catalase I11. Hemoglobin and Myoglobin Autoxidation and Other Reactions . . . . 1%'.Peroxidase Reactions . . . . . . . . . . . . . . . . . . . . . . . . V. CataIase Reactions . . . . . . . . . . . . . . . . . . . . . . . . VI. The Mechanism of these Hemoprotein Reactions . . . . . . . . . . . VII . Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . .

367 369 381 388 393 404 424 425

AUTHORINDEX . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

429

SUBJECTINDEX . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

440

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Chemical Concepts of Catalytic Cracking R. C. HANSFORD Socmy-Vacuum Laboratories, Paulsboro, New Jersey

CONTENTS Page

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13

111. Catalyst Evaluation 2. Physical Tests. ..

2. Mechanism of Catalytic Cracking.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3. Present Status.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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

8

21

27 29

I. INTRODUCTION It is the principal purpose of this paper to present to the reader a broad picture of the chemistry of catalytic cracking of petroleum, emphasizing the chemical nature of the catalyst in relation to the general reactions which differentiate this type of cracking from thermal cracking in the production of gasoline. A detailed review of the reactions of pure hydrocarbons in the presence of cracking cataIysts has been given elsewhere (1). In many respects, the term “cracking,” as applied to modern processes for converting various petroleum fractions to gasoline, is a misnomer. The term implies a simple decomposition of high molecular weight compounds to lower molecular weight products, without substantial rearrangement in the structure of the fragments. Thus, the decomposition of a long-chain paraffin molecule might be expected to give mainly shorter-chain paraffins and olefins by simple rupture of carbon-carbon bonds. Originally, the production of gasoline from the heavier fractions of petroleum was accomplished by a thermal destructive distillation process, and this was aptly called cracking, since the products were allowed to escape from the system as soon as they became sufficiently low in molecular weight. In other words, such a cracking process does not involve extensive secondary reactions, which strictly speaking are not 1

2

R. C. HANSFORD

cracking reactions. Most modern cracking processes, both catalytic and noncatalytic, depend to a considerable extent on other reactions besides true cracking t o produce high quality gasoline. Such reactions as polymerization and alkylation, which are the reverse of cracking, as well ~ h srearrangement reactions (isomerization), are involved in these processes. In particular, the superiority of catalytic cracking over thermal cracking is largely dependent on reactions other than those which involve simple bond rupture. As a matter of fact, catalytic cracking probably would not be able to compete economically with thermal cracking except for the higher quality of its products, made possible by reactions such as isomerization and hydrogen transfer. It is quite erroneous to conclude that the catalytic process owes its enormous success and rapid growth to the purely cracking functions of the catalyst, i.e., to a higher yield of gasoline over that obtainable by thermal cracking. In spite of the fact that the name “cracking” is none too descriptive and accurate, particularly in catalytic cracking, it is a well-established term, and its use seems appropriate enough if applied only to those processes which result in a substantial reduction in molecular weight of a considerable portion of the petroleum fraction being processed. In practice, this means the conversion of petroleum fractions, which boil principally above 210”C., t o gasoline, which boils in the range of O-21O0C., i.e., hydrocarbons from Cq to CI2,inclusive. This practical definition of cracking excludes such gasoline manufacturing processes as “reforming,” which convert petroleum fractions already boiling largely within the gasoline range t o a gasoline of higher quality, principally by dehydrogenation and isomerization reactions. A considerable amount of cracking may also occur in reforming, particularly in thermal ref0 rming, although generally about 80 to 90% of the product remains in the gasoline boiling range. The petroleum fraction which is normally converted1 t o gasoline by cracking is a broad cut, known as “gas-oil,” which boils mostly between kerosene and asphalt, i.e., between about 275°C. and 600°C. or higher. However, in catalytic cracking practices the boiling range of the charge stock may include all the kerosene (210-275°C.) and sometimes even the “straight-run ” gasoline (gasoline present in the natural crude oil). To some extent, the straight-run gasoline is “reformed” during catalytic cracking, a t least in some catalytic cracking processes. The general trend in modern catalytic cracking is to convert as much as possible of the crude boiling above the gasoline range, i.e., hydrocarbons boiling between gasoline and heavy asphaltic residuum (roughly 200-700°C.). The development of catalytic cracking as a major process for manufac-

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

3

turing motor gasoline has actually occurred largely since the end of World War 11, although during the war there was a very rapid expansion of catalytic cracking plants for the production of aviation gasoline. The first commercial unit to go into operation, employing the basic principles of modern catalytic cracking, was a 2000-barrel unit of the fixed bed type. This prototype of the Houdry process began commercial operation in 1936 and at the beginning of World War I1 in 1939 the total Houdry capacity in the United States was about 100,000 barrels or 15,000 short tons per day (Houdry, Burt, Pew, and Peters, 2). Although before the entry of this country into the war in 1941 much of the catalytic gasoline being produced was blended into premium grade motor (automobile) gasoline, a major interest in catalytic cracking by the Houdry process was in the production of aviation gasoline. The wartime demand for high grade military aviation gasoline greatly accelerated new developments in catalytic cracking processes, and two important improved processes were introduced early in the war. One process employs a moving bed of granular catalyst, which circulates continuously through the reactor and regenerator (kiln), and is known as the Thermofor catalytic cracking or TCC process (Simpson, Evans, Hornberg, and Payne, 3). The other employs a circulating “fluidized” bed of powdered catalyst and is known as the fluid process (Murphree, Brown, Fischer, Gohr, and Sweeney, 4). The rapid expansion of these processes under pressure of wartime needs for aviation gasoline brought the total catalytic cracking capacity of the American petroleum industry to about 1,000,000 daily barrels at the end of the war. Since 1945, this capacity has been doubled, so that the present processing capacity of 2,000,000 barrels per day represents about half of the total cracking (thermal plus catalytic) capacity of the American petroleum industry, and it is still expanding rapidly both in the United States and abroad. Catalytic cracking has thus become by far the largest catalytic process in all of American industry, producing some 30,000,000 tons per year of catalytic gasoline as compared with 10,000,000tons of sulfuric acid, the second largest product of commercial catalysis. Continued expansion of catalytic cracking a t the expense of thermal cracking is to be expected, because recent developments in both the fluid and TCC processes have made this the most economically attractive route to the production of the high antiknock fuels required by modern automotive engines. Furthermore, catalytic cracking has the advantage over thermal cracking in that it does not produce the very heavy residual fuels typical of thermal cracking, but, rather, more distillate fuels of considerably higher utility and value (domestic heating fuels and diesel fuels).

4

R. C. HANSFORD

I n spite of the fact that catalytic cracking of petroleum hydrocarbons rapidly attained prominent industrial importance after the introduction of the Houdry process in 1936, it was not until nearly ten years later that any significant progress toward a basic understanding of this field of catalysis became evident. Yet, much of the background, upon which present theories of catalytic cracking have been based, was published several years before 1936. The catalytic activity of certain activated clays towaxd hydrocarbons was observed as early as 1912, when Gurwitsch (5) in his studies of the adsorption of olefins on activated clays reported that, polymerization occurred. Herbst (6) in 1926 observed that the decomposition of hydrocarbons is accelerated by kieselguhr a t moderately elevated temperatures, and Kobayashi and Yamamota (7) obtained similar results with Japanese acid clays. Several patents (8) covering the use of floridin, pumice, and hydrosilicates of aluminum as cracking catalysts were issued in the period 1923-1932. I n 1933 Gayer (9) made a very significant observation when he reported that a silica-gel-supported alumina polymerization catalyst possessed acidic properties. About this same time, Whitmore (10) proposed his now famous carbonium ion theory of acid-catalyzed organic reactions, including polymerization and rearrangement reactions of hydrocarbons. Finally, in 1940 Frost (1 1) published an excellent review of the then known reactions of hydrocarbons over active aluminosilicates, calling attention t o the strong similarity of these reactions t o those catalyzed by acid catalysts like aluminum chloride, sulfuric acid, and phosphoric acid. It was not until some six years later, however, that anything like a unified picture of the chemistry of cracking catalysis began t o emerge from the many known facts concerning the catalyst and the reactions involved in catalytic cracking of hydrocarbons. Before discussing the recent advances in our understanding of the chemistry of catalytic cracking, it is felt that a brief review of the general nature of cracking catalysts, including methods of preparation, testing, and general application, would be of value in orienting those readers who may not be familiar with this important field of catalysis.

11. COMMERCIAL CRACKING CATALYSTS, Reference has already been made to the fact that act,ivated clays and synthetic compositions containing the hydrous oxides of silicon and aluminum were recognized, long before the development of a practical commercial process, to have catalytic activity toward hydrocarbons. Basically, these are the types of catalysts in commercial use today, although many modifications and improvements over the earlier catalysts

CHEMICAL CONCEPTS O F CATALYTIC CRACKING

5

have been made since the Houdry process was first developed. In recent years considerable attention has been given to silica-magnesia compositions, but this type of catalyst has not been widely used industrially as yet. Mention should also be made of silica-zirconia, silica-aluminazirconia, and alumina-boria compositions, which have been studied extensively on a laboratory or pilot plant scale but which have never been commercialized. The first industrial siliceous cracking catalyst was an acid-leached bentonite clay to which about 1% of manganese dioxide was added for the purpose of increasing the rate of burning of carbonaceous deposits during regeneration. Essentially the same type of catalyst, without the added manganese dioxide, is in wide use today. Preferred bentonite clays are those whose chief constituent is montmorillonite, a mineral of the composition corresponding to the empirical formula, 4SiO2.Al2O3.H20. The principal sources of raw clay for the manufacture of the presently most widely used natural catalyst (Filtrol Corporation) are deposits in Arizona and Mississippi. The clay from these deposits contains appreciable amounts of impurities, principally CaO, MgO, and Fe203, which replace part of the A1203in the ideal montmorillonite structure. The catalyst is prepared by leaching the raw clay with dilute sulfuric acid until about half of the alumina and associated impurities is removed. The resulting product is then washed, partially dried, and extruded into pellets, after which it is “activated” by calcination. A typical analysis of the finished catalyst is as follows (Mills, 12). Per Cent by Weight Dry (105°C.) basis: Ignition loss at 870°C. (mostly H 2 0 )

so4

Free HpSOa Ignited (870°C.) basis

SiOz A1203 MgO CaO Fez03 Alkali as NaZO CUO

8.5 4.3 0.83 73.9 18.0 4.9 3.1 2.1 0.3 0.005

While natural or activated clay catalysts are no longer employed in the fixed-bed Houdry process, they are still widely used in the fluid process and to a considerable extent in the TCC process. A natural bauxite catalyst is employed in the fixed-bed cycloversion process, developed by the Phillips Petroleum Company. This process is of greater importance as a naphtha reforming process than as a catalytic cracking process.

6

R. C. HANSFORD

A synthetic silica-alumina catalyst was introduceld commercially around 1940 t o replace the less satisfactory natural catalyst in the Houdry process. Although considerably more expensive than activated clay, th e synthetic catalyst is more stable toward deactivation by the effects of high temperatures and of sulfur (H2S). This increased stability permits much longer operation with a single charge of catalyst in fixed-bed units, which along with improved product quality makes the more expensive catalyst commercially attractive. The use of synthetic catalyst in moving-bed processes (TCC and fluid) is subject to the same considerations of thermal and sulfur stability. Synthetic cata1:yst is definitely preferred in operations which process fractions of high sulfur content, and better product quality may in many cases offset the higher initial cost of synthetic catalyst in the processing of low-sulfur stocks (Evans, 13). An important consideration in moving-bed processes is the mechanical stability or resistance to attrition of the catalyst particles. The advantages of greater catalytic stability of synthetic catalysts can be realized only if losses due to attrition are kept at a relatively low level. In practice, it is desirable that the necessary fresh catalyst make-up due to attrition will just balance the inherent loss in overall catalyst activity resulting from t he effects of operating conditions. This has been more nearly achieved with some synthetic catalysts than with clay catalysts. Generally the rate of attrition of clay catalysts is substantially greater than that of the best synthetic catalysts offsetting still further the relative advantage in cost of natural catalyst. In certain operations, particularly with petroleum fractions containing more than atbout 0.5% sulfur (as organic sulfur compounds), the greater inherent instability of clay catalysts requires excessive replacement of whole catalyst particles with fresh catalyst in order to maintain the desired activity level and product distribution. The preparation of synthetic silica-alumina catalysts is a relatively simple one, involving the coprecipitation or cogelation of the two hydrous oxides from mixed solutions of sodium silicate and aluminum sulfate. Depending on how the solutions are mixed and on the p H and concentration of the resulting mixture, the combined hydrous oxides will be formed as a coprecipitate, which separates from a greater part, of the aqueous phase, or as a true hydrogel, which embraces the entire solution volume. The formation of a hydrosol which sets t o a true hydrogel in a short time is the basis for an important recent advance in th e a r t of catalyst manufacture on a large scale. The gelation time of a silica-alumina hydrosol is a function of several variables (pH, concentration, degree of mixing of reagents, and temperature), the control of .which permits a uniform dispersion of the hydrosol into a n immiscible fluid, such as oil,

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

7

air, or steam, before gelation occurs. While still dispersed in the forming fluid the hydrosol droplets set t o fairly rigid and roughly spherical hydrogel particles, which may then be processed to a finished catalyst having the desired final particle size without going through the tedious operations of filtering and pelleting, molding, or grinding. A more detailed description of the method is given in the Marisic patents (14) and by Porter (15). It should be pointed out that the coprecipitation or cogelation of silica-alumina from sodium silicate and aluminum salts (sulfate, chloride, or sodium aluminate) results in the formation of a product having strong zeolitic properties. It is necessary t o remove the sodium ion by exchanging with another ion such as H+, NH4+, or A10+, and this is usually done by treatment of the precipitate or hydrogel with a dilute solution of ammonium chloride (or sulfate) or of aluminum sulfate. After the sodium is exchanged out, the material is washed free of electrolytes, dried, and calcined (700°C.). There are many other methods of preparing active synthetic silicaalumina catalysts. A fair catalyst can be made by impregnating dried silica gel with an aluminum compound which is easily converted t o the oxide by calcination, e.g, Al(N03)3. A preferred impregnation technique is t o soak a sodium-free silica hydrogel in a solution of an aluminum salt and t o follow this with an aqueous ammonia treatment to precipitate the hydrous alumina on the silica (Thomas, 16; Ryland and Tamele, 17). It should be noted that silica hydrogel can easily be freed of sodium ions by water washing, since it is not a zeolite. Exceptionally pure silicaalumina composites can also be prepared by the hydrolysis of mixtures of ethyl orthosilicate and aluminum alkoxides (Thomas, 18). Although only very small amounts of alumina (0.1%) are necessary to produce an active silica-alumina catalyst, more stable catalysts are obtained with larger contents of alumina. Commercial catalysts generally contain around 10% by weight of A1203 and 90% SiOz (dry basis). At the temperatures employed in catalytic cracking (ca. SOO”C.), synthetic catalysts contain about 1% by weight of “bound water” also. About half of this bound water is to be considered as being largely hydroxyl groups in the surface of the catalyst, and is therefore only potential ‘(water” of constitution, which is not removable except by splitting out a t high temperatures, with resulting collapse of the high surface structure. The particle size and form of cracking catalysts are varied, depending on the type of process in which they are used. Pelleted or extruded catalysts, employed in fixed- and moving-granular bed systems, are normally prepared in a cylindrical shape, with dimensions of about 4 mm.

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R. C. HANSFORD

in diameter and length. Similarly, spheroidal or bead catalyst particles are 2.5 to 4 mm. in diameter. Fluid catalysts (powdered and microspheres) range from about 100 to 300 mesh (0.15-0.04 imm.) in particle size.

111. CATALYST EVALUATION The laboratory evaluations of cracking catalysts may be divided into essentially two categories: (1) testing of activity for the conversion of a standard gas-oil to gasoline, gas, and coke, and (2) measurement of physical properties such as particle size, density, surface area, and poresize distribution.

1. Activity Tests A number of physical tests have been proposed for the estimation of catalytic activity, but except for purposes of control in catalyst manufacture none of these is an acceptable substitute for direct cracking tests. In general, such test,s as the measurement of the selective adsorption of an aromatic hydrocarbon from a standard binary mixture of aromatic and paraffinic components (Scheumann and Rescorla, 19), or the measurement of the heat of wetting with methanol (Mills, 20) merely reflect the extent of available surface. With a given catalyst coinposition, these methods may have some utilii,y in following the decline of activity due to the effects of temperature or steam (but not of sulfur) or as a rapid and approximate control method in the manufacture of a catalyst. The most satisfactory and meaningful catalyst activity tests are those involving the cracking of a standard cracking stock in small-scale staticbed reactors. Small laboratory-scale fluidized-bed test units have been described, but generally speaking these are more difficult to operate satisfactorily than fixed-bed units. Powdered catalysts are generally tested in pelleted form or in a nonfluidized static bed of powder. Alexander and Shimp (21) have described the widely used CAT “ A ” (catalyst activity test “A”) method for evaluating granular cracking catalysts. Figure 1 is a flow diagram of the apparatus employed in this test method. The reactor (cracking case) is made of Pyrex glass and is divided into two sections, The top section is filled with quartz chips and acts as the vaporizer and preheater. The bottoin section is the reactor proper and holds 200 to 220 ml. of catalyst. Thermocouple wells are located in each section, so that the temperature within the preheater and catalyst beds can be measured. Temperature control on the furnace is such that the temperature differential from the top t o the bottom of the catalyst bed is not greater than 10°F. under operating conditions. Figure 2 gives details of the construction of the reactor and furnace.

9

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

The cracking stock (standard light East Texas gas-oil) is charged as a liquid t o the vaporizer-preheater from the thermostated charging buret through a capillary tube. The charging rate is controlled by the pressure of nitrogen above the oil in the buret. At the end of the run (10 minutes), the catalyst is flushed with 900 ml. of nitrogen from the gas graduate. This nitrogen is collected, along with the gas produced in the cracking reaction, in the gas bottle. The gas yield is corrected for the measured nitrogen content so introduced. RACKING CASE ‘URNACE COMBUSTION

I / TUBE

8

1

N2

ASCARITE BULB

AIR

HOOD

4

REGULATORS WATER

I DRAIN

NTHETIC UDE TRAP

FIG.1. Flow diagram of CAT “A” apparatus.

The standard operating conditions of the test are as follows : Preheater temperature: 830-850°F. (443-454°C.) Catalyst furnace temperature: 800 f 5°F. (426 k 3°C.) Charging stock: light East Texas gas-oil (approximate A.S.T.M. boiling range, 220-380°C.; sp. gr. a t 15.6”C., 0.840) Charging rate: 5.0 f 0.3 ml. (measured a t 60°F. or 15.6”C.) per minute per 200 ml. catalyst Pressure: atmospheric Length of cracking period: 10 minutes Temperature of condenser and trap: 60°F. (15.6”C.)

Carbon on the catalyst is determined by burning in air, after the nitrogen flush and after replacement of the synthetic-crude trap (in which liquid cracked products are collected) with an empty one. The

10

R . C. HANSFORD

temperature of the preheater and catalyst is raised t o 950-975°F. (510524°C.) before air is introduced through the gas flowmeter. The regeneration gases contain appreciable amounts of carbon monoxide, in addition t o carbon dioxide. The combustion of carbon monoxide t o carbon dioxide is carried out over copper oxide a t 900-1000°F. (482538”C.),and the total carbon dioxide is absorbed in ascarite, weighed, and

IRON P I P E

25 MM.I.0. PREHEATER (CRACKED QUARTZ)

SS.SCREEN PLUG

5 MM. 1.0. CATALYST BED

GH TEMPERATURE

I N SU L A 1I0 N

SS SCREEN PLUG

II FIQ. 2. Details of reactor and furnace CAT “ A ” apparatus.

calculated t o equivalent carbon. Regeneration is carried out over a 2- t o 4-hour period, depending on the amount of carbon laid down. The liquid cracked products collected in the trap are distilled directly from the collecting trap through a low hold-up, 5- t o 10-plate column (see Fig. 3) a t a rate of about ml. per minute. The final cut point is 410’F. (210°C.) vapor temperature, giving the volume yield of 410°F. end-point gasoline.

CHEMICAL CONCEPTS O F CATALYTIC CRACKING

&!I I

I

INSULATION

11

REFLUX CONDENSER

v/1 v/1

( 3 0 ML?

SYNTHETIC CRUDE

U”

I-I-

ELECTRIC HEATER

FIG.3. CAT “ A ” gasoline distillation apparatus.

The data reported from the test are: Volume % 410°F. (210°C.) end-point gasoline Volume % 300°F. (149°C.) end-point aviation cut W-t. % gas Gas gravity (sp. gr. referred to air) Wt. % coke (as carbon)

All percentages are based on the gas-oil charged during the 10-minute cracking period. The following table shows some typical data for several catalysts (Mills, 12) : Catalyst Fresh Si02-A1203(synthetic) Normally aged” syn. catalyst “Sulfur-poisoned” syn. catalyst Fresh clay catalyst Normally aged“ clay catalyst Sulfur-poisoned clay catalyst

Gasoline, 410°F. Coke VOl. % Wt. % 45.0 3.2 32.0 1.5 40.7 3:l 39.3 4.3 29.3 1.9 16.2 10.1

Gas Gas Wt. % Gravity 9.6 1.60 5.0 1.50 8.3 1.50 7.4 1.48 3.5 1.44 6.9 0.37

a Aging equivalent t o t h a t encountered in processing of stocks containing less t h a n about 0.5 ?T h l f u r and no appreciable metal impurities.

12

R. C. HANSFORD

Other test procedures for the direct determination of catalyst activity have been proposed by Birkhimer, Macuga, and Leum (22), Shankland and Schmitkons (23), and Conn and Connolly (24). These are essentially similar to the CAT “ A ” method, except for the details of operating conditions and reporting of test data. In addition to the carbon and gas yields, the per cent recycle stock (“unconverted” charge stock boiling above 400”F.), the conversion (100% - % recycle stock), and a quantity called “D L” (per cent distillate to 400°F. plus distillation loss) are factors reported from the test data. The overall activity of the

+

(1

LIMI‘f SAFETY/’ TRIPS

THERMOREGULATOR

ICE WATERJACKETED POSITIVE DlSPLACEMEll -WATT HEATER Y

WTOMATIC DISTILLATION COLUMN

FROLN2 CYL!NDER

FIG.4. Apparatus for testing granular or powdered cracking catalysts in staticbed reactor (Shankland and Schmitkons, 23).

catalyst is expressed by the extent of conversion, while tJhe selectivity of the catalyst is reflected in the values of per cent by volume of D L and particularly in the spread between D L and conversion values. The method of Shankland and Schmitkons (23) (see Fig. 4 for details of apparatus) involves an additional factor of “relative activity,” that is, the weight of a standard reference catalyst which gives the same degree of cracking as a fixed amount of the test catalyst. For example, 2 lb. of catalyst of activity = 50 will produce the same extent of cracking as one pound of catalyst of activity = 100 at constant processing conditions. The determination of relative activity has a number of advantages, the most important being the cancelling out of variations in charge stock characteristics and in operating conditions.

+

+

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

13

A small-scale fluidized-catalyst test unit has been described by McReynolds (25). A standard gas-oil is vaporized and passed through a fluidized catalyst bed (400 g. of powdered catalyst) at 920°F. and at a rate of 925 ml. (800 g.) of liquid per hour for thirty minutes. The product is condensed, stabilized, and fractionated to 400°F. end-point gasoline. The activity of the catalyst is reported in terms of total conversion and per cent “D L.”

+

2. Physical Tests

A number of physical measurements on cracking catalysts are employed to complement the direct activity tests. Ries (26) has reviewed the principal methods employed in determining the physical properties of porous solids, and no attempt will be made to do so again in detail here. The determinations of surface area, bulk density, particle density, and real density are practically routine procedures in studying all types of cracking catalysts. From these data, the pore volume of catalyst particles, intergranular free volume, and average pore diameter can be calculated (Emmett and DeWitt, 27). The size of catalyst particles is of great importance, particularly in processes employing moving granular and fluidized catalyst beds. The flow characteristics of a moving granular catalyst bed and the fluidization quality of powdered or microspherical catalysts are affected by the particle size. Conventional screen analyses are used for particle size determinations, but with fluid (microgranular) catalysts special precautions are necessary for satisfactory analyses. Webb (28) states that humidification of microgranular catalysts is necessary to overcome the effects of electrostatic charges encountered with dry catalysts. Storage of the catalyst over 30% sulfuric acid for several hours (overnight) suffices to bring the catalyst to a satisfactory water content. Sedimentation methods are also useful in determining particle size distribution. Webb (28) describes the techniques employed in liquid sedimentation and gives data comparing the results with those obtained in mechanical screening. Sedimentation methods are particularly useful in measuring sub-sieve sizes. Air-elutriation methods (Roller, 29) are also useful, especially when used in connection with microscopic examination (Wiley, Deloney, and Denton, 30; Matheson, 31). The distribution of pore sizes in cracking catalyst particles is of considerable interest because of its bearing on diffusion limitations in the cracking and regeneration reactions (Wheeler, 32). A very considerable part of the internal pore volume of cracking catalysts is associated with pores having an average diameter of about 50 A. Pelleted or extruded catalysts have a relatively large proportion of their pore volumes in the

14

R. C. HANSFORD

macropore range (diameters above 200 A.), but exhibit distribution peaks a t around 50 A. and at considerably higher values of pore diameter (Ritter and Drake, 33; Drake, 34). S. Chemical Tests

Aside from chemical analyses for the major and minor constituents of cracking catalyst8s,there are few chemical tests which are normally made on these catalysts. A titration test, similar to the one described by Thomas (18), involving a correlation between the amount of potassium hydroxide reacting with a given quantity of catalyst and the cracking activity of the catalyst, has been used for control purposes (19). Unfortunately, the method requires calibration for each type of catalyst tested. For example, a different calibration curve (milliequivalents KOH per gram of catalyst us. gasoline yield) is obtained for clay catalysts and for synthetic silica-alumina catalysts. In reality, this method measures a quantity which is a function of total surface area (microparticle size us. rate of solution in aqueous KOH) rather than a chemical property. Tamele (35) has referred t o an unpublished titration method, which possibly would give a common correlation curve for different types of cracking catalysts, since it measures the ‘ I acidity ” of the catalyst surface. The method consists of measuring the surface acidity of the catalyst by titration with n-butylamine, using p-dimethylaminlobenzene as the indicator. The titration is carried out in a nonaqueous medium (powdered catalyst suspended in benzene) to the disappearance of the red color of the adsorbed indicator. A vapor phase “titration” of cracking catalysts with a volatile organic base (quinoline) has been described by Mills, Boedeker, and Oblad (36). The amount of quinoline irreversibly adsorbed (chemisorbed) at 315°C. was found to be a straight-line functioin of the CAT “ A ” activity of the catalyst. Such widely different catakysts as clay, synthetic silica-alumina, silica-magnesia, and silica-zirconia catalysts fall fairly closely on a common correlation curve (milliequivalents of chemisorbed quinoline per gram of catalyst us. CAT “ A ” gasoline yield). A possibly better correlation might be obtained by plotting milliequivalents of chemisorbed quinoline per milliliter of catalyst against CAT “ A ” gasoline yield, since the latter test is based on a constant volume of catalyst and not constant weight.

IV. CHEMISTRY OF CATALYTIC CRACKING I. Chemical Nature of the Catalyst Reference has already been made to the fact that cracking catalysts have definite and measurable acidic properties. Present concepts of the

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

15

mechanism of catalytic cracking are largely based on the observed acidic properties of the catalyst and on the nature of the reactions which occur in the presence of cracking catalysts. There are, however, some schools of thought (Taylor, Turkevich, Grosse, and others) which do not completely subscribe to the I‘acid theory ” of cracking catalysis, and there are widely divergent views on the nature and role of the acidic centers of cracking catalysts among those who accept the general cationic mechanism of catalytic cracking. The measurement of the ‘‘acidity” of cracking catalysts or of substances which possess considerable activity in cracking of hydrocarbons (various adsorbents derived from aluminosilicate or magnesium silicate clays, alumina-impregnated silica gel, etc.) may be carried out, at least in a qualitative way, with indicators. Recent investigations (WeilMalherbe and Weiss, 37; Walling, 38; Tamele, 35) have shown conclusively that synthetic silica-alumina cracking catalysts and activated clays have surfaces characterized by centers of high acid strength. It is important to note that the measurements of acidity reported by these investigators were carried out in nonaqueous media, eliminating any question as to the validity of the interpretation of acidity as an inherent surface property of the solid. Milliken, Mills, and Oblad (39) have criticized the interpretations of acidity derived from titrations with basic solutions (Thomas, 18; Grennall, 40). Their criticism is that “the quality being measured during titration is not ‘acidity.’ The property being measured is a capacity to react with a base that can be created a t the conditions of the experiment. Such a measure of ‘acidity’ may have little or no relation to the amount of ‘acids’ involved in catalytic reactions.” Such a criticism is justified only in the cases where the solvent employed, e.g., water, may lead to the creation of acids by solvolysis of an acid anhydride or other structure capable of being converted to an acid by the solvent. Certainly the work of Walling, which employs a completely nonpolar solvent (isooctane), cannot be criticized from this viewpoint, nor is there any reason t o suppose that the use of a solvent like benzene (Tamele, Weil-Malherbe) would create a capacity of the substance t o react with a base. If anything, the slightly basic properties of benzene might diminish the capacity of the solid surface to react with another base, although this would be insignificant with bases as strong as n-butylamine. The question as to whether the acid properties of cracking catalysts are due t o protons (Bronsted acids) or to electron-deficient atoms (Lewis acids) is somewhat more difficult t o answer. Before considering this question, the origin of acid centers in silica-alumina compositions should be discussed. It is generally believed that acid centers, either the

16

R. C. HANSFORD

Bronsted or the Lewis type, owe their existence in silica-alumina catalysts t o an isomorphous substitution of trivalent aluminurn for tetravalent silicon in the silica lattice (Hansford, 42; Thomas, 18; Tamele, 35). Such an isomorphous substitution would lead to a structure somewhat as follows :

Because the normally six-coordinated aluminum atom has been forced to assume a four-coordinated structure, there is a net unit negative charge created at this point in the catalyst surface, requiring neutralization by a cation such as a proton (Thomas, 18). A similar view (Tamele, 35) is that the aluminum atom in such a structure tends to acquire a pair of electrons to fill its p-orbital, creating a Lewis acid in the absence of water and a Bronsted acid in the presence of one molecule of water: -Si

I I

.. ..

.' I

: 0 : + - A h : 0 : Si-

1

.*

(Lewis acid)

I

:o: -Si-

I

-8i

I

..

I

H: 0 : H+

: 0 : e--Al-+

1

*.

:o: -Si-

I

.. :0 :

I

h-

.. I

(BrBnsted acid)

(Arrows indicate displacement of electrons toward the Si-0 group)

Although Milliken and co-workers (39) agree that an undried coprecipitate or hydrogel of silica-alumina probably has the structure of a Bronsted acid, where each aluminum atom in the hydrogel may create an acid or base-exchangeable center, they take a quite different view of the chemistry of the calcined catalyst. They argue that calcination essentially destroys the Bronsted acid by loss of water through reaction of the acid hydrogens from two alumina tetrahedra with an oxygen ion. Agglomeration of alumina and silica into a mixture of y-alumina and silica particles is believed t o occur, and evidence for this is presented in a study of the real density of heat-treated (7GO'C.) silica-alumina catalysts containing different amount of alumina. The real density increases linearly with increasing alumina content, giving a curve which, when

CHEMICAL CONCEPTS O F CATALYTIC CRACKING

17

extrapolated to 100% alumina from the 0 to 20% range, indicates that the added alumina has a density of 3.65 g. per milliliter. The real density of 100% y-alumina, prepared as a gel, was found to be 3.8 g. per milliliter. According to Milliken et al., the aluminum ions at the interface of y-alumina and silica micelles are believed to be in a three-coordinated structure corresponding to the anhydride of the acid, HA102. An anhydride of this structure is a potential acid of the Lewis type (see above example of Lewis acid structure), and this is believed to be the only type of acid which is present in the silica-alumina catalyst at cracking temperatures (500°C.). In addition to these conclusions, a new but entirely hypothetical concept of the nature and origin of the acid centers in cracking catalysts has been introduced by Milliken, Mills, and Oblad. They postulate that substantially all the alumina of the catalyst is in a six-coordinate structure at cracking temperatures, i.e., only potential Lewis acid structures are present. The aluminum ions closest to tetrahedral silicon are in a state of strain and can be induced t o assume tetrahedral coordination by the approach of a molecule having even weakly basic properties, e.g., even a paraffin hydrocarbon molecule. In other words, the acid centers of the catalyst are actually created only a t the time of approach by a base. The extent to which the catalyst surface becomes “acidic” depends on the amount of interfacial sharing of oxygen ions between silicon and aluminum, i.e., the degree of dispersion of alumina in silica, on the hydroxyl content of the alumina, and on the polarizing capacity of the base which is approaching the potential acid center. Weakly basic molecules (weak Lewis bases, e.g., paraffins), though having only slight ability t o polarize other molecules, will, according to Milliken et al., induce coordination shifts in the aluminum ions closest to tetrahedral silicon ions. Stronger bases, e.g., quinoline, may induce coordination shifts in aluminum ions further removed from the influence of silica. Thus, the acidity of the catalyst becomes a function of the basicity of the material used to measure it. Although this concept of the acid nature of cracking catalysts is novel and has some interesting aspects, it is completely unacceptable to this writer. First of all, the proponents of the concept have ignored a very considerable body of evidence that protons can and do exist on cracking catalysts at temperatures above which it is contended that a Bronsted acid cannot exist. Secondly, it is extremely difficult to believe that the polarizing ability of any but the most polar molecules, e.g., HzO, could exert an effect on the structure of a solid such as is postulated. The mobility of oxygen ions in such solids as measured by their exchange with a large excess of HzO1*(Mills and Hindin, 41) is no proof that such

18

R. C. HANSFORD

mobility exists under cracking conditions, where the amount of free water is usually small. Finally, it is not necessary t o assume that the possibility of coordination shifts provides the driving force for the desorption of product molecules, nor is it the only lpgical explanation why different bases give different values for the number of acid centers in the surface of a cracking catalyst. The still more radical suggestion that the principle of induced coordination shift can be applied generally to heterogeneous catalysis will be challenged by many iintil further proof of the universality of the concept is unequivocally esta,blished. With respect to the presence of protons even on calcined silica-alumina catalysts, the work reported by Tamele (35), by Hansford (42), by Hansford, Waldo, Drake, and Honig (43), and by the proponents of nonBronsted acids (Hindin, Mills, and Oblad, 44) strongly suggests that protons are essential to the activity of these catalysts. In the earlier work (Hansford, 42) it was shown that dehydration of a silica-alumina catalyst (presumably approaching a Lewis type acid) decreases the activity toward the cracking of n-butane. This has been repeatedly confirmed in unpublished work by the writer’s associates. It was further observed, by adding “activating ” amounts of water as deuterium oxide (ca. 0.25% by weight of the catalyst), that an exchange reaction occurs between the deuterated catalyst and the hydrogens of several hydrocarbons, including paraffins. This indicates that water in some way enters into the mechanism of activation of saturated hydrocarbons, at least. A detailed study of the exchange reaction with isobutane (Hansford, Waldo, Drake, and Honig, 43) has shown that the rate of exchange is dependent, among other things, on the amount of deuterium oxide adsorbed on the catalyst. There appears to be an optimum water (DzO) content, below and above which the rate of exchange decreases. The base water (DzO) content of the catalyst used in this investigation was 0.54 % (constitutional water not removable by evacuation at 760”C.), and the optimum amount of additional DzO required to activate the conditioned catalyst for exchange is about 0.25 % by -weight of catalyst. Hindin, Mills, and Oblad (44) observed a similar dependence of rate of exchange on the amount of adsorbed DzO. They found that drying the deuterated catalyst under vacuum at 525°C. destroyed practically all activity of the catalyst toward exchange with isobutane at 150°C. At this condition the catalyst contained about 0.7 % DzO (as total available water on high temperature ignition), but the adsorption of an additional 0.2% D 2 0 produced a remarkable increase in exchange activity. The results of these two investigations show quite conclusively that water promotes silica-alumina in its ability to activate a saturated

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

19

hydrocarbon molecule, and that i t is probably not the potential water built into the lattice as hydroxyl groups which is involved. This is t o say that the Lewis type acid centers require activation by water, creating a Bronsted type acid. It is believed that this activation is essentially equivalent to the activation of other Lewis type acid catalysts, such as BFI and AlCl,, by water or by hydrogen halides. Thus, the picture presented by Tamele (35) seems t o be quite adequate in explaining the origin of proton acidity through the coordination by water with the electron-deficient aluminum atom. The suggestion made by Hindin and co-workers that a second type of adsorbed water is involved, causing a “labilization” of hydrogen in the hydroxyl groups, is essentially an admission of the existence of protons in the catalyst surface, although it is not clear t o this writer how such labilization of hydroxyl hydrogen is brought about by adsorbed water. If protons exist in the catalyst at 150°C., then the contention by the co-workers of Hindin (39), that the existence of a Bronsted acid even at temperatures as high as room temperature is very doubtful, seems somewhat contradictory or a t least confusing to this writer. Although the origin of acidity in silica-alumina catalysts may be ascribed t o the existence of tetracoordinated aluminum atoms as pictured originally by the writer (42) or according t o Tamele (35), it may be that some other mechanism is also operating. For example, silica-zirconia gels are active cracking catalysts and possess acidity of about the same degree as do silica-alumina gels. Zirconium, being tetravalent like silicon, would not give rise t o a net negative unit charge if substituted for silicon in the silica lattice, unless possibly a difference in size of the ion (as compared to that of silicon) would create the requisite distortion in electron distribution. On the other hand, Thomas (18) considers zirconium t o be eight-coordinated in an active structure : 4

4H+

I

Such a structure would require four positive ions (protons) for electrostatic neutrality. However, while zirconium is one of the few elements which have been found t o form eight-coordinated compounds (Penney

20

R. C. HANSFORD

and Anderson, 45), it does so only with relatively small molecules, e.g., water and acetylacetone. I t would seem th at the structure suggested by Thomas would be too unstable to permit the preparation of a n active silica-zirconia catalyst suitable for high-temperature reactions. A sixcoordinated structure, having a net negative charge of two, appears more likely. Plank (46) has suggested an alternative explanation for the acidity of silica-alumina catalysts. From a study of the differences between silica and silica-alumina gels, he concluded th at alumina always becomes a terminal group in the micelle structure, and th a t therefore isomorphous substitution of silicon with aluminum does not occur. According to his hypothesis, t h e aluminum ions in th e terminal alurnina groups are coordinated with hydroxyl and water in such a way a.s to retain their normal octahedral coordination :

I

A --O-Si--O--A1-OH

I

d I

HOH

'

\ / OH

.oi'oz

AppIication of Pauling's electrostatic valence rule, distributing the 3f valence of aluminum among the six oxygen-aluminum bonds, would result in lability of the hydrogens of the coordinated water molecules, creating a Bronsted acid. A similar structure involving four-coordinated magnesium in silica-magnesia or six-coordinated zirconium in silicacatalysts would also represent a relatively strong Bronsted acid. Alumina can be practically completely extracted from a silica-alumina gel with dilute mineral acids, provided the gel is not calcined a t too high a temperature (ca. 300°C.) (Plank, 46). The physical structure of th e leached gel is not significantly different from that of th e unleached gel containing 10% Al2Oa. However, if the gel is calcined a t 700"C., leaching with 1 N HC1 for 420 hours will remove less than 10% of the Al2O3. This suggests, in agreement with Milliken et al., that high-temperature calcination leads to agglomeration of alumina into essentially a n acidinsoluble form of 4-A1203. It is not possible to decide a t present whether the acidity of calcined silica-alumina is due t o coordination of water with four-coordinated or with six-coordinated aluminum ions. I n summing u-p all the experimental evidence, the conclusion appears t o be that the active centers of

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

21

silica-alumina and related catalysts are located at the interfaces of the particles or micelles of the respective hydrous oxides, and that these active centers are regions of proton acidity created by coordination of electron-deficient points of the surface with water molecules.

2. Mechanism of Catalytic Cracking It is now generally agreed, with some exceptions, that the cracking of hydrocarbons over silica-alumina and related catalysts involves a cationic mechanism in which the acidity of the catalyst plays a decisive role. There is some disagreement among workers in this field concerning the actual part played by the acid centers of the catalyst in the initiating step of the mechanism. The writer (42) published the first attempt t o explain the mechanism of catalytic cracking in terms of an ionic mechanism, later modifying the initial concept, which involved both cationic and anionic complexes, to include only a cationic mechanism (47). Thomas (18) arrived a t essentially the same conclusions regarding a cationic (carbonium ion) mechanism somewhat earlier, although his very comprehensive paper on the subject was not published until 1949. Similarly, Greensfelder and his co-workers (48) came t o substantially the same general conclusions, The papers by Thomas and by Greensfelder, Voge, and Good have given an excellent treatment of the general mechanism involved in the cracking of hydrocarbons over silica-alumina catalysts. Essentially, the mechanism postulated by the writer and by these authors is based op the formation, reaction, and stabilization of cationic or carbonium ion intermediates. The formation of a carbonium ion from a paraffin or cycloparaffin mayebe brought about by the extraction of a negative hydrogen or hydride ion :

iLl

H H

H H H H

....

..

H : C : C : C : C : H -+

........

H .... .. H : C : C : 6 :?: H + : H........

H H H H

H H H H

On the other hand, hydrocarbons containing unsaturated bonds may form a carbonium ion directly by coordination of ?r-electrons with a proton: H H H

...... H : C : C : C:%: .... H H

H

H H H H H+

H-+

........ . . . . + ..

H : C’: C : C : C : H H H

H

Although many people write carbonium ion structures like the above in terms of free ions, it must be kept in mind that the catalyst (anion) is

22

R. C. HANSFORD

involved also. The structures as written merely represent a highly polarized part of a catalyst complex, and must be so treated. If the anionic portion of the complex is sufficiently electronegative, one might consider the cationic portion as being deficient in essentially two electrons. However, just as there are degrees of electronegativity in catalyst anions, so are there corresponding degrees of electron deficiency in the associated cation, i.e., the greater the electronegativity of the anion the more the cation will behave as a “free” carbonium ion. Since cracking catalysts have been shown t o be very strong acids, the hydrocarbon part of the catalyst complex can be treated subetantially as being deficient in an electron pair, but perhaps not as electron deficient as in t h e case of complexes involving aluminum chloride or boron fluoride. Neglecting for the moment the question of the mechanism of formation of the activated hydrocarbon cation, once it is formed, the associated electron deficiency constitutes the driving force for internal rearrangements and decompositions. A primary carbonium ion will tend t o rearrange t o a more stable secondary or tertiary structure. Evans and Polanyi (49) have given quantitative meaning t o this rule through their derivation of proton affinities from bond energies and ionization potentials of olefins and radicals formed from them by hydrogen atom addition. They define proton affinity, P , as follows: P = a - i3 + D - Z ionization potential of TI (312 kcal.), energy of the second half of the olefin double bond, = carbon-hydrogen bond strength, = ionization potential of the radical formed by addition of hydrogen atom t o olefin. The following table summarizes their conclusions :

where i

=

p D I

=

Proton Afiinities of Otefilzs Olefin CHz=CH2 (1) (2) CHZ=CHCH3 (1) (2) CHz=C( CHB)z (1) (2)

P 57.5

(Values in kilocalories) Dia Dz 11 97.5 97.5 200

PI

PZ

200

152

152

Iz

52.5

95

89

179

180

175.5

168.5

52.0

94

86

165

178

189

168

P I or

=

The subscripts refer t o carbon atoms 1 and 2.

CH2=CHy (1) (2)

(1) or (2) + I€+---+ CHI - &H2

152

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

23

Thus, t-butyl carbonium ion is more stable by 21 kcal. than isobutyl carbonium ion. For this reason, there is a tendency for all cationic hydrocarbon reactions to produce branched-chain structures, if possible. Both Thomas and Greensfelder consider decomposition of a carbonium ion to lead preferentially to a cleavage of the carbon-carbon bond which is in a /3-position with respect to the electron-deficient carbon atom (beta rule). This is to say, the electron-deficient carbon atom captures a pair of electrons from a beta bond, thus breaking that linkage. The act produces a new carbonium ion from the detached fragment, and decomposition may continue until there is finally formed a carbonium ion which cannot produce fragments of three or more carbon atoms. However, such ultimate decomposition may be interrupted by desorption of the reacting hydrocarbon (after ejection of a proton to attain a stable structure), or the ion may capture a hydrogen atom plus two electrons (hydride ion) from another molecule. This is the well-known hydrogen transfer reaction which produces saturated products, particularly from isoolefins. Voge, Good, and Greensfelder (50) showed that tertiary olefins undergo hydrogen transfer preferentially t o other olefinic structures, accounting for the high ratio of isobutane t o n-butane in commercial catalytic cracking. Further details of other reactions, such as dehydrogenation, polymerization, and dealkylation, will be found in the Thomas and Greensfelder papers. Returning now to the question of the initiation step of the cationic hydrocarbon reactions, there are essentially two schools of thought on the subject. One is that the catalyst, itself, has the intrinsic ability t o extract a hydride ion from the hydrocarbon or to distort a carbonhydrogen bond to such an extent that the net result is the same (Milliken et al., 39). The other mechanism is the formatiQnof a cationic complex by the simple addition of a proton to olefinic impurities, either initially present or formed in small amounts by thermal cracking or oxidation. This complex, however small its concentration on the surface of the catalyst, can then start a chain reaction through the hydrogen transfer mechanism, i.e., it can extract a hydride ion from a saturated hydrocarbon, which then becomes the propagating complex (Hansford, 43, 47; Thomas, 18).

24

R. C. HANSFORD

Possible variations of the first hypothesis (39) include the following: RH

+ Hi-+

R+

I

I

+

aI

1-O-ii-

0

I

I

I I

0

RH

+ HP

0

l

+ Ra+-Ha--Al--O-Si--- I I

The second hypothesis seems to be a straightforward conclusion (granted that protons are available) as far as olefin reactions are concerned. However, it is not so obvious with respect to reactions of saturated hydrocarbons. Nevertheless, there is experimental evidence that small amounts of olefins will greatly accelerate certain acid-catalyzed reactions of saturated hydrocarbons. Pines and Wack:her (51) showed that the isomerization of pure n-butane with AlCI, HCN does not proceed in the absence of added olefins (or other carbonium ion former) except under conditions where olefins or their equivalent are probably produced. Recent work reported by the writer and co-workers (43) has shown that the exchange between hydrocarbons and deuterium oxide is sensitive to small amounts of olefins, particularly tertiary olefins. Figure 5 shows a summary of the study of the rate of exchange of isobutane with DzO adsorbed on a silica-alumina catalyst. Although the “pure ” isobutane used in this investigation contained about 0.02% of unsaturated impurities, the rate of exchange was much less than that of the same isobutane to which had been added 0.5% isobutene. The apparent activation energy for exchange in both cases, however, is the same, 17 kcal. On hydrogenating the impurities present in the isobutane over platinum at room temperature, the rate of exchange dropped by a factor of about 3, showing the effect of trace amounts of olefins on the activation of saturated hydrocarbons. Similar work by Hindin, Mills, and Oblad (44) did not show this effect of added olefin. However, these authors used somewhat higher temperatures in their study (150°C.) which fact according to the results of Pines and Wackher with aluminum chloride-catalyzed isomerization may well account for the disagreement. Figure 6 shows a typical distribution curve for the deuterium in the exchanged isobutane, as reported by the writer and co-workers (43).

+

CHEMICAL CONCEPTS O F CATALYTIC CRACKING

25

OC

200

€'

150

100

50

I

I

I

I

B

+

ISOBUTANE 0.5% ISOBUTENE 17 KCAL.

I 3.5 FIG.5. Deuterium exchange rate us. temperature.

0.001

1 56

(c.

H 10)

60

62 64 66 MASS NUMBER

66 (CiO,o)

FIG.6. Distribution of deuterium-exchanged isobutanes.

26

R. C. HANSFORD

It is evident that all nine primary hydrogens are exchanged, but probably not the tertiary hydrogen. This was proved by the failure of 2-deuteroisobutane to exchange with H20. These results are in substantial agreement with those of Hindin et al. On the basis of the above facts and on the additional fact that the exchange reaction was found to be very nearly first order with respect to isobutane, the following scheme for the mechanism of the deuterium reaction was proposed : (1) RR’&-cH, (polarized olefin)

+ H+--A-

(catalyst)

$ [RIt’C-CH,I+-A

(initiating catalyst complex)

+ (CHa),&-H

(2) [RR’C-CH,]+-A-

i=? RR’CHCH3T

+ [(CH,),Cl+-A(pror’agatingcatalyst complex)

(3) [(CH,)3C]+-A-

(4) H+-A-

+

(CHs),6-eH2 +.H+-A(loosely adsorbed isobutene) DOD i=? D+-AHOD

(5) (CH3),&-CHz

$

+

+ IF-A-

s [(CH3),C-CH2D]+-A-

+ H+-A( 7 ) (CH3)&-EHD + zD+-A-, [(CD,H,),C]+-A- + (z - l)H+-A(8) [(CD,HJ3CI+-A- + (CHd3&-l? + (CD,H,)&HT + [(CH3)Cl+-A-

(6) [(CHs)aC-CHnD]+-A-

(CHI),&-fiHD

It should be pointed out that this mechanism accounts for the observed first order kinetics (based on the sum of all deuterated products), since the rate of the hydrogen transfer step (8) depends on the first power of the concentration of isobutane in the gas phase, i.e., tlne rate of adsorption of isobutane. Probably the hydrogen transfer step (8) is the slow step in the overall reaction, since the activation energy of the overall deuterium exchange reaction is not affected by the initiating step involving olefin. Although it is recognized that hydrogen exchange between hydrocarbon and catalyst-water system is not the same thing as catalytic cracking of the hydrocarbon, it seems very probable anld reasonable that both processes involve the same type of activation. The data of Greensfelder and associates on the cracking of pure hydrocarbons plainly indicate that the activation energy of catalytic cracking of saturated hydrocarbons is related to the relative ease of hydrogen transfer between hydrocarbons. Thus, hydrocarbons containing tertiary hydrogens crack at much the fastest rate, corresponding to the greater ease of hydride extraction in producing the activated propagating catalyst-

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

27

hydrocarbon complex. At 550°C., the relative rates of hydride ion removal from a primary, secondary, and tertiary carbon are 1, 2, and 20, respectively (Greensfelder, Voge, and Good, 48). A few words should be said about other mechanisms which have been proposed for catalytic cracking. Taylor and co-workers (52) have suggested that at least one C-H bond must be broken in order to bring the hydrocarbon into the sphere of influence of the catalyst. From a study of the exchange between CD4 and CH4 on silica-alumina at 345”C.,they concluded that the mechanism of catalytic cracking involves dehydrogenation as a primary step, followed by carbon-carbon scission leading to lighter fragments. These fragments evaporate from the catalyst as saturated or unsaturated molecules, depending on the hydrogen concentration a t the surface. Thus, the hypothesis is that dissociative adsorption of hydrocarbons is capable of explaining the results of catalytic cracking. However, this mechanism does not account for the tendency of cracking catalysts to cause chain branching in the products, e.g., isobutane us. n-butane, nor does it explain the experimental fact that addition of hydrogen to hydrocarbon vapors during cracking has the effect of any inert diluent, which is to produce more unsaturates rather than less. The primary postulate of Taylor’s mechanism, that a C-H bond must first be broken, is essentially in agreement with the conclusions of others, if viewed as hydrogen transfer. Turkevich and Smith (53) proposed a so-called unitary theory of hydrocarbon catalysis, based on hydrogen acceptor-donor centers in the catalyst. These centers (oxygen atoms of phosphoric acid, sulfuric acid, etc.) must be spaced close to the critical distance of 3.5 A. in order to produce the “hydrogen switch.” The theory has been successfully applied t o double-bond isomerization and to the exchange of deuterium during hydrogenation of olefins, but its application to cracking, particularly for explaining carbon-chain rearrangements, is not obvious.

3. Present Status In recapitulating the preceding discussion on the chemistry of catalytic cracking, the following conclusions are offered by the writer as to the present status of the mechanism of catalytic cracking, the role played by the catalyst, and the chemical nature of the catalyst. a. The general mechanism involved in catalytic cracking of hydrocarbons is almost certainly an ionic one, and the application of the carbonium ion theory has been extremely successful in accounting for the reactions which distinguish catalytic cracking from thermal cracking. b. The details of the mechanism include some points which may be difficult to resolve, particularly the initiating step of a probable chain

28

R.

C. HANSFORD

mechanism involving the reaction of saturated hydrocarbons. It has not been clearly proved that the general mechanism of catalytic cracking of hydrocarbons involves a chain reaction characteriried by initiation, propagation, and termination steps. However, there is strong evidence that such a chain feature is a necessary part of the mechanism of the reactions of saturated hydrocarbons. It seems more doubtful that a similar chain mechanism exists for reactions of unsaturated hydrocarbons (including aromatics), since it is completely unnecessary t o resort t o such an explanation for the activation and reaction of molecules containing easily polarizable multiple bonds. If a chain mechanism is involved in the reactions of saturated hydrocarbons, the question of the initiating step is not completely answered, although the evidence points strongly to the formation of a polarized catalyst complex from unsaturated hydrocarbons, present as impurities or formed by thermal cracking or oxidation. This complex may initiate the formation of a propagating complex by extracting a hydride ion from the saturated hydrocarbon (hydrogen transfer). It is not clear why the catalyst-hydrocarbon complex is more effective in extracting hydride ions than the catalyst, itself, but the experimental evidence from some of the work on deuterium exchange (43) indicates that this is the case. Further careful work may show that the catalyst is capable of extracting hydride ions from saturated hydrocarbons (H+ H- -+ Hz) and that both initiation mechanisms can operate simultaneously. c. The mechanism of hydrogen transfer between olefins and saturated hydrocarbons is not entirely clear. The indications are that the olefin forms a complex with the catalyst (polarized ester) and that in this state of activation (strong Lewis type acid) it is capable of polarizing C-H bonds (particularly tertiary) and extracting a hydride ion. Whether the donor molecule must be adsorbed or whether collision from the gas phase is a sufficient condition for hydrogen transfer cannot be decided on the basis of available data. d. All experimental data tend to show that silica-slumina catalysts which are most active in hydrogen transfer and crackiing reactions have high protonic acidity. There is no experimental evidence which indicates that protons are absent from such catalysts a t the high temperatures involved in cracking. This does not necessarily mean, however, that only proton acidity is involved in the mechanism of cracking catalysis. It is quite conceivable that electron-deficient centers of silica-alumina not connected with the presence of protons may be capable of activating unsaturated molecules by coordinating with the ?r-electrons of the multiple bonds of these molecules to form carbonium ion complexes. However, the present evidence indicates that probably such centers require

+

CHEMICAL CONCEPTS OF CATALYTIC CRACKING

29

coactivation by water in much the same way that other Lewis type acid catalysts require promoters like water and hydrogen halides. The actual structure of the acid centers of silica-alumina and of related catalysts is still largely unknown and in a speculative state. Much work remains t o be done to clarify this situation, and in particular the effect of temperature on the electron affinity or acid strength of these centers remains to be investigated. Needless to say, these are very difficult problems to attack, but until some reasonable solution to them is obtained there can be little further clarification of existing controversies.

REFERENCES 1. Hansford, R. C., Physical Chemistry of Hydrocarbons, Vol. 11, edited b y A. Farkas. Academic Press Inc., New York, 1952. 2. Houdry, E., Burt, W. F., Pew, A. E., Jr., and Peters, W. A., Jr., Petroleum Refiner 17, No. 11, 574 (1938). 3. Simpson, T. P., Evans, L. P., Hornberg, C. V., and Payne, J. W., Proc. Am. Petroleum Znst. 24 [111],83 (1943). 4. Murphree, E. V., Brown, C. L., Fischer, H. G. M., Gohr, E. J., and Sweeney, W. J., Znd. Eng. Chem. 36, 768 (1943). 5. Gurwitsch, L., Kolloid-2. 11, 17 (1912). 6. Herbst, H., Erdol u. Teer 2, 265, 411 (1926). 7. Kobayashi, K., and Yamamota, K., J . Sac. Chem. Znd. (Japan) 40, 54 (1927). 8. Thiele, F. C., and Cordes, C., German Patent 373,060 (1923); Erdoel u. Kohlenforschung Ges., British Patent 204,458 (1923); Erlenhach, E., U.S. Patent 1,671,573 (1928); Leamon, W. G., U.S. Patent 1,861,399 (1932). 9. Gayer, F. H., Znd. Eng. Chem. 26, 1122 (1933). 10. Whitmore, F. C., J. Am. Chem. Sac. 64, 3274 (1932); Znd. Eng. Chem. 26, 94 (1934). 11. Frost, A. V., J. Phys. Chem. (USSR) 14, 1313 (1940). 12. Mills, G. A., Znd. Ens. Chem. 42, 182 (1950). 13. Evans, L. P., Oil Gas J. 44, No. 47, 167 (1946). 14. Marisic, M. M., (to Socony-Vacuum Oil Co., Inc.), U.S. Patents 2,384,9422,384,946 (1945). 15. Porter, R. W., Chem. & Met. Eng. 63, 94 (April 1946). 16. Thomas, C. L. (to Universal Oil Products Co.), U.S. Patent 2,270,090 (1942). 17. Ryland, L. B., and Tamele, M. W. (to Shell Development Co.), U.S. Patent 2,469,314 (1949). 18. Thomas, C. L., Znd. Eng. Chem. 41, 2564 (1949). 19. Scheumann, W. W., and Rescorla, A. R., Oil Gas J. 46, No. 28, 231 (1947). 20. Mills, I. W., Oil Gas J. 46, No. 28, 237 (1947); Proc. Am. Petroleum Inst. 27 [III], 29 (1947). 21. Alexander, J., and Shimp, H. G., Natl. Petroleum News 36, No. 31, R-537 (1944); Alexander, J., Proc. Am. Petroleum Znst. 27 [III],51 (1947). 22. Birkhimer, E. R., Macuga, S. J., and Leum, L. N., Proc. Am. Petroleum Znst. 27 [111], 90 (1947). 23. Shankland, R. V., and Schmitkons, G. E., Proc. Am. Petroleum Znst. 27 [111], 57 (1947). 24. Conn, M. E., and Connolly, G. C., Znd. Eng. Chem. 39, 1138 (1947).

30

R. C. HANSFORD

25. McReynolds, H., Proc. Am. Petroleum Znst. 27 [III], 78 (1’347); Grote, H. W., Hoekstra, J.,$nd Tobiasson, G. T., Znd. Eng. Chem. 43, 545 (1951). 26. Ries, H. E., Jr., Advances in Catalysis 4, 88 (1952). [This volume] 27. Emmett, P. H., and DeWitt, T. W., J . Am. Chem. Soc. 66, 1253 (1943). 28. Webb, G. M., Petroleum Processing 2, 497 (1947). 29. Roller, P. S., Proc. Am. SOC.Testing Materials 32, 607 (1932). 30. Wiley, J. T., Deloney, J. E., and Denton, S. W., Proc. Am. Petroleum Znst. 27 [III],23 (1947). 31. Matheson, G. L., Proc. Am. Petroleum Znst. 27 [111], 18 (1947). 32. Wheeler, A., Advances in Catalysis, 3, 250 (1951). 33. Ritter, H. L., and Drake, L. C., Znd. Eng. Chem., Anal. Ed. 17, 782 (1945). 34. Drake, L. C., Znd. Eng. Chem. 41, 780 (1949). 35. Tamele, M. W., Faraday Soc. Discussion 8, 270 (1950). 36. Mills, G. A., Boedeker, E. R., and Oblad, A. G., J . Am. Chem. Soc. 72,1554 (1950). 37. Weil-Malherhe, H., and Weiss, J., J . Chem. SOC.2164 (1948). 38. Walling, C., J . Am. Chem. SOC.72, 1164 (1950). 39. Milliken, T. H., Jr., Mills, G. A., and Oblad, A. G., Faraday SOC.Discussion 8,279 (1950). 40. Grennall, A., Znd. Eng. Chem. 41, 1485 (1949). 41. Mills, G. A., and Hindin, S. G., J . Am. Chem. SOC.72, 5549 (1950). 42. Hansford, R. C., Znd. Eng. Chem. 39, 849 (1947). 43. Hansford, R. C., Waldo, P. G., Drake, L. C., and Honig, R. E, Joint Symposium on Use of Isotopes in Petroleum Chemistry, 118th. Meeting American Chemical Society, Chicago, September 1950 (see preprint of syniposiurn papers, p. 81). 44. Hindin, S. G., Mills, G. A., and Oblad, A. G., J . Am. Chem. SOC.73, 278 (1951) (Paper presented a t same symposium indicated in reference 43). 45. Penney, W. G., and Anderson, J. S., Trans. Faraday SOC.33, 1363 (1937). 46. Plank, C. J., J . Colloid Sci. 2, 413 (1947). 47. Hansford, R. C., Gordon Research Conference on Catalysis, New London, N. H., June 24, 1948. 48. Greensfelder, B. S., Voge, H. H., and Good, G. M., Znd. Eng. Chem. 41, 2573 (1949). 49. Evans, A. G., and Polanyi, M., J . Chem. Soe. 252, 1947. 50. Voge, H. H., Good, G. M., and Greensfelder, B. S., Znd. 1Sng. Chem. 38, 1033 (1946). 51. Pines, H., and Wackher, R. C., J . Am. Chem. Soc. 68, 595 (1946). 52. Parravano, G., Hammel, E. F., and Taylor, H. S., J . Am. Chem. SOC.70, 2269 (1948). 53. Turkevich, J., and Smith, R. K., Nature 167, 874 (1947); J. Chem. Phys. 16, 466 (1948).

Decomposition of Hydrogen Peroxide by Catalysts in Homogeneous Aqueous Solution J. H. BAXENDALE Chemistry Department, University of Manchester, England

CONTENTS I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Halides and Halogens.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Paw 31 35

1. Oxidation of Halides by Hydrogen Peroxide., . . . . . . . . . . . . . . . . . . . . . . 2. Reduction of Halogens by Hydrogen Peroxide.. . . . . . . . . . . . . . . . . . . . . 3. The Steady State . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

36 37 37

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

43

2. Active Intermediates. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

. . . . . . . . . . . . . . 62 VI. Copper Compounds.. . . . . . . . . . . . . . . . . . . . . . . . . . .

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

71

VII. Permanganat'e.,. . . . . . . . . . . . . .........

75

I. INTRODUCTION It is a measure of the complexity and variety of the chemical reactions in which hydrogen peroxide participates that they have attracted the attention of chemical kineticists for almost a hundred years, and that many of the earliest reactions to be investigated still provide matter for further investigations. Of these reactions the group involving the catalytic decomposition to water and oxygen by various agents has 31

32

J. H. BAXENDALE

probably been studied most extensively. Such agents are numerous and varied and include some elements and elementary ions-metallic and nonmetallic, more complex ions such as coordination compounds and ions derived from oxyacids, enzymes, and many materials in the solid and colloidal state. We shall be concerned here with the catalytic decomposition in true aqueous solution, and the catalysts which are active in this state are in the main inorganic compounds. From the many investigations on such systems over the past fifty years there have developed two theories t o account for the catalytic decomposition. First there is the intermediate product theory which assumes the presence of unstable compounds, usually peroxidic, which are formed from and are in equilibrium with the catalyst and hydrogen peroxide. Such compounds are presumed to decompose giving oxygen and simultaneously to regenerate the original catalyst, EIO that the reaction might be written in general outline as:

+

A(cata1yst) 2H202 A(HzOz)z+ A 0 2

A(H202)~

+

This idea has been attractive because of the similarity of such a scheme t o an enzyme reaction, the catalyst representing the enzyme and the intermediate compound being analogous t o the enzyme-substrate complex. The viewpoint has been taken by Spitalsky and co-workers (1) who have provided much evidence in support of it. The formation of various peroxide intermediates in many catalytic systems, e.g., chromate and molybdate, is easily demonstrated and t o take a less obvious case, the compound FeH02++ recently shown t o be present, in mixtures of ferric ion and hydrogen peroxide (2), might be considered as an intermediate in the catalytic decomposition which occurs in this system. On the other hand, there has been proposed the mechanism of “compensating reactions.” This is based upon the fact that, with few exceptions, the catalysts are capable of existing in a t least two oxidation states. Combined with the well-known behavior of hydrogen peroxide, as both an oxidizing agent and a reducing agent, this makes possible, a t least in principle, the following cycle of reactions : Reduced catalyst Oxidized catalyst

+ H20z + H202

-+

4

Oxidized catalyst Reduced catalyst

+ O2

So that the overall reaction will be simply 2H202+ 2Hz0

+

0 2

I n such a system there wil1 be dynamic equilibrium between the oxidized and reduced forms of the catalyst, the position of equilibrium being determined by the relative rates of the oxidation and reductilon reactigna.

DECOMPOSITION O F HYDROGEN PE RO X ID E BY CATALYSTS

33

This mechanism has been applied by Abel and Bray to the catalytic decomposition by the halide-halogen systems, in particular, and to other systems in more recent years by the former author. There is little doubt that in many cases the reactions follow in general outline the scheme given above. For example, it is well established that during catalytic decomposition by halides, both halide and halogen are present in the solution and that the oxidation of halide and reduction of halogen occur simultaneously. Some protagonists of these two views have had a tendency to account for all catalyses in terms of the one idea to the exclusion of the other. In actual fact it appears from the available data that with the possible exception of catalysis by molybdate, which appears to involve only the formation and decomposition of permolybdates, there is not one case which can be unequivocally accounted for in terms of one view only. Thus the chromate catalysis, which on the face of it is an example of the intermediate product mechanism, is more complex than the simple theory implies, and it is probable that in certain circumstances the reduction CrV1+ CrIT*and the reverse oxidation also occur, suggesting that compensating reactions are also important. On the other hand, the kinetics of the halide catalyses, which have been the main basis for the theory of compensating reactions, appear from more recent work to indicate the participation of intermediates probably of a peroxidic nature. All catalyses studied so far can be accounted for qualitatively by either the existence of intermediate peroxides or the occurrence of oxidation-reduction reactions. Of course the actual reactions concerned are in general found to be more complex than those given in the simple schemes written above. Thus in the compensating reactions scheme the overall reduction and oxidation of the catalyst usually involve two or more consecutive steps, and the same probably holds in general for the formation and dissociation of intermediate peroxides. The aim of kinetic studies is of course to elucidate these details, but the difficulties involved in this can be judged from the fact that, among the systems investigated and to be described subsequently, there is not one for which the kinetics of the catalysis in all conditions have been accounted for satisfactorily and quantitatively in terms of a detailed reaction mechanism. In both theories considered above, which have been current for many years, i t will be seen that the decomposition of hydrogen peroxide into oxygen and water is always accompanied by the stoichiometric formation of the peroxidic intermediate or an equivalent amount of oxidation and reduction of the catalyst. However, from recent work it appears that this need not always be the case. Thus Haber and Willstatter (3) in a speculative paper, which attempted to explain certain aspects of some

.

34

J. H. BAXENDALE

chain reactions in aqueous solution, proposed the participation of very reactive free radical intermediates. In particular, they suggested that, when acting as a reducing agent, hydrogen peroxide lost electrons in two separate steps, so that the radical HOz is formed in the first electron transfer as follows: A+++

+ H20,+

A++

+ HOz + H+

This radical may then initiate a chain decomposition of hydrogen peroxide as follows: HOz HO.

+ H 2 0 z h HO. + + HzO + HzOr HOZ + HzO 0 2

4

in which the hydroxyl radical takes part. With such presumably highly reactive entitites it can be supposed that these reactions are fast, and hence long chains are possible. This idea of reaction in single electron transfer steps was extended by Haber and Weiss (4) to the oxidation by hydrogen peroxide. I n this case the hydroxyl radical may be formed as follows : A++

+

H202-+

A+++

+ OH’ + HO.

and this may then take part in the above chain reaction. The combination of these steps results in catalytic decomposition and involves the compensating reaction principle. However, it is possible for the bulk of the catalysis to arise from the chain reactions which do not involve the catalyst directly and hence for the amount of decomposition to be many times greater than the turnover of the catalyst. This possibility was not envisaged in the earlier theories. As will be seen, there is now much evidence for the existence of radicals in some of the reactions of hydrogen peroxide, and their presence in others, although not yet proved, must be considered as a possibility. However, although the Haber-Willstatter chain reactions have been assumed to occur in certain catalytic systems, notably tbe ferrous-ferric ion system (4), more recent evidence to be described subsequently, does not support this assumption. On the other hand, such reactions appear to offer the most plausible explanation for the photochemical decomposition of hydrogen peroxide, although even here a satisfactory analysis of the kinetics has yet to be made. The variety in the mode of reaction of hydrogen peroxide makes it essential that a detailed kinetic study of any system should be made before any conclusions are drawn about a reaction mechanism and justifies the statement by W. D. Bancroft ( 5 ) that I ‘ . . the chemistry of hydrogen peroxide is a hopeless subject for the phenomenological or Baconian experimenter because misleading experiment is everywhere.”

.

DECOMPOSITION OF HYDROGEN PEROXIDE BY CATALYSTS

35

I n the following description of various systems an attempt has been made to point out the evidence for these various ideas, and although a qualitative interpretation is possible in most cases, it will be seen that the details of many systems are still in need of elucidation.

11. HALIDESAND HALOGENS Investigations on these systems were among the earliest t o be made in the field of chemical kinetics and the work of Abel, Bray, Livingston, Liebhafsky, and others has provided what must be one of the most extensive kinetic surveys of chemically related reactions yet carried out. The work up to 1931 has been reviewed by Bray (6), and there have been many investigations since that date. The experimental observations are that under suitable conditions halides in aqueous solution can decompose hydrogen peroxide catalytically according to the equation 2H202 + 2Hz0

+

0 2

Solutions containing originally only the halide, gradually build up a concentration of halogen until a “steady state” is reached in which the concentrations of halide and halogen are constant, so that decomposition of peroxide is the only overall reaction. If the halogen is the initial reactant, this is transformed into halide until, if the acidity is the same, an identical steady state is attained with respect to halide and halogen concentrations. The extensive investigations of Abel (7) on the iodineiodide system led him to conclude that the catalysis is the result of compensating reactions which can be written stoichiometrically as 2H+ 21’ HzOz ---t Iz + 2Hz0

+ + Iz + HzOz + 21’ + 2H+ + 02

and the steady state results when the rates of these two reactions are equal. Similar conclusions were reached by Bray and his school with regard t o the bromine-bromide and chlorine-chloride systems. Many investigations were undertaken with a view to proving that in the steady state the kinetics of the catalytic decomposition of hydrogen peroxide, and the concentration of halide, halogen, and acid could be accounted for by the kinetics of reactions (1) and (2) measured a t the steady state and also in conditions which are far removed from it. I n the cases of bromine and iodine (7,8) it has been possible to isolate reaction (1) or (2) from the other and obtain independent values of their rates. Many kinetic devices have been used, the results of which leave little doubt that the kinetics of (1) and (2) are the same at, or away from steady state conditions for all the halogens.

36

J. H. BAXENDALE

i. Oxidation of Halides by Hydrogen Peroxide Consider first reaction (1). In the case of iodine the reaction is stoichiometric in acid solutions, and the early work of Magnanini (9) and Noyes (10) established that the rate of iodine formation is accurately expressed by the equation d[Iz]/dt = ki"[HzOz][I']

+ ki[Hz0z1[I'I[H+]

(a)

where the value of k l o and kl are such that the H+ dependent term only becomes important in acidities greater than about 10V N (see Table I). For the bromide reaction, Bray and Livingston (8) showed that over the range of acid covered by their experiments, the kinetics are adequately described by the H+ dependent term only. The chloride reaction has not been isolated. Now it is clear that if (1) and (2) proceeding at equal rates account for the rate of decomposition of hydrogen peroxide in the steady state, then this rate must twice that of reaction (1) under steady state conditions. This has been established for iodide (7,11,12,13) where a t the pH's convenient for the measurement of the catalytic decomposition k l o >> k l [H+]and -d[HzOz]/dt = ko0[Hz02][I']

where k,"

=

2k10. I n the bromide case (8) the rate is given by

-d[Hz02]/dt

= k,[ H2OZ] [Br'] [H+]

where k, = 2kl. This correlation of the rates of (1) and the catalysis justifies the application of measurements of the latter t)o determine the former, and using this method Mohammad and Liebhafsky (14) have extended the measurements. They find that the kinetics of reaction (1) for all the halides can be expressed by the general E q . (a) above, and have determined klo, k,, and the constants in the Arrhenius equation k = A exp ( - E / R T ) in each case. These are given in Table I. The numerical values for chloride and bromide do not agree exactly with those obtained by other workers, the reason being that the rate constants vary with ionic strength. Livingston and Bray (8,15) have considered this salt effect in some detail and have shown that it is of the order expected on the Bronsted theory. The iodide reaction was thought to be anomalous in that no salt effect could be observed (16). However, recent measurements by Wynne-Jones (17) indicate that this is because most of the work has been done in the region of ionic strength where the activity coefficients concerned are almost constant. At lower ionic strengths the expected variation is observed.

37

DECOMPOSITION OF HYDROGEN PEROXIDE BY CATALYSTS

TABLE Ia (Units mole-' liters min.-l)

I' Br' C1' (1

k," (25°C.)

' k , (25°C.)

0.69 2.3 X 1.1 X

10.4 1.4 X 5.0 X

El" (kcal.) EL (kcal.) 13.4 21.1 23.6

10.4 16.7 20.7

A1

.

AI .

4.9 66 22

4.6 24 79

Data from reference 14.

2. Reduction of Halogens by Hydrogen Peroxide These have been measured directly for all the halogens under conditions where reaction (2) is isolated and it is found that they are stoichiometric according to Eq. (2). Using a flow method Bray and Livingston (18) established that the rate of reaction (2) for bromine is given by

over a range of acidities from 0.38 N to 2 x low4N . Makower and Bray (19) have found that the analogous equation for chlorine holds from 2 N to 5 N acid, but below 2 N discrepancies occur (see below). Liebhafsky (20), in extending the original observations of Abel on the iodine system, confirmed that in this case the dependence of the kinetics on the acidity was rather more complex than indicated by Eq. (b). I n the range pH 6 to pH 7 the rate varies inversely as [H+I2while from 0.01 to 0.31 M acid the variation is simply as inverse [H+]. Hence we have

Values of the rate constants and activation energies as far as they have been obtained are given in Table 11. For these reactions as for the halide reactions the specific rate constants are dependent on ionic strength. TABLE I1 k2

I2 (212

Br 2

(25°C.)

6.5 X 10-lo 5 x 10-3 1.8 X

E

k z o (25°C.)

Reference

12,000 16,000

1.18 X 10-l2 -

20 19 18

3. The Steady State

It is clear that if the kinetics of reactions (1) and (2) measured in isolation remain unchanged when the system is in the steady state, i.e.,

38

J. H. BAXENDALE

involves only the catalytic decomposition of hydrogen peroxide, then the concentrations in this state will be determined in the general case by equations similar to (a) and (c). Since for halogen Xz, tl[Xz]/dt = 0, we have

F has been called "the steady state function" and has been obtained in all cases by measuring the stationary halide and halogen concentrations. According to Eq. (d), in the general case F should be a function of hydrogen ion concentration, but it so happens that over the ranges of acidity convenient for its determination, F for each halogen when corrected for salt effects, is almost constant. This is to be expected from Eq. (d) since for chloride and bromide the acid range halppens to be such that kl[H+] >> kl" and kz[H+] >> kz". Hence in these systems it is to be expected that F = kz/kl and is independent of hydrogen ion concentration. For iodide the experimental range is from pH 4 to pH 8. Here k1" >> kl[H+] and for most of the range kzo >> kz[H+]. Thus from Eq. (d) we expect F = kz"/kl" and again independence of hydrogen ion concentration. There is also good numerical agreement between the directly determined values of F and those calculated from (d). Thus over a range of about 10l2in the product [H+]2*[I']2 Liebhafsky (13) found that F varied between 2.1 X 10-l2 and 4.9 X 10-l2 which agrees favorably with 1.7 X 10-l2 calculated from kzo/klo. Liebhafsky confirmed Abel's observation that F increased slightly a t the higher acidities which was attributed by Abel t o the fact that here kz is becoming appreciable. For the bromide-bromine system after correcting for salt effects (15) the value 1.7 is obtained for F. This is to be compared with 1.3 calculated for kz/kl after similar corrections for salt effects. For chloride-chlorine Makower (21) found F to vary from lo8 t o 2.7 X los for a range of HCI 2.84 N-5.25 N . This is paralleled by variations in kl and kz over the same range and can be accounted for by salt effects, and Makower has shown that the calculated and observed values of F are in excellent agreement. I n addition the observed temperature dependence of F is found to be given exactly by the difference in activation energies of k z and k l as required by (d). One further observation in this system is that the anomalous variation in kz at acidities below 2 N , which was mentioned above and which cannot be accounted for by salt effects, also appears in F . This has been investigated in detail by Connick (22) and is discussed below. Although there is general agreement on the existence and explanation of the steady state, Griffiths and McKeown (23) hold the view that it is a

DECOMPOSITION OF HYDROGEN PEROXIDE BY CATALYSTS

39

“pseudo steady state.” They report that for bromide-bromine there is a gradual increase in the “stationary” bromine concentration as the peroxide concentration decreases, and simultaneously the rate constant for the peroxide decomposition decreases slightly. There has been no confirmation or refutation of this work.

4. Reaction Mechanisms Although it has been well established by the work described above that the catalytic decomposition of hydrogen peroxide by the halides arises as a result of the compensating reactions (1) and (a), it is obvious from the kinetics of these reactions that the detailed mechanisms are more complicated than the stoichiometric equations would indicate. The early workers in this field recognized that the kinetics of reaction (1) as given by (a) could be explained by the sequence

+ kio + + +

+

Slow HzOz I’ + 01‘ HzO 01’ H+ F? HOI Fast HOI I’ H + - +HzO

(3)

+ Iz

together with the third order reaction ki

+ H+ + 1’-

Slow HzOz

HOI

+ HzO

(4)

Similarly the kinetics of the halogen reactions (Eq. b) are accounted for if the reacting molecule in the slowest step is the hypohalous acid produced by hydrolysis of the halogen

+ HzO HOBr + H+ + Br’ ks HOBr + H ~ 0 2 4 H+ + Br’ + + H20 Kh

Brz

0 2

(5)

In this case

and hence k z = k6Ka. For iodine where Eq. (c) applies, the additional inverse square dependence on H+ could arise if the ion 01’ reacted as well as HOI (7).

+ 01’ k 01’ + Hz02 I‘ + HzO + HOI

K O

H+

+

0 2

I n which case -

dt

=

kso[OI’][Hz02]= ks0K&

and kZ0 = ks”KaKn

*

[I~I[HzO~l ____ [H+12[I‘l

40

J. H. BAXENDALE

The reactions have been analyzed in this way and the constants kg and k6’ derived. However, although they will account for the kinetics the above reaction schemes are not of course unique in this respect, as was recognized by Bray and by Liebhafsky. There are in fact certain arguments against them. Thus the three-body collisi.on required by reaction (4) is unusual, although this objection might be overcome formally by assuming that the peroxide reacts through the ion HyO2+ formed rapidly by the equilibrium (24) : HzOz

+ H+ i=? H302+

Again in reaction (3) it is rather an involved process to transfer an oxygen atom from the structure HO-OH to the ion I+ since i t involves the breaking of two bonds and formation of two bonds in one step. A more probable alternative would seem to be the ion displacement reaction I’

+ HO-OH

-+

HOI

+ OH’

which is analogous t o the halogen hydrolysis step (25) HO’

+ CI-CI

+

HOCl f CI’

This objection on the grounds of complexity applies even more to reaction (5). Makower and Bray (19) have written the reaction schematically as: , I

I ’

H-j-0

/ I ; ~

O-;-H

,

I

,

I

H

II-;-0

~

I :

....

O--;-Cl or

O-/-Cl

,.-I.. . + ..-. 1 0 H 1 -+ €I+ + 0 I I

2

+ H i 0 + C1’

I

H

but it would seem desirable t o obtain additional evidence for such a complicated rearrangement of bonds before considering it as established. Such evidence would be for example to confirm that the reaction of the hypohalous acid itself is bimolecular according to Eq. ( 5 ) , and that the rate constant for this reaction measured directly is idsentical with t h a t calculated from kz and Kh. An investigation on these lines with hypochlorous acid has been done recently by Connick (22) in Bray’s laboratory. It was undertaken t o examine the deviations from equations (b) and (d) which occur in the chloride-chlorine system a t lower acid concentrations, and the results are probably important t o a consideration of the mechanism of reaction (2). Over a large range of hypochlorous acid, hydrogen peroxide .and chloride ion concentrations, Connick has found that the kinetics of reaction (2) are given more exactly by -_ d[HzOz] _ _- kaks[HzOs][Clz1 dt

- k?[H+l[Cl’]+ ka

(el

DECOMPOSITION O F HYDROGEN PEROXIDE BY GATALYSTS

41

At high concentrations of acid this expression leads to Eq. (b), as had been found by Makower and Bray (19), but at low acidities the rate is slower than that given by (b) and is found by Connick to conform t o the expression

-~ d[H2021= ka[H2Oz][ClZ]= dt

&

[H2O,][HOCI~[H+][CI’]

(f )

In the region of p H 4 this changes t o

--d[HzOz] = dt

[HzO21[Clz1 = [H+]2[Cl‘]

K h

[HzOzI[HOCll [H+]

(d

and in still more alkaline solutions the reaction is no longer stoichiometric according to Eq. (2). Observations over such a wide range of concentrations have not been made for the other halogens. However it may be significant that from 0.01 t o 0.31 N acid Abel and Liebhafsky (7,20) have found Eq. (b) to hold for iodine, while from p H 6 to p H 7 Eq. (g) holds. Unfortunately the intermediate conditions, where Eq. (f) might be expected to hold, have not been examined. The possibility is clearly present th a t iodine is exactly analogous t o chlorine in these reactions, with the difference that the kinetic expressions apply over different ranges of acidity. It is clear from this work th at for chlorine a t least, the original mechanisms for reaction (2) must be modified, and it seems probable that this may also apply t o the other halogens. Bray (22) has suggested the alternative scheme: ks + Clz + HOOCl + Hf + C1’ kr HOOCl + H+ + C1’ + ks

HzOz

(6) and ( 7 )

0 2

which leads t o the observed Eq. (b), (e), or (f), depending on the relative values of kT[H+][CI-] and kg. Connick explains Eq. (g) in terms of reaction through the hypohalite ion (reaction 5a), as proposed previously for the iodine reaction which shows the same kinetics a t high pH’s. Alternative schemes using Bray’s assumption of the intermediate HOOCl but involving the more probable ion displacement type of reaction and leading to the same kinetics would be: For Eq. (b), (e), or (f)

+ + + + HOz’ + ClOH HOOCl + OH’ HOOCl -+ H+ + + C1’ +

HzOz HOz’ H+ HOz’ Clz Ft HOOCl C1’ 0 2 C1’ HOOCl + Hi

+

For Eq. ( d

0 2

42

J. H. BAXENDALE

6. Radical Mechanisms As was mentioned in the introduction and is discussed more fully under catalysis by iron salts, it is possible for hydrogen peroxide t o react by single electron transfers thus giving rise t o the free radicals HO. and HOz. Mechanisms for the reactions of the halides and halogens based on these ideas have been proposed by Weiss (26) and by Abel (27). The oxidation of iodide may be formulated:

+ HzOz + I + OH’ + HO+ HO*+ I +OH‘

I’ I’

21-+ I*

These equations lead to the observed kinetics provided reaction (9) is rate determining. The observed dependence on hydrogen ion concentration (Eq. a) will be given if in addition t o this there exists the analogous reaction with H3 02+(see above) I’

+ HsOz+-+ I + HO. + Ha0

For the reduction of halogen it is suggested that there exist the equilibria,

+

If HzO @ HOI HOI H+ F? H20

+

involving the halogen cation.

+ H+ + I’ + I+

Together these are in effect: If

* I- + I+

(12)

By the electron transfer steps above we then have according to Abel,

while in addition to (12) and (13) Weiss suggests:

+

HOz HzOz-+ 0 HO. 1’- I I+I+It

+

+ HzO + HO. + OH’

2

Provided the equilibrium (12) is set up rapidly both these schemes will account for the earlier results on the chlorine and bromine reactions expressed by Eq. (b) above. For iodine an additional hydrogen ion dependent term is required (Eq. c) and Abel suggests that this arises through the reaction of peroxide as 02-in addition to HOz- in reaction (13). Although reactions such as these which involve radicals would conform to the pattern of hydrogen peroxide reactions in ot,her oxidizing and reducing systems, they fall short in certain respects. Thus consider the

DECOMPOSITION OF HYDROGEN PEROXIDE B Y CATALYSTS

43

oxidation of halide by the reactions-(9), (lo), (11). The reaction (9) has been calculated to be endothermic by 23 kcal. (28). Now the activation energy of a reaction must be at least equal to its overall endothermicity. This is clearly not satisfied here since we have seen in Table I that the activation energy for the iodide reaction is only 13.4 kcal. The situation for bromide and chloride is no better, since the corresponding endothermicities for reaction (9) are 38 and 52 kcal. respectively, which are far in excess of the observed activation energies 21 and 23.6 kcal. We may conclude therefore that reactions of the type (9) do not play a significant part in the halide oxidation reactions. Further although the above free radical reactions for the reduction of the halogens will explain the experimental results in certain concentration ranges, no radical mechanism has yet been suggested to include the more extended observations of Connick. On the other hand the experiments of Griffiths and McKeown (23,29) would indicate that in the photochemical bromide-bromine catalysis radicals and atoms do participate. The position at present in this field is that it is well established that the catalytic decomposition of hydrogen peroxide in halide-halogen systems is the result of compensating reactions which can be written stoichiometrically as Eqs. (1) and (23. For the same conditions the rates and kinetics of these reactions are the same both in the steady state and a t a distance from it. However, the detailed mechanisms of reactions (1) and (2) are still open to question, and it would appear that as a result of the more recent work, those suggested by the earlier workers are in need of extension or revision. I n considering the type of experimental approach which would assist in the clarification of the position, it is obvious that further evidence on the existence of intermediates of the type suggested by Bray would be useful. Also of importance would be the determination of the source of the oxygen which is evolved. This should be simple using HzO1sand it is surprising that it has not yet been done. It is clear from the various mechanisms outlined above that this knowledge will be significant only if all the oxygen does not come from the peroxide. If a large fraction originates in water this would be sufficient evidence to reject conclusively a number of the possible reactions discussed above. 111. IODATE

As another example of catalytic decomposition of hydrogen peroxide by compensating reactions, Bray and his school have studied the system iodine iodate-hydrogen peroxide. In solutions of moderate acidity (- 0.1 N ) iodate ion decomposes hydrogen peroxide and is itself unaffected at the end. Bray and Caulkins (30) observed that during the

44

J. H . BAXENDALE

reaction the solution is faintly colored with iodine. Subsequent work by Liebhafsky (31) showed that provided the iodine was removed as it forms, the iodate can be almost quantitatively reduced to iodine according t o the equation: 2103’

+ 5HZO2+ 2H++ Iz + 6Hz0 + 5 0 2

(1)

On the other hand, in acid solution iodine is slowly oxidized to iodate by hydrogen peroxide. The oxidation can be speeded up considerably by high acidities and also by the addition of iodate. I n these circumstances the reaction can be made to go quickly and almost quantitatively according to the equation Iz + 5H202 -+ 2103

+ 2Hf + 4H20

(11)

Reactions I and I1 occurring together will obviously lead to a catalytic decomposition of hydrogen peroxide similar in principle t o the halidehalogen catalysis. The interrelation between the reactions of iodide, iodate, and iodine with peroxide and the conditions for the isolation of each have been discussed by Bray and Liebhafsky (32) The catalysis itself has not received much attention, but the short investigation by Bray and Caulkins (30) showed it to ble rather complex. Measuring the evolution of oxygen with time for an iodate-peroxide-acid solution they found that the rate of evolution increased as the acid increased from 0.055 N to 0.110 N , but between these limits the evolution showed periodicity. The coloration due to iodine was also periodic, being formed during the slow evolution of oxygen and disappearing during the following rapid evolution. Whether this plzenomenon arises from a homogeneous reaction was questioned by Itice and Reiff (33), but the observations were later confirmed by Bray and Liebhafsky (32). Other observations on the system have been concerned mainly with the kinetics of reactions (I) and (11) in isolation. Liebhafsky (35) followed reaction (I) by removing the iodine with an organic solvent. After a short induction period the rate follows the equation: d[Izlldt

=

ki[Hz0zI[IOs’l

+ kzIH+I[HzO~l[IOs’l

where kl = 2.6 x LO4 and k2 = 129 x lo4 a t 50°C. that this equation arises from the slow reactions: H+

+

108’ 108’

+ H z 0 2 4 1 0 2 - + Hz0 + + HzOz -+ HIOz + HzO + 02 0 2

which are followed by the rapid reactions

+

+

HIOz Hf I’ -+ 2HOI HzOz + I‘ H+ Ha0 HOI HOI + I ’ H + + Hz0 Iz

+

+

+

+

+

+ 02

It was suggested

DECOMPOSITION O F HYDROGEN PEROXIDE B Y CATALYSTS

45

When catalyzed ‘by iodate and high acidity reaction (11) was found by Bray and Caulkins (34) to be first order in iodine concentration and independent of peroxide concentration. With increasing iodate and acid concentration the rate increases but ultimately becomes independent of these. It was considered that this dependence of the rate on the iodine concentration only, indicated that the reaction 12

+ HzO+

+ H+ +I’

HOI

became rate determining. However a subsequent investigation by Liebhafsky (35) shows this to be an oversimplification since the independence of acid, iodate, and peroxide concentrations only holds over a certain range of the latter. I n attempting to explain the rather complicated kinetics outside this range Liebhafsky maintains the slow hydrolysis of iodine and introduces another intermediate, H21203, previously suggested by Bray. His scheme is: Slow H20 + I * - + HOI I’ H+

+ + + + + + + + + + + + + + + +

Rapid equilibrium 103’ I’ 2H+ e H21203 Slow HzIz03 F? HOI HI02 HzIz03 HzOz + 2HI02 HzO HI02 HzOz + H+ 103‘ Hz0 HI02 I’ H + + 2HOI HOI HzOz+ H+ I‘ HzO 0 2

He considers the induction period t o be due to the slow build-up of HIOn, which is speeded up by any change which will lower the iodide concentration. In the catalytic decomposition of hydrogen peroxide it is suggested that the main oxygen evolution step, as in the halide-halogen system, is the reaction of HOI with hydrogen peroxide. Abel (36) has proposed a series of electron transfer reactions involving free radicals similar in principle to those suggested for the iodide-iodine kinetics. He suggests reaction through iodine cations formed for example in the equilibrium : 103’ e IOz+

+ 0”

For the reduction of iodate he writes: IOz+ IO+

+ HOz+ HOz-4

-+

102 I0

+ HOz + HOz-+

--t

+ H+ + + H+ +

102’ 10‘

0 2

0 2

and subsequent oxidation of HOI t o iodide and thence t o iodine. the oxidation of iodine to iodate:

+ Hz02 + HzOz Is+ + 3H20+ I+ 13+

--t

-+

+ +

+ +

OH’ HO. + 13+ 14+ OH‘ HO. 4 16+ HI08 5H’

I2+

+

+ 20H’ + 20H’

For

46

J. H. BAXENDALE

However, the experimental data are too fragmentary to support any particular mechanism, and there are many aspects of this very complex system, such as the catalytic periodicity and a certain amount of irreproducibility in the kinetics, which are still obscure.

IV. IRON SALTS The effectiveness of a mixture of iron salts and hydrogen peroxide as a n oxidizing agent has been known for a long time. Pis early as 1860 Schonbein found that the oxidation of iodide ion by hydrogen peroxide is considerably accelerated by the presence of iron salts. The important discovery by Fenton (37) in 1894 that a mixture of a ferrous salt and hydrogen peroxide could oxidize many hydroxylic orgmic compounds showed that the mixture possessed potent oxidizing properties not present in the separate reagents. These interesting reactions were the subject of many investigations around this period, but the first ;significant study of the reactions of iron salts and hydrogen peroxide alone was by Manchot and Wilhelms (38) in 1901. They found th a t oxygen was evolved from hydrogen peroxide not only by ferric salts but also by ferrous salts during their oxidation to ferric. Their observations werle later confirmed by Mummery (39) who found that ferrous salts led t o the evolution of oxygen only when they were added to the peroxide under conditions such th a t the latter was always in excess. However, if t'he peroxide was added t o the ferrous salt solution, so that ferrous ion was always in excess, then no oxygen was evolved and the reaction was quantitative according to the equation 2Fe++

+ HzOz + 2H+ = 2Fe++++ 2Hz0

(1)

The evolution of oxygen during the ferrous ion oxidation arises from the catalytic decomposition of hydrogen peroxide according to 2H202 = 2Hz0

+

0 2

(11)

In the same investigation Mummery found that after the initial evolution of oxygen (which occurred too rapidly for him to follow), there was a slow evolution which corresponded in rate t o that given by the same concentration of ferric salt. However, there was nct initial burst of oxygen in the ferric case. These observations have been confirmed by subsequent work and, contrary t o the assertions of v. Bertalan (40), Spita1sk.y and Petin (41), and Bohnson and ltobertson (42), it is well established that in suitable conditions the oxidation of ferrous ion to ferric by hydrogen peroxide can lead t o catalytic decomposition of the hydrogen peroxide. Unless the solutions are very dilute this decomposition is very fast and greatly

DECOMPOSITION OF HYDROGEN PEROXIDE BY CATALYSTS

47

exceeds the rate of decomposition by the same concentration of ferric salts under the same conditions. Because of this difference in rate it is usually possible to separate experimentally the catalytic decomposition by the two ions, and although, as will be seen later, it is probable that the same reactions are involved in both, it will be convenient to treat them separately. 1. Decomposition b y Ferrous Ion The earlier work was concerned only with measuring the total amount of catalytic decomposition when ferrous ion is oxidized and until quite recently (43), no attempt was made to follow the kinetics of the oxidation in these conditions. It is convenient t o express the amount of catalytic decomposition of peroxide in terms of what Haber and Weiss (4)have called the consumption ratio n given by

Thus in the absence of any catalysis, ferrous ion is oxidized according to reaction I above and n = 0.5. The evolution of oxygen indicates the occurrence of reaction I1 and leads to values of n greater than 0.5. Mummery measured n at constant ferrous ion with increasing peroxide concentrations and found that n increased a t first but tended to a limit at high concentrations. The quantitative value of his results is decreased by the fact that no free acid was present. Hence the acidity varied during the reaction and usually the final solutions were so alkaline that hydroxylated ferric salts were precipitated. Manchot and Lehmann (44) worked in acid solutions and found that n decreased with increasing acidity. They also claimed that their results indicated a value of n = 1.5 and adduce this as evidence for the existence of a higher oxide of iron Fez06. However, an examination of the data shows that this is only true for a very limited range of peroxide concentrations. Subsequently Haber and Weiss (4) reported that even with reagent molar, values of n of 15 or more are obtained concentrations as low as in “neutral solution.” These high consumption ratios, not observed by previous workers, were attributed to the use of more efficient mixing techniques. They confirmed Mummery’s observation that n increased with increase of peroxide concentration, but unlike Mummery they did not find any tendency for it to approach a limit. Their results could be fitted to the expression n = 0.5

+ K[H,O,]/[Fe++]

where K varied slightly with the acid concentration. Using a rapid mixing technique followed by quenching in permanganate after a few

48

J. H. BAXENDALE

seconds, they were able to measure the rate of ferrous ion. oxidation in the conditions where n = 0.5. The reaction rate was found to be independent of acid concentration and first order in both peroxide and ferrous ion with a bimolecular rate constant 23 mole-l/liter sec:-1 a t 20°C. and activation energy 8.5 kcal. Under similar conditions but with much lower concentrations so that reaction times were of the order of half an hour, Baxendale, Evans, and Park (45) confirmed the bimolecularity of the reaction and obtained a rate constant of 62 mole-l/liter sec.-1 a t 25°C. and activation energy 10.1 kcal.

1.00

n

0.75

0.50

I

I

I

2 .O 3.0 4.0 Concentrotion of hydrogen peroxide (Y:i 1.0

5.0

FIG. 1.

'ariationu of consumution ratio with oeroxide concentration. IFe++ (a) pH 1.5, (b) pH L8, (c) pH 1.5 with 0.04 M Fe+++. From 'Barb, Baxendale, George, and Hargrave, 43. = 0.04 M .

More recently the ferrous ion-hydrogen peroxide system has been examined in greater detail by Barb, Baxendale, George, and Hargrave (43). They have measured overall values of consumption ratios for a wide range of conditions, and in addition have followed the kinetics of ferrous ion oxidation in circumstances where there is catalytic peroxide decomposition. Contrary to the findings of Haber a,nd Weiss but in agreement with Mummery, they observed that at constant ferrous ion and acid concentrations n reaches a limit as the peroxide concentration is increased (Fig. 1). This limiting value increases with increase of ferrous ion or decrease of' acidity. Also, a t constant ferrous ion, acid and peroxide concentrations n is increased if ferric ion is added initially to the reaction mixture but ultimately reaches a limiting value a t high ferric ion concentrations. Cupric ion behaves similarly, but for the same con-

DECOMPOSITION O F HYDROGEN PEROXIDE BY CATALYSTS

49

centration is more effective than ferric.* Conversely n is decreased if a8 the ferric ion is formed, it is removed as the complex fluoride by having fluoride ion present. This effect of fluoride has also been reported by Humphrey and Weiss (46). These observations on n are paralleled by the kinetic measurements of Barb et al. Thus it is found that when n > 0.5 deviations occur from the previously established bimolecular rate of oxidation of ferrous ion. These deviations are in the sense that, although initially the oxidation is

0 TIME. seconds

FIG. 2. Rate of oxidation of ferrous ion. [Fe++]a = 1.4 X M ; [H202]0 29.5 X M . (a) 0.5 ilf HC104, (b) pH 2.65 (HCIOa), (c) pH 2.65 with 4.6 X M Fe+++, (d) pH 2.65 with 1.4 x loa8 M Cu++, (e) pH 2.65 with ( 0 ) 1.2 X 10-3 Fe+++,and (A)2.4 x 10-3 Fe+++. From Barb, Baxendale, George, and Hargrave, 43.

as fast as expected, it falls off more rapidly with time (Fig. 2). It is shown that this fall off is due to the ferric ion produced in the reaction since the rate can be slowed up initially if ferric ion is added, The deviations from bimolecularity follow exactly the variations in n described above, in that they increase with pH and they reach a limit at high ferric, cupric or peroxide concentrations (Fig. 2). The indications are that the evolution of oxygen which leads to n > 0.5 is accompanied by a decrease in the net rate of ferrous ion oxidation. Since ferric ion enhances both effects, the authors conclude that the oxygen is evolved simultaneously with a reduction of ferric ion to ferrous. The details of their conclusions will be given later. *The effect of cupric ion on the ferrous catalysis is also shown in the work of Spitalsky and Konovalova (51) but was not recognized as such.

50

J. H. BAXENDALE

2. Active Intermediates From the ready oxidation of many organic substances by solutions containing ferrous ion and hydrogen peroxide it was obvious that active intermediates were present in the reaction medium. Fenton (37) suggested that the activity arose from the formation of complexes between the iron and organic substrate. This was also concluded by Wieland and Franke (47) from an extensive investigation of the Fenton reaction using a wide variety of substrates. However many substances, such as benzene which gives phenolic products (48), or alcohol which yields aldehyde and acetic acid (49), show no tendency a t all t o complex with ferrous or ferric ion. Manchot and Wilhelms (38) considered the active intermediate t o be Fez05,but the experimental basis for this was shown t o be unfounded by Haber and Weiss. Mummery (39) suggested the existence of a ferrous salt-hydrogen peroxide complex [also proposed previously by Brode (SO)] t o explain both the catalytic decomposition and the Fenton oxidation. Mummery called it a “perhydrol” and wrote the reaction Fe

+ H z O F!~ Fe

\OH HzO

/

Fe

+ Fe

+ Fc

‘OOH SO4H

\eon

\

\OH

+ Hz0 OOH -+

2Fe-OH

\OH SO cH

+ H a 0 2 4 HzO + O2 + Fe/

\

OH[

For the Fenton oxidation, reactions involving the organic molecules instead of hydrogen peroxide were written for the lasit step. It can be seen that the competitibn between the ferrous salt and hydrogen peroxide for the complex determines the extent t o which there is catalytic decomposition. This is consistent with his observations t h a t n increases with increase of [H202]/[Fe++] but it does not explain the ultimate limit which n reaches. Manchot and Pflaum (52) wrote similar reactions involving the intermediate FeS04.(H202). Iron in the form of ferrate Fe04/’ was suggested as the seat of the activity by Bohrison and Robertson (42), who claimed t o have obtained spectroscopic evidence in its favor. As has been pointed out, however (53), ferrate is unlikely t o exist except in very alkaline solutions. More-

DECOMPOSITION OF HYDROQEN PEROXIDE BY CATALYSTS

51

over the colored solutions obtained by Bohnson and Robertson and on which they based their suggestion very probably had their origin in the stabilizer usually present in hydrogen peroxide (2). Another higher valence form of iron, the ferry1 ion FeO++, was proposed by Bray and Gorin (54). They suggested that it was in equilibrium with ferric ion as follows : 2Fe+++

+ HzO F? Fe++ + FeO++ + 2H+

It could also be formed from hydrogen peroxide and react according to: Fe++ FeO++

+ HZOZ+ FeO++ + HzO + HzOz-, Fe++ + + HzO 0 2

As was the case for many other solution reactions the picture was clarified somewhat by the free radical concept introduced by Haber and Willstatter (3). They accounted for the rapid decomposition of hydrogen peroxide by the enzyme catalase in terms of a chain reaction involving the free radicals HO. and HOT

+ H202-+ (Fe++enzyme) + HOa + H+ + HzOz+ HzO + + HO. HO. + HzOz HOz + Hz0

(Fe+++enzyme) HOz

Chain

0 2

-+

This idea was subsequently applied to the iron-peroxide reactions by Haber and Weiss (4) and in addition to the above chain, the initiation reaction Fe++

+ HzOz

HO.

--f

was suggested for the ferrous reaction. possible chain terminating steps: Fe++ Fe++

+ OH- + Fe+++

The following were proposed as

+ HO. + Fe+++ + OH-

+ HOZFe+++ + HOZ-+ Fe++ + HOz + HOZ

-+

Fe+++

+ HO,

+

Fe++

In the ferric reaction

Fe+++

+ 02 + H+

were considered important. We shall discuss later how far the details of the reaction schemes of Haber and Weiss can account for the experimental data. What appears to be conclusive evidence for the participation of free radicals in some hydrogen peroxide systems is contained in the work of Baxendale, Evans, and Park (45). They discovered that monomers such as acrylonitrile, styrene, and methyl methacrylate are rapidly polymerized if they are present when hydrogen peroxide is reduced by ferrous,

52

J. H. BAXENDALE

cuprous, titanous, and chromous ions in acid solution. Rapid polymerization of these compounds coupled with its easy inhibition by oxygen is generally accepted as evidence that the reaction proceeds by a free radical mechanism in the following way: Initiator -+ R. (free radical)

+ CHz=CHX RCHZ-cHX RCHAHX + C H ~ = C H X-+ R.CH~CHX--CH~.CHX. R.

-+

-+

Polymer

Hence it appears that the peroxide reduction reaction can supply the initiating radical R. The reduction by ferrous ion was studied in detail. I n concentration conditions which would normally lead to the catalytic decomposition of hydrogen peroxide by ferrous salts, it was observed that the presence of a monomer could completely inhibit the evolution of oxygen arising from this decomposition. It was also shown that glycollic acid, which is a typical substrate for the Fenton oxidation, decreases the extent of polymerization and can inhibit it completely at high concentrations. They concluded that the hydroxyl radical produced in the manner suggested by Haber and Weiss ii3 responsible for both the polymerization and Fenton oxidation reactions in the following way :

+ +

+ +

+

Fe++ H2O2-+ HO. OHFe+++ Fe++ HO. -+ Fe+++ OHHO. CHz=CHX -+ HOCH2-CHX + Polymer

+

or HO.

+ RH

-+

H20

+ R.

-+

Oxidation products

(0)

(1) (m) (8)

Thus reaction (m) can prevent the Haber-Willstatter chain reactions involving HO, H02, and H z 0 2and hence eliminate the catalytic decomposition of hydrogen peroxide as is observed. Competition for HObetween (m) and (s) accounts for the Fenton substrate, RH, inhibiting the polymerization reaction. Further quantitative evidence was obtained by examining the effect of monomer on the consumption ratio n. It was found that in acid solutions the value 0.5, obtained in the absence of monomer, gradually increases to 1.0 as the monomer concentration is increased. Moreover the concentration of monomer required to produce this effect at one concentration of ferrous ion is increased if the latter is raised. These observations are consistent with the competition for HO. between Fe++ and monomer in reactions (1) and (m) above. Further support for reaction (m) comes from work by Baxend.ale, Evans, and Kilham (55) who found that polystyrene produced by this method of polymerization contains hydroxyl groups and also that the molecular weights of such polymers are determined by the amount of hydroxyl radical supplied in reaction (0).

DECOMPOSITION O F HYDROGEN PEROXIDE B Y CATALYSTS

53

3. Mechanism of the Ferrous Ion Catalysis

The only data of any significance for the deduction of a mechanism are those of Haber and Weiss (4) and more recently Baxendale, George, and Hargrave (43). The former measured sumption ratio n under various conditions and put forward the scheme : Fe++ HO. HO. HO,

+ Hz02 HO. + OH- + Fe+++ + Fe+++ OH- + Fe+++ + HzOz+ HOz + HzO + HzOz-+ + Hz0 + HO. +

0 2

reaction of Barb, the conreaction (0)

(1) (2) (2’)

This combination of radical reactions leads to the equation n = 0.5

+ ka[H2Oz]/k,[Fe++]

(a 1

which is claimed to account quantitatively for their experimental observations. However a critical examination of the data reveals certain inconsistencies (43). Thus it is to be expected that equation (a) should hold throughout the reaction. For some experiments Haber and Weiss have obtained values of n for the initial part of a reaction and also mean values for the complete reaction. It is significant that although the mean values in these particular conditions can be as high as 2.0, in every case the initial values of n are 0.5. I n other words it appears from these data that there is no catalytic decomposition in the initial stages of the reaction. This was also found to be the case by Barb et al. I t is also apparent from the experiments in “neutral solution” that n is not given accurately by Eq. (a) in respect of hydrogen peroxide concentration. I n fact the data indicate a tendency for n to become independent of peroxide concentration as the latter increases. This agrees with Mummery’s earlier work and is confirmed by Barb et aE. Further, Haber and Weiss find that k 2 / k l varies with acid concentration. This cannot be accounted for on the above scheme but Weiss (56) has suggested that HOa reacting in its dissociated form Oz- in (2’) will explain this. However, it is clear that this is not the case since the reaction of HOz does not determine Eq. (a). A major objection to the mechanism of Haber and Weiss is that it predicts the rate of oxidation of ferrous ion t o be first order in ferrous ion and peroxide. The recent observations described above have shown that considerable deviations from this occur when there is catalytic decomposition of the peroxide. A reaction mechanism which they consider consistent with all the experimental observations is proposed by Barb et al. (43). This comprises the following combination of radical reactions :

54

J. H. BAXENDALE

+ + + + +

+ +

+

Fe++ H 2 0 2-+ HO. OHFe+++ Fe+++ HO. Fe++-+ OHHO. H202 -+ HOz Hz0 HOz Fe++-+ HOz- Fe+++ HOz Fe+++4 Oz. H+ Fe++

+ + + +

All the individual steps in this have been considered previously by Haber and Willstiitter or by Haber and Weiss. It will be seen, however, that step (2’) of the Haber-Willstatter chain has been omitted but the other, step (2), retained. Barb et al. reject reaction (2’) since this leads to values of n continuously increasing with H202 concentration which is contrary to their observations. Reaction (2) is supported by evidence obtained from experiments on the Fenton oxidation of dyestuffs. They find that the amount of dyestuff oxidized decreases as the hydrogen peroxide increases and interpret this to mean that the peroxide competes with dyestuff for hydroxyl radicals. From the above scheme the following conclusions can be deduced about the consumption ratio n and the kinetics of oxidation of ferrous ion, all of which have been observed experimentally: 1. At very low ratios of peroxide to ferrous ion concentrations (R1) only steps ( 0 ) and (1) occur and hence n = 0.5 and the rate of ferrous oxidation is given by -d[Fe++]/dt = k[Fe++][HzOz]

where k = 2k0. 2. As R 1is increased, reaction (2) becomes important and the HOz produced in (2) leads by (4) to oxygen evolution. Hence n increases and since ferrous ion is regenerated in (4) deviations front the bimolecular kinetics are produced in the sense that lc is decreased. However initially, when no ferric ion is present, (4) cannot occur, and the H 0 2 is removed by (3) in which case the initial value of n is 0.5 and k = 2k,. On the other hand, increase in ferric ion concentration favors (4) and hence leads t o an increased n and smaller k . At very high values of R 1effectively all the HO. radical reacts with HzOzin (2) and the dominating reactions are then (3) and (4). Since these do not involve peroxide, n and k are independent of the peroxide concentration. They are, however, increased and decreased respectively by increase in the ferric ion concentration since this favors (4). The effect of cupric ion in increasing n and decreasing k can be satisfactorily explained by assuming it to react as follows: Cu++ + HOz+ CU+ Hf OZ (5)

+ + + Fe++

CU+ + F e + + + - +Cu++

This amounts to a catalysis of reaction (4) and from the greater effectiveness of Cu++ compared with Fe+++ it is concluded that kg > k d .

DECOMPOSITION OF HYDROGEN PEROXIDE B Y CATALYSTS

55

At high values of R1 the above mechanism is reduced to reactions (o), (2), (3), and (4). Under these conditions the variation of n and the kinetics of the ferrous ion oxidation with reagent concentrations have been shown to be consistent with these reactions, and Barb et al. have obtained numerical values for k 4 / k 3 from both. They find that kq/ka increases with decrease in acid concentration, and this accounts for the increased catalytic decomposition observed in more alkaline solutions. Quantitatively k 4 / k 3is given by krlkr

=

Ki/(Kz

+ [H+D

(b)

K2 is numerically the same as the ferric ion hydrolysis constant Kp (21) Fe+++

+ H 2 0K&PFeOH++ + H+

Assuming K p and K z identical, various combinations of Fe+++, HOz, FeOH++, and 02-reacting in (3) and (4) can lead to the Eq. (b). It is suggested that the most probable reactions are Fe+++

+ HOz-+Fe++++ HOZKEOa

HOz Fe+++

These will give

+

H+ 0%4 Fe++ 0

+ ka/ka = kr' . K H O ~ / ( K+F[H+l)ka +

02-

(3)

2

(4')

at 25°C. and is almost Interpreting in this way k41KEO2/k3= 7 X temperature independent. Using cupric ion, reaction (4) can be replaced by ( 5 ) and it is found that k d k a = Ka/[H+l

Assuming a reaction (5') analogous to (4') then ks/ka = k~'K~oz/ka[H'l

which means that a t the acid concentrations used the cupric ion hydrolysis constant is negligible. k6/KBOL/k3 is about 0.1 at 25°C. and almost temperature independent. Using values of R 1 such that both reactions (1) and (2) occur, and a t the same time having high concentrations of ferric or cupric ion present, the reaction scheme reduces to (o), (l), (2), and (4) *or ( 5 ) . I n these conditions n was found to vary as expected with reagent concentrations, but the kinetics of the ferrous ion oxidation were considerably affected by the presence of traces of organic impurities which could not be removed. However, it was possible to make allowance for this and show that the kinetics agreed with those predicted. Values of h / k l calculated from n and from the kinetics show it to be independent of acid concentration,

56

J. H. BAXENDALE

but whereas n gave k J k , = 0.030 a t 25"C., the kinetics gave 0.011 a t the same temperature. It is suggested that this discrepanc.y may arise from the presence of another reacting species such as ferryl ion, in the conditions of high ferric ion concentration relative to peroxide concentration obtaining in the kinetic experiments. As mentioned above the ferryl ion FeO++ was proposed originally by Bray and Gorin, and as has been pointed out by Medalia and Kolthoff (53), it is possible to replace the hydroxyl radical by the ferryl ion throughout the IIaber and Weiss reaction scheme without altering the kinetics. The same is true for the scheme of Barb et al., and it is possible th at the above anomaly is due to the fact that hydroxyl radical, as well as ferryl ion formed by the reaction Fe+++

+ €10.

---t

FeOH++++ FeO++

+ Hi

both take part in reactions (1) and ( 2 ) . The evidence for the ferryl ion is a t present not very conclusive. Kolthoff and Medalia (58) have observed that there i:j no salt effect on the Fenton oxidation of alcohol and conclude that this argues against the participation of a charged reactant such as ferryl ion with ferrous ion in reaction (1). However, as we have seen for the iodide reaction, such evidence is of doubtful significance unless very low salt concentrations are used. On the other hand, some observations of Barb et al. on the oxidation of organic impurities as well as the above cliscrepancy in the determination of k2/kl can be interpreted in terms of the ferry1 ion. Abel (59) has criticized the experimental methods and the conclusions of IIaber and Weiss and of Barb, Baxendale, George, and Hargrave. He suggests t ha t reactions ( o ) , (I), and (4) together with Fe+++

+ HO?-+

Fo++ + 1301

are suffic+ientto account for the reactions of both ferrous and ferric ions with hydrogen peroxide. This scheme is obviously a t variance with the experimental observations, but since the experimental methods are questioned any confirmation must await further work

4. Decomposition by Ferric Ion It was first established by von Bertalan (40) that the rate of hydrogen peroxide decomposition in acid solution containing ferric ions is given by - d [ H z 0 2 ] / d t = k,[Fe+++][H202]/[H+]

(c)

This has been confirmed by many workers (4,6,41,42,60)when the peroxide concentration used has been high compared with the ferric ion concentration. However, a t low peroxide concentrations deviations from (c) have been observed. Thus Bray (6) found that k , tended to decrease

DECOMPOSITION O F HYDROGEN PEROXIDE B Y CATALYSTS

57

when the concentration became small. More recently Andersen (61), using lower ratios of peroxide to ferric ion concentration ( R z )than had been used hitherto, has found the rate to fit the equation -4HzOzI/dt = BIHzOzlz/([HzOzl

+

+ A)

(d)

A is proportional to [Fe+++][H+]/(Kh [H+]) and B to [Fe+++]/(Kh [H+]) where Kh is the ferric ion hydrolysis constant. At high values of Rz and [H+] Eq. (d) reduces to (c). Barb, Baxendale, George, and 1.0 used by Hargrave (62) have extended these observations from Rz Andersen, to R2 0.002. They confirmed that the reaction order in hydrogen peroxide is higher than unity and concluded that the rate expression at low Rz(< 1) is given by

+

-

-

+

-dlH~Ozl/dt = k[H202]$~[Fe+++1/([Hf] Kk)

(e)

They consider that Eq. (d) is observed because the value of Rz used by Andersen is between the high values which lead to the kinetics of Eq. (c) and the low values giving Eq. (e). An interesting observation which must be accounted for in any theory of the ferric ion catalysis was made by Bohnson and Robertson, who found that the catalytic decomposition by ferric salts is considerably enhanced if cupric salts are present (51,63,64). The resulting rate is much higher than that expected from the sum of the individual ferric and cupric ion rates. For a constant ferric ion concentration the enhanced rate a t first increases with added cupric ion but ultimately reaches a limit beyond which further cupric ion has little effect. Bohnson and Robertson assumed that in this promoted catalysis the rate was first order in peroxide concentration as had been found for the ferric ion alone, but there is as yet no experimental evidence for this. Analysis of their numerical data (43,62) shows that when the maximum promotion by cupric ion has been reached the rate of peroxide decomposition is proportional to [Fe+++]"/[H+]. 5. Mechanism of the Ferric Ion Catalysis

It was suggested by von Bertalan (40) that the catalysis occurs as a result of the alternating reduction of ferric and oxidation of ferrous according to the equations: 2Fe+++ 2Fe++

+ HzOz + H202

--t

--f

+ +

+

2Fe++ 2H+ 0 2Fe+++ 2H n0

2

This is very similar to the catalysis by compensating reactions as originally proposed by Abel for the halide-halogen-peroxide systems, and this author has discussed the iron catalysis in these terms (59). However, it

58

J. H. BAXENDALE

was recognized by von Bertalan that the rate determining step which leads to the kinetics of Eq. (c) above, is simpler than these equations indicate. He realized that reaction of ferric ion with the perhydroxyl ion H 0 1 , formed as a result of the acid dissociation of hydrogen peroxide KP

HI02 & H01-

+ H+

would give the required kinetics. However, he did not, suggest the production of the radical HOz in this reaction, as was done later by Haber and Weiss (4). The above compensating reactions are attractive because of the success of similar schemes in the halide catalysis, but proof in this case is more difficult. Thus it was possible to show in the halide systems that halogen and halide are present simultaneously. Elvidence for the presence of ferrous ion in the ferric catalysis would support a similar interpretation. Manchot and Lehmann (44) claimed to have proved that ferrous ion is formed from ferric ion in the presence of peroxide since the addition of a,a’-dipyridyl to the mixture resulted in the slow formation of the red ferrous tris-dipyridyl ion Fe(Dipy)3++. However, later work (65,66), which will be discussed when these systems are considered in more detail (IV,6), indicates that the ferrous complex ion may be formed by reduction not of the ferric ion, but of a ferric dipyridyl complex. Similar conclusions on the presence of ferrous ion were drawn by Simon and Haufe (67) from the observation that on addition of ferricyanide to the system Prussian blue is formed. This again is ambiguous, since peroxide is known to reduce ferricyanide to ferrocyanide and the latter with ferric ion will of course give Prussian blue (53). Although the evidence for the participation of ferrous ion in the reaction is not conclusive, it does seem probable that the active intermediates known to occur in the ferrous ion-hydrogen peroxide system are also present during the ferric ion catalysis. Thus although they are a great deal slower in reaction than ferrous salts, ferric salts and hydrogen peroxide can also oxidize organic compounds (68) and initiate the polymerization of substituted ethylenes. Moreover, during such oxidations or polymerizations the rate of the catalytic decomposition is decreased (69,70), which suggests that chain carriers for the catalysis are removed. Of the active intermediates previously considered, the ferryl ion and the free radical HO. appeared to be the most likely entities. The reactions of the ferryl ion (see p. 51) proposed by Bray and (forin will give the kinetics of Eq. (c), but we have seen that these are insiufficient to account for all the observations on the ferrous system. The ferryl ion itself cannot be entirely rejected as a reactant since it has not been possible to distinguish between it and the hydroxyl radical.

DECOMPOSITION O F HYDROGEN PEROXIDE BY CATALYSTS

59

The ferric ion catalysis has been considered by Haber and Weiss (4) in terms of t he reactions of radicals along the lines adopted to explain the ferrous ion reaction. They proposed the following mechanism:

+ + + +

+

Fe+++ HOz- + Fe++ HOz Fe+++ HOz + Fe++ H+ Oz Fe++ HzOz-+ Fe+++ HO. OHFe++ HO. 3 Fe+++ OH-

+ + + + +

(i) (4) (0)

(1)

It will be seen that on this scheme the catalysis is not a chain reaction, and provided (i) is the slow step, the rate of oxygen evolution is the rate of reaction (i) which will be given a t high acidities by Eq. (c). The scheme has been criticized (43,53,59) because it is not consistent with the set of reactions which Haber and Weiss adopted to explain the ferrous ion catalysis. Thus in the latter the reaction of HO with peroxide (reaction 2) is important because it leads to oxygen evolution via the HaberWillstatter chain. It has been omitted from the above scheme, although it is clear that i t is even more likely to occur here if (1) occurs, since the ratio of peroxide and ferrous ion concentrations is much higher. Later modifications by Weiss (71) do not remove this objection. Medalia and Kolthoff (53) have reconsidered the various combinations of reactions (i), (o), (l), (2), (3)) (4) and conclude that the required kinetics will be given by these reactions provided the rates of (1) and (3) are small in comparison with those of (2) and (4) respectively. To obtain the hydrogen ion dependence shown by Eq. (c) they assume that in (i) the reaction is via the hydroxylated ferric ion FeOH++. FeOH++

+ HOz- + Fe++ + HOZ+ OH-

These authors also criticize the mechanism,

+ HOZ+ HOz-+

Fe+++ 0

+

FeOH++ 0 OHOz

+

which was proposed by Andersen t o account for his observations (Eq. d). They point out that the very high positive free energy change associated with the production of the oxygen atom in the first step is likely t o mean that the rate constant is prohibitively low. Barb, Baxendale, George, and Hargrave (43) have applied the results of their analysis of the ferrous reaction t o the ferric catalysis. They conclude that for high ratios of peroxide to ferric ion concentrations (R2) the significant reactions are:

8cheme A

60

J. H. BAXENDALE

This choice of reactions is based upon the fact that a t high R 2 the predominant reaction of the hydroxyl radical will be with peroxide. Hence, as was the case in the ferrous ion reaction, chain termination by (3) is favored over (1). Using approximations which are justified by the numerical values of IC~/k,and k*/lc3,the above scheme gives:

+

-d[HtOtl/dt = 2(k,kok4'K,Krro,/k3)',/L . [Fe+++][H20t]/(Kh [€I+]) (f)

Equation (f) is equivalent t o (c) a t high acidities. Their own data and those of Andersen (extrapolated to high R z )show that (f) which involves the ferric ion hydrolysis constant Kh is a more exact expression than (c) for the p H dependence of the rate a t low pH's. Barb et al. have also developed a method for measuring the ferrous ion concentration during the ferric ion catalysis. They find that the complex ion Fe(Dipy)3+++is reduced extremely rapidly b y ferrous ion to Fe(Dipy)B++ but relatively slowly by hydrogen peroxide in acid solution. If the complex ion is added t o a mixture of ferric ion and hydrogen peroxide and the subsequent reduction (followed colorimetrically) extrapolated back to the time of addition, they find evidence for an initial rapid reduction which must arise from the presence of a reducing agent other than hydrogen peroxide. The concentration of this reducing agent measured over a range of reagent concentrations is found to be quantitatively consistent with the concentration of ferrous ion which can be calculated from the above reaction scheme. They therefore conclude that they are measuring the stationary ferrous ion concentration. From the values of IC, and IC4'KRO2/k3previously obtained and the rate constant for the catalytic decomposition given by Eq. (f), the constant ki is calculated to be k , = loz4exp. (-28,00O/RT) mole-' liters/sec.-l

The very high frequency factor lo2*is attributed to a high entropy of activation resulting from the charged reactants. It is pointed out that the entropy of activation involved (50 e.u.) is of the same order as the entropy change for the reaction Fe+++

+ H02- + FeHOz+'

which has been obtained by Evans, George, and Uri (2). Barb et al. have aIso considered the ferric ion cata1,ysis kinetics a t low values of R2 where, as stated above, deviations from the von Bertalan equation (c) occur. They conclude that with decreasing R z reactions (3) and (1) will become of comparable importance as chain-terminating reactions, since peroxide will no longer be of such a concentration as to eliminate the ferrous reaction in the competition for the hydroxyl radical. At

DECOMPOSITION O F HYDROGEN PEROXIDE BY CATALYSTS

61

sufficiently low I22, (1) will be the predominating chain terminator. They calculate that this will occur at R z < 1.0 at pH 1.6 or R z < 0.1 at 0.2 N acid. Under these conditions the react,iori scheme becomes (i), (o), ( I ) , (2), (4) (Scheme B), and this leads to the equation:

+

-d[HzO%]/dt = 2 ( k i k & X , / k ~ )~~[Fe+++]f’J[H,O,]la/([H+] KtJM

(g)

The dependence on peroxide concentration indicated by (g) was confirmed. It is suggested that Andersen’s data of Eq. (d) have been obtained a t values of R z which are not low enough to eliminate reaction (3) entirely, and hence something intermediate between first and three halves power in peroxide concentrations might be expected. However, both these sets of data show a first order dependence on ferric ion concentration and inverse first power in ([H+] Kh) which are at variance with Eq. (g). This discrepancy has not yet been resolved, and Barb et al. suggest that other entities might be present at these high ferric ion concentrations which play no part in the ferrous ion reaction or the ferric ion reaction at lower ferric ion concentrations. The accelerating effect of cupric ions on the ferric ion catalysis which was observed by Bohnson and Robertson is considered by Barb et al. to be due to reaction (5’), as was the analogous effect of cupric salts on the ferrous ion catalysis. For conditions in which Scheme A applies reaction (4’) is in effect catalyzed by (5’). At high cupric ion concentrations (5’) will eliminate (3), since effectively all the radical HOz will react in (5’). In these conditions the enhancement reaches a limit as was observed by Bohnson and Robertson. However (1) now becomes the operative chain-terminating step, and hence the kinetics of Scheme B should apply (Eq. g). Unfortunately no data on the peroxide dependence is available, but analysis of the data (43,66) shows the rate to be proportional to [Fe+++]15 as required by (g). There is the same discrepancy in hydrogen ion dependence as was found in the simple ferric reaction. In view of the long period for which the ferric ion catalysis has been under investigation it is surprising that it was only recently that deviations from the kinetic equation of von Bertalan have been examined in any detail. It would appear that the system is even yet not completely understood, and clearly more data at low peroxide and high ferric ion concentrations are required as well as a more extensive investigation of the copper enhanced reaction. One other factor which would seem to merit attention is the part played, if any, by the complex FeH02++, which has been shown by Evans, George, and Uri (2) to be formed from Fe+++ and NO2-. Barb et al. have pointed out that the accepted initiation step (i) could very well be replaced by

+

Fe+++-j- H02-

FeH02++3 Fe++

+ HO,

62

J. H. BAXENDALE

without changing the kinetics in the experimental ccinditions usually employed. Observations in conditions where the ferric ion is present to a considerable extent as FeH02++would show if the complex is an intermediate in this way, and also whether it plays a part in any of the subsequent reactions. 6. Iron Salts in the Presence of Complexing Agents

Manchot and Lehmann (44) observed that in the presence of the bases cup'-dipyridyl (Dipy) or 1:10 phenanthroline (Phen), ferric ion can be transformed quantitatively into the ferrous complex Fe(Dipy) 3++ or Fe(Phen)a++ by hydrogen peroxide. It was concluded that the ferric ion

TIME, minutes

FIG. 3. Katalasestoss. 0.66N. HzO2, 5 X Fe++, Fe-I++, or Fe(Dipy),++, 2.6 X Dipy; 40 cc. of solution at 17°C. in pH 4.1 acetate bufEer. I. Fe++orFe+++ Hz02. 11. Dipy added to Fe+++ H202. 111. H Z 0 2added to added to Dipy Fe+++ Dipy. IV. Fe(Dipy),++ added t o H z 0 2 . V. Fe(Dip;y),++ added to HzOz Fe+++. VI. Fe(Dipy)~++ added to HzOa Dipy. From Kuhn and Wassermann, 65.

+

+

+

+

+

was reduced to ferrous by hydrogen peroxide and this was followed by addition of the base to give the complex ion. The reactions were subsequently studied by Kuhn and Wassermann (65) who discovered that during the formation of the complex ion in this way thiere is considerable catalytic decomposition of the hydrogen peroxide. For example in acetate at pH 4.1 with 0.33 llil peroxide, 5 X M ferric ion and 2.6 X M a,a'-dipyridyl the maximum amount of Fe(Dipy)3++ is produced in about ten minutes. During this time oxygen is evolved a t a rate which is about one hundred times greater than that given by the ferric salt alone and about twenty times greater than t8hatof the complex ion which is formed finally (Fig. 3). After this initial burst, which is usually referred to as the "Katalasestoss" the rate of evolution falls to that due to catalytic decomposition by the complex ion. Experiments with various combinations of the reactants (see Fig. 3) established that

DECOMPOSITION OF HYDROGEN PEROXIDE BY CATALYSTS

63

this increased catalytic activity only occurs while the complex ion is being formed in the presence of hydrogen peroxide. It is the same for ferrous and ferric salts since in these conditions the oxidation of ferrous to ferric is much faster than the rate at which it forms the complex with the base. The magnitude of the initial burst increases with increase of pH and peroxide concentration, and also with the amount of iron salt and dipyridyl provided the latter are kept in the constant ratio 1:5. However increase in dipyridyl to decrease this ratio gives a slightly decreased effect. Observations were also made on the rate of formation of the complex ion in various conditions. In the absence of peroxide with ferrous ion and a,a'-dipyridyl concentrations about 10-3 to M , complex formation is almost instantaneous at pH 4. If a ferrous or ferric salt is added to dipyridyl plus peroxide the reaction may occur over a period of minutes and goes at the same rate whether ferrous or ferric salts are used. The rate of formation increases with increase of peroxide concentration and pH, and below pH 3.3 no complex ion is formed with the above reagent concentrations. In comparable conditions the complex forms about three times more quickly with a,a'-dipyridyl than with phenanthroline, and twice as much oxygen is evolved in the initial burst. Because of this difference between the two bases Kuhn and Wassermann concluded that the reduction of ferric to ferrous cannot be the ratecontrolling step in the formation of the complex. They suggested that after ferrous ion has been formed in this way, it can either be oxidized by peroxide or react with the base to give the complex. If the rate of the latter reaction is faster with dipyridyl than with phenanthroline this will lead to the observed difference between them. Simon and Haufe (67) do not agree with this and hold the view that the production of ferrous ion by the reaction Fe+++

+ HOz---+ Fe++ + HOa

(9

is the slowest step. They attempt to explain the difference between the two bases in terms of Fe(Phen)3++being more easily dissociated than F e ( D i ~ y ) ~ + +Experiments . are described showing that F e ( D i ~ y ) ~ + + forms at smaller concentrations of ferrous ion than does Fe(Phen)3++in comparable conditions, and it is concluded that a t very small stationary ferrous ion concentrations, which must obtain in the ferric ion-peroxide mixtures, the dipyridyl complex will be formed before that of phenanthroline. This is contrary to more recent determinations (72,73) of the rates of formation and instability constants of these complexes. Fe(Phen) a++ forms about lo3 times more quiekly than Fe(Dipy),++, and its instability constant is less by a factor of lo4. It seems probable that Simon and Haufe have confused the color intensity of the complexes

64

J. H. BAXENDALE

with their stability. However, the explanation of Kuhn and Wassermann is also untenable in view of the greater speed of formation of Fe(Phen)3++. An alternative considered by the latter authors was that the reaction with hydrogen peroxide occurs not with free ferric ion but with a complex of ferric ion and the base. Brown complexes of unknown constitution were reported to be formed in solutions containing ferric ion and dipyridyl or phenanthroline by Blau (74), and subsequently Simon and Haufe (67) isolated a compound Fe(Dipy)Cla by mixing the reagents in ether. Whether this is the same compound as that present in aqueous solution is in some doubt since Barb, Baxendale, George, and llargrave (66), by measurements of the optical density of the brown complex formed in aqueous solutions, conclude that the formula is probably Fe(Dipy) ,+++. However, Simon and Haufe report that the complex does not react with M FeC13 and 2 X lop2 hydrogen peroxide. For instance with 2 X M dipyridyl a t p1-I 2, hydrogen peroxide takes three days t o produce a visible amount of F e ( D i ~ y ) ~ +but + ; if the peroxide and dipyridyl are mixed first and then the ferric salt is added, F e ( D i p ~ ) ~ +appears + in about a minute. It is suggested that in the first case all the ferric salt is present as the brown complex and this is the reason for the inactivity. I n the second experiment some reduction to ferrous ion by reaction (i) occurs before the formation of the brown complex. Barb et al. were unable t o confirm these observations and found that it made little difference whether the ferric ion and dipyridyl were mixed first or not. They also find that the rate of formation of the brown complex is too high to measure and they deduce from their calculated value of k, (see p. 60) that in any case the rate of formation of ferrous ion by reaction (i) is far too small t o account for the rate of Fe(Dipy),++ formation by this mechanism. From measurements of the amount of the brown complex present in various concentration conditions they conclude that in the Katalasestoss, where the rate of Fe(Dipy) 3++ formation is considerable, almost all the ferric ion must be bound as this complex. The catalytic activity of these systems was interpreted by IIaber and Weiss (4) in terms of their reaction mechanism for the ferric ion catalysis. They consider that the reaction of ferrous ion with the base t o form the complex ion lowers the stationary concentration of ferrous ion and hence leads to longer reaction chains by decreasing the rate of the termination reaction Fe++

+ HO. + Fe++++ OH-

However, as has been pointed out by Kolthoff and Medlalia (53), ferrous ion reacts much more quickly with peroxide than with the base under the conditions necessary for the Katalasestoss, and hence the steady state

DECOMPOSITION OF HYDROGEN PEROXIDE BY CATALYSTS

65

ferrous ion concentration cannot be appreciably decreased by the presence of the base. Further evidence against this view is contained in the observations of Barb et al. (66) using the base a,a’,a’’-tripyridyl (Tripy). The complex Fe(Tripy)2++is formed directly from ferrous ion and the base much more rapidly than Fe(Dipy)o++ forms in analogous conditions. Hence on the above argument tripyridyl should reduce the stationary ferrous ion even farther than dipyridyl and show an increased catalytic effect. I n fact in the same conditions tripyridyl gives only one-tenth as much decomposition as dipyridyl. Kuhn and Wassermann (65) attributed the high catalytic activity to the presence of reactive intermediates arising during the formation of the final complexes. They suggest that F e ( D i ~ y ) ~ + is+ formed from ferrous ion by successive additions of the base a s follows: Fe++ Fe(Dipy)++ Fe(Dipy),++

+ Dipy + Fe(Dipy)++ + Dipy e Fe(Dipy)z++ + Dipy e Fe(Dipy),++

The intermediate complexes may react with hydrogen peroxide as does ferrous ion but give a n increased reaction chain length and hence show increased activity. The existence of Fe(Dipy)++, Fe(Dipy)Z++, and the corresponding phenanthroline compounds has been confirmed (73,75) by kinetic analysis of the reactions of ferrous ion with the bases. Simon and Haufe (67), although opposing the view that ferric complexes are involved in the initiating step, do not reject the possibility that other complexes may be active, but suggest that the ferrous-ferric ion equilibrium which is set up in the presence of hydrogen peroxide may be present a s Hz02

(Fe+++complex)

z? (Fe++complex) H202

The position of the equilibrium may be so changed from the normal by the complexing agent as t o lead to a bigger overall peroxide decomposition. Barb et al. (66) have found experimental evidence to support the reactivity of intermediate complexes. They measured the amount of Fe(Dipy),++ formed from ferrous ion and dipyridyl in the presence of peroxide. Assuming that the only reactions occurring are

+ + +

+ +

Fe++ H,Oa-, Fe+++ HO. HO. Fe+++ Fe+++ OHFe++ 3Dipy -+Fe(Dipy)s++

+ OH-

it is possible to calculate the expected amount of Fe(Dipy)a++ from the known kinetics of ( 0 ) and (1) and those of (C), which have been obtained independently (73). Since the rate of (C) varies with [Dipy], (73), the

66

J. H. BAXENDALE

amount of F e ( D i ~ y ) ~formed ++ should also be a function of [Dipy13. In acid solutions this is the case, and the data are therefore consistent with the occurrence of reactions (o), (l),and (C). However, a t higher pH’s, such as obtain in the Katalasestoss conditions, less, F e ( D i ~ y ) ~ +is+ obtained than expected, and the dependence on the dip yridyl concentration is between [Dipy] and [Dipy12. These observations can be interpreted by supposing that Fe(Dipy)++ and Fe(Dipy) 2++, which are known to be formed in reaction (C), react more quickly with HzOz than free ferrous ion, and these oxidations compete with the further addition of dipyridyl to give F e ( D i ~ y ) ~ + + . Barb et al. have also examined the complex formaiion and oxygen evolution in the catalysis conditions. They find that the reaction is complicated by destructive oxidation of dipyridyl, which occurs during oxygen evolution. Thus in order to get complete transformation of ferric ion into Fe(Dipy),++ it is necessary to have a tenfold excess of the base over ferric ion, if the reaction is carried out with 0.33 M peroxide at pH 4. The amount of destruction decreases as the peroxide decreases and can be partially inhibited by acrylonitrile which i,s known to react with hydroxyl radicals. Acrylonitrile also slows up the F e ( D i ~ y ) ~ + + formation. It is concluded that hydroxyl radicals are .involved in chain reactions similar to those present in the ferric and ferrous ion catalyses. Their measurements of the oxygen evolved during the Katalasestoss show that at pH 4 with constant ferric ion and peroxide concentrations, the amount increases when the dipyridyl is increased, but reaches a maximum at [Dipy]/[Fe+++]about 5 :1. The initial increase is interpreted as being due to the increase in the amount of brown ferric complex present which gives an increased initiation rate by the reaction

+ k02--+ Fe(Dipy)z+++ HOz

Fe(Dipy)2+++

The intermediate Fe(Dipy) z++maythen add more base to give Fe(Dipy),++ or be oxidized back by peroxide to give a hydroxyl radical in a reaction analogous to (0) above. The radical can then initiate reaction chains similar to those in the simple ferrous and ferric ion systems. The maximum and final fall off in the oxygen evolution may result from the fact that when all the ferric ion is combined as the brown complex further increase in dipyridyl will not increase the initiation step. However, since more of the intermediates will be removed as Fe(I)ipy)a++the number of decomposition chains will decrease, and hence also the a,mount of oxygen evolved. In addition, the breaking of reaction chains by the oxidation reaction of hydroxyl with dipyridyl will act in this direction. It is unfortunate that the complexity of the system makes it difficult to establish the details of this catalysis with any certainty, but the evi-

DECOMPOSITION OF HYDROGEN PEROXIDE BY CATALYSTS

67

dence so far available indicates the existence of radical reactions and would seem to support the idea of the participation of ferrous and ferric complex ions. These complexes react more rapidly than the free ions and apparently lead to greater reaction chain lengths. The possible importance of this in the biological reactions where the enzymes concerned contain iron-porphyrin complexes, is obvious.

V. FERROCYANIDE AND FERRICYANIDE The ready reversibility of the ferrocyanide-ferricyanide redox system makes it a potential catalyst for the decomposition of hydrogen peroxide by the mechanism of compensating oxidation-reduction reactions. Moreover, the well-known facts that in acid solution ferrocyanide is oxidized to ferricyanide, whereas in alkaline solution the reverse reduction occurs, seem a good indication that at suitable pH's both reactions might occur to give catalytic decomposition. But from the investigations to date it would appear doubtful whether any such catalysis occurs to a measurable extent, and that what seems to be ready reactions of ferro- and ferricyanides are in fact those of partial hydrolysis products of these ions in which water molecules replace the cyanide ions in the coordination shell. Kistiakowsky (76) investigated the catalysis by ferrocyanide-ferricyanide mixtures and observed that the rate of decomposition of peroxide increases considerably when the solutions are exposed to light. The reaction was found to be first order in peroxide concentration and the first order constant a t 25°C. with 0.1 M HzOz, 0.014 N K4Fe(CN)a, 0.002 N K3Fe(CN), increased from 5 x 10-4 min.-1 in the dark to 59 X min.-1 when illuminated. Moreover, this high value was maintained when the light source was cut off. Ferrocyanide shows the same effect without ferricyanide, but the latter was not examined alone. Kistiakowsky concluded that a catalyst was produced by the illumination and suggested that this might be of a colloidal nature. Winther (77) and Weigert (78) also explained the phenomenon in terms of a catalyst but were not very explicit on its nature. Srikantan and Rao (79) reported that the ferrocyanide catalysis is not first order in peroxide as found by Kistiakowsky but much more complicated. No details of the experiments were given. They also studied the ferricyanide-peroxide reaction and found that in the dark there was a pronounced induction period after which the decomposition rate was first order in peroxide concentration. Illuminating for five minutes in bright sunlight removed this induction period but the subsequent first order rate is somewhat less than the original dark rate. They suggest that the induction period was the time taken to build up on

68

J. H. BAXENDALE

intermediate compound of the form [Fe(CN)6H202]”which is analogous t o the aquopentacyanoferrates formed by hydrolysis of Ferricyanide : [Fe(CN)a]”’

+ H 2 0 F? [Fe(CN)5H20]”+ CN‘

(1)

These aquopentacyanoferrates and t,he corresponding ferrites have been prepared by Hofmann (80) and they were shown by Briggs (81) and Iimori (82) t o be formed by the action of heat, acid, or light on the hexacyanides. Srikantan and Rao consider it possible for hydrogen peroxide to enter the coordination shell and for oxygen to be evolved by the reaction 2[Fe(CN)aHzOzl”4 2[Fe(CN)5H20]”

+

0 2

The removal of the induction period by illumination was attributed to colloidal Prussian blue which is known to be formed on prolonged illumiriation of ferricyanide. La1 (83) in agreement with Kistiakowsky found the ferrocyanide reaction t o be first order in peroxide concentration, although there was a certain irreproducibility in the results which was attributed to traces of impurities. He found that colloidal ferric hydroxide, a product of prolonged illumination of ferrocyanide, did not give the increased rate observed on illumination and therefore rejected this as the catalytically active material. EIxperiments using sodium aquopentetcyanoferrite gave results consistent with the assumption th at this is the compound responsible for the enhancement. Thus a solution of N / 6 hydroxen peroxide, M/64.2 K4Fe(CN)6, and 1.7 X M Na3Fe(CN)6H20 showed about the same rate of decomposition as the same solution without the aquopentacyanoferrite exposed t o sunlight for two minutes. This rate mas about twenty times Iarger than the ferrocyanide dark rate. An increase of threefold on the dark rate was produced by the presence of as little as 6.8 X lop6 M aquo salt. However, in the absence of ferrocyanide 1.7 X M aquo salt decomposed peroxide only a t about the same rate as the ferrocyanide alone in the dark, and judging from the color of the solution the salt is present mainly as the aquoperitacyanoferrate. Also of significance is the observation that if a ferrocyanide solution is irradiated and then kept in the dark for a while before adding peroxide, the catalytic activity is much less than if the irradiated solution is used immediately. Further, if the separately irradiated ferrocyanide solution is mixed immediately with the peroxide, the time of irradiation has little effect on the reaction rate, but if the ferrocyanide-peroxide mixture is irradiated, a n increase in the time of irradiation increases the decomposition rate. La1 explains these observations as follows. I n an irradiated ferrocyanide solution there exists the equilibrium

DECOMPOSITION O F HYDROGEN P E R O X I D E BY CATALYSTS

[Fe(CN)el””

+ H?O

light

+

[Fe(CN)5H10]”’ CN‘ dark

69 (1)

This is rapidly established from the left-hand side, and therefore the position is independent of the time of illumination. The reverse reaction occurs slowly in the dark and in the absence of hydrogen peroxide. Aquopentacyanoferrite formed in this way decomposes peroxide catalytically by a reaction which is not specified. At the same time it is oxidized by peroxide t o aquopentacyanoferrate 2[Fe(CN)6H20I‘”

+ H202-+

+

2[Fe(CN)5H20]” 20H’

as is shown by the experiment with the ferrite alone. The ferrate, however, is not as active catalytically as the ferrite, but there is a rapid regeneration of ferrite through the reaction [Fe(CN)#’”

+ [Fe(CN)6H20]”+ [Fe(CN)e]”’ + [Fe(CN)JLOl”’

which is shown to occur by independent experiments. The difference between the irradiation of ferrocyanide alone and when mixed with peroxide is that the ferrite produced in reaction (I) is oxidized by peroxide and hence more will be produced t o maintain the equilibrium. Subsequent work (84) confirmed these ideas. The addition of oxidized aquo salt to ferrocyanide-peroxide mixtures has the same effect as the same concentration of the reduced form, as it should. Further, addition of cyanide, nitrite, or nitrosobenzene to both illuminated ferrocyanide solutions and to ferrocyanide containing aquopentacyanoferrate, suppresses the catalytic decomposition. This effect is due to the reversal of the equilibrium (I) by a higher cyanide concentration, and to the removal of aquo salt by nitrite and nitrosobenzene to form the complex ions [Fe(CN)E,NO~]’”’and [Fe(CN),ONPh]”’, respectively, reactions which are known to occur readily. It is also suggested (85) that even the catalytic decomposition by ferrocyariide in the dark is due t o a small amount of aquopentacyanoferrite in purely thermal (as opposed to photochemical) equilibrium with ferrocyanide according to (I). La1 and Singhal (86) have examined the catalysis by ferricyanide. Using much smaller concentrations (M/300) than those of Srikantan and Rao they find that the dark rate is about one hundred times slower than with the same concentration of ferrocyanide, and that mixtures of the two salts give about the same rates as ferrocyanide alone. Initial irradiation of ferricyanide alone gives an increase of about tenfold over the ferricyanide dark rate, but if new ferrocyanide is added to this solution the rate rises t o about ten times that of the ferrocyanide dark rate. These observations are explained in terms of the reactions given above.

70

J. H. BAXENDALE

Thus ferricyanide hydrolyzes photochemically to aquopentacyanoferrate which is a slightly better catalyst for decomposition than ferricyanide. On addition of ferrocyanide the ferrate is reduced to the ferrite which is more active. Some confirmation of these reactions is given by the observations that addition of the ferrate to ferricyanide increases the dark rate, and also that cyanide, nitrate, and nitrosobenzene inhibit this as well as the catalysis induced by irradiation of ferricyanide. This work of La1 clearly establishes that the ferro and ferricyanide catalyses in light arise not directly from these ions but from the hydrolysis products. The same is true of catalysis by nitroprusside observed by Qureshi (87) and also studied by La1 (88). It also seems probable that the aquo salts are the active species in the dark reactionc; and support for this is provided by some recent experiments by the author and Mr. E. Walker. While measuring the rate of oxidation of ferrocyanide by hydrogen peroxide it was observed that the same stock solution of ferrocyanide gave gradually increasing rates and also that the rate was very light sensitive. Suspecting the hydrolysis to be responsible for this the effect of addition of potassium cyanide was examined. In a phosphate buffer pH 6.1 with 4.3 X M H 2 0 2 and 1.3 X M K4Fe(CN)a freshly made in the dark, the half life of the oxidation (also in the dark) was about thirty minutes at 25°C. In the same conditions with the solution also 0.01 M in KCN there was no reaction in several hours. Increasing the peroxide to 3 M ,the half life was still about two hours. The reduction of ferricyanide by hydrogen peroxide was analogous. With 1.3 X M K3Fe(CN)aand an equivalent amount of hydrogen peroxide in 0.01 N KOH the reduction was complete in a few seconds in the dark. Potassium cyanide decreases the rate and thle reaction is first order in peroxide in these conditions. A t 0.01 M KCN the half life is about six minutes and this does not change when the cyanide is increased to 0.05 M . These observations are readily explained in terms of reactions of the aquo salts and it would appear that the usually accepted rapid oxidation of ferrocyanide by hydrogen peroxide is misleading. The fact that this is a very slow reaction means that although the ferricyanide reduction is fairly rapid there is no possibility of appreciable catalytic decomposition of the peroxide by the compensating reactions mechanism. There is not sufficient data available as yet to say what reactions are involved in the rapid catalytic decomposition by the aquo salts, and clearly a separate study in the absence of ferro- and ferricyanides is required. However Baudisch (89) and Petow and Kosterlitz (90) have shown that like ferrous and ferric ions they can behave peroxidatically with suitable organic substrates. Also Baxendale et aE. (45) have found

DECOMPOSITION OF BYDROGEN PEROXIDE BY CATALYSTS

ii

that ferrocyanide and hydrogen peroxide can initiate polymerization-a reaction which probably goes via the aquo salt since, as pointed out above, ferrocyanide itself does not react readily with peroxide. Moreover the oxidation of aquopentacyanoferrite to ferrate by hydrogen peroxide in acid solution is reversed on making the solution alkaline. It seems probable therefore that reactions involving free hydroxyl radicals play some part in the catalysis and the reactions may well be very similar to those concerned in the ferrous and ferric ion systems. It is of some interest to inquire why there is this big difference in catalytic activity between the ferro-ferricyanide system and that of the corresponding aquo salts. There is very little difference in the redox potentials of the systems for Davidson (91) has obtained 0.198 v. and 0.209 v. respectively against the N calomel electrode. Of course it is not possible in general to compare overall free-energy changes with reaction rates but since presumably both oxidation and reduction reactions are involved in the catalysis, some sort of similarity might have been expected here. For a more rigorous analysis, a knowledge of the individual heat and entropy changes for the two redox systems would be more valuable but these unfortunately are not available. It is of course possible that there is something specific about the aquo salt system and there are in fact indications to this effect. Thus Davidson (91) and Michaelis (92) find that the redox titration curves of these salts do not follow the course of a simple one electron redox system as might be expected. Davidson explains his results in terms of an association complex having the constitution 2[Fe(CN)gH20]”’.3[Fe(CN) 6H20]”1 while Michaelis found that the curves corresponded to a four electron system having two pairs of overlapping single electron steps. He suggested that the ions are polymerized through hydrogen bonds to give 4[Fe(CN) gH20]””. There is also the possibility, as suggested by Srikantan and Rao, that hydrogen peroxide may be coordinated instead of water either in the single ions or in the association complexes. This is an interesting speculation for there would be a close analogy between such compounds and the substrateenzyme complexes of biochemical reactions, including the decomposition of peroxide by catalase.

VI. COPPERCOMPOUNDS

I. Cupric Ion Kiss and Lederer (93) have made the only significant quantitative study of the decomposition of hydrogen peroxide by cupric ions alone and this is not very extensive. The reaction is much slower than the ferric ion catalysis at the same concentrations and acidities. For three

72

J. H. HAXENDALE

acid concentrations around 5 the reaction rate is given by

x

10k4 N and two cupric ion concentrations

- d [ H 2 0 2 ] / d L = k[C~++].[H202]S~/[~1+]

in which lc has the value 1.5 X lo-" mole-1 min.--l a t 40°C. There are deviations from this expression in more alkaline solutions, and a t about N acid the reaction becomes first order in peroxide. The order in peroxide decreases further in the absence of added acid. Euler and Jansson (94) made a few measurements on the copper catalysis in solutions made alkaline with caustic soda, but it is doubtful whether cupric ion can be considered as a reacting entity in these conditions. It seems probable that the decomposition occurs by a mechanism involving radicals analogous to the ferric ion catalysis, for in acid solutions cupric ion with hydrogen peroxide oxidizes fatty acids such as palmitic, down to the lower homologues, aldehydes and carbon dioxide (95). Another analogy with the iron system is that cuprous salts and peroxide cause rapid oxidation of organic compounds (96). From the limited amount of information available, any conclusioiis on the detailed mechanism of the reaction would be mere conjecture, but it would seem safe to conclude that the redox system Cu++-Cuf is involved possibly in a manner similar to the ferric system. It is interesting to note that the above kinetic expression of Kiss and Lederer is almost exactly the same form as that obtained by Barb el al. (62) for the ferric ion catalysis when the ratio of ferric ion to peroxide concentration is very high. It was suggested that in these conditions the reaction

+ 02-

Fe+++

4

Fe++

+ Op

(4')

eliminates all other reactions of the radical 110, and thus changes the chain termination mechanism. This is likely to happen a t much lower concentrations of cupric ion, since the corresponding lo2- reaction with cupric ion has been shown to be about twenty times a s fast as the ferric ion reaction. It seems possible therefore that the reaction mechanisms are similar in the two cases. As has been mentioned above the ferric ion reaction in these conditions still requires elucidation. An interesting enhancement of the cupric ion catalysis by manganous ion has been reported by Bobtelsky (97). It was found that in neutral solution with 0.05 N manganous sulfate the rate of decomposition increases with increase of copper sulfate but passes through a maximum at about 0.05 N . Here the rate is about ten times th a t of copper sulfate alone. Zinc and cadmium sulfates are also reported to enhance the Cuff-Mn++ catalysis.

DECOMPOSITION O F HYDROGEN PEROXIDE B Y CATALYSTS

73

2. Copper Complexes

Catalysis by Cu(NH3)J304 has been examined by Nikolev (98) who found a maximum rate a t pH 8.5 to 9.0. The reaction is strictly first order in peroxide up t o 0.25 M peroxide but decreases to zero order a t 0.5 M . Addition of ammonium hydroxide a t first increases the rate but ultimately depresses it, an effect also noted by Bobtelsky and Kirson (99). Unfortunately the experiments are not very extensive, and it is impossible to separate the effect of ammonia in complexing from the effect of the accompanying pH increase. Information on the relative activities of the complexes with various amounts of coordinated ammonia would be interesting. The only indication from this work is that the depressing effect of high ammonia concentrations supports the authors’ view that the hexammine probably formed in these conditions is less active than the tetrammine. One other point of interest is that it seems probable that radicals are formed in the reaction, possibly in conjunction with a cupric ammine-cuprous ammine redox system, since organic substances can be oxidized by the cuprammonium-peroxide mixtures. Nikolev also measured the relative catalytic properties of a series of complexes of cupric ion with aliphatic amines. These are less effective than the ammonia complexes and among themselves show a variation of about a hundred-fold in their efficiencies. However no account was taken of the relative stability of the complexes under the experimental conditions, and since also oxidation of the organic amine compounds probably occurs it is doubtful whether the results give a true picture of the relative reactivities of the compounds. VII. PERMANGANATE

I n strong acid solution it is well known that permanganates react stoichiometrically with hydrogen peroxide according to the equation :

+

~ H z O Z 2MnOd-

+ 6H+ = 2Mn++ + 8HzO + 502

In solutions initially alkaline or slightly acid manganese dioxide is produced, and in certain conditions more oxygen is liberated than corresponds to the above reaction or to the following one: 3H202

+ 2Mn04- + 2H+ = 2Mn02 + 302 + 4Hz0

Probably because of the high velocity of the reaction, very little kinetic work has been done on it, and it is still doubtful whether the catalytic decomposition of peroxide in these conditions arises solely from the manganese dioxide which is formed, or whether in addition other compounds of manganese participate in a homogeneous catalytic reaction.

74

J. H. BAXENDALE

The catalysis by manganese dioxide either in bulk or as a colloidal solution is well known and has been the subject of much work. The only experimental study of the permanganate reaction which gives some indication that a homogeneous catalysis might be present, is a recent one by Fouinat (100). The velocity of and extent to which 0.056 M hydrogen M permanganate was measured peroxide is decomposed by 1.15 X at different acidities and alkalinities. For concentrations of caustic soda greater than 0.012 N the peroxide is completely decomposed, the time for half decomposition is about twenty-five minutes and is independent of the alkali concentration. With decreasing alkalinity the velocity N of decomposition increases and reaches a maximum a t 8.3 X caustic soda when the half life of the peroxide is about two minutes. At the end of these experiments manganese dioxide is formed either as a precipitate or as a colloidal solution. At a slightly greater acidity than 1.30 X which gives a final pH of 6.7 to 7.0, there is no visual evidence of manganese dioxide but the hydrogen peroxide is still completely decomposed and has a half life of about twelve minutes. At this point the reaction becomes very sensitive to small changes in acidity. Thus at M , 80% and 30% respectively of the M and 1.38 X 1.34 X peroxide are decomposed. The absence of manganese dioxide in these more acid solutions might indicate a homogeneous catalysis, although it is of course possible that colloidal material formed initially is dissolved during subsequent reaction. Also the maximum in the rate could be interpreted either in terms of a homogeneous reaction increasing in rate as the pH increases, until at the higher pH’s the active manganese intermediates are removed as MnOz; or alternatively in terms of surface catalysis by MnOz, the surface area being decreased by coagulation at the higher pH’s. It would assist, in the understanding of any catalytic reaction in this system if the details of the stoichiometric reaction in ,acid solution could be elucidated. This is also extremely complex as has been shown by the investigations of Riesenfeld (101,103) and Bailey and Taylor (102). An idea of some of the oxidation states of manganese which might be important in the reactions is given by the kinetic studies on the oxidation of organic compounds by permanganate (104,105). In these, the entities, Mn04”, Mn03’, Mn4+, Mn3+ have been invoked to explain the kinetics. Mn4+ can be regarded as the ion which on hydration leads to MnOz, while MnO4” and Mn3+ can be prepared and are relatively stable. The latter is reduced to Mn2+ by hydrogen peroxide and the reaction is too fast to measure even in high dilution and high acid concentrations (106). The existence of so many intermediate redox states is of course highly favorable to the existence of compensating reactions in which hydrogen

DECOMPOSITION OF HYDROGEN PEROXIDE BY CATALYSTS

75

peroxide behaves first as a reducing agent and then as an oxidizing agent, and it seems very likely that a homogeneous catalytic decomposition occurs in the slightly acid conditions obtaining in Fouinat’s experiments discussed above. However, the experimental data are too meager to give any details of the steps involved, and although plausible reactions can be suggested, as for example those of Abel (107), these must at present be mainly speculative.

VIII. CHROMATE

*

When hydrogen peroxide is added to a potassium chromate solution no visible change occurs but oxygen is slowly evolved and ultimately the hydrogen peroxide is all decomposed leaving the chromate unaffected. On the other hand, if the chromate solution is made acid with mineral acid there is an immediate formation of the well-known blue perchromic acid which decomposes fairly quickly with the evolution of oxygen, leaving the chromium reduced completely to Cr+++. Subsequent decomposition of the hydrogen peroxide is very slow. Spitalsky (108) has described the changes which occur between these extremes. I n slightly acid solutions or alternatively using dichromate, the solution quickly goes violet (sometimes described as brown-red) on the addition of peroxide, and a steady evolution of oxygen occurs more quickly than with neutral chromate. As the reaction proceeds, the violet color gradually disappears and when all the peroxide has decomposed the original dichromate color reappears. With slightly more acid or using chromic anhydride, addition of hydrogen peroxide immediately gives the blue compound which disappears in a few minutes to give place to the violet compound. This remains unchanged for some time and catalytic decomposition of the peroxide proceeds until suddenly the solution darkens and then quickly reverts to the color of the CrOs solution. At this stage all the peroxide has been decomposed. If now more peroxide is added, the violet and not the blue compound is formed initially, and the reaction proceeds as before. The reason for this is that some of the CrOa is reduced to Cr+++, and this reaction increases the pH to a value at which the blue compound does not form. Spitalsky has investigated the kinetics of peroxide decomposition in these various conditions at 25°C. With only CrOq” or only Cr20;’ the reaction is first order both in peroxide and chromium concentrations, the dichromate reaction being about three hundred times as fast as that of chromate. Additions of Cr04” to Crz07” solutions decrease the rate constant at first but with more Cr04” it passes through a minimum and then increases steadily. This behavior is attributed by Riesenfeld (109)

76

J. H. RAXENDALE

to the equilibrium Cr&”

+ I1?0

2Cr04”

+ 211+

and he found that taking this into account the observed rates are given by the sum of individual contributions by G O 4 ” and Cr207”. The kinetics in more acid solution or using CrOa are much more complicated. I n the first place, the reaction is not a pure catalysis in that some or all of the Cr03 is reduced to Cr+++. Spitalsky (110) has measured the extent of this reduction in various conditions. I n solutions containing only CrOa and hydrogen perolride the final amount of Cr+++ produced increases as the peroxide is increased up to a concentration of about 0.02 M . Beyond this, even up to 8.5 M , the amount of Cr+++ stays constant a t 28.3 yo of the total chromium. This maximum amount increases with the acidity of the solution and a t high acid concentrations the reduction is complete. These observations sug:gest a dynamic equilibrium between the two oxidation states of chroinium during the catalysis reaction, and this was shown by Spitalsky to be the case. Thus a solution containing 3 i Yo of the chromium initially as Cr+++ was found t o have the “equilibrium” 28.3% when it had decomposed a large amount of hydrogen peroxide. Spitalsky considered that the production of Cr+++ occurred through decomposition of the blue compound which is formed initially but which disappears relatively quickly compared with the time for the complete decomposition of h,ydrogen peroxide. The catalytic decomposition rates with CrOa solutions vary in a rather unusual way with the peroxide concentration (11I). They are approximately first order i n the total chromium and almost independent of the peroxide concentration until the decomposition is nearly complete. Then a t about 0.02 M peroxide, there is a sharp increase in rate and a subsequent rapid decrease until all the peroxide has disappeared (Fig. 4). This type of variation occurred for a range of CrOJ concentrations and was independent of the initial peroxide concentration.. The transition from the first order behavior of CrzO,” t o that of C r 0 3 was followed by using Crz07)’with increasing amounts of nitric acid. Curves of the type shown in Fig. 4 were obtained. It can be seen that with increasing acid there is a decrease in the initial rates and a maximum appears a t a certain peroxide concentration. At very high acids (curves g and f ) the decomposition is very rapid, but does not go to complet,ion. This region corresponds t.o complete reduction to Cr+++and the decomposition occurs during the disappearance of the blue compound. Further evidence that Cr+++ is in dynamic equilibrium with CrV1 during cat’alysis is that the rate is independent of the extent of oxidation of the chromium. Thus a solution containing 38% as Cr+++ and the rest

DECOMPOSITION OF HYDROGEN PEROXIDE BY CATALYSTS

77

CrOI follows exactly the same course as one with 100% CrOs (110). Spitalsky suggests t h a t the catalysis arises as a result of the compensating reactions

+

+

+ l l H & + 502 + 7H20 + 402 + 8H+

CrzO7” 7H#& 8H+ + 2Cr+++ 2Cr+++ llHzOz+ CrzO?”

+

but as is obvious from the appearance of the solutions, both of these reactions probably proceed through intermediates. An attempt was made to detect these at various stages in the reaction by measuring the

180

is0

120

90

60

30

0

[H2OZ] x i03

FIG. 4. Chromate catalysis a t different acidities. [K2Cr207]= 0.00192 M in each experiment. a b c d e f g 1O3[HNO,3] 0 0.405 1.01 2.51 5.03 7.61 15.2 From Spitalsky, 111.

conductivity of the reaction mixture (11‘2). It was observed t h a t in acid Cr207” solutions pretreated with peroxide so t h a t the further reduction t o Cr+++ is absent, there is an immediate decrease in conductivity which is attributed t o the removal of hydrogen ions. As the reaction proceeds, the conductivity remains steady until in the region of the velocity maximum there is a sharp increase. By measuring this initial change, and from it calculating the consumption of acid a t various Cr20,” and peroxide concentrations, it was concluded t h a t along the steady portion of the velocity curve a compound KH6Cr4018’’is in equilibrium according t o 2Cr201“

+ 2H202+ K + + H+ & KH6CrrOls”

78

J. H. BAXENDALE

The catalysis was attributed to the formation of this substance and its dissociation with the evolution of oxygen. When the peroxide concentration becomes small the above equilibrium reverses liberating H+ which causes the formation of other more active intermediates giving a maximum in the velocity curve. The same general picture differing only in the details was given by Kobosev and Galbreich (113). They observed that the Crz07” catalysis found by Spitalsky to be first order a t 25”C., changes t o second order a t 56°C. At 0” and 56°C. the velocity is given more exactly by the equation:

+

V = Kk[Cr~Oj’l[H,0~]~/(1 K[Hz0zIz)

Variation of K with temperature changes the reaction order and fortuitously leads to first order a t 25°C. The above kinetic (equation is interpreted by the following reactions: CrzOr”

+ 2H202 k

K

Cr204’

Cr20/’-+ CrnOj’

+

+ 2Hz0 0 2

The experimental data lead t o the evaluation of k and K which are given by k = 2.1 X 1Olo exp (-15,78O/RT)

and K

= exp

(2725/RT)

I n more acid solutions they consider HCrz09’t o be the reactive intermediate and the constants for the corresponding reactions H+

+ Crz07” + 2 H 2 0 2Ki F? HCrOo’ + H 2 0 ki HCrZOs’ -+ CrzO7” + H+ + OZ

are given as kl =

lOI7

exp (-23,00O/RT)

and K 1 = 5 X lov7exp (17,50O/RT)

but the origin of and method of calculation from the data are somewhat obscure. It appears that the two complexes Cr2O9’’ and IICrz09’ are t o be identified with the violet and blue intermediates respectively, since these are visible in the less and the more acid solutions. Bobtelsky et al. (114) also consider these to be kinetically important and they have found that the decomposition velocity follows the concentration of these compounds. However, these workers have used alcoholic solutions: and i t is open t o question whether their results can be applied in detail t o aqueous solutions in view of the possible interference by oxidation of the alcohol. They consider the mechanism of the catalysis t o be CrVI

+ HZOZ-+

Blue or violet intermediate -+ Decomposition products

+

0 2 -i

Regeneration of CrVI

DECOMPOSITION O F HYDROGEN PEROXIDE B Y CATALYSTS

79

It is evident that the details of the reactions occurring in these catalytic decompositions are still somewhat obscure. It is clear that the blue and violet perchromate intermediates play an important part, but in what way cannot be surmised with any certainty until their constitutions have been established. Direct analytical investigations on these have given conflicting results (115~116,117). Of the formulae used in the interpretation of the kinetics Spitalsky’s KH6Cr4OI8’appears too complex to be probable and the evidence supporting it is not very strong. Kobosev’s HCr209’and Cr209” are based on erroneous deductions from his observed kinetics. He ignored the fact that at the dichromate concentrations (- 10-3 M ) used by him (and also by most other workers), the chromium is present mainly as HCr04- or CrOl’, depending on the acidity (for HCr04’ pK, = 7). This follows because the equilibrium constant for Crz07”

+ HzO & 2HCrOa’

a t 25°C. is (118) K = [HCr04’]2/[Crz07’’]= 0.023. Hence with an initial chromate concentration of M , 90% of the chromium is in the form HCr04’. At these dilutions the actual dichromate concentration in solution will vary more nearly as the square of the total concentration of chromium present, Since Kobosev and the other investigators have shown the reaction rates in all conditions to be proportional to the total chromium concentration, it would appear from these considerations that the kinetics are determined by the concentration of HCr04’ or of Cr04” and not that of Cr207”. Thus HCrOa’ and H2CrOs would be consistent with Kobosev’s kinetic data, but this is by no means conclusive evidence. It is not clear whether catalysis can occur purely by formation and decomposition of these perchromate intermediates as suggested by Kobosev or whether a t all acidities there is reduction of CrV1to Cr”’ and subsequent regeneration of CrV1 along the lines of the compensating reactions mechanism. It appears from Spitalsky’s work described above that in the more acid solutions Cr+++is in dynamic equilibrium with CrV1. Whether this is also the case in less acid solutions where no Cr+++ is produced finally cannot be concluded from the present evidence, but such a possibility is clearly present since Cr+++ is readily oxidized to CrO4” by hydrogen peroxide in neutral solution (111). If this is the case then the participation of the CrVand CrIV states of oxidation found to be present in other reactions of chromates (119) is possible. These active oxidation states are considered to be responsible for the induced oxidations which occur in certain chromate reactions, for example the oxidation of manganous ion during the chromic acid-arsenious acid reaction.

80

J. H . BAXENDALE

Such oxidation reactions may be responsible in part for the enhancement of the chromate catalysis which is produced by Mn++, Co++, Cu++, Ce+++, Ni++ (120,121). Alternatively this promotion may arise from the reaction of these ions with perchromate compoundsy and it is possible that chain reactions may occur similar t o those in the ferric ion catalysis with the perchromate replacing the peroxide. Uri has suggested such a scheme for promotion in the molybdate and tungstste catalyses (see Sec. IX.3). However the data are too fragmentary for any definite conclusions to be drawn.

IX. MOLYBDATE AND TUNGSTATE 1. Molybdate

The addition of’ alkali molybdates to hydrogen peiroxide in aqueous solution can give two permolybdates which are the salts of the acids HzMoOs (yellow) and HPMoOs (red). A third compound HzMo05 can be obtained under certain conditions (115). According to Jahr (122) HMoOs’ is formed with only a slight excess of hydrogen peroxide or in acid solution, and HMoOs’ is given by a large excess of peroxide in neutral solution. He found no evidence for H2M005in these conditions. When the compounds are formed in neutral solution oxygen is evolved and ultimately they disappear when all the hydrogen peroxide has decomposed t o leave the molybdate apparently unchanged. The kinetics of the catalytic decomposition have been investigated by SpitaIsky and Funrk (123). At high acid concentrations where a yellow permolybdate (whether HMo06’ or H P M o 0 6has iiot bt3en established) is formed, and also a t high alkali concentrations where there is no visible evidence of any permolybdate the rate of decomposition is extremely slow. In slightly acid (- 0.005 N ) and slightly alkaline (- 0.01 N ) solutions the decomposition is zero order with respect, t o peroxide concentration until the latter is small (- 0.01 M ) when the rate falls off with further decrease of peroxide. With 1.4 X M molybdate the zero mole/liter/min. to about 10-3 order rate increases from about 5 X mole/liter/min. in going from 0.0055 N acid to 0.0096 N alkali, but increases of alkali to 0.075 N do not affect the rate further. In neutral solution, however, two odd features appear. First, the rate can only be reproduced with the same catalyst provided it is first pretreated with hydrogen peroxide and used fairly soon after this. Secondly, with this pretreated catalyst the reaction rate shows a maximum at about 0.01 M peroxide which is not present in acid or alkaline solutions. This is similar t o the behavior of chromate in slightly acid solutions. Unlike the chromate system there is no reason to believe that any

DECOMPOSITION O F HYDROGEN PEXOXIDE B Y CATALYSTS

81

intermediates are concerned in these reactions other than the permolybdates and i t is very probable that the formation and decomposition of one or more of these is responsible for the catalytic decomposition. Spitalsky suggests that three are present. The yellow compound formed in acid solution is considered inactive. Two others formed in neutral and alkaline solution are thought to be catalytically active in these conditions and lead to the maximum observed in neutral solution. This could be the case if the equilibrium constants for the formation of these compounds are such that the more easily decomposed peracid is formed preferentially a t lower peroxide concentrations, but Jahr (122) could not detect any compounds other than HMo06’ and HMo08”. It seems possible that these are the two active substances and the inactive yellow compound formed in acid solution is undissociated HzMoOs. The zero order rates observed by Spitalsky and Funck are readily explained in terms of a single species, e.g., HMo06’forming and decomposing as follows:

+ 2Hz02 F?K KMoOB’ + H2O + OH’ k HM006’4 M004” + + H+

M004” ,

0 2

If the equilibrium is attained rapidly the reactions lead to

+

-d[HnOi]/dt = kK[M00a”][HzOz]~/(l K[HzOzI2)

Hence a t high peroxide concentration when K[H,O,I2 >> 1 the rate is zero order. The fall off in rate a t the lower peroxide concentrations also follows from this equation. Analogous equations can be written for HMoOs’. Kobosev and Sokolev (124) have isolated the permolybdates Mo06” and Moos” and investigated the kinetics of their decomposition (125, 126). That of Moos” is first order until the concentration is small, when deviations occur which are accounted for by the increased dissociation of the compound into peroxide and molybdate. The decomposition of MOO;’ is more complex and it is suggested that it occurs through the intermediate formation of MOO;’. Both MOOS” and Moo6’’ have an activation energy of 17.4 kcal. for decomposition, and the rate of Moog” is about 4.5 times that of Moos“. Equilibrium constants K have also been obtained for the two compounds from the kinetics but the method of derivation is open to question. Until accurate data on these equilibrium constants and their variation with acidity are available the relative importance of two permolybdates in the catalysis cannot be assessed. However, it seems probable that the anomalous behavior in neutral solution and the variation of rate with acid and alkali concentrations might be understood in terms of

82

J. H. BAXENDALE

these equilibria. Since this is one of the few systems where there is apparently nothing more involved in the catalysis than the combination of substrate and catalyst and the subsequent decomposition of this complex the clarification of the kinetics would be of some interest.

2. Tungstate Tungstates are very similar to molybdates in their catalytic decomposition of hydrogen peroxide. They form analogous peroxy compounds HWOe’ and WOE’’ which are more stable than most peracids but decompose slowly with the evolution of oxygen (115). It seems probable that the catalysis occurs via these compounds. Bogdanov (127) found that with acid more concentrated than 0.0067 N the catalysis is insignificant but that below 0.00168 N it is fast. The variation of rate with peroxide concentration in these acid solutions has exactly the same form as that for molybdate. To 0.22 N peroxide the rate is zero order but becomes first order and ultimately second order as the concentration decreases. The velocity in the first order region is proportional to the tungstate concentration and varies inversely as the acid concentration. Bogdanov assumes the existence of two intermediates formed from WO,”, OH’ and one or two molecules of hydrogen peroxide, which in a scheme similar to that given above for molybdates, will account for the kinetics. However there is no other evidence for a compound with one molecule of hydrogen peroxide and it seems probable that both the first and second order kinetics can be explained by one intermediate since the equation given above could simulate first order a t certain peroxide concentrations. Comparing these observations with those of Spitdsky (123) on the molybdate catalysis, it is found that in the zero order region the latter reaction is about twice as fast, which is in keeping with the greater stability of the pertungstates. Another quantitative difference is that the deviations from zero order for tungstate occur at much higher peroxide concentrations, which seems to indicate that pertungstates are more easily dissociated into tungstate and peroxide than are permolybdates. Here again, however, data on these equilibria are required for a quantitative interpretation of the kinetics. 3. Promotion of Molybdate and Tungstate Catalyses

I n acid solution where the catalyses by molybdate and tungstate alone are slow, Fe+++, CU++, Mn++, Co++, Ni++, and I’ cause a pronounced increase in rate. Konovalova (128) found similar kinetics for the ferric ion promoted tungstate reaction as for the tungstate alone, and also showed that ferric ion increases the rate of decomposition of the pertungstate WOE”. Similar observations were made by Bogdanov

DECOMPOSITION O F HYDROGEN PEROXIDE B Y CATALYSTS

83

(129) and Bogdanov and Petin (130) on the molybdate system. The latter also found that molybdate and tungstate exert a mutual enhancing effect, the rate with both present being much greater than the sum of the two individual rates. Bobtelsky (131) observed a weak activity of Mn++, Co++, and Ni++ on the tungstate system but found Cu++ and I’ to be very effective. Uri (132) also found Cu++ to be the most active of the transition element ions in the tungstate and molybdate catalyses and also observed that the chloro and citrate complexes of copper were very effective. The kinetic measurements have not been detailed enough for the mechanism of this promotion to be deduced with any certainty, but two views have been proposed. The Russian workers suggest an intermediate complex of metal ion, peroxide and tungstate or molybdate which decomposes more quickly than the simple peracid. On the other hand Uri considers that the peracids behave like hydrogen peroxide in their reactions with the metal ions, so that a chain reaction occurs along the lines proposed by Haber and Weiss for hydrogen peroxide :

+

+ + + +

CU++ HWOS‘ @ CU+ HWOe HWOe Hz0a-t 02 HzO HWO6 HWOa Ha02 --* HzO HWOe

+ +

It is pointed out that since HzWQs and the other peracids are much stronger acids than hydrogen peroxide the higher concentration of HW06’ compared with HOz’ will favor a faster reaction of HzW06 than H 2 0 2 with the metal ion. REFERENCES 1. 2. 3. 4.

5. 6. 7.

8. 9. 10. 11. 12. 13. 14. 15. 16.

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Structure and Sintering Properties of Cracking Catalysts and Related Materials HERMAN E. RIES, JR.* The Research and Development Department, Sinclair ReJining Company, Harvey, Illinois CONTENTS

Page ........................ , . . . . . 88 I. Introduction. . . , . . , . , . . . . . . . . tion . . . . . . , . . . . . . . . . . . . . . . . . . . . . 90 11. Experimental Procedures and In 1. Apparatus and Technique for Gas Adsorption Measurements. . . . . . . . . . 90 2. Sintering Treatments.. . . . . . , , . . . . . . , . . . , . , . , . . . . , . . , . . . . . . . . . . . , . 92 3. Helium-Mercury Pore Volumes.. . , . . , . . . . . . . , . . . . . . . . . . , . . . . . . . . , 92 4. Materials . . . . . . . . , . , . . . . . . . . . . . . , . . . . . , . . . . . , , . . . . , . . . . . . , . . . . . . . 94 a. Cracking Catalysts.. . . . . , . . . . . . . . , , , . , . . . , . . . . . , , . . . . . . . . . . . . . 94 b. Silicas and Related Materials.. . . . . . . . . . , . . . . . . . . . . . . . . . . , . . . . . . . 94 5. Interpretation of Adsorption Isotherm D a t a . . . . . . . . . . . . , . . . . . . . . . . . . 95 a. Surface Area . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 95 b. Pore Volume.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 97 . . . . , . . . . . . . . 98 c. Pore Radius . . . . . . _ . . . . . . . . . . . . . . . d. Hysteresis Loops and Pore Structure.. . , , . . , . . . , , . . . . . . . . . . . . . , . . 99 111. Structure and Sintering Properties of Representative Cracking Catalysts. . . 99 1. Silica-Magnesia Catalysts. . . . . . . . . . , . . . . . . . , . . . . , . . . . . . . . . . . . . . . . . 103 2. Silica-Alumina Catalysts, . , . . . . . . . . , . . . , , , , . . . . . . . , , . . , , . . , , . . , , . . 104 a. TCC Beads 105 b. Aerocat Microspheres., . . . . . . . . . . , , . , . . . . . . . , , , , . . . . . . . , . . . . , , . 114 c. Diakel.. . . . . . . . . . , , . . . . . . . . . . . . . . . . . . . . . . . . , . , . , . . , , . , . . . . . . . 118 3. Clay Type Catalysts ....................................... 119 IV. Structure and Sintering Properties of Various Forms of Silica and Related ................................................... 124 ,

I

,

,

2. Silica Aerogel.. . . . . . . . . . , . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . a. Aerogel “ S ” . . . . . . . . . . . . . . . . . . . . . . . . , , . . . . . . . . . . , b. Bead Aerogel.. . . . . , . . . . . . . . . , . . . , , . . , . . . . . . , , , , . , . . . . . . . . . . . . . . . . . . 131 c. Santocel “ C ” . . . . . . . . . . . . . . . . . . . . . . . . . . . d. Wetted Aerogel . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 131 3. Nonporous Silica . . . . . . . . . . . . . . . . . . . . . . . . . . . _ . _ _ . . . . . . . . _ .132 ._ 4. Adsorption per Unit Surface for Three Forms of Silica., , , , , , . . . . . . . . . 133 5. Sintering Curves for Silicas.. . . . . . . . . . . . . . , . , . . . , , , . , , . . . , . . . . . . . . . 135 . 137 6. Alumina . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

* Present address: The Research Laboratories of the Standard Oil Company of Indiana, Whiting, Indiana. 87

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HERMAN E. RIES, JR.

Page

7. Titania . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8. Diatomaceous Ehrth.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . V. Summary and Conclusion.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

139

140 146 147

I. INTRODUCTION Within the past few years the large scale commercial use of catalysts in the petroleum industry has undergone a remarkable expansion. The most important of these catalysts are those employed in cracking processes. Principal interest has been focused on thiree types: silicaalumina, silica-magnesia, and clay type catalysts. The manufacture of catalysts has in itself become a major industrial activity. Nevertheless it is still true that not one catalytic reaction involving solid catalysts can he explained in rigorous detail. Considerable effort, however, has recently been directed toward the interpretation of heterogeneous catalytic phenomena and some advances may be chimed even in a field as complex as that of catalytic cracking of hydrocarbons (Bloch and Thomas, 4; Greensfelder, Voge, and Good, 24; Hansford, 26; Mills, Boedeker, and Oblad, 41 ; Parravano, Hammel, and Taylor, 43; Thomas, 60; Turkevich and Smith, 61). One of the obvious reasons for slow development in this direction is that there are very few tools available for the study of the chemistry of the surface layer of a solid and for the observation of chemical changes that take place at the surface. The techniques that have been developed involve many theoretical and experimental difficulties. It has become increasingly clear, however, that while these efforts in the direction of surface chemistry are of critical importance, knowledge of the physical structure of catalysts is just as essential for a complete understanding of catalyst function. Furthermore, the physical properties are valuable for following catalyst deterioration on use and regeneration, the effect of processing variables on the catalyst and the relative importance of various factors in catalyst preparatio:n and treatment. Attention is currently directed toward two aspects of physical structure; namely, surface area and pore geometry. In general the activity of a given cracking catalyst is proportional to its surface area providing that there is no poisoning and no major alteration of pore structure or surface composition. One of the most useful techniques for observing changes in physical properties such as surface area, pore radius and pore volume is that of low-temperature gas adsorption developed principally by Emmett (Emmett, 15, 16; Emmett and Brunauer, 1’8). For example, one of the earliest and most valuable applications of area ‘measurement, in addition to that of following and anticipating catalyst deterioration,

STRUCTURE AND SINTERING

PROPERTIES OF CRACKING CATALYSTS

89

is in the interpretation of effects attributed to catalyst poisoning. If the activity of a catalyst decreases more rapidly than its surface area, poisoning may be suspected; whereas, if activity and area decrease in direct proportion, thermal deactivation without poisoning is indicated. Furthermore, area values provide a means for determining whether promoters and supports improve catalysts by increasing or maintaining area or by qualitative changes that increase catalyst activity per unit surface area. The effects of temperature and steam on the deterioration of catalysts employed in cracking processes are under continuous surveillance, and these effects may be studied in a rather straightforward manner by physical property measurements independent of the chemistry of the surface. For example, as will be shown later, if a significant increase in pore radius accompanies loss in area, steam deactivation is probably indicated. Moreover the pore structure as determined by adsorption techniques is undoubtedly related t o the ease of admission of reactant molecules and the diffusion out of product molecules as well as to the regeneration properties when carbonaceous deposits must be removed. The material presented in the following discussion represents one phase of a program designed todetermine, for catalytic cracking processes, the relationship between catalyst structure and catalyst function and the various factors involved in catalyst deterioration. Considerable emphasis is given to the effect of steam since it is believed that this is one of the principal factors promoting structural deterioration of cracking catalysts. Sintering in the presence of steam not only accelerates catalyst deterioration but also apparently produces catalyst properties very similar to those resulting from commercial use. The use of steam in accelerated aging procedures is of current practical interest. Steam stripping and the hydration of catalysts to combat poisoning are important commercial operations (Conn and Brackin, 11; Davidson, 13; Mills, 40). Furthermore, steam is present to some extent throughout the entire cracking process. Moreover, cracking catalysts of current interest are prepared in aqueous media. It has also been demonstrated that a certain amount of surface water is necessary for cracking catalyst activity (Hansford, 26). The present discussion is limited for the most part to low temperature nitrogen adsorption studies and sintering experiments. An adsorptiondesorption isotherm yields at once surface area, pore volume, average pore radius and an approximate pore size distribution. Such an isotherm is thus an excellent fingerprint of the physical structure of the catalyst. Sintering curves, or temperature-area plots, obviously demonstrate relative thermal stabilities of these structures under the conditions

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HERMAN E. RIES, JR.

of the sintering experiments. These techniques are t:herefore employed in the study of virgh, vacuum-sintered, steam-sintered, and commercially used cracking catalysts as well as in the study of certain related materials.

11. EXPERIMENTAL PROCEDURES AND INTERPRETATION 1 . Apparatus and Technique for Gas Adsorption Measurements

General procedures for the determination and interpretation of adsorption isotherms have been described by Emmett and Brunauer (Emmett, 15, 16; Emmett and Brunauer, 17; Brunauer, 5). The particular apparatus and techniques employed in tbe present study have been reported earlier (Ries, Van Nordstrand, Johnson, and Bauermeister, 50: Ries, Van Nordstrand, and Kreger, 51; Ries, Van Nordstrand, and Teter, 52). A photograph of one of the a,dsorption systems used by M. F. L. Johnson in the Sinclair Laboratories is presented in Fig. 1. The glass fabrication for this system was performed by J. A. Glover. It is essentially a low-temperature, high-vacuum apparatus for the measurement of gas adsorption by the volumetric method. An exceedingly useful discussion of this type of apparatus and procedure has been prepared recently by Joyner (34). Many procedures and precautions important in this type of study have been discussed in detail by Frankenburg (22). In general, the temperature of the liquid nitrogen bath in which the sample bulb is immersed varies less than 0.4"C. during a single adsorptiondesorption hysteresis experiment. An experiment of this type usually extends over a period of eight to ten days. During the portion of the isotherm measurements used for calculating surface area by the BrunauerEmmett-Teller (BET) method the temperature of the bath rarely varies more than 0.03"C. A nitrogen vapor pressure thermometer is used continuously and readings are made to within 0.2 mm. of mercury. The relative pressure term, p/po, which is used throughout as a parameter, is simply the ratio of the measured pressure, p , in the a,dsorption system to the liquefaction or saturation pressure, PO. The liqu'efaction pressure, p o , given by the vapor pressure thermometer is the vapor pressure of the adsorbate at the temperature of the bath. The use of relative pressures tends to correct for small temperature fluctuations. Occasionally a slight separation of the adsorption and desorption curves is observed at low relative pressures, but probably little significance should be attached to this since a leakage error accumulated throughout the lengthy experiment is undoubtedly indicated.

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

91

Special care is exercised in the determination of adsorption-desorption hysteresis curves since significant differences in structure are thus established. Time intervals varying from ten minutes to two hours are required for equilibrium, depending on the portion of the isotherm concerned and the amount adsorbed. After apparent equilibrium is reached the system is allowed to stand for an appropriate period to confirm the

FIG.1. Adsorption isotherm apparatus for catalyst studies.

equilibrium. At certain critical points for both adsorption and desorption, equilibration is verified by allowing the system to stand for sixteen hours. Unless otherwise indicated the catalysts studied are degassed a t 350°C. for sixteen hours by means of a mercury diffusion pump system. The helium used for dead space measurements is 99.9% pure and is obtained in special Pyrex flasks from the Ohio Chemical Company. The nitrogen which is over 99 % pure as received from the Ohio Chemical Company is purified further by passage over copper gauze a t 500°C. and phosphorous pentoxide with a final distillation a t liquid nitrogen

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HERMAN E. RIES, JR.

temperatures. Periodic mass spectrometer analyses of the purified nitrogen indicate that 100.0% purity is achieved. 2. Sintering Treatments

Vacuum sintering curves are determined by successive heat treatments performed during evacuation. The sample under study is contained in a small quartz bulb sealed to the adsorption isotherm system through a graded seal. A heat treatment a t a given temperature extends for approximately twelve hours. Sintering curves therefore do not represent the final state of sintering for the indicated temperatures. At certain points on these plots the heat treatments are repeated at the designated temperatures; the consequent decrease in area and possible approach to a final state for infinite time are indicated. However, it may well be true that a stable state is never reached at any temperature at which sintering is observed. Steam treatments are performed for the most part on larger samples of catalyst in a conventional metal reactor type system. The catalyst and the system are thoroughly purged with steam andL then maintained in an effectively static steam atmosphere at the pressure and temperature and for the time specified in Table I. Small Pyrex bulbs are used for atmospheric pressure steam treatments. The latter treatments are performed by means of a slow stream of steam generated and preheated below the catalyst sample. The effectiveness of the steam in reducing catalyst area increases somewhat with increasing pressure. However, the particular treatments selected for this discussion were chos’en primarily for purposes of convenient comparison with vacuum sintered and commercially used samples. Most of the steam treatments at elevated pressures were performed by R. I. Lindberg of our laboratories. 3. Helium-Mercury Pore Volumes

A helium-mercury displacement method is sometimes used to check the pore volume obtained from the adsorption-desorption isotherm at the saturation pressure (Ries, Van Nordstrand, Johnson, and Bauermeister, 50). The helium measurement which gives the solid volume and solid density is obtained by means of a miniature isotherm type apparatus. This part of the determination is essentially the same as that described by Smith and Howard (58) and Schumb and Rittner (55). Mercury displacement yields the pellet volume and pellet density and is measured volumetrically by means of a mercury buret attached directly to the same sample bulb in which the helium measurement is performed. The difference between the pellet volume and the solid volume is obviously the pore volume.

TABLE I Area, Pore Volume and Average Pore Radius for Representative Cracking Catalysts an the Virgin, Steam-Treated, Vacuum-Sintered, and Used States V,, cc. Adsorbed Experiment in Monolayer on 1 g. (ig. wt.) Number (Hysteresis) Adsorbent (BET) Nalco Nalco: Nalco Nalco:

silica-magnesia microspheres used equilibrium (U O.P.) steam-treated 621'C. (1150°F.) 400 psig., 24 hr. steam-treated: 621OC. (1150'F.): 400 psig., 24 hr.

DA-5, silica-magnesia DA-5, silica-magnesia (calcined) DA-5. silica-magnesia calcined), steam-treated, (1175"F.), 400 psig., 246r. '

635'C.

BET Area, sq. m./g. ig. wt.

Per PI, Average Cent Ve. Gas Pore Volume Pore Volatile Adsorbed a t (from V , ) Radius. A. PO, cc./g. Matter cc./g. (r = 2P,,/A)

2575 3537 2810 2858

144.7 72.7 72.8 73.8

630 317 317 322

4.6 8.9 4.4 4.9

292 2 10 193 183

0.451 0.325 0.298 0.283

14.3 20.5 18 8 17.6

461 1 2674

151.0 71.8

656 312

20.9 3.4

236 134

0.365 0.207

11.1

302

34.7

151

5.0

68

0.105

13.9

1.3 7.2

310 290

0.479 0.448

21.8 33.2

13.3

Socony TCC Beads, silica-alumina Socony TCC Beads used Socon TCC Beads: steam-treated, 538'C. (lOOO°F.), 28 hr.; 566&. 1050'F.) 10 hr atm. Socon T&C Bead;, steal;;-treated, 538°C. (lOOO°F.), 28 hr.; 566%. (1050°F.), 22 hr.; 5930C. (llOO°F.), 40 hr.; atm. Soconv TCC Beads. vacuum-sintered, 970'C. (1778"F.), 12 hr.; see reference.31

745 716

101.3 62.2

440 270

3665

70.3

306

2.5

282

0.436

28.5

3690

55.2

240

0.1

274

0.424

35.3

754

43.2

188

1.0

135

0.208

22.2

Aerocat Aerocat: Aerocat, Aerocat,

2620 2828 315 319

160.9 19.6 95.0 32.8

700 85 413 143

27.4 0.6 25 9 25.9

375 326 237 96

0.580 0.504 0.366 0.148

16.6 119.0 17.7 20.7

Diakel silica-alumina Diakel: silica-alumina, used Houdry 5-46, silica-alumina pellets Houdry Porous Beads

2021 2035 327 3658

94.2 29.9 69.6 57.0

410 130 303 248

6.7 0.5 11.2 2.2

284 191 313 334

0.439 0.295 0.484 0,516

21.4 45.4 31.9 41.6

Fluid Fluid Fluid Fluid

2790 217 4059 2857

77.9 31.3 32.5 31.1

339 136 141 135

21.0 3.7 2.5 1.9

265 250 215 240

0.410 0.386 0,332 0.371

24.2 56.8 47.1 55.0

4052 282 1 2636 3696 2611

38.7 25.9 63.5 37.6 52.2

169 113 276 164 227

3.1 2.8 7.5 17.5 1.0

248 244 235 200 230

0.383 0.377 0.363 0,309 0.356

45.4 66.1 26.3 37.7 31.4

.

silica-alumina microspheres steam-treated 621°C. 1150DF.),75 psig., 24 hr. vacuum-sinteied. 900°&. (1652'F.), 12 hr. vacuum-sintered. 975'C. (1787'F.). 12 hr.

Filtrol Filtrol: Filtrol Filtrol:

clay used, equilibrium (A) used equilibrium (E) ured: equilibrium (GI

Fluid Filtrol steam-treated 570OC. (105S°F.), 75 psjg., 24 hr. Fluid Filtrol: steam-treated: 627°C. (1160°F.), 75 pslg., 24 hr. TCC clay pellets Fluid Filtrol SR Houdry Type I

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HERMAN E. RIES, JR.

Q. Materials a. Cracking Catalysts. The cracking catalysts studied are for the most part well-known commercial products. Approxi mate compositions of these materials are tabulated below. Catalyst DA-5 Nalco Microspheres

Composition 35 % MgO, 65 % SiOz 25% MgO, 75% SiOz

Manufacturer Davison Chemical National Aluminate

TCC Beads Houdry Porous Beads Aerocat Microspheres Diakel Houdry synthetic Houdry S-46

11% A1~03,89% SiOz 10% AlzOa, 90% SiOz 11% A1~03,89% SiOz 11% A1~03,89% SiOz 13% A1203, 87% SiOz 12% Al2O3,88% SiOz

Socony-Vacuum Houdry Process American Cyanamid Davison Chemical Houdry Process Houdry Process

TCC clay pellets

16% A1203,74% SiOz (MgO, CaO, Fez03,SO,’) 15% i%Oa, 72% SiOl (MgO, CaO, Fez03, Sod’) 42% Al2O3,56% sioz, 2% Fez03 18% AIZO3,78% SiOi(Mg0)

Filtrol

Fluid Filtrol Filtrol SR Houdry Type I

Filtrol Filtrol Houdry Process

b. Silicas and Related Materials. The various forms of silica and some related materials investigated are also principally commercial products. These adsorbents may be described as follows. The high area small pore silica xerogel used in these studies is manufactured by the Davison Chemical Corporation and is labeled Davison Silica Gel 12966-120. It has a silica content of 99.7%, on the dry basis. The silica aerogel, Aerogel “S,” was prepared by J. L. Gring of our laboratories following the general method of Kistler (36). The procedure consists essentially of alcohol replacement of the water in the hydrogel and subsequent removal of the alcohol above its crihical temperature. The silica content of this preparation is 99.8 % on a dry basis with 0.03 % A1203,0.02% Fe, 0.06% Na, and 0.01% SO*‘. Aerogel “S”has a bulk density of approximately 0.14 cc./g. The silica bead aerogel was similarly prepared by Henry Erickson of our laboratories from a silica bead hydrogel obtained from the Socony-Vacuum Laboratories. Santocel “ C ” is a silica aerogel produced by Monsanto Chemical Company. Its silica content is approximately 96% on the dry basis, and its impurities include small percentages of NazS04, A1203, and Fe2O3. Linde silica, according to its manufacturer, the Linde Air Products Company, has a purity of 99.9% on the dry basis and a maximum of 8%

STRUCTURE AND SINTERING PROPERTIES O F CRACKING CATALYSTS

95

adsorbed water. A special method of preparation has produced an extremely finely divided nonporous material. The alumina was made available by the Harshaw Chemical Company and is designated “Activated Alumina.” Titanium dioxide was obtained from the Titanium Pigment Corporation and is referred to as “Titanox-A MO” (anatase). The diatomaceous earth discussed here and in studies reported earlier is designated as Celite 337 by the Johns-Manville Corporation. Because of the wide variation in moisture content, areas and related properties are expressed per gram of ignited weight. Ignition losses are obtained by heating the samples to constant weight at 1000°C. 5. Interpretation of Adsorption Isotherm Data

Since nitrogen adsorption-desorption isotherms provide considerable information on the physical structure of cracking catalysts, the methods used in interpreting isotherm data will be briefly considered. In this connection it will be convenient to refer to the isotherms of some representative cracking catalysts and certain related materials presented in Fig. 2. More detailed plots of the isotherms may be found in subsequent sections. In all cases of hysteresis presented in Fig. 2 the upper portion of the hysteresis loop represents desorption and the lower curve represents adsorption. Catalyst characteristics of principal significance obtained from these plots are listed in Table I. As pointed out above the adsorption-desorption isotherms yield at once surface area, pore volume, and pore radius information. a. Surface Area. Area values are calculated by means of the Brunauer-Emmett-Teller (BET) method (7). For a number of somewhat similar adsorbents studied earlier in our laboratories and by others both the BET and Harkins-Jura (HJ) methods were used and reasonably satisfactory agreement obtained (Harkins and Jura, 28; Ries, Johnson, and Melik, 47). There is of course a variety of additional support of the BET method (Cassie, 8; Emmett, 16; Hill, 30). Although the BET method is widely used and generally accepted for relative area measurements, the method has been criticized for calculating areas of solids having very small pores (Pierce, Wiley, and Smith, 44). Nevertheless it is believed that the BET areas for the small pore catalysts discussed here are quite reliable as relative values. This viewpoint is supported by a recent study in our laboratories in which a large pore aerogel was shrunk to a small pore system without any significant change in BET area and without any indication of nonuniformity or sintering of the structure (Johnson and Ries, 33; Ries, 46). This point will be discussed in a later

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HERMAN E. RIES, JR.

section in connection with aerogel and xerogel structure comparisons (Sec. IV.2.d). It is also in order to point out that care must be observed in the application of the BET area method when nitrogen adsorption is employed

FIQ.2. Nitrogen adsorption isotherms.

on metallic surfaces. Beeck (3a) has reported that nitrogen chemisorption occurs on certain metals at liquid nitrogen temperatures. Therefore such determinations would give area values too great by an amount equivalent t o the chemisorbed nitrogen. No complication of this type should be involved in the present study since no metallic adsorbents are

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

97

considered. All the catalysts and related materials investigated in this series are in the form of oxides. The BET area method simply involves the application of the following equation : p =-+ 1 - (c - U P Vdp, - p )

vmc

vmcpo

in which V , is the volume adsorbed at the measured pressure, p ; V , is the volume adsorbed in the monolayer; po is the liquefaction pressure of the adsorbate gas; c is a constant related to the heat of adsorption and the heat of liquefaction of the adsorbate. Since the equation is linear a plot of p / V , ( p ~- p ) against p / p o will yield a value for V,. The Vm values thus obtained are translated into area units, square meters per gram, by using 16.2 sq. A. as the area value for the nitrogen molecule (Emmett and Brunauer, IS). The V , values are included in Table I because they represent relative areas according to the BET equation without the assumption of a specific molecular area for nitrogen or, more correctly, they represent the relative number of molecules adsorbed. These calculated V , values are generally in reasonably good agreement with those of the so-called B point method recommended in the earlier work of Emmett and Brunauer (18). The B point is the lower extremity of the central linear portion of the adsorption isotherm. b. Pore Volume. The pore volumes listed in Table I are calculated from V s data. The V , value is the volume adsorbed a t the liquefaction pressure, po, the point at which the relative pressure equals 1.0 and the pores are filled with liquid nitrogen. This is obtained by determining the point at which the desorption isotherm breaks away from the po line; in other words, the point a t which the pores begin to empty (Holmes and Emmett, 32; Ries, Van Nordstrand, Johnson, and Bauermeister, 50). It is apparent that for certain isotherms of Fig. 2 the selection of a break point is a straightforward procedure since these isotherms meet the po line almost perpendicularly whereas for others great care in the determination of the desorption branch is necessary. In some cases an inflection point at a relative pressure of approximately 0.9 may be chosen as V , since there is reason for believing that adsorption above this point may be principally interparticle condensation (Van Nordstrand, Kreger, and Ries, 62). Due to adsorption forces the density of liquid nitrogen in very small pores probably differs from the density of the bulk liquid which is used in calculating pore volume. An uncertainty is therefore introduced in the pore volume calculated by this method. Nevertheless in some earlier work in our laboratories on other materials the pore volumes determined by the helium-mercury displacement method were in satisfactory agreement with the isotherm values.

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HERMAN E. RIES, JR.

c, Pore Radius. Adsorption-desorption isotherms represent a combination of thin-film or multilayer adsorption and capillary condensation. Nevertheless, average pore radius values, which are used throughout this discussion, may be obtained simply by the application of geometry to the pore volume-area relationship. This relationship is

where r is the average pore radius, V the pore volume, and A the surface area. Long or open end cylindrical pores are assumed (Emmett and DeWitt, 19). For a platelet structure the average separation of the platelets, d, is also equal t o 2V/A or d is equal to r. Approximate pore radius values may also be obtained for a given point on an isotherm by application of the Kelvin equation to the corresponding relative pressure. The Kelvin equation, which relates the vapor pressure of liquids in small capillaries to the radius of curvature, indicates that the vapor pressure decreases with decrease in capillary radius. In the Kelvin relationship In p l p ~=

-2uv cos 8

rRT

where p is the equilibrium pressure, po the liquefaction pressure, u the surface tension of the adsorbate, V the volume of one mole of the liquid adsorbate, 0 the angle of contact of the adsorbate on the adsorbent, r the capillary radius, R the gas constant, and T the absolute temperature. The many assumptions involved in the application of the Kelvin equation have been discussed earlier (Foster, 21 ; Ries, Van Nordstrand, Johnson, and Bauermeister, 50). However, a relatively long and steep desorption portion of an isotherm clearly indicates a narrow distribution or a predominance of pores in the indicated size range. Such Kelvin values when corrected by adding the thickness of two adsorbed layers, or approximately 8 A., frequently are in good agreement with average pore radius values. Examination of the hysteresis isotherms presented in this study discloses that in general the desorption branches rejoin the adsorption curves above 0.4 relative pressure. The significance of this observation is probably that below 0.4 relative pressure, in the region of two statistical monolayers and less, the nitrogen molecules are adsorbed and desorbed in a manner independent of the capillttry condensation forces. Hysteresis phenomena are generally attributed to capillary condensation and evaporation effects (Cohan, 9, 10; Emmett and DeWitt, 19; Foster, 20, 21; Gleysteen and Dietz, 23; Kraemer, 37; McBain, 38; Ries, Van Nordstrand, Johnson, and Bauermeister, 50; Ries, Van Nordstrand, and Teter, 52).

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

99

Refined methods for combining the Kelvin relationship with adsorption equations for obtaining pore size distribution have recently been suggested (Barrett and Joyner, 2; Shull, Elkin, and Roess, 57; Wheeler, 64, 64a). With some experience, however, simple visual inspection of the isotherms proper provides a reasonable impression of the relative pore size distributions. It must be borne in mind when considering pore structure and particularly the structure of high area gels that the shape of the ultimate particles and thus the shape of the pores is not known. It remains to be established whether the particles are platelets, fibers, spheres, or complex combinations of many 'structures. d. Hysteresis Loops and Pore Structure. The following qualitative observations on hysteresis loops may be of some significance in structure and adsorption theory. For small pore systems the hysteresis loops are extremely narrow, probably because adsorbed layers fill the pores without the benefit of capillary condensation in the classical sense; that is, deposition and removal may involve similar forces. The hysteresis loops then broaden and increase in area for adsorbents composed of larger pores (Ries, Melik, and Johnson, 49). Condensation and evaporation from such structures involve different surface contours and consequently different forces. With systems having very large pores, such as aerogels, a reversal of the trend or a narrowing of the hysteresis loop takes place. The latter isotherms approach those of nonporous powders which show no hysteresis (Harkins and Jura, 28; Ries, Johnson, Melik, 47; Ries, Johnson, Melik, and Kreger, 48). In other words the extreme of large pores is no pores at all or perhaps a system in which the solid adsorbing surfaces are sufficiently separated that there are no cooperative force effects. It should also be pointed out that the Kelvin relationship between relative pressure and pore radius discussed above accounts for some of the narrowing of the hysteresis loop for large pore systems due to the logarithmic nature of the expression.

111. STRUCTURE AND SINTERING PROPERTIES OF REPRESENTATIVE CRACKING CATALYSTS A preliminary overall picture of cracking catalyst structures is available in the first three horizontal rows of the composite plot of Fig. 2 and the corresponding data of Table I. Isotherms presented in the lowest row are discussed in Sec. IV. Only the general features of these representative types of cracking catalysts are indicated here, since the detailed plots of individual isotherms will be considered in subsequent sections on sintering. Cracking catalysts of principal interest are represented by three types : silica-magnesia; silica-alumina; and activated clay.

100

HERMAN E. RIES, JR.

The silica-magnesia catalysts, DA-5 and Nalco, in the virgin state, along with Davison silica gel have practically their entire area and pore volume contributed by the very smallest of pores that are encountered in catalyst structures; that is, pores in the 10 to 15 A. radius range. It is apparent in Fig. 2 that for these materials there is no appreciable adsorption at the high relative pressures. This indicates the absence of large pores. One and one-half monolayers according to the BET theory effectively fill the pore volume of the DA-5 and the Davison silica gel, and only two monolayers are required for Nalco. Ver,y little hysteresis is observed for any of these three materials. The synthetic silica-alumina catalysts, TCC Beads, Aerocat, and Diakel, are composed of pores appreciably larger but with average pore radius values almost exclusively in the small pore range of 15 to 25 A. (Ries, Johnson, Melik, and Kreger, 48; Shull, Elkin, and Roess, 57). The complete absence of large pores is indicated for the TCC Beads and the Aerocat Microspheres. In the case of silica-alumina catalysts, the hysteresis loops have considerable breadth and area in contrast to the silica-magnesia isotherms. The clay type catalysts such as TCC clay and Fluid Filtrol have a considerably wider pore size distribution which inclu.des pores having radii much greater than those of the synthetic silica-alumina and silicamagnesia. There is appreciable adsorption in the high relative pressure region, and the hysteresis loops are broad (Oulton, 42; Ries, Johnson, Melik, and Kreger, 48; Ritter and Drake, 53). It is thus clear that the isotherm contours are strikingly different for the three types of representative cracking catalysts considered. The isotherms for the members of each group are, however, remarkably similar. Extensive surface area is certainly characteristic of the currently important cracking catalysts. The surface areas of the synthetic materials are in general considerably greater than those of the clays. It is also indicated that internal area contributed by the small pore structure functions catalytically since some of the catalysts, such as the bead type, consist almost exclusively of small pores acid have an entirely negligible external area. Furthermore, it is generally observed that for a given cracking catalyst, activity decreases with diminution of area. The importance of sintering studies of cracking catalysts as pointed out earlier is evident at once since sintering or surface area loss generally represents an irreversible deterioration of the catalyst. An investigation of the various factors involved in sintering is not only of value in commercial applications of catalysts but will also improve our understanding of solid structure as related to the basic reorganizations that occur during treatments of various types. Changes in surface area, pore volume, and

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101

pore radius for known compositions subjected t o various types of heat treatment will provide clues for the mechanism of the deterioration of structure as well as for the changing catalytic functions of the solid. The following discussion is limited to a consideration of sintering effects obtained in vacuum, in steam, and in commercial use. Similarities and differences for silica-magnesia, silica-alumina, and clay catalysts are indicated. The various forms of silica are considered separately in Sec. IV. A major portion of the sintering study presented at this time has been devoted to the synthetic silica-alumina catalysts. The reasons for this emphasis on silica-alumina are twofold; in the first place, during World War I1 this material was of great industrial importance in the production of high octane aviation gasoline; secondly, the uniformity of both the physical structure and the chemical composition of the synthetic silica-aluminas provides obvious advantages in such an investigation. A few introductory statements on the general relationships observed in sintering studies will help to orient the more detailed discussion. It appears that in general the sintering of silica-alumina cracking catalysts in vacuum causes a decrease in area and a corresponding decrease in pore volume without appreciable alteration of the pore radius or the structure of the pores remaining. The presence of steam, however, not only accelerates the sintering process but also causes the area t o decrease more rapidly than the pore volume and consequently produces an increase in pore size. Steam tends to maintain the total pore volume during sintering and therefore the external dimensions of the particles are changed much less than in vacuum sintering. Isotherms of the used catalysts indicate that an effect similar to that of steam treatment accompanies commercial processing. The presence of steam during heating accelerates the sintering process for silica-alumina catalysts to a much greater degree than it does in the case of silica-magnesia preparations. Steam apparently interacts intimately with silica-alumina and reorganizes its structure in a rather profound manner. Evidence for this interaction is in the form of increases in pore size and decreases in activity and area. Isotope exchange reactions between the oxygen of water maintained at 100°C. in contact with the catalyst and the oxygen of the catalyst have been reported by others (Hindin and Mills, 31; Heinemann, 29) and also support the intimate nature of this interaction. I n the case of silicamagnesia, however, steam evidently does not have as important an effect on the pore or surface structure. In Fig. 3 are presented sintering curves for representative cracking catalysts. Surface area values are plotted against temperature of

102

HERMAN E. RIES, JR,

treatment. The heat treatments for this series were performed in vacuum at the designated temperatures and extended for twelve hours unless otherwise indicated. Because of the limited time period of the treatment, the area values do not represent the final state of sintering at the I 700

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FIQ.3. Effect of temperature on surface area of cracking catalysts. 12 hours in vacuum at indicated temperatures.

Treatments:

indicated temperatures. Perhaps no truly stable state is reached at any temperature at which sintering is observed. Furthermore, two other considerations must be borne in mind when comparisons are made with commercial stability data. In the first place commercial cracking and

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

103

regeneration operations are in the approximate range of 480 t o 570°C. (896 to 1058°F.). This range is well below the steep fall-off portions of the vacuum sintering curves. Secondly, steam which is nearly always present in commercial units may profoundly alter the sintering relationships as is indicated below (see also Conn, Meehan, and Shankland, 12; Richardson, Johnson, and Robbins, 45). The DA-5 silica-magnesia catalyst studied in the sintering series demonstrates good thermal stability up to 700°C. (1292°F.) and then loses area rapidly to reach zero area at 800°C. (1472°F.). Other samples of DA-5 indicate that this one may have been overcalcined and thus have abnormal stability in the lower temperature range. The Nalco silica-magnesia catalyst with a greater initial area also falls to zero area at 800°C. Nalco, however, begins to suffer appreciable area losses in the region of 600 to 700°C. (1112-1292°F.). The silica-alumina catalysts, Aerocat, Diakel, and TCC Beads, maintain relatively large areas at considerably higher temperatures and fall to zero area in the neighborhood of 1000°C. (1832°F.). The TCC Beads show a negligible area loss up to 800°C. The clay type catalysts, Fluid Filtrol, TCC clay and Houdry Type I, fall to zero area in an intermediate region, 850 to 900°C. (1562-1652°F.). The Houdry Type I catalyst demonstrates good stability up to 8OO"C., whereas the Filtrol clays lose considerable area between 700 and 800°C. 1. Silica-Magnesia Catalysts

The effect of steam sintering the silica-magnesia catalysts, Nalco and DA-5, is demonstrated in Figs. 4 and 5 and Table I. In the case of Nalco the steam treatment at 620°C. (1150°F.), which reduces the area from 630 to 322 sq. m./g., has very little effect on the contour of the hysteresis isotherm. The curves for the virgin and steam-treated Nalco have similar hysteresis loops and relatively flat portions above 0.5 relative pressure. In the higher relative pressure regions the isotherms curve up slightly. The pore volume has been decreased to a lesser extent than the area so that the pore radius increases from 14.3 to 17.6 A. It is of interest that the steam treatment results in an area of 322 sq. m./g., practically identical to that of the used equilibrium sample, 317 sq. m./g. The lower portions of the isotherms for the used and the steam treated samples almost coincide but the higher relative pressure region of the used catalyst shows the presence of pores larger than those produced by the steam treatment. The average pore radius of the used equilibrium catalyst is 20.5 A., somewhat greater than that of the steam treated material, 17.6 A.

104

HERMAN E. RIES, JR.

In Fig. 5 the striking similarity of the isotherms for a virgin DA-5 and a steam-treated DA-5 is apparent. The pore volume and area of the steam-sintered sample are far below those of the virgin, but the isotherm contours are almost identical. A small increase in pore radius is observed. Thus the presence of steam during sintering does not alter significantly the pore structure of the silica-magnesia, catalysts studied. It should also be noted here that the isotherms for the DA-5 catalyst

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FIQ.4. Nitrogen adsorption-desorption isotherms for virgin, steam-treated, and used Nalco silica-magnesia microspheres.

exhibit practically no hysteresis and are more like Davison silica gel in adsorption characteristics than are those for Nalco. IDavison silica gel isotherms are designated as Type I isotherms according to the classification of Brunauer, Deming, Deming, and Teller (6).

2. Silica-Alumina Catalysts The adsorption-desorption isotherms for synthetic silica-alumina catalysts are somewhat similar to those of silica-magnesia in that there is very little adsorption in the higher relative pressure region. However, the average pore radius of the silica-alumina preparations, generally in

OF CRACKING

STRUCTURE AND SINTERING PROPERTIES

CATALYSTS

105

the range of 15 to 25 A., is greater than for the silica-magnesia catalysts, whose pore radii are in the range of 10 to 15 A., and the hysteresis loops of the silica-alumina isotherms are much broader. a. TCC Bea.ds. A detailed isotherm plot of the virgin Socony TCC Bead catalyst and a comparison of a steam-treated with a commercially used sample are presented in Fig. 6. Area values for the steam-sintered and used materials are respectively 240 and 270 sq. m./g. while that of

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FIG.5 . Nitrogen adsorption-desorption isotherms for virgin and steam-treated DA-5 silica-magnesia.

the virgin sample is 440 sq. m./g. The structural similarity of the steam-treated and the used is immediately apparent (see also Drake, 14). The steep portions of both isotherms are displaced considerably to the right of the virgin catalyst isotherm. In addition the average pore radii are similar; 35.3 A. for the steam-treated and 33.2 A. for the used, compared to 21.8 A. for the virgin. The isotherms also demonstrate that the steam treatment of the catalyst and its use in actual processing have produced very little if any of the large pores whose presence would be indicated by adsorption near PO. It appears likely that sintering in the

106

HERMAN E. RIES, J R .

FIG.6. Nitrogen adsorption-desorption isotherms for virgin, steam-treated, and used Socony TCC Beads.

presence of steam may be the principal factor in the physical deterioration of this type of catalyst during commercial processing. Additional evidence supporting this viewpoint and a detailed comparison of vacuum and steam sintering are presented below. Complete adsorption-desorption isotherms obtained throughout a series of vacuum sintering treatments of the TCC Elead catalyst are

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

107

plotted in Figs. 7 and 8 (Van Nordstrand, Kreger, and Ries, 62). These isotherms are good examples of Type IV according to the classification of Brunauer, Deming, Deming, and Teller (6). In Fig. 7 isotherms for sintering treatments at 660, 880, and 970°C. are presented. The data for Fig. 8 were obtained by heating at 970°C. for different periods of time. During the sintering series shown in Figs. 7 and 8 the area is reduced stepwise from a value of 440 sq. m./g. to approximately 10 sq. m./g. It is apparent a t once that the general contour of the

RELATIVE PRESSURE. P/R

FIG.7. Nitrogen adsorption-desorptionisotherms for Socony TCC Bead catalyst sintered in vacuo.

isotherms remains unchanged during the sintering sequence. I n other words adsorption throughout the entire relative pressure range decreases proportionately. This means that the structure of the pores that survive the heat treatment is the same as their initial structure. Pore volume that disappears on vacuum sintering is lost completely, and the pores that remain are effectively unchanged (62). Throughout the sintering series shown in Figs. 7 and 8 the isotherms appear to intersect the p o line perpendicularly, and thus there is no indication of the development of large pores during the sintering process. Moreover, the adsorption and capillary condensation processes that fill

108

HERMAN E. RIES, JR.

the pores are effectively complete at 0.6 relative pressure. Average pore radius values as calculated from pore volume and area data obtained throughout the series vary only slightly from the original pore radius of 21.8 A. Inspection of the isotherm plots indicates that the steep portion of the isotherm near 0.5 relative pressure is displaced o'nly very slightly t o the left during the sintering. This also demonstrakes that there is very little change in the size of those pores that contribute almost the entire pore volume and area. It may be of some interest that the average I

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FIG.8. Nitrogen adsorption-desorption isotherms for Socony TCC Bead catalyst sintered in vacuo at 970°C. pore radius values calculated from pore volume and area show a very slight increase on sintering whereas the displacement of the steep portion which is related to the Kelvin type calculation of pore size indicates a very slight decrease in pore radius. However, neither method is sufficiently accurate nor are the differences sufficiently great to attach any significance at this time to these opposite trends. I t will also be noted that in this series the desorption branches of the isotherms in general rejoin the adsorption curves above 0.4 relative pressure. This probably means that below this point, in the region of two statistical monolayers and less, the adsorption and desorption of nitrogen molecules are independent of capillary condensation forces. Capillary condensation

STRUCTURE AND S I N T E R I N G P R O P E R T I E S O F CRACKING CATALYSTS

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109

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FIG.9. Nitrogen adsorption-desorption isotherms for virgin, steam-treated, and vacuum-sintered TCC Beads.

effects are generally considered to be responsible for hysteresis phenomena (see Sec. 11.5.~). I n Fig. 9 two steam-treated TCC Bead samples are compared with a vacuum-sintered material. Although the heat treatment in vacuum reduces the area from 440 to 188 sq. m./g., the pore radius shows only the insignificant increase 21.8 t o 22.2 A. (Van Nordstrand, Kreger, and Ries, 62; Milligan and Rachford, 39). This is in sharp contrast t o the pore

110

H E R M A N E. RIES, JR.

radius increase, 21.8 to 35.3 A., accompanying a steam treatment which effected a much smaller area decrease, 440 to 240 sq. m./g. The steam not only accelerated the sintering since a much lower temperature was used with steam (Table I), but it enlarged the pores (Ries, Johnson, Melik, and Kreger, 48; Ries, 46). However, the isotherms obtained after the steam sintering treatments also intersect the p o line sharply. This indicates that very large pores have not developed during the

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FIG.10. Schematic geometry of extreme cases of catalyst sintering.

sintering. It might be said that the pore structure enlargement is due principally to the effect of steam in maintaining pore volume. In the more severe of the two steam treatments only about 12% of the pore volume was lost whereas more than 50% of the pore volume disappeared with the vacuum sintering. An idealized platelet picture of catalyst structure as related to possible mechanisms for sintering in vacuum and in steam is presented in Fig. 10. Extreme cases of sintering are shown schematically. Structure representations of this type may aid in the ultimate interpretation of sintering

STRUCTURE AND SINTERING PROPERTIES O F CRACKING CATALYSTS

11 1

which will involve both chemical change and particle growth. The schematic geometry of Fig. 10 illustrates the following: (I) The top row represents a set of four parallel square platelets spaced one unit apart and each one square unit in area. The platelets are assumed to be the ultimate particles of a catalyst. Some cracking catalysts such as certain clays do possess platelet structures. Only the internal area is considered here, since in a high area catalyst system the external surface is negligible; thus: A (area) = 6 V (pore volume) = 3 d (platelet separation) = 1

It will be recalled that d

= T = 2 V / A for pore systems in which r is the average radius of long cylindrical pores. (11) Vacuum sintering may cause the outer platelets to move in and fuse with the inner platelets giving up area and pore volume in direct proportion and not affecting the radius or platelet separation of the remaining pore or pores. In this case, illustrated in the second row,

A = 2

V = l d = l

A threefold decrease in area and pore volume is thus represented with no change in average pore width. In other words, pores that disappear are lost completely, and pores that remain retain their initial structure. This also implies an overall shrinkage of the apparent volume and such has been observed in a number of cases. (111) Steam tends to maintain pore volume and to increase pore radius as illustrated in the third row. If the inner platelets move out to join the outer platelets, then A = 2 v=3 d = 3

This represents a threefold decrease in area, the same decrease as in the vacuum sintering example above. However, in this case of steam sintering, the pore volume does not change and, consequently, there is a threefold increase in pore width. (111’) Steam may also maintain pore volume by moving the inner platelets together to give A = 4 v=3 d = 14

Thus the area of the catalyst is reduced by one-third, the pore volume again is not changed, and the pore width is increased by 50%.

112

HERMAN E. RIES, JR.

Although the above picture is greatly oversimplified, it is not unreasonable t o think of certain cracking catalyst systems as platelet structures (Davidson, 13; Webb and Ehrhardt, 63). Perhaps other catalysts may be idealized in this manner. The mechanisms of various types of sintering, nevertheless, must eventually be solved in terms of molecular structure and crystal growth (Shapiro and Kolthoff, 56). Another observation, however, may be of significance in the general picture of sintering. With respect to the vacuum sintering of the TCC Beads it might be suggested that the fact that the pore structure remaining after sintering is unchanged could be explained by a nonuniform sintering of the bead as a whole. In other words complete sintering or total loss of pore volume might take place from the outside of the bead to the inside or from the inside out. This process at intermediate stages would give, respectively, shells of zero area or cores of zero area. Although this mechanism seemed rather unlikely in view of the long time periods of the heat treatments, an ingenious method for disproving this type of nonuniform sintering was devised by J. S. Melik of our laboratories. The method provides a technique for ph,ysically separating shell and core material at the desired point. The sintered beads are immersed in water and allowed to remain a short time until visual observation indicates that the water has radially permeated a shell section of the desired thickness. The beads are then quickly transferred to liquid nitrogen which freezes the water. Subsequent removal of the shell structure which has been severely weakened by the process is easily accomplished. I n the case of the vacuum-sintered beads it was found that the core areas were in good agreement with areas of the whole sintered beads. Nonuniform sintering with respect to the whole bead is thus not indicated. Other applications of this separation technique are obvious. Additional information related to sintering may be obtained from grinding experiments performed with the TCC Beads (Van Nordstrand, Kreger, and Ries, 62). I n brief it has been observed that manual grinding of TCC Beads to a rather fine powder causes an appreciable loss in area and pore volume, but in addition the appearance of an asymptotic type curve a t po. This probably indicates sintering induced by the localized frictional heat and appears to be somewhat similar to the vacuum type of sintering effect. It is believed that the small particles produced by grinding probably are responsible for the asymptotic approach to p o . Both the adsorption and desorption isotherms for the beads before grinding are almost perfectly horizontal as they approach the po line. A relatively recent catalyst preparation, Houdry Porous Beads, is compared with the virgin and steam-sintered TCC Beads in Fig. 11.

STRUCTURE AND SINTERING

0

0.2

PROPERTIES OF CRACKING CATALYSTS

113

0.4 0.6 0.8 RELATIVE PRESSURE, 9'Fb

FIG. 11. Nitrogen adsorption-desorption isotherms for virgin and steam-treated Socony TCC Bends and Houdry Porous Beads.

The Houdry Porous Bead catalyst has an average pore radius of 41.6 A., a value about twice that of 'the TCC Beads studied. The Houdry Beads were prepared with a system of large pores presumably to improve regeneration characteristics. Although a small amount of adsorption a t high relative pressures indicates the presence of some very large pores, the long steep portion of the isotherm preceding this indicates a con-

114

H E R M A N E. RIES, JR.

centration of pores in the region of the average pore radius. The area of the Houdry Porous Beads, 248 sq. m./g., and that of the second steam treated TCC Bead sample, 240 sq. m./g., are very similar. The Houdry Porous Beads may be compared with an earlier type of Houdry pelleted catalyst, S-46, in Fig. 12 and Table I. The isotherm contours for these 4 0 0 ,

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1 04 RELATIVE

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FIG. 12. Nitrogen adsorption-desorption isotherms for Houdry silica-alumina catalysts.

two materials bear a general resemblance. Since the pore volume of the 5-46 is smaller and its area greater than the corresponding values for the Houdry Porous Beads its average pore radius, 31.9 A., is considerably smaller than that of the Porous Beads, 41.6 A. Moreover 31.9 A. is a rather large pore radius value for a synthetic silica-alumina. b. Aerocat Microspheres. The effect of steam sintering is strikingly demonstrated in the case of the silica-alumina Aerocat Microspheres.

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

115

In Fig. 13 is illustrated the sharp contrast between the isotherms for the virgin and for the steam-treated Aerocat. The steam sintering treatment reduces the area from 700 to 85 sq. m./g., a decrease of almost 90%, whereas the pore volume has fallen only from 0.580 to 0.504 cc./g., a loss of approximately 13%. The average pore radius increases from 16.6 to 119 A., a sevenfold increase. The isotherm plots indicate that the pore size distribution is quite narrow in both cases with steep portions

0

0.2

04 06 RELATIVE PRESSURE. P/R.

08

1.0

FIG. 13. Nitrogen adsorption-desorption isotherms for virgin and steam-treated Aerocat silica-alumina microspheres.

of the desorption isotherms a t relative pressures that correspond to pore radius values close to the calculated averages. Furthermore, the PO line is met quite sharply by both the adsorption and desorption branches. The severe steam sintering has evidently produced relatively few pores that are much larger than the average. The narrowing of the hysteresis loop with increase in pore size is also demonstrated in Fig. 13. I n addition it is noted that there is a certain resemblance between the isotherm for the steam-sintered Aerocat and that for the silica aerogels discussed in See. IV. The hysteresis isotherms presented in Fig. 14 illustrate very clearly for the Aerocat catalyst the differences between sintering in vacuum and

116

HERMAN E. RIES, JR.

sintering in steam. Successive vacuum heat treatments reduce the area from 700 t o 413 and then t o 143 sq. m./g. The pore volumes which correspond t o adsorption values a t po decrease in almost direct proportion t o the decrease in area which corresponds t o the adsorption in the neighI

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FIG.14. Nitrogen adsorption-desorption isotherms for virgin, vacuum-sint red and steam-treated Aerocat silica-alumina microspheres.

borhood of 0.1 relative pressure. Thus the vacuum sintering does not change appreciably the average pore radius. The isotherm contours remain very similar during the vacuum sintering and the steep portion of the desorption branch is not displaced by a significant amount. How-

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

117

ever, in contrast, as indicated above, the isotherm for the steam treated material shows a major loss in area, a small loss in pore volume and a great displacement of the steep portion of the isotherm toward p,. The sintering curves, in which surface area is plotted against temperature of treatment, are presented in Fig. 15. These curves also demonstrate the wide difference between vacuum and steam sintering of Aerocat. The area value of 85 sq. m./g. obtained in steam at 621°C. (1150°F.) corresponds approximately to a vacuum sintering value in the

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FIG.15. Effect of temperature and steam on surface area of Aerocat Microspheres.

neighborhood of 600 sq. m./g. The vacuum sintering curve falls to zero area at about 1000°C. (1832'F.) whereas the steam curve extrapolates to zero area in the region of 700°C. (1292°F.). The area-temperature curves and the isotherms make clear the accelerated sintering produced by steam and the profound reorganization of pore structure effected in the case of silica-alumina catalysts. This reorganization or intimate interaction between steam or adsorbed water and the catalyst is not too surprising when it is considered that the catalysts are prepared in aqueous media, that surface water is necessary for activity (Hansford, 26) and that steam deactivates catalysts. This

118

HERMAN E. RIES, JR.

intimate interaction is further exemplified by the exchange that takes place between the oxygen of the catalyst and the oxygen of water in contact with the catalyst as recently demonstrated with oxygen isotopes at the Houdry Laboratories (Hindin and Mills, 31; Heinemann, 29). The latter experiments were performed a t 100°C. and extended for a period of one month. This temperature is obviously much less severe than that used in the steam sintering treatments reported above, although

c: 3

0

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I 08

LO

FIG. 16. Nitrogen adsorption-desorption isotherms for virgin and used Diakel.

the time periods for the isotope exchange experiments were much greater. Nevertheless in practically every case in which extensive oxygen exchange is found there is also observed a structural transformation in the form of area and pore size changes. c . Diukel. The structural differences between virgin and used Diakel are indicated in Fig. 16 and Table I. A virgin Iliakel, which had an area of 410 sq. m./g., fell t o 130 sq. m./g. during comlmercial use in a fluid cracking unit. The average pore radius increased from 21.2 to 43.7 A., a twofold enlargement. The isotherms indicate again that the increase in pore radius has been a uniform one since the hysteresis loop is displaced to the right and both branches are steep. The plots shown

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

119

in Figs. 17 and 18 are included for two reasons: in the first place they demonstrate the high degree of reproducibility obtained in duplicate determinations of complete adsorption-desorption isotherms, and secondly they illustrate in the insets the expanded plots of the high relative pressure region used for estimating V. for pore volume computation. The three sets of data plotted in Fig. 17 practically coincide throughout as do the two determinations of Fig. 18. The expanded I

0

0.2

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0.6

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RELATIVE PRESSURE. P/P-

FIG.17. Nitrogen adsorption-desorption isotherms for virgin Diakel.

plots in the inset curves show clearly the isotherms breaking away from the PO line. It will be noted that the Diakel isotherms do not approach PO as sharply as do those for the TCC Bead and Aerocat Microspheres. This asymptoticity may be caused by the presence of some very small particles or by some exceptionally large pores in Diakel (see Sec. II.5.b). 3. Clay T y p e Catalysts

A comparison of virgin, steam-treated, and used Fluid Filtrol catalysts is presented in Fig. 19 and Table I. The hysteresis isotherm for the virgin material has been discussed in some detail by Oulton (42). Virgin Filtrol has an area of 339 sq. m./g., which is considerably smaller than

120

HERMAN E. RIES, JR.

the values for most of the synthetic silica aluminas discussed above. Although its isotherm indicates the presence of some large pores and a rather broad distribution in the high relative pressure region it has a sufficient concentration of small pores t o yield an average pore radius of 24.2 A. The very broad hysteresis loop is probably characteristic of

L

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0.4 0.6 0.8 RELATIVE PRESSURE. P/PD

I

FIG.18. Nitrogen adsorption-deaorption isotherms for used Diakel.

this type of pore size distribution, but might possibly indicate more of the bottleneck type of pore geometry than is present in other catalysts studied (52). Commercial use of a Fluid Filtrol catalyst effected more than a twofold decrease in area, 339 to 141 sq. m./g., in the case of the Equilibrium E sample and an approximately corresponding increase in average pore radius, 24.2 to 47.1 A. The isotherm for the steam-treated sample

STRUCTURE AND SINTERINQ PROPERTIES OF CRACKINQ CATALYSTS

121

shown in Fig. 19 is remarkably similar to that of the used material, both in contour and calculated structure characteristics. Its average pore radius is 45.4 A. compared with 47.1 A. for the used catalyst. Again it is clear that steam sintering alters the physical structure of the catalyst, that is, its area, pore volume, and pore size, in approximately the same manner as accomplished by commercial use. Steam is used extensively in catalytic cracking with Filtrol clay catalysts for stripping and for

,

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0.2

04

06

48

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RELATIVE PRESSURE. P/pO

FIG.19. Nitrogen adsorption-desorption isotherms for virgin, steam-treated, and used Fluid Filtrol.

hydration of the catalyst. A considerable quantity of steam is also produced during the regeneration process. Consequently it appears that steam may be the principal factor in the physical deterioration of the clay catalysts as well as for the synthetic silica-aluminas. The significance of steam in catalyst deterioration is supported indirectly in Fig. 20. Plotted here are three equilibrium used Filtrols from three different refineries whose charging stocks differ rather widely, for example, with respect to poisonous contaminants. The three refineries have in common the use of steam and approximately the same tem-

122

HERMAN E. RIES, JR.

perature range. The three adsorption-desorption isotherms practically coincide giving areas of 135, 136, and 141 sq. m./g., Val-ues almost within experimental error. The physical deterioration obviously has been almost identical and may be attributed tentatively to steam and temperature factors. Steam sintering at a higher temperature, 627°C. (1180°F.) is compared with that which occurs a t 570°C. (1058°F.) i n Fig. 21. The higher

02

04

06

0-8

19

RELATIVE PRESSURE. P/Po

FIQ. 20. Nitrogen adsorption-desorption isotherms for virgin and three used Fluid Filtrols.

temperature treatment lowers the area from 169 t o 113 sq. m./g. and introduces a curve in the desorption branch between 0.9 and 1.0 relative pressure that approaches closely that of the used catalyst. The lower temperature steam treatment increases the average pore radius from 24.2 to 45.4 A. and the higher temperature treatment increases it further to 66.1 A. radius. A comparison of Filtrol sintering curves obtained in vacuum and in steam is available in Fig. 22. Area values are plotted against temperature of treatment. It is apparent that steam accelerates the sintering

STRUCTURE AND SINTERING PROPERTIES OF CRACKING

CATALYSTS

123

process. The curves indicate that in the range of about 600-700OC. (1112-1 292’F.) the area values for the steam sintered samples are approximately 200 sq. m./g. smaller than those which result from corresponding treatments in vacuum. The steam sintering curve falls sharply in this range whereas the vacuum sintering curve shows only a small fall-off. Another Fluid Filtrol cracking catalyst designated “Filtrol SR ’’ and claimed t o be sulfur resistant has been introduced into commercial use

oo 02

0.4

06

08

10

RELATIVE PRESSURE. P/Po

FIQ.21.

Nitrogen adsorption-desorption isotherms for steam-treated Fluid Filtrol.

recently. It is presumably prepared from a different type of clay, and its physical properties as shown by adsorption measurements vary considerably from those of the other Filtrol catalyst discussed above (Fig. 23 and Table I). The area of Filtrol SR, 164 sq. m./g., is about one-half the value found for the earlier Filtrol product, 339 sq. m./g. Its average pore radius is 37.7 A. compared t o 24.2 A. for the ordinary Filtrol. However, of more interest, is the observation that the hysteresis isotherm plotted in Fig. 23 shows no steep portion in the 0.4 to 0.6 relative pressure range, Both the adsorption and desorption branches of this isotherm are unusually smooth and free of any apparent indication of inflection

124

H E R M A N E. RIES, J R .

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re*m.LU

-1

i I

1

FIG.22. Effect of temperature and steam on surface area of Fluid Filtrol.

IV. STRUCTURE AND SINTERING PROPERTIES OF VARIOUS FORMS OF SILICAAND RELATEDMATERIALS Silica is present to an important extent in all of the cracking catalysts discussed in the preceding sections. The state or form of the silica may vary considerably. For many of these catalytic materials silica undoubtedly provides the basic physical structure and in add:ition is intimately involved in the surface chemistry of the catalyst. It therefore seems

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

125

important to consider carefully the structure of various forms of silica. Furthermore the relative simplicity. of the chemistry and physics of silica compared with the more complex catalyst systems will facilitate understanding the relationship between structure and adsorption characteristics. Principal emphasis in the following discussion is placed on

t I

I

1

I

I

I

I

1 I

TABLE I1 Area, Pore Volume and Average Pore Radius for Various Forms of Silica and Some Related Materials

V,, cc. Adsorbed Experiment in Monolayer Number on 1 g. (ig. wt.) (Hysteresis) Adsorbent (BET)

BET Area, sq. m./g. ig. wt.

Per p,, Average Cent V., Gas Pore Volume Pore Volatile Adsorbed at (from V , ) Radius, A . Matter PO,cc./g. cc./g. ( r = 2P ,/A)

x

$

z

z Davison silica xerogel Silica Aerogel “ S” Silica bead aerogel Santocel “ C ” Silica Aerogel “S,”wetted Linde nonporous silica Alumina Titania Diatomaceous earth (Ci-:it~337)

173 223 3697 2778 4056 2758 215 2482

2507

185.8 182.9 115.8 49.7 186.8 41.4 32.5 2.4 ”5. -1

809 796

504 216 813 180 141 10

22

7.5 11 .o 7.2 6.1 5.1 1.7 5.4 0.5 5.3

285 2525 3700 2550 4-25 193 -

0.44 3.91 5.71 3.94 0.66 0.30 -

11 98 227 365 16 42 -

-

M

a M m Y

?

STRUCTURE AND SINTERINQ

PROPERTIES OF CRACKING CATALYSTS

127

to the classification of Brunauer, Deming, Deming, and Teller (6). The Davison gel has an area value of 809 sq. m./g. (Table 11). Its hysteresis isotherm as shown in Fig. 24 has a horizontal portion above 0.5 relative pressure which intersects the p o line a t a right angle. This indicates an almost complete absence of large pores. Furthermore, very little hysteresis is observed in the entire adsorption-desorption plot (see also Joyner, Weinberger, and Montgomery, 35). It is also apparent

0

02

0.4 RELATIVE

FIG. 24. in vaeuo.

0.6

08

10

PRESSURE. P/P.

Nitrogen adsorption-desorption isotherms for silica xerogel sintered

that approximately one and one-half monolayers eff ectively'Jfil1 the pores. An average pore radius obtained by means of the relationship, T = 2 V / A , is 11 A. According to pore radius values calculated by means of the Kelvin equation and not corrected for adsorbed layers, one must attribute the surface area to pores having radii below 13 A., the value which corresponds to a relative pressure of 0.5. The upper limit of pore radius corrected for two adsorbed layers is approximately 21 A. Adsorption-desorption isotherms obtained during a series of vacuum sintering treatments are also presented in Fig. 24. It is apparent that the contour of the isotherm is changed only slightly as sintering proceeds.

128

HERMAN E. RIES, JR.

Adsorption decreases proportionately throughout the isotherm. The pore volume-area ratio thus remains effectively constant and there is no appreciable change in average pore radius even after sin tering to an area of less than 100 sq. m./g. (Van Nordstrand, Kreger: and Ries, 62). Somewhat similar observations have been reported by Milligan and Rachford (39) and by Tamele, Byck, Ronay, and Vinograd (5’9). Throughout the sintering series of Fig. 24 the isotherms remain almost flat above 0.5 relative pressure and intersect the p o line at right angles. There is thus no indication of the development of large pores. However it appears that sintering reduces, almost to the point of elimin,ation, the small hysteresis loop observed in the initial isotherm. This could be related to a disappearance of some bottleneck type pore structure or possibly to the disappearance of some of the slightly larger pores. Nevertheless the general picture remains very similar to that observed in the sintering studies on synthetic silica-alumina catalysts, TCC Bead and Aerocat Microspheres (Secs. III.2.a, III.2.b) ; that is, vacuum sintering reduce8 the pore volume in proportion to the area decrease arid the pores that remain are essentially unchanged. In preliminary steam sintering experiments with Davison silica gel it appears that sintering is accelerated considerably by the presence of steam a t atmospheric pressure. Furthermore, there is a small but significant increase in pore radius. The small hysteresis loop that almost disappears on vacuum sintering is apparently maintained or perhaps slightly broadened on sintering in steam. Detailed data must await the completion of the steam series. 2. Silica Aerogel

a. Aerogel ‘‘8.” The structure of the large pore Aerogel “S” stands in sharp contrast to that of the small pore Davison silica gel. Nevertheless, the BET area of the aerogel is 796 sq. m./g., a value almost the same as that of the Davison gel, 809 sq. m./g. The contrast in pore structure is strikingly apparent in a comparison of the isotherm for the aerogel plotted in Fig. 25 and that of the Davison xerogel in Fig. 24. The low relative pressure region of the aerogel isotherm is similar to that of the xerogel and accounts for the similarity of a.rea values. However, the aerogel isotherm continues to rise throughout the middle relative pressure region and then rises very steeply in. the neighborhood of 0.9 relative pressure, whereas the Davison gel rshows a relatively negligible adsorption above 0.5 relative pressure. Approximately one and one-half monolayers effectively fill the pores of the Davison material, but many times this amount is required to fill the large pores of the aerogel. At a very high adsorption value, 2525 cc.,/g., the desorption

STRUCTURE AND SINTERING

PROPERTIES OF CRACKING

CATALYSTS

129

isotherm of the aerogel breaks sharply from the pa line and provides a value for the computation of pore volume. The aerogel thus gives a Type IV isotherm or possibly a combination of Type I1 and Type IV (Brunauer, Deming, Deming, and Teller, 6).

I

0

I 0.2

I

1 0.4

I

I

1

06

I

I

0.8

RELATIVE PRESSURE. P/Pe

FIG.25. Effect of wetting on the nitrogen adsorption-desorption isotherm of a silica aerogel.

An average pore radius value obtained from the pore volume-area relationship is 98 A. for the Aerogel “S” compared with 11 A. for the Davison xerogel. The steep portion of the aerogel isotherm occurs in the region of 0.9 relative pressure and corresponds, according t o the Kelvin equation, to a rather narrow distribution of pore radii in the neighborhood of 100 A. The pore volume of the Aerogel “S” is 3.91 cc./g., and that of the Davison silica gel, 0.44 cc./g. The larger pores of the aerogel compared with those of the xerogel probably account for the aerogel’s greater fragility, its greater light scattering to give a white rather than a

130

HERMAN E. RIES, JR.

clear transparent appearance, and its lower hygroscopicity attributable to pores too large for effective capillary condensation. The large pore structure picture of aerogels is also supported by elect(ron micrographs obtained by J. A. Brown in our laboratories. Micrographs were made by lightly shadow casting the specimen with gold. Ultimate

RELATIVE PRESSURE. P/p.

FIG.26.

Nitrogen adsorption-desorption isotherm for silica, bead aerogel.

particles and “pores” of the correct order of magnitude were observed a t the thin edges of the aerogel structure. No structure of this type was found in similar electron microscopic examinations of the small pore Davison gel. b. Bead Aerogel. An adsorption-desorption isotherm which is quite similar to that of the Aerogel “S” is obtained for a silica bead aerogel and is shown in Fig. 26. The aerogel bead was prepared from a Socony-

STRUCTURE AND SINTERING PROPERTIES OF CRACKING

CATALYSTS

131

Vacuum silica hydrogel bead by Henry Erickson of our laboratories. The area of this material, 504 sq. m./g., is somewhat lower than that of Aerogel “S,”796 sq. m./g., and its pore volume, 5.71 cc./g., obtained from its V , value, is greater than that of Aerogel “S,”3.91 cc./g. Thus its average pore radius, 227 A., is considerably larger. The steep portion of the bead aerogel isotherm is consequently displaced to the right toward the greater pore radius region. The long narrow hysteresis loop for this material is of interest because of the large pore structure. The desorption and adsorption curves are in good agreement at the lower relative pressures, and thus negligible errors have been introduced during the lengthy determination. c. Santocel “ C.” Consideration of a commercial aerogel, Santocel “C,” is also of interest. However, the sample studied has an area of 216 sq. m./g., a value well below those of the aerogels discussed above. Nevertheless its isotherm is very similar in contour to those of the above .aerogels and it has the large pore structure typical of aerogel preparations. Its pore volume as obtained from V,, the volume adsorbed at the saturation pressure, is 3.94 cc./g., practically identical t o that of the Aerogel “S,”3.91 cc./g. However, a greater average pore radius, 365 A., is obtained for Santocel “ C ” since its area, 216 sq. m./g., is smaller than that of the Aerogel “S,” 796 sq. m./g. The steep portion of the Santocel isotherm occurs a t a higher relative pressure and corresponds to its greater pore radius. d. Wetted Aerogel. I n a study designed to determine the effect of shrinkage on the pore structure and area of aerogels, the Aerogel “S”was wetted with water (Johnson and Ries, 33). This was accomplished simply by immersion of the aerogel in water for one and one-half hours at room temperature followed by drying for twelve hours a t 110°C. (230°F.) with a final drying for two hours at 593°C. (1100°F.). The adsorption-desorption isotherm for the resulting material is plotted with that of the original aerogel in Fig. 25. The wetting treatment has obviously produced a xerogel type structure. The pore volume suffers a sixfold contraction, falling from 3.91 to 0.66 cc./g. The BET area remains essentially the same, 813 sq. m./g. for the wetted material compared to 796 sq. m./g. for the original aerogel, while the average pore radius decreases from 98 A. to 16 A. The isotherms demonstrate rather narrow pore size distributions for both materials. Uniformity of both structures is also indicated by optical microscope observations. Apparently the shrinkage is uniform and simply involves a tighter packing of the ultimate particles (fibers, platelets, etc.) without crystallization or sintering and with point contacts of the particles retained. Since the BET area method is generally accepted for large pore structures, this

132

HERMAN

E. RIES, JR.

study tends t o support the applicability of the BET method to small pore systems. 3. Nonporous Silica Linde silica is of special interest because, for a nclnporous material, it has an unusually high area, 180 sq. m./g. A very small particle size obviously accounts for its extensive area in spite of its nonporous structure. The duplicate isotherm determinations plotted in Fig. 27

0

02

a4

06

0.8

10

RELATIVE PRESSURE. P/Po

FIG.27. Kitrogen adsorption-desorption isotherm for Linde silica.

are in good agreement throughout and yield area values of 179.8 and 180.2 sq. m./g. Linde silica presents a typical Type I1 isotherm with its sigmoid curve and its asymptotic approach to po. It is an excellent example of a nonporous adsorbent showing effectively no adsorption-desorption hysteresis and no steep portions in the intermediate relative pressure region. The Linde silica isotherms are very similar t o those obtained for nonporous titania of relatively low surface area (Harkins and Jura, 28; Ries, Johnson, and Melik, 47). The surface area of the Linde material is so much greater than that of the titania, whose area is 10 to 12 sq. m./g., that it should be more useful for a variety of studies such as the important relative molecular area determinations for various adsorbates.

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

133

A general similarity in isotherm contour observed for Linde silica and the aerogels discussed above may be significant. The principal difference is that the Linde isotherm approaches Po asymptotically whereas the aerogels show limits in adsorption at Po sufficiently well defined for pore volume calculations. The interpretation of the isotherm similarity between large pore and nonporous structures is undoubtedly related to the fact that in the large pore structure of the aerogel the solid surfaces or “ultimate particles” are separated in a manner somewhat similar to that found in the loosely aggregated independent particles of the nonporous preparation. Appreciable overlapping of the solid surface forces in aerogel pores is unlikely and the solid surfaces probably behave almost as independent particles. Nevertheless the aerogel has a measurable pore volume filled at Po, whereas the Linde material apparently does not have a well-defined pore volume for its loosely formed aggregates. Thus the Linde isotherm exhibits an asymptotic approach to the saturation pressure, po. Adsorption studies on packed or pelleted Linde silica should be informative in this connection. In further distinguishing the aerogel structure from the aggregates of nonporous type powders, perhaps one should think of the aerogel ultimate particles or fibers as joined in point contacts by chemical bonds that may be weak but that give a certain physical stability to the structure. There is, of course, some basis for considering a large gel particle as simply a large polymer type molecule.

4. Adsorption per Unit Surface for Three Forms of Silica The V J V , or adsorption per unit surface area plots are presented in Fig. 28 for a silica aerogel, a silica xerogel, and a nonporous silica. These plots are of particular interest in the silica series since undoubtedly the surface compositions of the three forms of silica are very similar. Certainly the small differences in surface chemistry that may exist should not be of appreciable significance in the low-temperature adsorption of nitrogen. Consequently appreciable differences observed in the V,/V, curves may be attributed to differences in pore structure characteristics since any differences in surface forces or adsorbate-adsorbent interaction effects should be practically negligible for such a series. In general the effect of pore structure on adsorption in the low relative pressure or monolayer region has been considered a minor one and the small differences observed have been ascribed t o a certain specificity of the van der Walls adsorption (Ries, Van Nordstrand, and Teter, 52). Apparently for the silicas studied here (Fig. 28) the adsorption differences at low relative pressures are not of great significance. Physically adsorbed monolayers in general are statistically complete in the neighborhood of 0.1 relative pressure. At this pressure Davison

134

HERMAN E . RIES, JR.

silica gel exhibits the most adsorption in terms of oaonolayers, 1.1, whereas the large pore aerogel has adsorbed a smaller fraction of a monolayer; namely, 0.9. A statistical monolayer is completed (Va/Vm = 1) on the Davison gel at approximately 0.07 relative pressure and at 0.14 relative pressure for the aerogel. The small pores of the Davison gel in this case show at low relative pressures a slightly greater adsorption

i

0

0.2

0.4 0.6 RELATIVE PRESSURE. P/pO

0.8

/'

I

FIG.28. Isotherms for nitrogen adsorption per unit surface.

per unit surface than the large pores of the aerogel. However the Linde nonporous silica data are intermediate in this region. It is also clear in Fig. 28 that 1.5 monolayers are sufficient to fill completely the pores of the Davison gel, but many more layers are required for the aerogel structure and for the interparticle spaces or loosely bound aggregates of the Linde silica. It thus appears that pore structure may be a factor affecting to a slight extent adsorption at low relative pressures but is a

STRUCTURE AND SINTERING PROPERTIES O F CRACKING CATALYSTS

135

factor to a really significant extent in the high relative pressure regions (Ries, Van Nordstrand, Johnson, and Bauermeister, 50). 5. S‘intering Curves for Silicas

Four silica samples, Aerogel “S,” Santocel “C,” a Davison xerogel, and Linde silica have been subjected to vacuum sintering treatments. I

-c

100

I

400

I l l

IW

800

I

700

I

MD

I

900

I

ID00

I

I100

.E

TEMPERATURE

FIG.29. Effect of temperature on surface area of various forms of silica. ments: 12 hours in vacuum a t indicated temperatures.

Treat-

The sintering curves, in which area values are plotted against temperature of treatment, are presented in Fig. 29. Unless otherwise indicated the heat treatments extended for twelve hours. As pointed out in Sec. I11 such area values do not represent final states of sintering a t the indicated temperatures because of the limited time period of the treatment. According to the data of these plots the Davison xerogel demonstrates relatively good thermal stability compared with a wide variety of high

136

HERMAN E. RIES, JR.

area catalytic materials (compare with Fig. 3). Its ‘high area small pore structure appears to be quite stable up to 600°C. (1112°F.). It then sinters a t a slightly greater rate a t successively higher temperatures. Zero area is reached in the neighborhood of 1050°C. (1922°F.). The structure change which takes place on vacuum sintering #of this particular silica gel is of considerable interest as pointed out above (Sec. IV.l). Apparently the pore distribution remains essentially unchanged throughout the sintering to zero area. Pores disappear but the radii of those remaining evidently neither increase nor decrease significantly. Silica Aerogel “ S ” demonstrates by far the greatest thermal stability of all the materials of this series. Not only has this aerogel sample the greatest initial surface area, but throughout its sintering curve its area remains well above those of the cracking catalysts and the other forms of silica. It retains a good percentage of its surface up to 1100°C. (2012”F.), and its area falls only to 200 sq. m./g. after four successive twelve-hour treatments a t this temperature. No other material sho’wnon the sintering charts retains significant area in the neighborhood of 1100°C. The large pores of the aerogel may be related to its remarkable thermal stability; that is, the solid surfaces of the “pores” are so widely separated that their approach and “fusion” on heating is difficult and slow. It seems somewhat paradoxical, however, that in an abbreviated sintering series for a silica aerogel, in which hysteresis isotherms were determined, indications are that the larger pores in this structure disappeared first (Van Nordstrand, Kreger, and Ries, G2). Vacuum sintering of the aerogel produced a pronounced decrease in the average pore radius. Complete interpretation must await more extensive studies. Somewhat more unexpected than the relatively low area of the Santocel “C,” the other aerogel studied in this series, are its sintering data plotted in Fig. 29. The sintering curve shows its area to fall sharply above 800°C. (1472°F.) and to reach zero area a t 900°C. (1652°F.). Although its range of thermal stability may be sufficient for most practical purposes, a significant difference between the tlvo aerogels is indicated. Perhaps this difference may be related to impurities or possibly to basic differences in composition. Composition diifferences or contamination may of course account for sintering differences in many cases. The nonporous finely divided Linde silica demonatrates remarkable thermal stability up to 900°C. (1652°F.) and loses very little area a t 1000°C. (1832°F.). Following a twelve-hour treatment a t lOOO”C., in addition to the many preceding treatments a t lower temperatures, the Linde area is still above 125 sq. m./g. Superior thermal stability was anticipated for this material since an individual nonporous particle in

STRUCTURE AND SINTERING PROPERTIES OF CRACKING CATALYSTS

137

general can lose very little area itself and the loose contacts between such particles should make particle fusion relatively difficult. The fact that the silicas discussed above sinter t o zero area far below the silica “melting point ” range, 1500-1700°C. (2732-3092”F.), is of significance in solid structure and surface diffusion considerations (see also Shapiro and Iiolthoff, 56). Another relationship that must be clarified in further sintering studies is the effect of previous treatment or the history of the silica sample on its sintering characteristics. For example it has been found in severaI cases that an initial heat treatment a t a high temperature sinters the sample to an area considerably lower than that obtained with the same treatment which has been preceded by several treatments a t lower temperatures. This apparent anomaly may well be related to the degree of hydration and the rate of dehydration; that is, there may be an effective accelerated “steam” sintering in the case of samples not previously heated stepwise to the higher temperature levels. In other words removal of water from a sample a t low temperatures may protect the sample from excessive sintering effected by this “inherent steam” if present a t higher temperatures. However, there may also be a stabilization process occurring at lower temperatures which strengthens the structure. 6. A lumin a

An alumina adsorbent is included in this discussion since not only is alumina an important constituent of the synthetic silica-alumina catalysts discussed a t length above but because it is a major constituent of many other catalysts (Haensel and Sterba, 25) and an important catalyst in its own right. Although the structure of alumina in combination with silica may be entirely different from that of alumina alone, the structural properties of free alumina must be studied for a complete understanding of its various combinations. For example, it may be possible t o detect free alumina by adsorption techniques if the alumina structure is sufficiently well defined. A Harshaw activated alumina was chosen for the preliminary investigation. An adsorption-desorption isotherm for this material is presented in Fig. 30. It has a surface area of 141 sq. m./g. and a pore volume of 0.30 cc./g. The contour of the isotherm is rather unusual in that both the adsorption and desorption branches have two steep portions above 0.5 relative pressure. Apparently there is a tendency for two hysteresis loops to develop in the 0.5 to 1.0 relative pressure region. Hysteresis curves with similar contours have been obtained for active magnesias of the same area range by Zettlemoyer and Walker (65). Isotherms of this type indicate a concentration of two different pore sizes or pore size ranges. An average pore radius based

138

HERMAN E. RIES, JR.

simply on total pore volume and surface area is 42 A. Very little hysteresis is observed in the smaller pore range; that is, in the region below 0.6 relative pressure. A relatively large area value, 70 sq. m./g., has been obtained for this alumina by the stearic acid adsorption technique (Ries, Johnson, and Melik, 47), but in view of studies with other compositions having

280

c

240

i

-

0

ta

20 0

0

ffa

0:

160

G is

0

0.2

OL)

RELATIVE

0.6 0.6 PRESSURE (P/P.\

FIG.30. Adsorption-desorption isotherm for alumina.

large pores it now appears that this is related more to the basicity of alumina than to its large pore structure. The adsorption of stearic acid from solution by alumina may be considered t o be a type of chemisorption (see also Russell and Cochran, 54). Preliminary vacuum sintering data on this alumina sample indicate fairly good thermal stability up to approximately 900°C. (1652°F.). In comparison a magnesia adsorbent which has an area of 195 sq. m./g. exhibited relatively poor thermal stability. Its area began to fall sharply in the neighborhood of 500°C. (932°F.).

T

7. Titania The hysteresis isotherm for titania, to which earlier reference has been made, is plotted in Fig. 31. Although there is some interest in titania as a catalyst it is included here principally because in adsorption research it may be considered the classical example of a nonporous finely I

70

2 -

100

60

Gt L

50

160

n 0 a P w

d

40

v)

0 0

d W

m

30

0 0

W

5

20

100

2482

-I

>

{

0 ADSORPTION

0 DESORPTION

90

I0 0.996 ID IPmI

0

0.2

a4

0.6

0.8

1.0

RELATIVE PRESSURE (PIP.)

FIG.31.

Nitrogen adsorption-desorption isotherm for titanium dioxide.

divided solid (Harkins and Jura, 27). It has been employed, therefore, for the determination of molecular areas or oriented cross-sections of various adsorbates (see Beebe, Beckwith, and Honig, 3). The titania used here corresponds closely to that investigated by Harkins and Jura (28). However, as pointed out above, the nonporous finely divided Linde silica has certain advantages principally because it has a greater area. The surface area value for the titania is 10 sq. m./g. compared with 180 sq. m./g. for the Linde material. The isotherm for titania shows practically no hysteresis and provides an excellent example of a

140

HICRMAN E. RIES,

JR.

Type I1 isotherm with an extended asymptotic approach to p o . The similarity of this isotherm and that of the Linde silica is apparent a t once in a compa.rison of Fig. 27 with Fig. 31. 8. Diatomaceous Earth

The structure of diatomaceous earth is of interest for several reasons. In the first place it is a form of silica quite different from those discussed 280

P

240

5s

z 2

-

3

ZOO

0 K W L

f

16C

v)

d

0

0W

p

12c

0

I0 n

a W

5

ec

-1 0

2107

{

0 ADSORPTION 0 OESORPTION

4c

0

0.2

QI RELATIVE

06 od PRESSURE (PIP.)

FIG.32. Nitrogen adsorption-desorption isotherm for diatomaceous earth.

above. Diatomaceous earths are the siliceous remains of microscopic marine organisms. Secondly diatomaceous earth is widely used as a catalyst support. In the third place its adsorption-desorption isotherm demonstrates a rather unexpected hysteresis effect. Finally an electron microscope study of this material illustrates the ap:plication of another technique t o the study of catalysts in general and pore structure in particular. The diatomaceous earth studied was obtained from the JohnsManville Corporation and is known as Celite 337. I n earlier studies in

STRUCTURE A N D SINTERING

PROPERTIES

141

OF CRACKING CATALYSTS

our laboratories it was indicated that diatomaceous earth supports of low area, relatively large pores, and an irreguIar particle geometry are responsible for the development and maintenance of the high area, small pore structure of certain supported catalysts (liies, Johiison, and Melik, 47). The area of Celite 337 as determined by low-temperature nitrogen adsorption is 22 sq. m./g., a value far below those obtained for the silica gels and the rioriporous silica discussed above. The hysteresis isotherm for the diatomaceous earth is plotted in Fig. 32. An asymptotic approach to PO, similar tjo that for nonporous or large pore systems, is apparent.

-

2 a2

I

a4 RELATIVE

I

I a6 PRESSURE

0.8

1I LO

(PI81

FIG.33. Nitrogen atlsorptioii-desorption isotherm for diatomaceous earth.

The scale of Fig. 32, however, is such that the definition of the curve in the lower relative pressure region is not satisfactory. Therefore, this portion of the curve is plotted on a greatly expanded scale in Fig. 33. Curiously enough, in the latter figure, the diatomaceous earth shows a hysteresis effect in the 0.4 to 0.6 relative pressure region, a n effect that is typical of so many synthetic adsorbents. The desorption curve rejoins the adsorption branch rather sharply near 0.5 relative pressure. The presence of some small pore structure is thus indicated. Considerable information on the structure of the diatomaceous earth is obtained by means of electron micarowopic examination. I t was hoped that the electron micrographs might indicate the presence of the small pore structure suggested by the adsorption-desorption isotherm. Immediately apparent, however, is a large pore structure with pore

142

HERMAN E. RIER, JR.

FIG.34. IClcrtron micrograph of di;ttomac~ouscartli sliowirig widc writsty partick size and shape.

iri

FIG.35. Electron micrograph of tlirttomaceoiis rarth; an i?xample of regularity in the circular pore pattern.

STHUCTITIZE r \ N D S I N T E R I N G I’ROPEItTIES O F CRACKING CATALYSTS

143

diameters from somewhat less than 1000 A. to approximately 20,000 A. An irregular part,icle geometry is also observed. Pore distribution studies on a caloscly related diatomaceous earth, .Johns-hhiville Celite 296, by ltitter and I h k e ( 5 3 ) show a predominance of pore diameters i n the same range. A mercury porosimeter developed by Ritter and Drake was employed ill their study. The micrograph showir i n Fig. 34 is one of relatively low magnification :mi thus gives L: rather representative picture of the various particle

FIG.36. 15lectron niirrogrnpli of dintomacrous earth; nn example of irregularity in thr rircii1:ir pore pattern.

types present. It was obtaiiied in 1941 through the kiudness of E. F. Carman and E. Fullam, then with the Interchemical Corporation Laboratories. All the remaining micrographs were obtained in our own laboratories by .J. A. Bro\i711, using a RCA-EMli instrument. In Fig. 34 it is immediately apparent that there is a wide variety of particle sizes and shapes and that there is great irregularity in the external contours of the particles. T o this irregularity may tie attributed the poor packing properties of tliatomac.coris earth and possibly as a (.onsequence a considerable dispersion of catalyst material deposited on such structures.

144

H E R M A N IG. RIES, J H .

I t may he observed in Fig. 34 that, a major port,ion of the centmlly located particle having large circbular pores is transp.~rentto electrons. This indicates an extremely small partlick thickness probably in t,he neighborhood of 400 A. An approximate range of circwlar pore sizes observed in this investigation is from several hundred t o several thousand angstroms. Approximately 75 yo of the particles observed were porous and had structures of the general type illustrated i n the micrographs.

FIG.37. Idlectron micrograph of tlintomnceoris earth with hexagonal structures containing small pore formations.

A recent study of diatomaceous earth by Anderson, McCartney, Hall, and Hofer (1) invludes many electron micrographs of low magnificat,ion, which, of murse, provide more information on particle size and shape than do the high magnification pict,ures. A striking example of a regular spatial distrihution of small ('ircult~r pores is presented in Fig. 35. Nature has produced all most^ perfect, order in the pore structme of this particle. Furthermore many of bhe pore cross sectioiis approximate perfect, circularity. I n contrast to t,his pattern Fig. 3G presentasan extremely irregular arrangement of circ.ular pores. Careful examination discloses no periodicity or regularit,y what-

S T R U C T U R E A N D R I N T E R I N G P R O P E R T I E S O F CItACKING CATALYSTS

145

soever in the pattern. The pores are approximately 1000 A. in diameter. I n Fig. 37 is illustrated a multiple hexagon system c*ompletelyfilled with a structure caontaining many small circular pores. According to this micrograph the hexagon riblike structure could he either a denser material or a thicker material. Application of the shadow casting technique would probably decide the question since thick or ridgelike structures

FIG.38. Electron micrograph of diatomaceous earth showing pores containing pores which in turn contain smallcr pores.

mould yield shadow effects, whereas a flat denser material obviously would not. The electron micrograph of Fig. 38 dcrnonstrates the presence of pores containing other pores which in turn contain smaller pores. I n the central pore strurture, which would appear as a single pore a t low magnification, are nine smaller pores, and these cwntain still smaller pore structures. The smallest of these pores is several hundred angstroms in diameter. The obvious implication is that such pores may contain still smaller structures, not visible perhaps because of the low density of their framework to electrons or because the width of the solid portion is below the limit of resolution of the eleetron microscope. Smaller pores wonld

146

HEHM.\N

13:.

RIKS, J H .

awount for the adsorption-desorpt ion hysteresis effecat observed in the intermediate relative pressure region and for some of the area of the tliatomacwus earth as discussed above.

v.

SIJMMARY

AND

CONCLIJSION

Stru(turLt1 similarities and differences for representative cwtc*king vatalysts and caertain related materials are clearly demonstrated by low temperature nitrogen adsorption techniques. Surface area, pore volume, and an average pore radius have been determined for a series of c.rac+king vatalysts d1ic.h inrludes Xalco and DA-5 silira-magnesia cbatalysts, Socony T C C Beads, Aerocat Micmqheres, Diakel, Houdry Porous Beads and Houdry synthetic silicaa-alumina catalysts, Fluid Filtrol and pelleted d a y catalysts, as well as various forms of silica. Struc*ture characteristics of the silica-magnesia, silica-alumina, and clay catalysts studied are widely different. Practically the entire area of the virgin silica-magnesia cstdysts is cwitrihuted by the very srrtallest of pores i n the 10 t o 1.5 A. average pore radius range. The syntjheticb silic~n-:tluminn vatalysts are composed of pores appreciably larger, h u t with a \ ~ r : t g e pore radius values almost excslusively in the small pore range, 13 to 23 A . The clay ('atalysts have a much broader pore size distribution whivh includes pores (witriderably greater in radius. With respect to extent of surfwe, the w e a vstlues generally arc greater for the synthctir m:itei*ials t h s n for the clays. Vacuum sintering and steam sintering properties, as well as the eff ect of cwmmercial use, have been investigated for typicd catalysts. The sintering curves obtained in vacuum are quite different for the silicamagnesia, silira-alumina, and clay catalysts. In the preseiwe of steam the sintering process is accelerated for the three types of catalysts. Apparently the silica-magnesia preparations are somewhat, less sensitive to steam than the silica-alumina and clay catalysts. Steam sintering effects a reorganization of the silica-alumina and d a y structures in a relatively profound manner to increase pore size, whereas the pore structure of the silica-magnesia catalysts is not changed significantly. For the vacuum-sintered catalysts the structure of the pores remaining is essentially the same as their initial structure. l'se of the cat,alyst,s i n cxommercial cracking processes causes an increase in pore radius and related structure changes that are strikingly similar t o those produced by small scale steam treatments. It thus appears that steam a t elevated temperatures may he the principal factor in the de1,erioration of the physicd striwture of cracking caatalysts during commercial use. Adsorpt ion-desorption isotherms demonstrate the sharp (Boiitrstst between a small pore silica xerogel and a large pore aerogel structure.

S T R U C T U l t E .tND S I N T E R I N G PROPERTIES O F CRACKING C.4TALTSTS

147

The hysteresis isotherm for a nonporous Liiide silica presents a n excellent example of the low temperature adsorption properties of a high area notiporous solid. An investigation of sintering properties reveals that a silica aerogel has the greatest thermal stability of all the materials in the series. Hysteresis isotherms determined throughout a vacuum sintering study of Davison silica gel indicate that the pores which do not disappear maintain their original structure. The use of electroil microscopy as an aid i l l structure studies of catalysts is illustrated i n the case of another form of silica; namely, diatomaceous earth. Many other tec*hniques such as chemisorption for surface composition and surface acidity, x-ray arid electron diffraction, thermal analysis arid magnetic susceptibility are gaining wider applic*ation in catalyst research. Elevated temperature adsorption studies of yea(*tants and producats are still in their infancy. The thorough irivestigatioii of adsorption properties along with information from other sources will improve our uiiderstandiiig of catalysts and will eventually provide a means for the design and preparation of catalytic materials which possess the desired properties.

ACKNOWLEDGMENT The author rxprcsses his appreciation to J. A . 13rowri, hl. F. 1,. Johnson, LV. E. Kreger, J. S. hlelik, and It. A. Van Nordstrand for their collaboration throughout these studies. He also thanks J. W.Teter, B. S. Friedman, and L. 15. Olson of our lahoratorics and S. H. I h u w of C~ornell University for helpful suggestions, 11. D . Duncan for the preparation of the figures, and the management of the Sinrlair Refining Company for permission to publish this work.

REFERENCES 1. Anderson, 11. U., hlcCartnry, J . T., Hall, IV. K., anti Ifofrr, L. J. E., I d . Enq. Cherii. 39, 1618 (1!)47). 2. Sarrett, I<;. P., and Joyncr, I,. G., paper presented before tlir Division of l’rtroleuiii Chemistry, h i . Cheni. so(,., Atlaritic City, kkptoinhr, 1849. 3. Beehe, It. A,, Uerkwitli, J. U., and Honig, J. XI., J. ,4nr. Chettt. SOP.67, 1554 (1945). 3a. Beeck, O., Advances z n Catalyszs 2, 151 (1950). 4. Blorh, H. S., and Thomas, C. I,., J . A n r . C h e w . Soc. 66, 1589 (1944) 5. Brunauer, S., T h r Adsorption of Caws and Vapors, Vol. I. 1’rinc.eton University I’ress, Princeton, N.J., 1048. 6. Srunauer, S., Drining, I,. S., I ~ r i n ~ i i I\’. g , L,and Trllrr, E., J . L 4 t t ~ .Chetir. SOC. 62, 1723 (1!)40). 7. Brunauer, S., I’hnic-tt, 1’. H., and Teller, E., J. ,4171. Chettr. Soc. 60, SO9 (1938). 8. Cassie, A. 13. D . , ‘/‘ram. Farads!/ Soc. 41, 450 (1!)45). 9. Cohan, I,. €I., J . Am. Cheni. Soc. 60, 433 (1938). 10. Cohan, I,. H., J . A m . Chent. Sor. 66, 98 (1944). 11. Conn, A. I,., a n d I3rarkin, C. LV., I d . 13’ng. Chettt. 41, 1717 (1948).

148

IIERBIAN l3. RIES, JI1.

12. C h n n , A. L , hlrehan, \V. I ? , :ind Shankland, R. V., (‘hem. 1Sny. I’ioqiess 46, 176 ( 1950). 1;3. Davidson, I{. C., I’ctrolrurtr It‘cjttter 26, 663 (1947). 14. Drake, 1,. C , I d . ICirg. Cherir. 41, 780 (1949). 15. I<mmett, 1’. H., Aduairces Z?L Cutulrlsis 1, 65 (1948). 16. l h m e t t , 1’. H , Advances z n Collozd Scz. 1, I (1012). 17. fCinmett, 1’. H., arid 13ruri:iiier, S., .I. .I ttt (’hem. Sac. 69, 310 (1937). 18. b h n i e t t , 1’. I I . , and I3runuuw, H.,J . .I)// (‘hrtti. Soc. 69, 1563 (1937). 19. 15mmett, 1’. I I . , and De\Vitt, T. W., J . < 4 t t r . Chetti. Soc. 66, 1253 (1943). 20. Foster, A. G., ’/’tans. Faraday Soc. 28, ti45 (1032). 21. Foster, A . G., Discussions of thr I’atadu!/ Soczrti/, No. 3, 41 (l!)48). 22. Ihtnkerihurg, \I. G., J . A t ) ) . (“hcttt. Soc. 66, 1827 (l!)44). 2:3. Glrystrcn, I,. 11‘ , arid Divtz, V . It., J . lirseorth N n f l . N u r . S l ~ i r l ~ 36, rd~ 28.5 (1945). 24. Greensfeldcr, 13. S., Vogc, H. f I . , and Good, G . M., fwl. Birg. CheJtl. 41, 25i3

( I 949). 25. IIacnsr~l,V., and S t w h , 11. ,J., f t d . h’nq. C‘hetrr. 40, 1660 (1948). 26. Hansfortl, 11. C., I d . Etiq. C h r w . 39, 84!) (1947). 27. Harkins, \V. D., and Jula, G., J . A t t 1 . Cheut. Soc. 66, 1362 ( 1941). 28. Harkins, \i’. D., and Jura, G., J . L 2 / ~ iChettr. . Soc. 66, 1366 (1944). 29. Heincmann, H., A.A.A.S. Gordon Conference on Catalysis. New I,oiidon, NPWHampshire, 1950. 30. Hill, T. I,., J . C h e m Phys. 14, 263 (1946). 31. IIiridin, 8. G., arid Mills, G . A., paper prcsrntcd before the Division of I’tiysical and in organ it^ Chemistry, Am. Chein. Soc., Detroit, April, 1950. 32. Holnics, J., and lhim e tt, 1’. H., J . Phys. & Collozd Chem. 61, 1262 (1947). 3 3 . Johnson, hl. F. I,., arid lties, H. 15., Jr., J. .4??i.C h e w Soc. T2, 4289 (19.50).

31. Joyner, I,. G., Srimtific, :inti 1ndristri:tl Glass I3lowirig and 1,ahoratory Trt~liniqrit~s, cditcd hy \V. Rarr and V. .J.Anhorn. Instrrimcnts l’u1)Iishing C‘oiiijmiiy, I’ittshurgh, I%., 1949. 3.5. Joyner, I,. G., \Vt,inbcrger, 15. H., tirid Ilontgomery, C. \V., J . .Z t i / . ( ‘ h ~ t / t . Soc. 13;.

67, 2182 (1015). 36. Kistler, 8.S., J . 1’hUs. Chetti. 36, 52 (1932). 3i. Kraerner, I<. O., A Treatise on Physical Clicriiistry. E d i t d by 11. S. Taylor. 1).Van Nostraiid Company, N t v York, 1931. 38. McBain, J. \V., J . A1n. Cheni. Soc. 67, 699 (1935). 39. Milligan, \V. O., and Ilachford, H. H., Jr., J . f’hus. & Collocd Cherri. 61, 338 (1947). 40. hfilla, G. A., I n d . I h q . Chrttr. 42, 182 (1950). 41. Mills, G. A , Hoeticker, 1:. It., and Ohlad, A. G., J . .1 ttL. (‘hem. Soc. 72, 1554 ( 1950). 42. Oulton, T. I)., J . l’hys. CC Colloid Chem. 62, 1296 (1948). 43. Parravano, C., H~trriiiiel, 15 F.,and Taylor, H. S., J . A t n . Chrtti. Soc. 70, 2260 (1948). 44. t‘iert*e, C., \Yiley, J. W., anti Yiritth, It. S . , J . I’hUs. tC Colloid Chem. 63, titi0 (I!)40). 45. Itit~hardson,li. \V.,Johnsori, 14’. U., and Rolhins, I,. V., Jr., I n t l . B n y . C‘herti. 41, 1i 2 9 ( IO49). 16. Ilies, H. I<;., Jr., paper presented heforc thc Division of Potrolerini Chciiiistry, Am. Chem. Soc., Houston, >larch, 1950. 47. Ities, €I. lC., Jr., Johnson, M. F. I,., and Melik, J. S., J . Phys. 6 Collozd Chem. 63, G38 (1949).

STRUCTURE AND SINTERINQ

PROPERTIES OF CRACKINQ CATALYSTS

149

48. Ries, H. E., Jr., Johnson, M. F. L., Melik, J. S., and Kreger, W. E., paper presented before the Division of Petroleum Chemistry, Am. Chem. SOC.,Atlantic City, September, 1949. 49. Ries, H. E. Jr., Melik, J. S., and Johnson, M. F. L., communication to the General Discussion of the Faraday Society on “Heterogeneous Catalysis,” Liverpool, April, 1950. 50. Ries, H. E., Jr., Van Nordstrand, R. A., Johnson, M. F. L., and Bauermeister, H. O., J . Am. Chem. SOC.67, 1242 (1945). 51. Ries, H. E. Jr., Van Nordstrand, R. A,, and Kreger, W. E., J. Am. Chem. SOC. 69, 35 (1947). 52. Ries, H. E. Jr., Van Nordstrand, R. A., and Teter, J. W., Znd. Eng. Chem. 37, 310 (1945). 53. Ritter, H. L., and Drake, L. C., Znd. Eng. Chem., Anal. Ed., 17, 782 (1945). 54. Russell, A. S., and Cochran, C. N., Znd. Eng. Chem. 42, 1332 (1950). 55. Schumb, W. C., and Rittner, E. S., J . Am. Chem. SOC.66, 1692 (1943). 56. Shapiro, I., and Kolthoff, I. M., J. Am. Chem. SOC.72, 776 (1950). 57. Shull, C. G., Elkin, P. B., and Roess, L. C., J. Am. Chem. Soc. 70, 1410 (1948). 58. Smith, R. C., Jr., and Howard, H. C., Ind. Eng. Chem. 34, 438 (1942). 59. Tamele, M. W., Byck, H. T., Ronay, G., and Vinograd, J. R., paper presented before the Division of Colloid Chemistry, Am. Chem. SOC.,New York, September, 1947. 60. Thomas, C. L., Znd. Eng. Chem. 41,2564 (1949). 61. Turkevich, J., and Smith, R. K., J. Chem. Phys. 16, 466 (1948). 62. Van Nordstrand, R. A., Kreger, W. E., and Ries, H. E., Jr., paper presented before the Division of Petroleum Chemistry, Am. Chem. SOC.,New York, September, 1947. 63. Webb, G. M., and Ehrhardt, C. H., Petroleum Processing 2, 5 (1947). 64. Wheeler, A., A.A.A.S. Gordon Conference on Catalysis. Gibson Island, 1945. 64a. Wheeler, A., Advances in Catalysis 3, 250 (1951). 65. Zettlemoyer, A. C., and Walker, W. C., Znd. Eng. Chem. 39, 69 (1947).

This Page Intentionally Left Blank

Acid-Base Catalysis and Molecular Structure R. P. BELL Balliol College, Oxjord, England

CONTENTS Page

I. [ntroduction.

. . . .

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

11. The Empirical Law atalysis.. . . . . . . . . . . . . . . . . . . . . . . . . . . ....... 1. Salt Effects.. . . . . . . 2. General and Specific ..................................... 3. Acid-Base Catalysis in Nonaqueous Solvents. . . . . . ......... 4. Relations between Acid-Base Strength and Catalyt 111. The Molecular Mechanism of Acid-Basc Catalysis 1. The General Nature of Acid-Base Catalysis.. . . . . . . . . . . . . . . . . . . . . . . . . 2. Examples of the Mechanism of Acid-Base Catalysis.. . . . . . . . . . . . . . . . . . a. Some Reactions of Enolisable Ketones and Similar Substances. . . . . . b. Nitro-Compounds and Nitramide. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . c. Reversible Addition of Hydroxy-Compounds to the Carbonyl Group.

153 153 157

164 165 165 169 171

d. Esterification and Hydrolysis of Carboxylic Esters, . . . . . . . . . . . . . . . . 173 3. Kinetic Steps in Acid-Base Catalysis., . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 174 a. Reactions Involving a Single Proton Transfer.. . . . . . . . . . . . . . . . . . . . 174

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

1. Pseudo-Acids and Pseudo-Bas

V. The Importance of Molecular Structure.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1. The Structure of t h e Substrate.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2. The Structure of the Catalyst.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

201 201 204 207

I. INTRODUCTION The early study of catalysis by acids and bases was concerned chiefly with the use of catalyzed reactions for investigating general problems of physical chemistry. For example, the first correct formulation of the kinetic laws of a first order reaction was made by Wilhelmy in 1850 in connection with his measurements of the catalytic inversion of cane sugar by acids (Wilhelmy, 1). Catalytic reactions also played an important 1 F1

152

R. P. BELL

part in the foundation of the classical theory of electrolytic dissociation toward the end of the nineteenth century. The parallelism between the electrolytic conductivity and the catalytic activity of solutions of acids received a ready explanation in terms of the high mobility and catalytic activity of the hydrogen ion (Ostwald, 2), and measurements on catalyzed reactions (notably the hydrolysis of esters) were used widely for investigating the state of solutions of electrolytes. The classical theory of acid-base catalysis assumed that hydrogen and hydroxyl ions were the only effective catalysts, that the reaction velocity was proportional to the concentration of the catalyzing ion, and that the degrees of dissociation of the electrolytes involved were given directly by their electrolytic conductivities. These assumptions, although giving a good general description of the facts, led to a number of discrepancies when applied quantitatively. The next phase in the study of acid-base catalysis, especially associated with the name of J. N. Bronsted, dealt mainly with the clearing up of these discrepancies, partly by the application of modern views on electrolytic solutions, and partly by the deduction from experiment of new laws governing catalytic phenomena. In this way the systematics of acid-base catalysis were largely established in the decade 1920-1930, and little has been added later to this aspect of the subject. This part of the story is well known l(see, for example, Bell, 3), and Sec. I1 of this review therefore contains only a brief summary of the empirical laws of acid-base catalysis, with few references. The aspect of acid-base catalysis which will be mainly dealt with in this article is its bearing on the molecular mechanisms of the reactions concerned and the structure of the molecules taking part. This side of the subject is of comparatively recent development. Most reactions catalyzed by acids and bases involve fairly complicated organic molecules, and although physical chemists used these reactions widely as tools for investigating their own problems, there was a general reluctance to speculate as to the detailed reaction mechanisms. Catalyzed reactions were of course included in some of the early attempts of organic chemists to devise reaction mechanisms, but the catalyst was often regarded only as an auxiliary influence which facilitated an uncataly zed mechanism, a view which is now known to be incorrect. Closer analysis shows that in most catalyzed reactions in solution the rate-determining step is chemically a simple one, though the overall reaction may be fairly complex. Catalyzed reactions are in fact often very suitable for studying general problems in the field of reactivity and structure, and much modern work on the effect of substitution on reactivity has in fact dealt with catalyzed reactions. Sections 111 and IV describe how the nature of the ratedetermining step can be ascertained in many reactions, and Sections

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

153

V gives examples of how the reaction velocity depends on the molecular structure of the catalyst or substrate. 11. THE EMPIRICAL LAWS OF ACID-BASECATALYSIS 1. Salt Efects

The subject of salt effects in one which arises in all reaction-kinetic problems involving electrolytes and has no special relevance to acid-base catalysis. However, much of the early work on salt effects was in fact carried out with catalyzed reactions, and a neglect of these effects is still the commonest cause of misinterpretation of data on acid-base catalysis, so that a brief account will be given here. It is convenient to include under the heading of “salt effects” the various ways in which the assumptions of the classical theory have been modified by modern views on electrolytic solutions. Since the catalyst itself is commonly ionic, the same problems often arise even when no other electrolyte has been added to the system. The classical theory regarded all electrolytes as being incompletely dissociated in solutions of moderate concentration, the degree of dissociation being given by A/A,, the conductivity ratio. The present view is that those electrolytes commonly classed as strong (most salts, and a few acids such as HC1, HBr, HI, HC104, and sulfonic acids) are effectively completely dissociated in aqueous solution, even at concentrations where the conductivity ratio indicates a considerably smaller degree of dissociation. The decrease in equivalent conductivity with increasing concentration is attributed to the electrostatic forces between the ions rather than to a decrease in the degree of dissociation. This assumption of the complete dissociation of strong electrolytes (for which there is, of course, much evidence from various sources) simplifies considerably the interpretation of catalysis by strong acids or strong bases, since it is often found that the reaction velocity in such solutions is approximately proportional to the total concentration of acid or base rather than to the conductivity of the solutions. * The same is true in nonaqueous media such as methyl and ethyl alcohol, in spite of the more powerful interionic forces in these media. It is not strictly true to say that the velocity of a reaction catalyzed by a strong acid or a strong base is universally or exactly proportional to the catalyst concentration. In the first place, this statement ignores the primary salt effect (see below), though deviations attributable to this cause are unlikely to exceed a few per cent in 0.1 N solution. In the second place it is once more becoming fashionable to describe such salts

* For examples, see Bell, 3, Chapter 11.

154

R. P. BELL

as “incompletely dissociated,” especially those containing multiply charged ions or the anions of certain organic acids, though the degrees of association attributed to them are much smaller. than those allotted by the classical treatment (Davies, 4). As far as acid-base catalysis goes, the most important examples are certain metallic hydroxides such as Ca(OH)2, Ba(OH)2, TlOH, whose aqueous solutions are supposed to contain appreciable concentrations of the species CaOH+, BaOH’, and TlOH. There has been some difference of opinion about the status and usefulness of the concept of incomplete dissociation in these cases (cf. Owen, 5 ) , but it is supported by kinetic measurements using solutions of the above hydroxides (Bell and Prue, 6; Bell and Waimd, 7). However, in spite of the above qualifications it still remains a good approximation to say that the catalytic effect of a strong acid or base is proportional to its total concentration, provided that high concentrations and certain types of catalyst are avoided. One of the main anomalies encountered in app1.ying the classical theory of catalysis was the large accelerating effect produced by the addition of neutral salts to catalyzing solutions containing weak electrolytes: for example, the addition of 0.1 M K N 0 3 to 0.05 M acetic acid increases by 30% its catalytic effect in the reaction of diazoacetic ester with water (Bronsted and Teeter, 8). Such an effect is now termed a secondary salt efect and is independeht of any kinetic considerations, being due to the increase in the degree of dissociation of the acetic acid caused by the increased ionic concentration. This; increase can be detected by other means (e.g., indicators, conductivity measurements) and its theoretical basis is now well understood. If the hydrogen ion concentration of a solution is controlled by a dissociation equilibrium of the type HA G H+ A-, then the concentration dissociation constant K , is given by the expression

+

where K is the thermodynamic dissociation constant, dependent only on the solvent and the temperature, and the f’s are activity coefficients. To a good approximation f H A is unity, independent of salt concentration, and the ionic activity coefficients can be predicted by the interionic attraction theory of electrolytes. For aqueous solutions at 25” a good approximation is (Guggenheim, 8a),

where z is the valency of the ion and I the ionic strength defined by I = fZnLizi2

(3)

155

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

mibeing the molality of an ion of species i, and the summation being taken over all the ions present in the solution. Applying this t o Eq. (1) we find log,, Kc = log,, K

+ Z$4/(1 + X W )

(4)

which shows that the Concentration dissociation constant (and hence the degree of dissociation) is increased by the addition of salt and serves t o estimate the magnitude of the effect. When a neutral salt is added t o a solution of a weak acid, the ions of the added salt are the main contributors t o the ionic strength. If, on the other hand, measurements are being made in a buffer solution (e.g., sodium acetate), then the constituents of the buffer itself acetic acid contribute largely t o the ionic strength, and there may be considerable secondary salt effects even without the addition of other salts to the system. For example, the classical dissociation theory predicts that the hydrogen ion concentration in a buffer solution of acetic acid sodium acetate should depend only on the ratio of the buffer constituents. Actually, because of the secondary salt effect, the hydrogen ion concentration depends also on the total buffer concentration, increasing with increasing salt concentration. These variations can be allowed for approximately by using Eq. (4),but it is usually simpler t o carry out kinetic experiments a t constant total salt concentration by adding appropriate amounts of a neutral salt. For example, the following series of solutions are to a good approximation “isohydric,” since the ionic strength is throughout 0.1 :

+

+

0 . 1 N acetic acid 0.075 N acetic acid 0 . 0 5 N acetic acid 0 . 0 2 5 N acetic acid

+ O . 1 N sodium acetate +O. 075 N sodium acetate +0.05 N sodium acetate +0.025 N sodium acetate

+O ,025 N NaCl + O . 05 N NaCl +0.075 N NaCl

This principle of maintaining a constant ionic strength is of great value in simplifying the comparison of kinetic data and should be followed whenever possible. Equations (l), (3), and (4) apply only when the hydrogen ion concentration is controlled by the dissociation of an uncharged acid molecule. I n many instances different types of equilibria are involved. For example, in a buffer solution of ammonia ammonium chloride the relevant equilibrium is NH4+ $ NH3 H+, while in a phosphate buffer it is HzP04HP04= H+. It is easy t o generalize the expressions for the secondary salt effect to cover all these cases. If the equilibrium is A B H+, where A bears z positive charges and B z - 1 (z can of course be a positive or negative number), then the application of Eq. 2 gives

+

+

+

+

156

H. P. BELL

Equation (4) is a special case of Eq. ( 5 ) with z = 0. I t may be noted that if z = +1 (as in a buffer of ammonia ammonium chloride), the secondary salt effect is to a first approximation zero, so that the hydrogen ion concentration of this type of buffer is relatively insensitive to changes in salt concentration. An analogous treatment applies to the hydroxyl ion concentration in solutions containing weak electrolytes, where it must be remembered that the value of the ionic product [H+][OH-] will increase with increasing ionic strength. The primary salt e$ect deals with the effect of salt concentration on reaction velocity when the reacting system involves no equilibria which can be displaced by a change in ionic environment. This effect can be very large when both the reacting species are ions, but it is of less importance in acid-base catalysis, where the substrate is almost always an uncharged molecule. To avoid complications due t o secondary salt effects, the primary effect is best studied in catalysis by solutions of strong acids and bases, and there exists a large body of experimental data. Some of the main conclusions are as follows: (i) For a given reaction and a given added salt the percentage change in velocity is a linear function of the salt concentration. (ii) The magnitude of the effect depends upon the individual nature of the reaction and of the added salt, but it rarely exceeds 4 to 5% in an 0.1 N solution of a uni-univalent salt. (iii) The addition of salt invariably causes an increase of velocity in reactions catalyzed by hydrogen ions, while for hydroxyl ion catalysis there is sometimes an increase and sometimes a decrease. It is possible to give a formal theoretical treatment of the primary salt effect, though it leads to little in the way of quantitative prediction. For a reaction involving the hydrogen ion and an uncharged substrate S, the theory gives for the effect of environment on the velocity of the reaction v ,

+

=

k[H+][S]fa+fs/fx+

(6)

where X+ represents the transition state (or critical complex) between the two reactants, and k depends only on the temperature. This expression was originally advanced by Bronsted (9) on not very clear theoretical grounds, but would now be regarded,as a special cas,e of the transition state theory of reaction velocities. It should be noted that it is never justifiable in a reaction of this type to introduce the factor fE+fB into the rate expression, omitting fx+, This assumption was made in the so-called activity rate theory, but it is correct only for reactiolns between ions of

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

157

equal and opposite charge. Unfortunately the ratio fH+/fx+ is not accessible experimentally (except from the kinetic measurements themselves), nor is it amenable to theoretical treatment, since theoretical expressions such as (2) predict the same value for the two activity coefficients. The most useful approach is that due to Hammett (lo), who pointed out that in the equilibrium of a simple basic indicator B, the ratio of the two forms is given by

where the activity factor has just the same form as in (6). This factor can be determined experimentally and is found to have approximately the same value for different indicators. Moreover, it gives a good account of the catalytic power of solutions of strong acids, with and without the addition of salts, in many cases up to high ionic concentrations. However, in view of the absence of any satisfactory theory of concentrated electrolytic solutions it would seem desirable t o confine kinetic measurements as far as possible to more dilute solutions, in which primary salt effects are small and secondary ones can be reliably estimated. 2. General and SpeciJic Catalysis The classical ,theory of catalysis supposed that the hydrogen and hydroxyl ions were the only effective catalysts in solutions of acids and bases. In a few instances early attempts were made to remedy some of the discrepancies encountered by attributing some catalytic power to undissociated acid molecules, but these attempts were mostly based on incorrect values for degrees of dissociation, and they did not take into account the possibility of primary or secondary salt effects. However, later work has shown definitely that species other than hydrogen and hydroxyl ions often can exert a catalytic effect, and the development of these ideas was closely linked with a closer understanding of the nature of the hydrogen ion in solution, and with the clarification of acid-base definitions (cf. Bell, 11). As long as the hydrogen ion was regarded as a bare proton, H+, it seemed reasonable to suppose that it had a unique catalytic power (perhaps in virtue of its powerful electrostatic field), and this idea seemed t o fit in with its abnormally high ionic mobility. However, evidence soon accumulated to show that there could be no significant concentration of free protons in any solvent, but that the “hydrogen ion” is H30+ in water, C2H,0H2+ in ethyl alcohol, NH,+ in liquid ammonia, etc. This realization destroyed the unique position of the hydrogen ion, and it was soon found (Lowry and Smith, 12; Bronsted and Guggenheim, 13) that

158

R. P. BELL

both the ammonium ion and the hydrogen ion acted a s catalysts in the mutarotation of glucose. At about the same time the Bronsted-Lowry definition of acids was put forward, according to which a n acid i s a n y species which has a tendency to give u p a proton. This definition makes no mention of the charge in the species, and in fact we now regard H30+, NH4+,etc., as cation acids, completely analogous to uncharged acids like HC1 and CH,COOH. The first clear demonstration that uncharged acid molecules are catalytically active was given by the work of Dawson (14) on the reaction between acetone and iodine, though some of his quantitative conclusions need modification in the light of later work on electrolytes. Uncharged acids were also found to be catalysts in the mutarotation of glucose, and subsequently in many other reactions. The position is similar in basic catalysis. The hydroxyl ion has no strong claims to uniqueness, being merely the anion of a weak acid. According to the Bronsted-Lowry acid-base definition, a base i s a n y species which has a tendency to accept a proton. This obviously includes anions like OH-, CH3COO-, HPO,=, as well as uncharged basic molecules like ammonia and the amines. Catalysis by all these species was first found in the decomposition of nitramide (Bronsted and Pedersen, 15), and subsequently in many other reactions. General catalysis by acids and bases is now recognized as a very common phenomenon. Since a reaction may exhibit both acid and base catalysis, many different species may contribute to catalysis in the same solution. For example, in an acetate buffer the most general expression for the observed velocity would be 2,

=

Uo

+

f ~ H + [ H$ +.]~oH-[OH-] ~EAO[HAC] k ~ o - [ A c - ]

(8)

The term v o represents what is commonly described as the “spontaneous” reaction, though this is really a misnomer, since it it3 actually due to catalysis by solvent molecules, acting as acids or bases. It appears th a t all reactions catalyzed by acids or bases can be arrested completely in solvents having no acidic or basic properties (e.g., hydrocarbons), and an apparently uncatalyzed reaction of this type can usually be traced to the presence of some adventitious acidic or basic catalyst. I n aqueous solution it is not always easy to establish with certainty the existence of catalysis by species other than hydrogen or hydroxyl ions, and some of the early conclusions in this field werle based on insufficient evidence. The safest method is to use as catalysts a series of buffer solutions of equal ratios but varying concentrations, using the principle of constant ionic strength to eliminate secondary salt effects, as described in the last sub-section. If the observed reaction velocity increases with increasing buffer concentration in such a series, this is proof that one or

159

ACID-BASE CATALYSIS A N D MOLECULAR STRUCTURE

both of the buffer constituents is exerting a catalytic effect. The evaluation of the separate catalytic constants, e.g., kH+, kOH-, kHAo,and kr.- in Eq. (8), demands an extensive series of carefully planned experiments (see for example, Bell and Baughan, 16). Although many reactions exhibit general acid-base catalysis in the sense described above, there remain a few in which no catalysis by species other than hydrogen or hydroxyl ions can be detected. This behavior is known as specific catalysis. We shall see later that it is doubtful whether these reactions differ in principle from those exhibiting general catalysis. Most probably the failure to detect general catalysis is due to the quantitative relations between the catalytic effects of different species rather than to any pecularity in the reaction. However, specific catalysis is important in practice, since it provides a means of measuring the concentrations of hydrogen or hydroxyl ion in a solution without disturbance from catalysis by other species. Examples of specific catalysis by hydrogen ions in aqueous solution are the decomposition of diazoacetic ester, and the hydrolysis of acetals, while specific catalysis by hydroxyl ions is exhibited by the depolymerization of diacetone alcohol and the decomposition of nitrosotriacetonamine. 3. Acid-Base Catalysis in Nonaqueous Solvents

In general catalytic measurements in aqueous solutions are easier to interpret with certainty than those in other solvents, since our knowledge of the properties of solutions (especially of electrolytes) is very limited outside water. However, it is often necessary to use nonaqueous solvents for practical reasons e.g., solubility and chemical inertness, and the use of different solvents has elucidated a number of points of interest in the general theory of acid-base catalysis. The solvents which differ least from water are the lower alcohols. For example, in ethyl alcohol the analogues of the hydrogen and hydroxyl ions in water are C2HsOH2+ and C2H50- respectively. These ions cannot exist in appreciable concentration in aqueous solution, since the reactions CzH50Hz+ HzO CzH50H H30+ and C2H50H2O + C2H50H OH- go completely from left to right. This shows that C,H50H2+ is a stronger acid than H30+, and C2H50- a stronger base than OH-. Correspondingly, alcoholic solutions will often catalyze transformations which cannot be effected in water. Similarly, certain acids which are strong in water become weak in alcoholic solution. For example, nitric acid dissociates almost completely in water according to NOa-, while it is only slightly the scheme H N 0 3 H2O 3 &of so that catalysis by undissociated HNOl molecules dissociated in alcohol, can be studied in alcohol, but not in water. On the other hand, the

+

+

+

+

+

+

160

R. P. BELL

quantitative interpretation of kinetic data is much more difficult in alcoholic than in aqueous solutions, partly because of the lower dielectric constant. For this reason salt effects (both primary and secondary) are much greater in magnitude and less well investigated. Thus in the activity coefficient expression (2), the numerical factor which is 0.5 in water would be about 2.8 in ethyl alcohol, and the concentration range over which such expressions are valid is much reduced. In solvents which differ more radically from water, the chief new feature is the range of catalytic species which can be effective. This may be illustrated by reference to the solvent anhydrous acetic acid, which is highly acidic, but which has very weak basic properties. For this reason, solutions of strong acids like HCl, HBr, and HClO, are very little dissociated in this solvent and differ considerably in their catalytic and other properties, whereas solutions of the same acids in water are converted completely into hydrogen ions, and therefore exhibit almost identical properties. Correspondingly, all bases which in water are stronger than aniline behave as “strong l 1 bases in anhydrous acetic acid, reacting completely with the solvent according to the equation B CH3COO- + BH+ CH3COO-, and thus: yielding solutions of identical basic properties. Special interest attaches to solvents such as the hydrocarbons, which exhibit neither acid nor basic properties, being unable t’o lose or to gain a proton. They are often described as aprotic and are typified by the hydrocarbons and their halogen derivatives. Although there are no analogues of the hydrogen and hydroxyl ions in these solvents, and even the strongest acids and bases remain undissociated, inevertheless their solutions possess catalytic activity, often exceeding that of any aqueous solution. This provides very direct evidence of the possibility of catalysis by undissociated acids and bases, and in principle the study of catalysis in aprotic solvents should be much simpler than in solvents of other types. For example, if acetic acid is dissolved in water, the sohtion contains the species HzO, H30+,OH-, CH3COOII, and CH3COO-, all of which may be catalytically active (cf. Eq. ( S ) ] , while in a solution of acetic acid in benzene the only active species is the acetic acid molecule itself. In practice this advantage is to some extent counterbalanced by complications which are found to arise in the kinetics of catalyzed reactions in these solvents. These are due to the low dielectric constants of the media concerned, which favor strong interactions between polar molecules, leading to association of reactants and catalysts and to kinetic abnormalities analogous to salt effects. Nevertheless, measurements in aprotic solvents often provide information of interest and have been increasingly used.

+

+

ACID-BASE CATALYSIS A N D MOLECULAR STRUCTURE

161

Since acid-base catalysis can take place in solvents which play no part (by promoting dissociation or otherwise) in the reaction, it is natural to inquire whether such catalysis can take place in the complete absence of a solvent, i.e., as a gas reaction. Practical difficulties connected with the instability or low volatility of the reactants often hinder an investigation of this point, but it has been shown in a few instances (Bell and Burnett, 17, 18; Wassermann, 19, 20) that catalysis by acids can certainly take place in the absence of a solvent, though in most cases the reaction takes place predominantly on the walls of the reaction vessel. Only for one reaction, the dimerization of cyclopentadiene, has homogeneous acid catalysis been observed in the gas phase. The preference for a heterogeneous reaction can be reasonably accounted for in terms of the type of reaction mechanism which we shall discuss later in this article. These mechanisms involve a considerable charge separation, which will be favored by the proximity of any polarizable material, such as the wall of the vessel (cf. Bell, 21). In this connection it may be of interest to mention that occasionally acid-base catalysis may be effected heterogeneously at the surface of a solid catalyst. For example, ion exchange resins have been used as acid catalysts in the esterification reaction (Haskell and Hammett, 22; Levesque and Craig, 22a), and it is possible that some of the oxide catalysts used industrially in various organic reactions may be functioning as acids or bases (Walling, 22b). However, little is yet known about acid-base catalysis of this type.

4. Relations between Acid-Base Strength and Catalytic Power Since there is a qualitative correlation between the acid-base properties of a species (ion or molecule) and its ability t o act as a catalyst, it is reasonable to expect that there may also be a quantitative relation between the acid-base strength of a species and its effectiveness as a catalyst for a given reaction. Such relations become still more likely when we consider (as in later sections of this article) the actual mechanism of catalysis, but they were in fact originally derived empirically from experimental data without reference to mechanism and will therefore be described briefly here. The catalytic power of a particular species is most conveniently specified by the catalytic constant (kA or kB), as shown in Eq. (8). The acid strength of a catalyst is usually described by its dissociation constant in water, K A . The basic strength of a catalyst may be measured by the conventional basic dissociation constant, but it is more convenient to use the reciprocal of the dissociation constant of the corresponding acid :thus for the acetate ion we write KB = l/KHAD = [HAc]/[H+][Ac-], and for ammonia, K , = 1/KNH4+= [NH4+]/[H+][NH3]. Strictly speaking,

162

R . P. BELL

these should be converted into thermodynamic constants by the inclusion of activity coefficients, but the distinction is often unimportant in the present context. The following equations are then found to hold approximately for acid and basic catalysis respectively kA =

G A K A ~ ,kB

=

GB(~/KA)O

(9)

where G A , G B , a,and p depend only on the solvent, the temperature, and the nature of the catalyzed reaction considered, a and p being always positive and less than unity. A different choice of constants for specifying the strength of acids or bases would not affect the -form of the equations or the values of a or p, but would of course modify the parameters G, or G,. This type of relation was first established by Bronsted and Pedersen (15) for the decomposition of nitramide and is usually known as the Bronsted relation. Later sections of this article will deal with the exact range of validity of the Bronsted relation and with its theoretical interpretation, and we shall give here only some of the salient facts. A relation of this kind has been found to hold, a t least approximately, for every reaction in which a series of related catalysts has been investigated, the range of k or K frequently covering several powers of ten. The values of G, a, and p vary from one reaction to another. G also varies markedly with solvent and temperature, a and p much less so. A drastic change in the nature of the catalyst (in particular a change in the charge which it bears) necessitates a change in the parameters of the equation. The same type of relation governs catalysis in nonaqueous solvents, even when these are aprotic. In the latter case it is not possible to specify the strengths of acids and bases by dissociation constants in the same solvent, but use may be made of equilibria with an added acid-base system, e.g., an indicator. Frequently such equilibrium measurements are not available, and dissociation constants in water are commonly used as a basis of comparison with catalytic measurements in other solvents. Since the relative strengths of acid-base systems of the same charge type vary little from one solvent to another, the use of aqueous dissociation constants will not alter the form of Eqs. (9), though it will of course change the value of G. If it can be assumed that the Bronsted relation remains approximately valid for the ions and moIecules of the solvent, it is possible to make some interesting deductions about the possibility of detecting general acid-base catalysis. The point is best illustrated by taking a particular example, e.g., acid catalysis in an aqueous solution 0.1 N with respect to both acetic acid and acetate ions. We assume that the catal.ytic effect of acids (without reference to charge or structure) is given ttpproximately by

ACID-BASE CATALYSIS A N D MOLECULAR STRUCTURE

163

Eqs. (9). In applying this to the species H30+and HzO we use the acid constants

The quantitative significance of these figures is doubtful, since they involve the concentration of water molecules in pure water (taken as 55.5 moles/liter), but experience shows that they give a roughly correct estimate. The following figures are then obtained for the proportion of catalysis due to the three acidic species in the solution, for different values of the exponent Proportion of Catalysis due to

Exponent 01 = 0 . 1 01 = 0 . 5 a = 1.0

H,O

H30+ 0.002y* 3.6% 99.8%

98 % 0.01%

5

x

10-’0%

CHECOOH 2% 96.4% 0.2%

When a = 0.1, most of the catalysis is due to the solvent, and the reaction would in practice be regarded as uncatalyzed, since the rate is but little increased even in solutions of strong acids. When a = 0.5, the rate in the buffer solution is largely due to the undissociated acetic acid, as could be verified by varying the buffer concentration and keeping its ratio constant. On the other hand, the catalytic effect of OH3+ could be measured independently in solutions of strong acids, and the “spontaneous” (H20-catalyzed) reaction in solutions sufficiently alkaline to repress the effect of the hydrogen ions. This case is thus a favorable one for the study of general acid catalysis. Finally, if a = 1, the catalytic effect of the buffer is almost entirely due to the hydrogen ions which it contains, and it is clear that no experiments could detect with certainty the small effect of the acetic acid molecules. Moreover, the watercatalyzed reaction is so slow that it would be impossible to detect it. The reaction would thus be classed experimentally as an instance of specific catalysis by hydrogen ions. This treatment is easily generalized and leads to the conclusion that in water or similar solvents general acid catalysis will be observable only for intermediate values of the exponent a. If a is too small, the catalytic effect of acids will be swamped by that of the solvent, while if a approaches unity, the effect of all other acids will be masked by that of the hydrogen ion. Similar conclusions apply to basic catalysis. It is possible, therefore, that those reactions in aqueous solution which appear to show specific catalysis by hydrogen or hydroxyl ions (cf. preceding sub-section) do not constitute a special class of reaction, but are actually

164

R. P. BELL

examples of general acid-base catalysis in which the exponent of the Bronsted relation approaches unity. These considerations do not, of course, affect the practical application of these reactions for measuring hydrogen and hydroxyl ion concentrations. In aprotic solvents, on the other hand, there are no limits to the value of a! observable in practice, and values as low as 0.2 and as high as unity have in fact been observed (Bell and Brown, 23).

111. THEMOLECULAR MECHANISM OF ACID-BASECATALYSIS I. T h e General Nature of Acid-Base Catalysis Early views on the nature of catalysis regarded it as an indefinite influence of some kind. Somewhat later a rather more definite picture was formed of catalysis by the hydrogen ion (regarded as a bare proton), which was supposed to attract the reactants together in virtue of its powerful electric field. This explanation did not seem especially appropriate to hydroxyl ion catalysis and obviously would not apply to catalysis by uncharged molecules. These early views envisaged reactions which could take place in the absence of a catalyst, but which were facilitated by itt3 presence. Evidence gradually accumulated to show that many of the reactions subject to acid-base catalysis could not take place at all in the complete absence of catalysts, apparently “spontaneous reactions” being often due to catalysis by acidic or basic solvent molecules, or by some adventitious acidic or basic impurity. This seemed to indicate that the catalyst took a fundamental part in the reaction, possibly in :a chemical sense. It was soon realized that the essential property of acilds and bases was their power respectively to lose and to add on a proton, and enquiry also showed that substrates involved in acid catalysis could always be supposed to have some basic properties, while those in base-catalyzed reactions could always in principle act as acids, though the acid-base properties of the substrates were often so weak as to elude detection by ordinary means. This led to the suggestion that acid-base catalysis always involves an acid-base reaction between the catalyst and the substrate. Such a reaction is also often termed a protolytic reaction, since it involves the transfer of a proton between the two reacting species. This view of acid-base catalysis is now generally accepted, and specific mechanisms have been proposed for a large number of types of reaction. For the purpose of illustration a few of these are given in the next subsection, but no attempt has been made at comple-teness. I n many instances the mechanism involves two successive proton transfers, and it may be a matter of some difficulty to decide the relative rates of the two successive steps. This question is considered in Sec. 111.3.

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

165

a. ExampEes of the Mechanism of Acid-Base Catalysis a. Some Reactions of Enolizable Ketones and Similar Substances. The earliest reaction in this class to be studied in any detail was the halogenation of ketones and allied compounds. It was found by Lapworth (24) that the rate of reaction of ketones with iodine is independent of the iodine concentration and was in fact the same for bromine as for iodine under the same conditions. This shows that the process being measured is not the halogenation reaction at all, but some change in the ketone itself. The rate of halogenation is increased by addition of strong acid, and halogenation also takes place rapidly under alkaline conditions, though further reactions often take place (e.g., the iodoform reaction with acetone and similar substances). These facts led Lapworth to suggest that the process being measured was in fact the enolization of the ketone, since enols are known to react very rapidly with halogens, and the interconversion of keto-enol isomers is catalyzed both by acids and by bases. The ordinary tests do not show any detectable amount of enol in simple ketones, and the supposed preparation of these enols by indirect means has been disputed (Kohler and Thompson, 25). These and other arguments have led some authors to reject the enolization mechanism for the halogenation of simple ketones (Arndt, 26). However, the presence of an appreciable amount of enol at equilibrium is not a necessary condition for the correctness of Lapworth’s mechanism, and in any case it has been shown recently that simple ketones do contain a very small proportion (of enol (Schwarzenbach and Wittwer, 27). Even before this demonstration the enolization mechanism was generally accepted, for example in the extensive work of Dawson (14) on the iodination of acetone. Dawson found that the reaction was catalyzed not only by hydrogen or hydroxyl ions, but also by undissociated acid molecules and by the anions of weak acids, and similar demonstrations of general acid-base catalysis were given later for the halogenation of other carbonyl compounds (Pedersen, 28, 29; Bell and Lidwell, 30). It is of interest to note that the same kind of kinetics can be observed for the acid-catalyzed halogenation of acetone in hydrocarbon solvents, though here it is necessary to use N-halogen compounds as halogenating agents rather than free halogens, so as t o avoid the formation of undissociated hydrogen halides, which have an enormous catalytic effect (Bell and Tantram, 31). We must now consider the mechanism by which acids and bases can effect the transformation of a keto to an enol form and in doing so shall have occasion to modify slightly the enolization view of halogenation. All compounds containing a carbonyl group have some basic properties.

166

R . P. BELL

These are too weak t o detect in aqueous solution, but ar’erevealed in very acid solvents such as concentrated sulfuric acid (e.g:., Hammett, 32; Flexser, Hammett, and Dingwall, 33). Similarly, every enolizable ketone containing the group

\

giving the ion

/

/

C:C.O-.

\ /

/

C H C : 0 can in principle act as an acid,

@-Diketones and similar substances are in

fact acids of measurable strength in aqueous solution, blut simple ketones are such weak acids that their acidic properties can o,nly be studied in inert solvents (Conant, 33a). Reasonable mechanisms for the formation of the enol are then as follows, where A represents any acid, and B any base : Acid catalysis

+A

‘CH.C
/

\ i +

/

/+ CHC:OH

+B

\

/

C:C.OH

/

+A

(10)

(1)

Basic catalysis

\

/

CH.C:O+B=

/

\

/

/

C:C.O-+A*

\

/

C:C.OH + B

/

(11)

Each mechanism involves a two-stage reaction, and there is no direct transfer of a hydrogen atom from the carbon atom t o the oxygen. I n fact, the hydrogen of the -OH will probably not be the same atom as Each mechanism also requires the participation of that of the -CH. both an acid and a base, though the order of attack is different. It is not, however, necessary deliberately to add both an acid and a base in order to bring about the change. In water or a similar medium the solvent molecules themselves can act either as acids or as bases, and even in an aprotic solvent the first stage of the reaction will convert an acid into a base, or vice versa, so that both types of molecule will be present although only one was added initially. This point will be considered further in the next sub-section in connection with tlhe kinetics of the successive steps. The above mechanism for the acid-catalyzed formation of the enol accords with Lapworth’s interpretation of the halogenation, but for basic catalysis some modification is necessary. The proposed mechanism involves the anion (11) as an intermediate in the formation of the enol,

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

167

and it would be anticipated that this anion will react very rapidly with halogens. This is in fact found to be the case for ketones which are sufficiently acidic to be converted completely into the enolate ion. This means that in presence of halogen no enol will be formed, and the measured rate of halogenation under conditions of basic catalysis is therefore the rate of ionization of the ketone, rather than its rate of enolization. No corresponding change is needed in the mechanism for acid catalysis, since the cation (I) will not be reactive toward halogens. The conversion of a ketone into its enol or enolate ion can be observed by other means than by reaction with halogens. For example, if the ketone is dissolved in water (or a similar solvent) containing a considerable proportion of deuterium, the formation of enol or enolate ion will lead to the exchange of hydrogen isotopes between the ketone and the solvent. We should therefore expect that this isotope exchange would be catalyzed by acids and by bases, and this is in fact found to be so. Similarly, if an optically active ketone R1R2CH*C0.Ris prepared, the formation of enol or enolate ion will lead to racemization, and correspondingly the racemization of such a substance is found to be catalyzed by acids and bases. The close relation between the mechanisms of halogenation, isotope exchange, and racemization is further confirmed by quantitative comparisons of the rates of these processes. Thus the rate of racemization of the ketone CsH&OCH(CH3)C2HS by sodium hydroxide in a dioxane-water mixture was found to be identical with its rate of interchange with sodium deuteroxide under the same conditions (Hsii, Ingold, and Wilson, 33b). Similarly, the rates of bromination and racemization of the ketone

were found to be identical when catalyzed by acetate ions in aqueous acetic acid solution (Hsii and Wilson, 34). These two comparisons refer to basic catalysis, but the position is similar for catalysis by acids. For example, the rates of acid-catalyzed halogenation and racemization have been found to be identical for both the optically active ketones mentioned above (Ingold and Wilson, 35; Bartlett and Stauffer, 36), and the acidcatalyzed bromination and isotope exchange of acetone also proceed at equal rates (Reitz, 37). It has just been assumed that the formation of an enol or enolate ion from an optically active ketone automatically leads to racemization, and this is clearly to be expected on the basis of the planar structures

168

R . P. BELL

Ri

\

7%

R2/c=c\K

R1 \

0/'

,c=c\, R2

R

However, it is possible that the actual structures derive partly from the forms R,

\ - //

/c-c\R

R2

which would be nonplanar and might or might not racemize readily. For this reason some differences of opinion have been expressed about the optical stability of carbonium ions (cf. Hammett, 38). The rate correspondences mentioned above show that this stability is very small for monoketones, but it may well depend upon the actual compound concerned, and perhaps upon the ability of the solvent to solvate the ion. One piece of evidence is sometimes quoted which seems to conflict with the mechanism given above for the bromination and racemization of the cyclic ketone studied by Ingold and his collaborators-the observation of Leuchs (39) that this ketone can be brominated without complete loss of optical activity. However, this bromination was carried out in a solvent differing greatly from that employed by Ingold, and in any case it is doubtful whether the claim of Leuchs can be accepted, since he isolated only a small quantity of product of low optical activity and doubtful purity. The enols and enolate ions of ketones and similar substances will react not only with halogens but with many other substances, and in particular the many addition and condensation reactions undergone by carbonyl compounds in presence of basic catalysts can be referred t o an ionization of the type

\ / CHC:O /

---t

\ / C:C.O-. /

An example of such a

reaction is the reversible aldol condensation of aldehydes and ketones, for which the reaction scheme is:

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

169

Alternative mechanisms have been proposed (Nelson and Butler, 39a), but the above scheme seems the most probable (Walters and Bonhoeffer, 40). b. Nitro Compounds and Nitramide. Nitroparaflns which have at least one hydrogen atom on the carbon atom to which the nitro group is attached behave rather like enolizable ketones in their halogenation reactions. The reaction is of zero order with respect to halogen and is catalyzed by anion bases such as the acetate ion (Junell, 41; Pedersen, 42: Reitz, 43). This was originally interpreted in terms of conversion to the aci-nitro form, which is present at equilibrium only in minute concentrations (Turnbull and Maron, 44) ; however, the arguments already given for the halogenation of ketones indicate here also that the reactive species in halogenation is the ion I in the following scheme: \ CH.N +/ (a)

/

\

0-

0

+ B e \ C:N +/

0-

/

\

+.)

0(13)

The nitroparaffins differ from the ketones in that their halogenation in aqueous solution is not catalyzed by acids: this is because the nitro group is much less basic than the carbonyl group, having in fact hardly any tendency to add on a proton. As would be expected, deuterium exchange between nitromethane and DzO is catalyzed by acetate ions, and the rate of exchange is equal t o the rate of bromination under the same conditions (Reitz, 43). Similar correspondences would be expected in the rate of racemization of optically active nitro compounds, but here the position has been complicated by reports (Kuhn and Albrecht, 45; Shriner and Young, 46) that these compounds do not lose their activity completely on conversion to the ion. However, it has recently been shown (Kornblum et al., 47, 48) that the residual activity in these observations was due to the presence of alkyl nitrate as an impurity. The mutarotation of a-nitro camphor represents a similar type of reaction: it is catalyzed both by acids and bases (Lowry, 49), and a quantitative study of acid catalysis in chlorobenzene solution has been made (Bell and Sherred, 50). Lowry and most subsequent writers have supposed that the observed change of rotation is due t o conversion to the aci-nitro form (11),i.e.,

..;/'

170

R. P. BELL

co

'

-+

CaH~I/I

HbN02

0-

\ C: N +/

O 'H

(1)

(11)

but it is much more likely that what is observed is a change of configuration about the carbon atom marked with a n asterisk. Since the molecule contains a second asymmetric center this will lead to miitarotation rather than racemization. I n catalysis by bases this will ta,ke place through the ion

c:o

\

0-

with a charge distribution somewhere intermediate between the two structures shown. I n acid catalysis th e intermediate may be the aci form (11) above, or it may be the enol (111). C.0H

(111)

There is some evidence th at the mi form is involved, but the question is not settled. The nitroparaffins differ from the ketones in that t:hey are sufficiently strong acids (pK = 8-10) to be completely neutralized hy aqueous solutions of sodium hydroxide. It is well known that this neutralization takes place a t a measurable rate (Hantzsch and Veit, 51). and Hantzsch and others have supposed that the nitro compound is f'rst converted slowly into the aci form, which then reacts rapidly with hydroxyl ions. However, according to modern views, the aci form cannot be produced without the intermediate formation of the anion, and the slow process must therefore be 0-

0

\CH.&/

/

\

+OH--+

0-

\

/

+/ C:N

\

+HzO 0-

a particular case of the reaction with bases (13a). In this case, therefore, the slow protolytic reaction can be observed directly, without having

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

171

recourse to indirect processes such as halogenation, racemixation, or isotope exchange. The decomposition of nitramide has been studied extensively as an example of general basic catalysis, but its mechanism has been little discussed. Nitramide is a weak acid (pK 7), but the loss of a proton cannot be the step concerned with the decomposition, since neutralization and deuterium exchange take place rapidly and without decomposition. There is still some doubt about the structure of the nitramide molecule. It is usually written NHZN02, as suggested by its discoverers (Thiele and Lachmann, 52). Hantasch (53) maintained for many years tjhat it was a stereoisomer of hyponitrous acid, HON:NOH, but eventually abandoned this structure in favor of NH: NO.OH (Hantzsch, 54) in view of optical evidence on the structure of the anion of phenylnitroamine (Cambi and Szego, 55). Other optical evidence (Kortiim and Finck, 56) suggested that in aqueous solution the equilibrium NH2NOz N H :NO.OH is set up, with the form NH :NO.OH predominating, but some kinetic evidence (Bell and Trotman-Dickenson, 57) favors the view that NHzNOz is the predominant form. Measurements of dipole moment were inconclusive (Hunter and Partington, 58) but led to the

-

- +

suggestion of N:N(OH)z, a structure not supported by any other evidence. The most likely mechanism of decomposition seems to be that suggested by Pedersen (59) B

+ NH:NO.OH -+

A

+ N10 + OH-

(13a)

which will be applicable independent of how much of the form N H : NO.OH is present in solution. c. Reversible Addition of Hydroxy Compounds to the Carbonyl Group. Many reactions involving addition to a multiple bond (or the reverse process) are catalyzed by acids and bases, and we shall take here as an example the addition of hydroxy compounds to a carbonyl group. The egneral equation is

\

/

OH

C=O

+ R.OH

’

G! 3 (‘’

‘OR

and the simplest case (R = H) is the hydration of carbonyl compounds. The earliest direct investigation deals with the reversible hydration of carbon dioxide, COz HzO H&03 which exhibits general acid-base catalysis (Booth and Roughton, 60). The same is true for the hydration of acetaldehyde in aqueous solution (Bell and Darwent, 61) and for the dissociation of the same hydrate in aqueous acetone (Bell and Higginson,

+

172

R. P. BELL

62). The carbonyl group of ketones is hydrated only to a very small extent, but the kinetics of the process can be studied by observing the rate of exchange of the isotope 0l8between acetone and water (Cohn and Urey, 63). This exchange is also catalyzed both by acids and by bases. When R is an ethyl group, Eq. (14) represents the reversible formation of a semi-acetal. This reaction is known to be catalyzed (e.g., Dieckmann, 64, 65), but most of the quantitative information available relates to slightly more complicated reactions. I n particular, the ring structures now attributed to the simple sugars involve a semi-aoetal link between the carbonyl and hydroxy groups, and the interconversion of different configurations of the ring involves the breaking of this link. This interconversion is just what is observed in the mutarotaizon of glucose and similar substances, already quoted as one of the best established examples of general acid-base catalysis. An exactly analogoils ring formation occurs in the reversible dimerization of a-hydroxy aldehydes and ketones, according to the scheme

and it is in fact found that the depolymerization of climeric dihydroxyacetone (Bell and Baughan, 66) and glycolaldehyde (Bell and Hirst, 67) exhibit general acid-base catalysis with kinetics very similar to those of the mutarotation of glucose. In this reaction the observed kinetics provide strong support for the structure of the dimer given in Eq. (15). It had been suggested on chemical evidence (Bergmann, 68) that only a loose association took place between the two monomer molecules, though later work (Bergmann and Miekeley, 69) provided an alternative interpretation of this evidence, and confirmed the structure (15) for the dimer. The catalyzed exchange reactions which take place between esters and alcohols (Schaefgen, Verhoek, and Newman, 70) are also explained by a reversible addition to the carbonyl group, i.e.,

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

173

and the mutarotation of certain a-keto esters (McKenaie and Mitchell, 71; McKensie and Ritchie, 72) in presence of alcohol is probably due t o the asymmetric addition of alcohol to the keto group. All these reactions thus fall under the general scheme (14), and in view of the acid character of the hydroxyl group and the basic properties of oxygen reasonable mechanisms for acid and basic catalysis are : Acid catalysis OH

\c/

/

/ \OR

R+B+

\o/

\/ /

+ROH+A

+‘H

Basic catalysis OH

\c’

0

OH +A$\C/

(16)

0-

’

+ B A /

/ ‘OR

+A=

\OR

\c/

0

/

+ROH+B

We shall discuss later the relative rates of these successive steps, and also the question of whether they can be split up into simpler consecutive reactions. d. Esteri3cation and Hydrolysis of Carboxylic Esters. The general reaction is: 0

//

R‘,C

+ HzO

R’.COOH

+ ROH

\OR

being catalyzed in both directions by acids. The hydrolysis of esters can of course be also effected by alkali, but this reaction is not reversible. Moreover, alkaline hydrolysis is a bimolecular replacement of OR- by OH-, rather than an acid-base reaction and will not be considered here. The usual acid catalyst employed is the hydrogen ion, and it is doubtful whether catalysis by other acid species is established with certainty for hydrolysis in aqueous solution (Dawson and Lowson, 73). On the other hand, undissociated acid molecules certainly act as catalysts for esterification in alcoholic solution (Rolfe and Hinshelwood, 74; Hinshelwood and Legard, 75), and any mechanism must therefore be consonant with general acid catalysis. A great variety of mechanisms have been proposed, but many of them are in fact equivalent, differing only in the extent to which the reaction is dissected into steps, and in the assumptions made about the relative rates of the consecutive steps. The following version is that given by Day and Ingold (76):

174

R . P. BELL

R’.C

//

0

0 +A=

\OR

+

R e R ’ .C=O

R’.C

+ R.OH

\O’

+ H ‘

0

Under different conditions different steps may become rate determining, and the further discussion of this mechanism will not b,e attempted here. 3. Kinetic Steps in Acid-Base Catalysis

In all the mechanisms described above the initial step involves the transfer of a proton between the catalyst and the substrate, but it is only rarely that the observed reaction consists solely of this transfer. It is commonly followed either by a second proton transfer from another part of the substrate molecule (thus regenerating the catallyst), or by some other reaction of the species initially formed. We shall now inquire under which conditions it is possible, from kinetic or other data, to obtain information about the relative velocities of the consecutive steps involved in such a mechanism. The following sections follow largely the treatment given by Pedersen (77). a. Reactions Involving a Single Proton Transfer. In any truly catalytic reaction the initial proton transfer must be followed by a transfer in the reverse direction, so that the catalyst may be regenerated, but in many reactions this second transfer does not atiect the kinetics of the reaction. For example, the mechanism (13a) suggested above for the decomposition of nitramide will be followed by the simple acid-base reaction A OH---+ B HzO, regenerating the basic catalyst B, but this will take place rapidly and does not affect the decomposition of the nitramide. We shall take as an example the reaction of a substrate S to

+

+

ACID-BASE CATALYSIS A N D MOLECULAR STRUCTURE

175

give products X, catalyzed by an acid A in aqueous solution. The reaction scheme is (a)

S+A=

ki

k -1

(b)

SH+

SH++ X

+B

(18)

ki

where we suppose for the sake of simplicity that the reaction is irreversible. If the concentrations of A and B are kept effectively constant, the kinetics of the system are described completely by the three first order and kz. The values of k , and k-1 depend upon the constants k l , Ll, nature of A and B and are proportional to their respective concentrations. Information can be obtained about their ratio by considering what happens if the equilibrium (18a) were set up without interference from the reaction (18b). If asterisks denote equilibrium concentrations we then have

where KBH+ measures the acid strength of SH+ (being inversely proportional to the basic strength of S). It should be noted that Eq. (19) does not imply that equilibrium is actually set up in the first stage of the reaction, but only that kl and k-1 have the same values which they would have at equilibrium. [Hf] in Eq. (19) is strictly speaking the hydrogen ion concentration which would obtain if reaction (18a) were a t equilibrium, but this differs inappreciably from the actual value of [H+]in the system for all the cases considered below with the exception of (iib). The state of affairs now depends on the relative values of kl, k-1, and kZ,and the different possibilities can be classified as follows. this will be so in most reactions, since S is (i) First assume k l << L1: normally a much weaker base than B. The amount of SH+ will then remain very small throughout the reaction, and we can equate its rates of formation and destruction. This shows that the formation of the products X will take place with a first order velocity constant k given by

There are now three possibilities, according to the relative values of kz and k-1. (ia) k z >> Ll. Equation (20) becomes k = k l = T A . ~ [A], where T ~is ,a constant ~ characteristic of the proton-transfer reaction between the catalyst A and the substrate S. If the solution contains a number

176

R. P. BELL

of different acid catalysts, this can be generalized to k =

ri[Ai) i

Under these conditions the observed reaction velocity is determined by the rate of the initial proton transfers, and general acid catalysis will be observed. (ib) kz << k-1. (20) becomes k = Iclkz/k-l, and hence from (19)

+

+

In this case the system S A SH+ B is in equilibrium throughout the reaction, and the rate-determining step is the further reaction of SH+. Moreover, the velocity is proportional to the hydrogen ion concentration, although the initial proton transfer takes place from tQe acid A, and it would be classed experimentally as an instance of specific hydrogen ion catalysis. It is easily seen that Eq. (22) is still valid if the solution contains a number of different acid catalysts, and the same conclusion holds. (ic) k z k-1. This gives

-

The velocity now depends partly upon actual proton transfers, and partly on an equilibrium controlled by the hydrogen ion concentration. The reaction would appear to be catalyzed by acid species other than the hydrogen ion, but the quantitative behavior would be complex, and it is doubtful whether this case has been observed in practice. (ii) As an alternative assumption let k-I k l . This will be so if the substrate is a base of moderate strength or if the ratio [A] :[B] (and hence the hydrogen ion concentration) is high. There are, as before, three sub-cases. (iia) k2 >> k-1. The concentration of SH+ will remain small, as in case (i), and the rate is determined by the proton tra-nsfers (lsa), giving a first order velocity constant k = si[Ai], i.e., general acid catalysis.

-

i

+

+

(iib) k z << Ll. The system S A S SH+ B will be in equilibrium throughout, the rate being determined by k z . The fraction of the total amount of substrate present in the form SH+ is k l / ( k , k-l), which is appreciable, since k-1- k l . Unless the substrate concentration is

+

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

177

much smaller than [A] and [B], the two latter concentrations (and hence also [H+]) will not be the same as they would be in the absence of substrate and may tend to vary during the course of the reactions. We shall take [A], [B], and [H+] to represent the actual concentrations in the reaction mixture, and assume that they are maintained a t constant values throughout the reaction. Under these conditions the first order constant for the disappearance of the substrate is given by

+ [SH+]dt (IS1 + [SH+ll -1

= (S]

d

from (19). The reaction velocity is thus a function of hydrogen ion concentration only (specific catalysis) and will be directly proportional to [Hf] for low values of the latter. However, as soon as [H+] becomes + velocity will increase less rapidly than the comparable with K N Hthe hydrogen ion concentration, and for very high values of [H+]it will reach a limiting value corresponding to the complete conversion of S into SH+. (iic) k2 k-I. In this case the variation of substrate concentration with time is given by an expression w’ith two exponential terms (Rakowski, 78) and the reaction will not follow first order kinetics. This behavior is encountered only in isolated cases (Bartlett, 79; Zucker and Hammett, 80). An exactly similar treatment can be applied t o basic catalysis. The important general result of these considerations is that if general acidbase catalysis is observed in a reaction involving only one proton transfer, then this proton transfer is rate determining. However, it is not safe to assume that the converse is true, i.e., that the substrate and catalyst are effectively in equilibrium in reactions found experimentally t o be specifically catalyzed by hydrogen or hydroxyl ions. This is because (as shown in Sec. 11.4) catalysis by species other than H+ or OH- may frequently escape observation, giving a false impression of specific catalysis. The existence of a preequilibrium between catalyst and substrate (leading to a genuine specific catalysis) can, however, be shown unequivocally if Ll kl in the above scheme, i.e., if the substrate has sufficiently marked basic (or acidic) properties, and the concentrations of hydrogen ions (or hydroxyl ions) are sufficiently high. Under these conditions the reaction velocity in solutions of strong acids or bases will increase less rapidly than [H+] or [OH-], corresponding to the conversion of an appreciable proportion of the substrate into its cation or anion [cf. case iib and Eq. (24) above]. The test cannot often be applied, since the

-

-

178

R. P. BELL

necessary values of [H+] or [OH-] are usually so high that measurement becomes impossible. However, a few cases are known in which the data provide good evidence for the existence of a preequilibrium. Thus in the hydrolysis of acetamide by aqueous solutions of strong acids (Euler and Olander, 81) the apparent catalytic constant of the hydrogen ion decreases by 28% between 0.1 N and 1 N , and by 58:!& between 0.1 N and 3 N . Although the magnitude of this change is little greater than might be expected for a primary salt effect, it is significant that such effects appear always to be positive for hydrogen ion catalysis (cf. Sec. II.l), and a positive salt effect has in fact been observed in the hydrolysis of acetamide (Taylor, 82). There is thus little doubt that the initial step in the hydrolysis is the setting up of the equilibrium CH3CONHt

+ H80++ CH3C06H3+ H 2 0

followed by a slower step such as

+

CH3CO&H3 H10 --* CH3COOH

+ &H
b. Reactions Involving T w o Proton Transfers. This is the case most frequently met with in practice, and in particular it includes the so-called prototropic isomerizations represented by the scheme HX.Y:Z

X:Y.ZH

(25)

or the ring-chain equivalent

&)’I

Z

/ *(A) ‘YXH \

ZH (26)

Y:X

where the atoms X and Z may be carbon, nitrogen, 0-r oxygen, and Y is carbon or nitrogen. (Other atoms or groups are of course attached to X, Y, and Z to satisfy their normal valencies.) The best known case of prototropic isomerization is keto-enol tautomerism, already discussed in Sec. III.2a1 and other examples are lactam-laction, nitroso-isonitroso, nitro-acinitro, and three-carbon tautomerism. By analogy with the discussion in Sec. 111.2 the mechanism of basecatalyzed prototropic isomerization will be

where the anion will in general be a hybrid of the two structures shown, the actual charge distribution depending on the nature of the atoms

179

ACID-BASE CATALYSIS A N D MOLECULAR STRUCTURE

X, Y, and Z. We are assuming here that the two proton transfers take place successively rather than simultaneously. The possibility of two simultaneous transfers (the so-called ternary mechanism) will be discussed in the next section. For brevity the scheme (27) will be written ki

ke

k-i

k-z

HR ;=:R- eRH

(28)

Provided that the concentrations of catalyst acids and bases remain constant during the reactlion (ie., the concentration of R- is small throughout) the kinetics of the process can be described by first order velocity constants k l , k-l, k z , k-2, whose values depend upon the catalysts present. We shall also assume that the reaction goes completely from left to right (i.e., kL2 = 0) : this can be ensured if necessary by moving R H as quickly as it is formed. Under these conditions the kinetic analysis follows the lines of the last section. There are three cases: (iiia) k 2 >> k-l. The ion R- is transformed into R H as soon as it is formed, and the observed first order velocity constant is k l * k l represents the rate a t which the various basic species present can remove a proton from HR, and will therefore be of the form zai[Bi]; i.e., general basic i

catalysis is in principle observable. (iiib) k z << Ll. The equilibrium HR

+

c

R-

Bi

+ 2 Ai is praci

i

tically undisturbed, and the velocity is determined by the rate at which R- adds on protons to give RH. The expression for the velocity can therefore be written in the form v = k z [ R - ] = [R-] ?ril [A& However,

i

since H R and R- are in equilibrium with each acid-base pair Ai - Bi, we have [Bi][HR]/[Ai][R-] = Ki’, and hence v = [R-]

c i

ri’[Ai] = [HR]

c i

ni’[Bi]/Ki’

c

(29)

The first order constant for the transformation of HR is therefore ?ri‘[Bi]/Ki’ which is experimentally indistinguishable from ~i[Bi]

1 i

i

found in case iiia. The reaction thus appears experimentally as a general basic catalysis of HR, though the rate is in fact determined by proton transfers from acid catalysts to the ion R-. of the ions R- formed (iiic) k z k-1. Only a fraction kz/(lcz iLl) from HR is transformed into RH, and the first order velocity constant is

-

+

180

R. P. BELL

If only a single acid-base pair has any catalytic effect klkz/(kz+ L1). we can write ki

=

r[B], k2 = r’[A],

k-i = r ” [ A ]

and hence for the observed velocity constant

which is indistinguishable from the results of iiia or iiiib. However, the matter is more complicated if several acid-base pairs axe effective simultaneously. We shall then have

giving k =

2

’ ri‘[Ail

i

*

+ 2 ait’[Ail i

This expression can only be simplified if it is possible to write

where /3 is the same for all catalyzing acids. Equation (31) then becomes

which again corresponds to the usual Iaw for general basic catalysis. However, there is no reason why (32) should be generally valid, especially for catalysts of widely differing strength or structure. When (32) fails, the reaction will show the qualitative characteristics of basic catalysis of HR, but will not follow the usual quantitative laws. In particular, the effect of several basic catalysts present simultaneously will not be additive. There is no clear evidence that this behavior haa been encountered in practice, though it may account for the kinetic anomalies observed in some reactions (King, 83; King and Bolinger, 84). Similar considerations apply to acid-catalyzed prototropy, the analogous scheme for the transformation being HX.Y:Z

+A

{HT*y:in) + HX :v.ZH

lj

1

(34)

181

ACID-BASE CATALYSIS A N D MOLECULAR STRUCTURE

In this case, however, the representation of the intermediate ion by a hybrid structure is not always correct. For example, in keto-enol

l l + I

tautomerism [cf. Eq. (lo)] the cation can only be written as HC.C:OH, since the alternative form would give carbon a covalency of five. In the

I l l I I

I l l I I

so-called three-carbon tautomerism, H C C : C S C: GCH, and the Schiff

I

l

l

1

base isomerism, HC*N:C G? C:N.CH, neither form shown in (34) is , I I I I I

I l l

allowable, and the actual structures of the cations must be H C . C C H / + I

I

1

and H C - N C H respectively, with a sextet of electrons on the central

l + l

atoms. The formulation given in (34) is correct only when the atoms X and Z each have an unshared pair of electrons in their normal valency

I 1

I I

state, e.g., in the tautomerism of the amides, H N C :0 s N :G O B , the

I 1

I 1 +

cation can be written as HN :C.OH or H N C :OH.

+

In either event (34) can be abbreviated as kz

ki

HR

k-1

(HRH)+

RH

(35)

k-2

by analogy with (28)) and the further treatment is exactly analogous. If we again assume that the concentration of (HRH)+ is always small, and that the formation of R H is irreversible, our conclusions can be stated briefly as follows. If k , >> k-1 the rate is determined by proton transfers from acid catalyst species to the molecule HR. If k z << k-l the ion (HRH)+ is throughout in equilibrium with HR, and the rate is determined by proton transfers from this ion to basic catalyst species. In either case general acid catalysis of HR will be observed. If k z k M 1 the rate is determined by both of the consecutive reactions, and the effect of several catalysts acting simultaneously may not be additive. The results of the above analysis of reactions involving two proton transfers may be compared with those for reactions in which only a single transfer is kinetically relevant (Sec. III.3.a). For a single transfer the existence of an equilibrium between catalyst and substrate always produces the appearance of specific catalysis by hydrogen or hydroxyl ions, and the detection of general acid-base catalysis therefore excludes

-

182

R. P. BELL

the existence of such an equilibrium. This is not the case when two proton transfers are involved, since here the existence of .B preequilibrium does not affect the observation of general catalysis by acids or bases. It is easy to show that the observation of a quantitative :relation between catalytic effect and acid-base strength is similarly unaffected. For example, consider case iiib above, where there is preequilibrium between the substrate R H and the basic catalyst Bi, the rate being determined by proton transfers between R- and the acid A;. From (29), the catalytic constant observed for the base Bi will be rjl/K,', where IG' is the equilibrium constant [Bi][HR]/[Ai][R-].Ki' is equal to Ki/KH Rl where K, and K H Rare the acid dissociation constants of Ai and HR, and ri' will be related to Ki by an expression of the form ril = GKp, where G is the same for a series of similar acids and a is less than unity. The observed catalytic constant for the base Bi is therefore equal t o GKHR(l/Ki)l-a, and since KHR is the same for a series of catalysts this is exactly the same as the relation normally found between catalytic power and acid-base strength [cf. Eq. (9)]. c. Kinetic Mechanism in Nonaqueous Solvents. The analysis given in the two preceding sections will apply in principle not only to aqueous solutions but also to those in similar solvents such as the alcohols, though its application in practice would be more difficult because of our incomplete knowledge of acid-base equilibria in solvents other than water. The position is different] however, in the so-called aprotic solvents (cf. Sec. 11.3) which are devoid of both acidic and basic properties. In the first place the solutions contain no analogues of the hydrogen or hydroxyl ion, so that there is no meaning in speaking of' specific catalysis by these ions, and in the second place the strengths of acids and bases can no longer be defined by the usual dissociation constants] which involve reference to a reaction with the solvent. Consider acid catalysis of a reaction involving only a single proton transfer, which we can write as S

+ Ai k-1ki

SH+-+ X ka

SH+

+ Bi-

I

We shall assume that the amount of SH+ present is throughout much smaller than [S] or [Ail. If the first step is rate determining] then the reaction will be first order with respect to both catalyst and substrate, with a catalytic constant depending on the nature of Ai. It is reasonable to suppose that, for a series of similar catalysts the catalytic constants will run parallel with their acid strengths, but since acids do not dissociate in the aprotic solvents being considered, some other reaction must be

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

183

used. This may be either reaction with a standard base (e.g., an indicator) in the solvent being considered, or dissociation in some other solvent (usually water) possessing basic properties. Experience shows that the relative strengths found for a Aeries of acids are approximately independent of the solvent, and either method may therefore be used for comparison with the catalytic effect. If the second step (b) is rate determining, the rate is proportional t o the concentration of SHf, which is governed by the equilibrium [SH+][Bi-]/[S][Ai] = Ki. In contrast with aqueous solutions, [Bi-] is not now determined by the reaction of A; with the solvent. If none of the base Bi- has been added, and if the reaction product X has no basic properties, then Bi- can be formed only by the reaction of Ai with s, and we have [ B i ] = [SH+]. Under these conditions the reaction velocity will be proportional to [S]’*[Ai]’, i.e., the reaction will no longer be first order. This result will be modified if the product X has some basic properties, but we shall not pursue the various possibilities, as it is doubtful whether they have been encountered in practice. The position is simpler if two proton transfers are involved. For example, consider an acid-catalyzed prototropic change, proceeding according to the scheme HR

+ Ai

ki

ki

(HRH)+ k-1

+ Bid+

RH

+ Ai

(37)

assuming as usual that the concentration of (HRH)+ is always small, and that the formation of R H is irreversible. If the first step (a) is rate determining, we shall have general acid catalysis following the usual laws. If (a) is in equilibrium and (b) rate determining, then the velocity will be proportional to [HRH+][Bi-], and hence by the law of mass action t o [HR][Ai]. The reaction thus again shows general acid catalysis with normal kinetics, and (as in aqueous solution) the question of whether (a) or (b) is the rate-determining step cannot be directly decided from kinetics. In the above treatment it has been assumed that the ions (HRH)+ and B; exist in the free state, as they undoubtedly do in aqueous solution. However, solvents which are aprotic, e.g., hydrocarbons, normally have such low dielectric constants that the majority of ions present will exist at least as i9n pairs, and partly as larger aggregates. This circumstance may well affect the kinetics. For a reaction involving only a single proton transfer the concentration of ion pair will be directly proportional to the concentrations of catalyst and substrate, so that the kinetics will be simpler than when free ions are present. On the other hand, in reactions involving two proton transfers, if the first product of

184

R. P. BELL

reaction between HR and Ai is the ion pair (HRH)+(B,)-, there are several possibilities for the subsequent reaction. If the second proton transfer to give R H and A, can take place within this ion pair, then we shall still observe normal kinetics, as described in the last paragraph. However, it may well be that the orientation of the components in the ion pair is unsuitable for the formation of RH, and reaction will then be completed only by collision of the ion pair with anothier species, e.g., a second ion pair or a second molecule of HR. This will obviously lead t o more complex kinetics. To sum up, acid-base catalysis may often follow normal kinetics in nonaqueous solvents, which may then be used for investigating reaction mechanisms in the ordinary way. On the other hand, more complicated kinetics may sometimes occur, especially in aprotic solvents, and caution must be exercised in drawing conclusions without investigating the kinetic behavior experimentally. d. The Use of Deuterium in Determining Kinetic .Mechanisms. We have already seen (Sec. III.2.a) that observations on the rate of hydrogendeuterium exchange can be used to elucidate the mechanism of other chemical processes. It is also possible t o obtain useful information by studying the velocity of a catalyzed reaction when the hydrogen in the catalyst or the substrate (or both) is replaced by deuterium. From a theoretical point of view it is certain that the transfer of a deuteron will always take place more slowly than the transfer of a proton, because of the higher zero point energy in the link containing the lighter isotope. On the other hand, if there is a pre-equilibrium such as S H30+$ SH+ HzO, the effect of isotopic mass will involve the relative values of the two constants KH = [H3O+][S]/[SH+]and KD =- [D3O+J[S]/[SD+], i.e., the dissociation constants of the acids SH+ and SD+. It is not possible to predict the relative magnitudes of KH and KD on theoretical grounds, but experiment has shown that in aqueous solution KH > KD. This means that under comparable conditions the con centration of SD+ will be greater than that of SH+, and the pre-equilibrium will therefore tend to make the deuterium compound react faster than the hydrogen one. Examples are in fact known of both kD/kE> 1 and kD/kH < 1, where kE and kD are the velocity constants for the hydrogen and deuterium compounds respectively. If b e find experimentally kD > kH in a catalyzed reaction, then an acid-base equilibrium must always be involved in the kinetic scheme, If the reaction involves only one proton transfer, then the converse is also true for catalysis by hydrogen ions, i.e., if k H > h D , then there is no pre-equilibrium. On the other hand, in a reaction involving two successive proton transfers the pre-equilibrium and the subsequent proton

+

+

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

185

transfer may be affected in opposite directions by isotopic substitution, SO that the observation that kH > kD does not indicate whether or not there is a pre-equilibrium. The isotope effect is therefore not so helpful as was at one time supposed in elucidating kinetic mechanisms, but examples will be given in the next section of how it can be used in conjunction with other data. It is also possible in principle to obtain information by studying rates in a series of solvents varying in isotopic composition from 100% HzO to 100% DzO, but the method is not a sensitive one and there has been some difference of opinion about the interpretation of the results (Gross, Steiner, and Suess, 85; Orr and Butler, 86; Nelson and Butler, 87; Brescia and La Mer, 88; Brescia, 89; WynneJones, 90; Gross, 91; Reitz, 92; Hamill and La Mer, 93). e. Examples of Kinetic Analysis. The preceding paragraphs have given in some detail the possibilities which exist for determining the nature of the rate-determining step when the mechanism is qualitatively established. There are not many cases where such an analysis has been carried through in practice, and in this section we shall give only a few examples in which the conclusions are fairly well established. In the halogenation of ketones and allied reactions (e.g., isotope exchange, racemization) we have seen in Sec. III.2.a that a distinction must be made between acid and base catalysis. In catalysis by bases the reactive species is the anion formed in the first step of the reaction, and the reaction thus involves only a single proton transfer. The independence of halogen concentration and the identity of rates of halogenation, racemization and isotope exchange prove that this proton transfer is rate determining. This conclusion is borne out by the fact that all these reactions exhibit general basic catalysis and is consonant with the fact that kH > k D for catalysis by acetate ions (Wilson, 94). I n acid catalysis, on the other hand, the above reactions depend upon the actual formation of the enol, involving two successive proton transfers. There is good reason to believe that the first step,

\ / \ /* CH.C:O + A + CH.C:OH / /

+B

is in equilibrium, the subsequent removal of a proton from the action being rate determining. This is most convincingly shown by the isotope effect, where kD > k , for hydrogen ion catalysis (Reitz, 94a), showing the existence of a preequilibrium. Further it is found that the addition of acids t o optically active ketones in nonaqueous solvents causes an instantaneous change of rotation, followed by a slow racemization. The instantaneous change is presumably due to reaction (38), and this view

186

R. P. BELL

receives confirmation from cryoscopic studies (Bell and Caldin, 95; Bell, Lidwell, and Wright, 96). Indirect evidence on the same point comes from a comparison of the rates of acid-catalyzed halogenation for a series of ketones. If (38) were rate determining, there should be a parallelism between the reaction velocity and the basic strength of the ketone. If, on the other hand, (38) is in equilibrium, this parallelism may be absent, since the second step of the reaction involves the removal of a proton from a different part of the molecule. Hammett and his co-workers have measured the relative basic strengths of a number of substituted acetophenones by observing their ultraviolet absorption spectra in sulfuric acid-water mixtures (Flexser, Dingwall, and Hammett, 97 ; Flexser and Hammett, 98) and have also measured the rates of halogenation of the same ketones under acid conditions (Zucker and Harnmett, 99). No parallelism was found between the two sets of measurements again confirming the view that reaction (38) is in equilibrium. If it is possible to measure the actual rate of formation of an enol from its keto isomer (or vice versa), then of course two successive proton transfers are involved in catalysis by both acids and bases. It is rarely feasible to study this interconversion directly, but it is sometimes possible to obtain indirect evidence about its kinetics. The typical scheme for the base-catalyzed conversion of keto to enol is

\ / / \ / C:C.OH + B C H C : O + l3z \C:C.O- + A + / k-i / kl / where we assume that the enol is removed as soon as it is formed. If the ketonic substance is a sufficiently strong acid, it will be possible to prepare a solution in which it is chiefly present as the anion. If this solution is acidified, the anion will be converted into keto and enol in the proportion k-l:kz. This experiment has been carried out with acetoacetic ester (Pedersen, 100) where the product of acidification is 100% enol, i.e., in this case kz>> k-I. This means that practically every anion formed is converted into enol, and the rate of enolization by bases will be identical with the rate of ionization (or halogenation). This result contrasts with that obtained for a Substance of similar constitution, the menthyl ester of a-phenylacetoacetic: acid. By comparing the rates of mutarotation and enolization of this substance, catalyzed by a solution of piperidine in hexane, Kimball (101) concluded that for every anion which goes on to form enol, approximately two revert to the keto form. There is no obvious explanation of this difference between two apparently similar compounds, but it may be that the interpretation of the data in hexane solution is invalidated by some of the kinetic complications described in the preceding section.

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

187

It is also possible to draw some conclusions about the detailed kinetics of the aldol condensation, which also goes by way of the anion of an aldehyde or ketone. We shall assume that the sequence of reactions is that given in Eq. (12). The condensation of two molecules of acetaldehyde to acetaldol in presence of hydroxyl ions is, surprisingly, of the first order with respect t o aldehyde (Bell, 102), demonstrating that the first step (a) must be rate determining, followed by two faster steps, (b) and (c). This means that the reaction should in principle exhibit general basic catalysis, but this has not been detected experimentally. There are no direct kinetic observations on the condensation of acetone to diacetone alcohol, but the reverse depolymerization reaction is well known to be of the first order, so that the condensation reaction must be of the second order in acetone. This suggests that for this substance reaction (a) is in equilibrium, with reaction (b) as the probable rate-determining step. Under these conditions the reaction should exhibit specific catalysis by hydroxyl ions, as is in fact observed, the catalysis exerted by amine molecules depending on a different mechanism (Westheimer and Cohen, 103; Westheimer, 104; Westheimer and Jones, 105). Further evidence comes from a study of the rate of deuterium interchange with acetone (Walters and Bonhoeffer, 40). The rate of Condensation of two acetone molecules can be calculated from the known rate of depolymerization of diacetone alcohol, and the measured equilibrium constant. I t is found to be very much slower than the observed rate of isotope interchange. Since interchange will take place every time the anion is formed, this result supports the view that in the condensation reaction (a) is in equilibrium, the rate being determined by the slower reaction (b). There is less evidence available for the reuersible addition of hydroxy compounds to the carbonyl group (cf. Sec. III.2.c). If we accept the mechanism in Eq. (16), it is likely that the reactions on the left are in equilibrium both for acid and base catalysis, since they represent simple protolytic reactions with no change of bond structure (cf. the discussion in Sec. V.2). The only reaction of this class for which the isotope effect has been investigated is the mutarotation of glucose (Pacsu, 106; Moelwyn-Hughes, Klar, and Bonhoeffer, 107 ; Moelwyn-Hughes, 108; Hamill and La Mer, 109). It is found that deuterium substitution causes a decrease in velocity for catalysis by water, acetate ion and hydrogen ion, but we have just seen that in a reaction involving two proton transfers this does not exclude the possibility of a preequilibrium. There are two instances in which the value of kD/kHdepends upon the conditions under which a reaction is carried out. In the acid-catalyzed hydrolysis of ethyl orthocarbonate Wynne-Jones (110) finds kD/k, > 1 for catalysis by hydrogen ions, but k,/k, < 1 for catalysis by acetic acid.

188

R. P. BELL

Wynne-Jones takes this to mean that preequilibrium is set up in hydrogen ion catalysis, but not in catalysis by acetic acid; however, another explanation seems preferable. In hydrogen ion catalysis the value of ko/ka depends upon the ratio of the two equilibrium constants, [SU+]/ [S][D,O+]and [SH+]/[H30+][S],which is normally greater than unity. For catalysis by acetic acid the corresponding equilibrium constant is [SH+][CH3COO-]/[S][CH3COOH],which may be written as the ratio of two constants [SH+]/[H30+][S]and [CH3COOH]/[H30+][CH3COO-]. Each of these last two constants will be increased by replacing hydrogen by deuterium, and their ratio may well decrease, leading to k D / k H< 1. The other example is the hydrolysis of acetarnick in slolutions of strong acids, where it has been found that k D / k H= 1.50 in 0.1 N acid, 1.00 in 2.3 N acid, and 0.86 in 4.0 N acid (Reitz, 111,112). This accords excellently with the conclusions reached in Sec. III.:3.a, according to

+ which a large proportion of the acetamide is present as the ion CH3CONH3 in the strongly acid solutions. In the more dilute solutions the value kD/kH > 1 depends upon the greater proportion of acetamide converted t o cation in the deuterium system. In concentrated mid, on the other hand, most of the acetamide will be present as cation in both systems,

+

and kD/k,

< 1 because the ion CHICOND3 reacts

further more slowly

+ than CH3CONH3. f. The Question of Ternary Mechanisms. In all the above treatment it has been assumed that when two proton transfers occur they do so successively in two distinct steps. It is equally possible to imagine that both transfers take place simultaneously by the approach of an acid and a base to different parts of the reacting molecule; for example, in a prototropic change HX*Y:Z -+ X :Y.ZH the reaction would be

Such a mechanism has been advanced many times (Lowry and Faulkner, 113; Lowry, 114) but has not generally been accepted. However, recent work has shown that some of the arguments advanced against it are of questionable validity, and we shall therefore discuss t'he question here. The kinetic consequence of a ternary mechanism is t8hatin any mixture of acids and bases the reaction velocity should be given by an expression of the form

ACID-BASE CATALYSIS A N D MOLECULAR STRUCTURE

189

instead of the usual expression

which we have shown to follow from the assumption of two consecutive proton transfers. In the ternary mechanism there is no distinction between acid and base catalysis, and any acid can cooperate with any base in the “push-pull” mechanism represented by Eq. (39). However in water or a similar solvent it is not easy to distinguish experimentally between (40) and (41), since a large proportion of the observed velocity will be due to terms in which the solvent (present in high and constant concentration) is acting either as the acidic or as the basic partner. There is in fact only one set of data in aqueous solution which appears to show the existence of a ternary term. Dawson and Spivey (115) found that the rate of iodination of acetone in acetic acid-acetate buffers could be represented by the expression 10% = 0.006

+ 560[H30f] + 7 X 1OB[OH-]+ 1.3[CHaCOOH] + 3.3[CH3C00-] + 3.5[CH3COOH][CH3COO-]

(42)

The quantitative interpretation of the data may be open to some doubt, since the experiments were carried out at high and variable ionic strengths, but the existence of a measurable contribution from the last term in (42) is probably established. Nevertheless, it has been usually held that the reaction cannot go exclusively or even chiefly by a ternary mechanism, using the following argument due to Pedersen (116). Omitting the terms in [HaO+] and [OH-], the remaining terms of (42) will be represented as follows on the ternary scheme: (a) (b) (c)

(4

Acid HzO HAc

Base

+ acetone + HzO + acetone + HzO H 2 0 + acetone + AcHAc + acetone + Ac-

Relative Velocity 0.006 1.3 3.3 3.5

By comparing (a) and (b) it will be seen that on changing the acid from HzO to HAc (the base being in each case H20) the rate is increased by a factor of 220. We should therefore expect a similar increase in going from (c) to (d), where the acid again changes from HzO to HAc and the base is in both cases Ac-. The observed value for reaction (d) thus appears to be 200 times smaller than might be expected on the termolecular hypothesis. However, it has been recently shown by Swain (117) that the above criticism of the termolecular picture is not valid. He points out that there is an ambiguity in interpreting the fourth and fifth terms of Eq. (42)

190

R. P. BELL

on the termolecular hypothesis. Because of the dissociation equilibria of acetic acid and water it is impossible to distinguish kinetically between the systems HAc acetone HzO and H30+ acetone Ac-, or acetone H20 and OHacetone -4- HAc. Swain between Acmakes the reasonable assumption that the relative reactivities of different acids are independent of the base with which they cooperate (and similarly for different bases); i.e., that Eq. (40) can be rewritten as

+

+

+

+

+

+

+

thus reducing the number of constants. On this basis it can be shown that on the ternary hypothesis the term 3.3[Ac-] in (42) arises mainly from the reaction OHacetone HAc, and not from Acacetone HzO. Under these conditions the above criticism of the ternary mechanism is no longer valid, and it can in fact be shown that the term in [Ac-][HAc] corresponds within a factor of two to the predictions of (43). Swain has also shown that the corresponding term would not be detectable experimentally in the mutarotation of glucose. Less direct support for a ternary mechanism come:3 from measurements by Ingold and his collaborators (118,119) on the base-catalyzed interconversion of azomethines in alcoholic solution. It was possible to measure both the rate of racemization of one isomer and its rate of conversion to the other isomer. These two rates were found to be equal. If the mechanism involved the formation of a free anion, the rate of racemization should be greater than that of conversion, since the known position of equilibrium shows that an appreciable fraction of the anions would revert to the original compound instead of going on to form the isomer. The experimental findings thus suggest that rto kinetically free anion is formed, and support a ternary mechanism for this reaction. It should in principle be easier to test the ternary hypothesis when there is not present a constant excess of acidic or basic solvent molecules, i.e., by using solutions of acids and bases in inert solvents. There are very few experiments of this kind, but an observation which is often quoted in favor of a ternary mechanism is that of Lovvry and Richards (120), who found that the mutarotation of tetramethylglucose took place very slowly in dry pyridine (possessing no acid properties) or in dry cresol (possessing hardly any basic properties), but was rapid in a mixture of the two solvents, or in either solvent when moist. This observation shows that the reaction is facilitated by the simultaneous presence of an acidic and a basic substance, and can readily be interpreted in terms of the ternary picture.

+

+

+

+

ACID-BASE

CATALYSIS AND MOLECULAR STRUCTURE

191

However, in spite of these various pieces of evidence in favor of a ternary mechanism, it is doubtful whether it can be accepted as valid for acid-base catalysis in general, as there are a number of arguments in the reverse direction. Although further work on the acetone-iodine reaction (Bell and Jones, 120a) has confirmed the magnitude of the term in [HAc][Ac-] found by Dawson and Spivey (115), the same authors show that no product term is detectable in glycollate buffers, although calculation by Swain’s method predicts a much larger contribution than in acetate buffers. Further evidence adverse t o the ternary mechanism comes from a study of the hydration of acetaldehyde, a reaction which is chemically very similarly to the mutarotation of glucose (cf. Sec. III.2.c). Bell and Clunie (121a) failed to detect any product term in a kinetic study of this reaction, even under conditions where Swain’s treatment predicts that 75% of the total velocity should derive from such a term. In nonaqueous media, it is difficult to interpret with certainty the observation of Lowry and Richards (120), since such drastic changes of medium are involved in going from pure pyridine to pure cresol. Other observations on acid-base catalysis in nonaqueous solvents speak against the ternary hypothesis. The interconversion of the two forms of mesityl oxide oxalic ester (an example of ring-chain tautomerism) in the aprotic solvent chlorobenzene is catalyzed by dilute solutions of acids or bases, but the velocity in a solution containing both acid and base is no greater than the sum of the velocities due to the acid and the base separately (Bell and Rybicka, 121). This result is in direct contrast to that obtained by Lowry and Richards for the mutarotation of tetramethylglucose, and the system studied by Bell and Rybicka has the advantage that dilute solutions in the same solvent were used throughout, showing definitely that the ternary hypothesis cannot apply to their reaction. The same conclusion may be drawn from the fact that solutions of acids or bases alone in aprotic solvents frequently show a reproducible catalytic effect with simple kinetics. In conclusion it should be pointed out that the ternary picture cannot be dismissed merely on the grounds that a triple collision is an improbable event. Exactly the same result is achieved kinetically if two of the molecules concerned associate together loosely (e.g., by hydrogen bonding) before the approach of the third. Although the modern tendency is to split up complex reactions into consecutive bimolecular steps (Hinshelwood, 122) there appear to be cases in which this cannot be done (Bell and Darwent, S l ) , and the ternary hypothesis must certainly be considered as one of the possible mechanisms for reactions catalyzed by acids and bases.

192

R. P. BELL

IV. THEVELOCITY OF ACID-BASEREACTIONS 1. Pseudo-Acids and Pseudo-Bases

I n the preceding sections we have frequently postulated reactions which take place a t a measurable speed, which are formally acid-base reactions involving the transfer of a proton from one molecule t o another. Since such reactions are commonly thought of as being very fast, we must examine how far the idea of acid-base reactions of measurable speed is consistent with information outside th e field of acid-base catalysis. It is common experience that many acid-base reactions (e.g., dissociation, hydrolysis, neutralization) take place too rapidly to be followed by ordinary methods, even when very weak acids and bases are involved. However, a few instances are known in which a slow reaction can be observed, the best known of which is th e reaction of nitroparaffins with hydroxyl ions. For example, nitromethane, CHaN02, is neutralized at a measurable speed by hydroxyl ions t o given an anion whose structure 0-

+/

is presumably CH2:N

\

.

This phenomenon was discovered by

0Hantzsch (123), who introduced the term pseudo-acid to describe th e nitroparaffins and similar substances. Hantzsch supposed that the slow change was t he transformation of the nitroparaffin into the aci form, e.g., OH

+/

CH2:N

\

,

which then reacted rapidly with hydroxyl ions.

How-

0ever, we have already seen (cf. Sec. III.2.b) that the interconversion of two tautomers of this kind is now believed to involve the anion as a n intermediate, and we should now regard the actual loss of a proton from the methyl group to the hydroxyl ion as a slow process. A similar reaction has been investigated recently by Lewis and Seaborg (124). The substance tri-(p-nitropheny1)-methane reacts with sodium ethoxide t o give a colored anion with the structure

-

0

-/ \ d = ~ =c(p.C,H,No*). 0

in which the negative charge will be equally distributed between th e three nitro groups. If this alcoholic solution is acidified at -80" with

ACID-BASE

CATALYSIS AND MOLECULAR STRUCTURE

193

various weak acids, the color changes at a measurable rate owing to the slow regeneration of the original compound. This case is of particular interest, since there is an approximate correlation between the rate of reaction (which was not measured accurately) and the strength of the acid used for neutralization, showing an obvious analogy with Bronsted’s relation between catalytic power and acid-base strength (cf. Sec. 11.4).* In the examples quoted above there is a considerable difference in electronic structure between the pseudo-acid and its anion, and it may be surmised that the relative slowness of the reactions observed is connected with this electronic reorganization. The change in electronic structure is also associated with a change in absorption spectrum, and Hantzsch has proposed that this latter change should be used as a criterion for pseudo-acids. On the basis of slight optical changes on ionization he has concluded (126,127,128) that almost all acids (e.g., halogen hydrides, nitric acid, sulfuric acid, carboxylic acids) are pseudo-acids, but his interpretation has been challenged by a number of authors (Fajans, 129; von Halban, 130; Ley and Hunecke, 131). I n any case, such a wide extension of the term would destroy its usefulness, and it is best to reserve it for the more extreme types of behavior, though we shall see later that it is impossible to draw any sharp line of demarcation between pseudo-acids and “true ” acids. In many instances it is not practicable to investigate directly acidbase reactions of the ordinary type, since if very weak acids or bases are involved the extent of reaction is very small. Under these conditions measurements of rates of racemization or of deuterium exchange may serve as an indirect method of measuring rates of ionization. For example, hydrogen attached to the a-carbon atom of ketones, carboxylic acids, esters, nitriles, and similar substances exchanges very slowly (if a t all) in neutral aqueous solution, but a t a measurable speed in presence of hydroxyl ions (Bonhoeffer, 132). Although the amount of anion present is very small, even in strongly alkaline solution, the rate of H-D exchange can be used to measure the rate of its formation. Similarly, the rate of C6Hs racemization of the ion

\ CDCOO/

in alkaline solut,ion is

CH3CsH4 equal to its rate of isotope exchange (Ives and Wilks, 133), both rates being determined by ionization at the asymmetric carbon atom.

* Lewis and Seaborg (124) give a different and more complicated explanation of the observed facts, but there are good arguments for rejecting their interpretation (cf. Kilpatrick, 125).

194

R. P. BELL

All the instances of slow acid-base reactions mentioned so far involve a considerable structural change in one of the acid-base pairs taking part in the reaction. This statement rests partly on theoretical considerations (notably the improbability of structures with a charge on a carbon atom), and partly on direct observations of the change in absorption spectrum accompanying the reaction. These reactions are of exactly the same type as those occurring in the schemes proposed.in Sec. 111.2 for various acid-base catalyzed reactions, and the behavior of pseudo-acids thus supports the conclusions already reached about these mechanisms. In catalyzed reactions involving two steps it is frequently the case that one of these involves a structural change and the other does not. For example, in the acid-catalyzed prototropy of ketones

\ / CH*C:O + HA /

\

A

CH*C:OH

+ A-

/ \ / -+ C:C.OH + H A /

(b)

step (a) involves no rearrangement of valencies, while in step (b) there is a shift of the double bond. As we have already seen, there is evidence that the equilibrium (a) is set up rapidly, while (b) is a slow reaction. This again accords with the picture of pseudo-acids given above. In extreme cases the valency rearrangement may involve the break-up of a molecule into two separate pieces, or the reverse process. For example, in the reversible reaction between an aldehyde and a hydroxy compound (cf. Sec. III.2.c) the rate-determining step is believed t o be

R-CHO

+ HA + R’*OH

/OH RCH R

+ A--

\O/ +H ‘

and

/ R.CHO + B + R’.OH S R.CH \

OR

+ 13H’ 0-

for catalysis by acids and by bases respectively. Although on paper each of these processes can be dissected into two conseciitive bimolecular reactions in a number of ways, closer examination shows that none of these dissections is consonant with the facts (Bell and Darwent, 61).

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

195

The rate-determining steps thus involve simultaneously a proton transfer and a change in the number of molecules, thus showing analogy with the usual picture of a pseudo-acid. The term pseudo-base has been used less widely and less consistently than pseudo-acid. Logically it should be applied t o a species which undergoes a change of structure when it adds on a proton. This would include the anions of the pseudo-acids discussed above, and a few uncharged species such as coloring matters (e.g., anthocyanins and flavones) derived from y-pyrone, where the addition of a proton involves the reaction CH.CH

/CH:CH \ C:O 0

\

/

+ H + e +O/

\

CH: CH

%.OH CH: CH

/

However, the term pseudo-base has been more widely used to describe compounds which can lose a hydroxyl ion with change of structure, and hence undergo slow reactions with acids, which are reversed by the addition of base. Examples are the “carbin01 bases” of various triphenylmethane dyes, such as crystal violet, and derivatives of pyrazine and acridine (Hantzsch and Kalb, 134; Aston, 135, 136). However, this kind of change is not strictly analogous to an ordinary acid-base reaction, and it is probably better not to use the term pseudo-base in this connection. I n any case, processes involving transfer of hydroxyl ions are not believed to play any important part in catalyzed reactions. We have seen that acid-base reactions proceeding a t a measurable rate frequently involve structural changes, but this is not always the case. If an acid-base reaction involves an acid or base which is extremely weak, it will be considerably endothermic in one direction, and the activation energy for reaction in this direction must be at least as great as the endothermicity, independent of any structural factors. For example, acetylene and water do not exchange deuterium under neutral conditions, but do so in presence of a considerable concentration of hydroxyl ions (Reyerson, 137). This process must depend upon the slow reaction CzHz OH- -+ CtHHzO, which does not involve structural change, but which is presumably markedly endothermic. The same applies to the slow ionization of aliphatic amines in alcohol observed by Ogston (138), where the slow stage may be the ionization of the alcohol itself; similarly, in the zero order halogenation reactions studied by Hughes (139) the rate-determining step is probably the endothermic protolysis HOCl H 3 0 + - +[H20Cl]+ HzO (cf. Bell and Gelles, 140) rather than the production of C1+.

+

+

+

+

196

R . P. BELL

2. The Molecular Interpretation of the Bronsted Relation

If it is accepted that reactions catalyzed by acids or bases involve at some stage a slow acid-base reaction, then the Bronsted relation (cf. Sec. 11.4) assumes a reasonable aspect. The constants used to express the strengths of the catalyzing species are usually defined with reference to an equilibrium with some standard acid-base system such as the solvent, but they could in principle be defined in terms of the (hypothetical) protolytic equilibrium between the catalyst and the substrate. The Bronsted relation then amounts to a parallelism between the rates and equilibrium constants of a series of similar reactions. The general form of the relation can in fact be inferred without rtny reference to a molecular interpretation. Suppose that we have any acid-base equilibrium Ai+Bz=Az+Bi governed by the equilibrium constant [AZI[BII/[AII[BZ]= K

and let the velocities of proton transfer in the two direc.tions be

If K1 and K z are the strength constants of A1 and A2 on any scale, then clearly =

m,&?,i

K = Ki/Kz

Now let A2, B2 be a fixed acid-base pair (corresponding to the substrate) while A1, B1 varies through a series of substances (catalysts) of increasing acid strength in A, and hence decreasing acid strength in B1. As we pass along the series it is reasonable to suppose that ~ 1 will , increase ~ steadily while 7r2,1 will decrease. Since, however, the ratio ~ 1 , 2 / 7 r 2 , 1 must be increases less rapidly than proportional to K1, we must suppose that rl,:, K1, while r 2 , decreases 1 less rapidly than l/Kl. This inay be expressed by writing d log ~ - d log ~

= ad log K1 = (1 - m ) d log K I

1 . 2

2 , i

1

(45)

where 1 > a > 0. If a is effectively constant over some range of acid strengths these equations can be integrated to give ~ 1 . 2=

gBZKP,

TZ,I

= g.4z(l/K1)'-a

(46)

where gA2and gB2 are characteristic of a given substrate, solvent, and temperature. If the measured reaction velocities are i:n fact determined

ACID-BASE CATALYSIS A N D MOLECULAR STRUCTURE

197

by q 2 or 7r2,1, then these equations are identical with the Bronsted relation as usually expressed. The molecular picture behind this reasoning can be demonstrated by the use of potential energy curves. Subject t o certain reservations (cf. Bell, 3, Chapter VIII) the transfer of a proton can be represented by the passage of a point along the full curves in a diagram like Fig. 1. If the reaction is X H Y X HY, then curve I represents the energy of XH for varying separation of the proton, and I1 is the corresponding curve for HY. E represents the energy of activation for the reaction

+

+

~

\

I

I

DISTANCE OF PROTON

FIG.1.

from left to right, and e is the total energy change for the same reaction. If now the nature of Y is changed a little by chemical substitution, curve I1 will be replaced by a slightly different curve 11', and the new energies of activation and reaction are respectively E' and e'. If the change in Y involves only a vertical displacement of curve I1 and not any change in its shape or horizontal position, then it is clear from the geometry of the figure that 6E

= ci6t

(47)

where a is a quantity less than unity such that a / ( l - a) is the ratio of the slopes of the two curves a t their point of intersection. Equation (47) is clearly equivalent t o (45) if we can write 6 log P =

RTsE,

6 log K = RT6t

(48)

198

R. P. BELL

Equation (48) is of the expected form for relations between reaction velocities and activation energies on the one hand, and between equilibrium constants and heats of reaction on the other. However, there are difficulties in the way of using (48) as a quantitative basis-for the Bronsted relation. In the first place, the quantities E and B in the diagram refer strictly to the behavior of the system at absolute zero r3ince no account is taken of the internal thermal energy of the molecules. In the second place experiment shows that even in a series of similar reactions the observed velocities and equilibrium constants often involve variations in entropies of activation and of reaction, and not only energy changes. These difficulties are not yet fully resolved, but there seems little doubt that diagrams such as Fig. 1 represent the essential molecular basis of the Bronsted relation. This presentation makes clear several additional points arising from such relations. Figure 1 [also Eqs. (44)and (45)] makes no distinction between substrate and catalyst, and we should theref'ore expect t o find (for a given catalyst) relations between the reaction velocity and acidbase strengths of a series of substrates. There is little direct experimental evidence for this, mainly because substrates are iisually such weak acids or bases that the ordinary methods for measuring acid-base strength are not applicable. However, there is limited evidence for the validity of such relations, and it is probably safe to use them for estimating the acid-base strength of very weak substrates. For example, data on the rates of ionization of a series of ketones and similar substances have been used to estimate that acetophenone has an acid dissociation constant in water of about lO-'9, in agreement with estimates based on independent evidence (Bell, 141; McEwen, 142). On the other hand, this parallelism is only present for substrates of closely similar structure, and when comparing substances of different chemical types there may be wide divergencies. For example, acetoacetic ester, and nitromethane have similar acid dissociation constants (2 x 10-l' and 6 >< l O - l l ) , and their rate of halogenation in presence of a basic catalyst is 'believed to depend on the loss of a proton to the catalyst. Nevertheless, the rate of bromination in presence of acetate ions is about 6000 times as great for the ester as it is for the nitroparaffin (Pedersen, 29,42; Reita, 43). The general question of the effect of structure will be treated in Sec. V. 3. Statistical Efects The velocities of proton transfer in a series of similar acid-base systems may be affected by purely statistical features, and when this is the case some modification is necessary in the simple Bronsted relations expressed by Eq. (46). The point can be best explained by means of

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

199

an example. Suppose that we have as catalyst a carboxylic acid CH, (CHz),COOH whose catalytic effect is given by the equation k~ = GAKA"

(49)

and that we wish to compare with it a dibasic acid COOH (CH2),COOH, where n is so great that the mutual effect of the carboxyl groups is negligible. The tendency of the carboxyl groups t o lose a proton will be essentially the same in the two acids, but the first dissociation constant of the dibasic acid (KA')will be twice as great the monobasic acid ( K A ) , since the ion COOH-(CHz);COO- can be formed by losing a proton from either end of the chain. Similarly, the catalytic effect of the dibasic acid (kA') will be twice as great as that of the monobasic acid, since in the former case the substrate molecule can approach either end of the acid molecule. This is not, however, what is predicted by Eq. (49) as it stands, which gives ka'lka = (KA'/KA)"= P The error can be removed if we reckon both the acid strength and the catalytic power per carboxyl group, giving the observed result ik~'= GA(+KA')== GAKA" = k A

A simila'r problem arises if we compare the two acids and where again n is great. In this case the tendency to lose a proton from the carboxyl group (and hence the catalytic power) will be the same for two acids. On the other hand, the dissociation constant of (I) will be only half that of (11),since the conjugate base COO-.(CHz);COO- has two equivalent points at which a proton can be added, while COO-. (CH2),*COOCH3 has only one such point. Here again the straightforward use of Eq. (49) will not predict correctly the observed catalytic behavior, and for the acid I it is necessary to multiply the observed dissociation constant by two before inserting it in Eq. (49). These arguments can easily be generalized. Thus if we have a conjugate acid-base pair A-B in which A has p dissociable protons bound equally firmly, while B has q equivalent points at which a proton can be attached, then the catalytic power of A is related to its acid strength by the equation

200

R. P. BELL

A similar treatment for basic catalysis by B gives

Analogous equations can be developed (Westheimer, 143) for the case in which the various protons or points of attachment are not all equivalent, though it is then necessary t o know the relative tendencies of losing or gaining a proton at the different points. The idea of a statistical correction was originally put forward in an incomplete form by Bronsted and Pedersen (15) and later stated correctly by Bronsted (144). Although it is undoubtedly correct in principle, there is very little experimental evidence which can be quoted in support of it, almost the only clear-cut example being the catalysis of the nitramide decomposition by the anions of polycarboxylic acids (Bronsted and Pedersen, 15). I n the majority of cases various ambiguities arise in determining the correct values of p and q t o use in Eqs. (50) and (51). It is often doubtful whether several protons attached t o the same atom should be reckoned independently or not, e.g., is it correct t o take p = 4 for the acid NH4+? The hydrogen atoms are so close together t h a t the chance of proton transfer on collision is probably less than four times the chance for an analogous ion containing only one hydrogen atom. I n any case, the concept of a n analogous ion is a somewhat nebulous one, since in a system of this kind it is impossible t o substitute some of the hydrogen atoms by other groups without affecting the nature of the remaining hydrogens. Difficulties also arise in deciding upon the correct values of q, and are often related t o the particular views as t o the electronic structure of the species concerned. For example, if the structure 0 of the carboxylate ion is written as -C

// , we should have q \

=

1, but

0-

the mesomeric formulation with the charge shared equally between the two oxygens would suggest q = 2. Similar problems arise in the ions of oxy acids such as HsPOd and H2S04. The mesomeric structures for these ions are presumabIy the correct ones, but it is again doubtful how far oxygen atoms attached t o the same central atom can be regarded as kinetically independent. It is difficult t o subject any of these points t o experimental test, since (as we shall see in Sec. V) the presence of mesomerism is likely t o cause deviations from the Bronsted relation much greater than those depending upon statistical corrections. The above considerations have been applied t o iche comparison of different catalysts, but are also relevant in the comparison of differ-

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

20 1

ent substrates. For example, in comparing reactions of the ketones CHBCOCH3 and C6HBCOCH3,we must take into account the different numbers of methyl groups. The same difficulties arise in considering several protons bound to the same atom, e.g., in comparing the groups

-CH3,

\ \ CH2, and -CH, / /

though these ambiguities are frequently of

minor importance compared with the large effects which can be produced by chemical substitution.

V. THE IMPORTANCE OF MOLECULAR STRUCTURE 1.

The Structure of the Substrate

The variation of velocity in a series of similar substrates will be related t o their acid-base properties, and there are some instances in which kinetic measurements represent the only experimental approach t o these properties, e.g., when dealing with acids or bases which are so weak that no equilibrium measurements are feasible. For very weak bases equilibrium measurements often become possible if a solvent of sufficiently high acidity is employed, but for very weak acids no analogous strongly basic solvents are available, and rate measurements (e.g., of base-catalyzed deuterium exchange or halogenation) represent the only method of comparing acidities. Data of this kind can thus be used for testing theoretical predictions about the effect of chemical structure on acidity. As an illustration, Table I shows data for the base-catalyzed halogenation of a number of ketones and similar substances. I n this table, R is the catalytic constant (in liters/mole/minute) of the anion of a hypothetical acid of dissociation constant obtained by interpolating data for carboxylate anions. In computing R a statistical correction has been made for the number of equivalent hydrogen atoms in the substrate, counting as independent atoms attached t o the same carbon. 0 is the exponent in Eq. (51). The last column of Table I contains values for the pK of the substrate in water, estimated from the kinetic data. The range of velocities covered is very wide (about lo9), and the value of p shows a steady decrease as the rate increases. In integrating the relation d log R = pd log K , it is therefore not allowable to take ,f3 as a constant, and the usual Bronsted equation [cf. Eq. (51)] must be replaced by log K , =

1f

d log R

+A

where A is an inteqation constant and p is now a function of R.

(52)

The

202

R. P. BELL

values of pK, in Table I are derived from this equation, the value of A being obtained by using the directly measured value pK, = 10.5 for acetoacetic ester. Values of pK, have also been measured directly for acetylacetone (pK, = 8.9) and for benzoylacetone (pK, = 9.4), in fair agreement with the table. It is difficult to check the values for the very weak acids a t the top of the table, but fortunately McEwen (142) has TABLE I Substrate

logio R

P

PK*

CHsCOCHs CHsCOCeHs CHsCOCHzCHzCOCHo CHsCOCHzCl CH3COCH2Br CHsCOCHClz CHZCOCHCOOCzHr ‘(CH213CH3COCHzCOOCzH6 CHzCOCHCOOCiHs 4CH2&CHsCOCH2COCeHa CH3COCHBrCOCaHh CH&OCHBrCOOCzHs CH3COCHzCOCHs CH~COCHB~COCHI

-6.78 -6.07 -5.46 -3.51 -3.25 +2.00 -0.98

0.88 0.89 0.82 0.82 0.82 0.64

20.0 19.2 18.7 16.5 16.1 14.9 13.1

+0.72 +1.18

0.59 0.58

10.5 10.0

+1.33 +1.33 +1.35 +1.54 +2.04

0.52 0.52 0.48 0.42

9.7 9.7 9.7 9.3 8.3

I

I

~~~~~

~

Source

a

d

~~

Bell and Lidwell, 30. * Bell. 141. c Bell and Goldsmith, 145.

Bell, Smith, and Woodward. 146. Bell, Gelles, and Molbsr, 147.

4

estimated the acid strength of acetophenone by a series of displacement reactions and finds pK = 19, in excellent agreement with our value of 19.2. Neither value would be expected t o be more certain than a power of ten, but the agreement strengthens our confidence in the general method employed in estimating K,. The values of K , thus derived can in many cases be interpreted in terms of molecular structure. The increase in acidity (- 4 pK units) produced by introducing a halogen atom into acetone can be classified as an inductive effect, depending upon the stabilization of the anion by interaction between its charge and the carbon-halogen dipole, e.g., a+

8-

CH8.C: CHCl

A-

This effect is of course paralleled in the effect of halogen substitution on the strength of aliphatic carboxylic acids, though here it is of smaller

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

203

magnitude (- 2 pK units for a-substitution) because of the greater distance between the dipole and the negative charge. The still larger increase in K , which occurs when we pass from simple ketones to P-diketones or keto esters demands a different type of explanation, and this is found in the occurrence of mesomerism in the ion, resulting in the distribution of negative charge between two oxygen atoms. For example, the anion of acetylacetone can be given the equivalent structures CH3.COCH:CO-.CH3 and CHa.CO-: CH*CO.CHs,and it is well known that in this situation the species concerned is more stable than would be expected for a single structure. The substitution of a halogen atom in the 0-diketones and keto esters produces a small increase in velocity, but much less than in the simple ketones. It is not clear why the effect of halogen substitution should be so much reduced, since the inductive effect would be expected to operate in roughly the same way in all these compounds. Another feature of the table which is difficult to explain is the large difference (a factor of over 100) between the rates for 2-carbet hoxy cy clopentanone and 2-carbet hoxycyclohexanone , though there are other similar anomalies in the reactivities of such ring systems, Interest also attaches to the variation of the exponent p among this series of compounds. The range of velocities covered is very much greater than is normally investigated when studying a series of catalysts for a given substrate, and constancy of p is therefore less likely. Since p is related to the slopes of two intersecting energy curves (cf. Fig. l ) , its change could be interpreted as due to a change of slope when traversing a considerable part of the energy curve. This view was put forward by Lidwell and Bell (148), but it was found difficult t o account for the relatively large changes observed. It seems more likely that the variations of B are associated with changes in the shape of the potential energy curves, and not merely in their position. Such changes of shape are likely when considerable changes of structure are involved (especially those involving mesomerism), and will be considered further in the next section in connection with the effect of structure of catalysts. An extreme instance of changes in the shape of energy curves is met with when the chemical nature of the substrate is drastically altered. As already mentioned, although acetoacetic ester and nitromethane have similar acid dissociation constants, their rates of ionization are in the ratio 6000 :1 (Pedersen, 29,42; Reitz, 43). Although the proton is lost from a carbon in both compounds, the structures of the resulting ions are very different, and the energy curves would be expected to differ considerably in shape. Another point of interest arises when considering the series CH3N02,C2HsN02, CH3CH2CH2N02,(CH3)2CHN02. The acid strengths of these increase in the order given (Turnbull and Maron,

204

R. P. BELL

149; Wheland and Farr, 150), but their rates of reaction with hydroxyl ion are in the reverse order, so that the dependence of velocity upon acid strength is the exact opposite of t h a t usually found. Moreover, the effect of alkyl substitution in increasing the acidity is 1,he opposite of the usual inductive effect (e.g., in the carboxylic acids). It seems likely t h a t the sequence of acid strengths results from the stabilization of the undissociated molecule by a hyperconjugative effect which depends on the number of hydrogen atoms on the carbon adjacent to the nitro group, e.g., in nitromethane we can write three structures of the type H+

H-C=N

H

/

0-

+ / ~

\

0-

This type of structure is not possible for the anion (or for carboxylic acids). The sequence of velocities is then explained on the basis of the ordinary inductive effect (Cardwell, 151).

2. T h e Structure of the Catalyst This problem is in principle the same as that of the structure of the substrate, but there are some differences in practice. I n the first place, the acid-base strength of the catalyst is usually me:asurable by direct means, and a large part of the variations in reaction velocity merely parallel these variations in acid-base strength. I n the second place, it is rarely practicable t o vary the strength of the catalyst over a very wide range (because of the catalytic effect of the molecules or ions of the solvent). Moreover, most measurements on catalyzed reactions have been confined t o a series of very similar catalysts, and little evidence is available as t o how far catalysts of widely differing structure conform t o a single Bronsted relation. From what has already been said, it might be expected that a pseudoacid would have a smaller catalytic effect than a “normal” acid of the same dissociation constant. This may be made cleairer by reference t o Fig. 2. Curve (a) is the potential energy curve for the removal of a proton from a pseudo-acid, and (b) is the corresponding curve for a “normal” acid of the same strength. For the sake of definiteness we can suppose t h a t (a) refers t o a nitroparaffin, and (b) t o a phenol. For small displacements of the proton the nitroparaffin curve will be essentially that for a C-H ionization, and hence much steeper than (b) for the 0-H ionization of the phenol, but for greater separations the anion of the nitroparaffin is stabilized by the shift of charge t o the oxygen atoms, and the final energy is t h e s a m e for the two curves. If the proton is

ACID-BASE CATALYSIS A N D MOLECULAR STRUCTURE

205

transferred t o a base (e.g., the substrate in a catalyzed reaction) the activation energies in the two cases will correspond t o the intersections of (a) and (b) with a curve such as (c), and it is clear that the pseudoacid will react more slowly than the ordinary acid of the same strength. The magnitude of this difference will depend upon the position of the proton in the transition state (corresponding to the intersection of the curves in Fig. 2), and may vary widely from one reaction to another. Similar considerations apply to proton transfers in the opposite direction,

1 POSITION OF PROTON

FIG.2.

.

and to the converse case where the undissociated state is stabilized by an electronic rearrangement. The same kind of explanation will obviously account for the variation of the exponent in the series of ketones referred to in the last section, and also the large disparity in rates between substrates of similar acid strength but widely differing chemical nature. I n its application t o catalysts, the preceding paragraph suggests that pseudo-acids or bases will be ineffective catalysts compared with other acids or bases of the same strength; i.e., they should exhibit negative deviations from the Bronsted relation. There are a number of isolated observations which confirm this idea. Thus Bronsted and Pedersen (15) found that the nitrourethane ion had an unexpectedly small catalytic

206

R. P. BELL

effect in the decomposition of nitramide, which they attributed to the pseudo-acid nature of nitrourethane. Similarly, the abnormally low catalytic effect of the picric acid molecule in a number of reactions (Bronsted and Bell, 151 ; Bell, 152) may be related t o the spread of negative charge to the nitro groups in the picrate anion. More systematic evidence is provided by the results of Bell and Higginson (62) for the reaction CHaCH(OH)2+ CHaCHO H 20, the mechanism of which has already been discussed (Sec. 111.2.12). A detailed study of acid catalysis was made, and it was found that 45 carboxylic acids and phenols, with acid strengths ranging over 10 powers of ten, obeyed a Bronsted relation with a maximum deviation of 0.3 logarithmic unit, and a mean deviation of 0.1 logarithmic unit. On the other hand, catalysts of other chemical types exhibit large deviations from the relation which is valid for carboxylic acids and phenols. Some of the deviations are shown in Table 11. The left-hand side of the table shows that nitro-

+

TABLE I1 Dehydration o j .4cetaldehyde Hpdtate-Deviations from the Briinsled Relation Negative Deviations

Positive Deviations

Logarithmic deviation

Catalyst Benzoylacatone enol 1,3 diketo-5-dimethylcyclohexane cnol (dimedone) Nitromethane 1-Nitropropane Nitroethane 2-nitro propane

-1.4 -1.1 -1 . 4 -1.5 -1.7

Logarithmic deviation

Catalyst Benzophenone oxime Acetophenone oxime Diethyl ketoxime Chloral hydrate Water

+1.2 +1.4 +2.1

+0.7 +1.6

-1.9

paraffins and the enols of @-diketonesare 10 to 100 times less efficient as catalysts than carboxylic acids or phenols of the same dissociation constant. The nitroparaffins have already been discussed as examples of pseudo-acids, while the ionillation of 8-diketone enols involves the distribution of a negative charge between two separated oxygen atoms, i.e., -C-CH=C-

I\

-C-CH=C AH4

8

-C=CH--C

b-

b-

!I

+ 11

+

The right-hand side of Table I1 shows that certrtin catalysts exhibit. large positive deviations from the Bronsted relation. This can be interpreted if me remember that both the carboxylic acids and the phenols undergo some structural rearrangement on ionization. The carboxylate

ACID-BASE CATALYSIS AND MOLECULAR STRUCTURE

207

anion contains two equivalent carbon-oxygen bonds in place of one double and one single bond, while the high acid strength of phenols relative t o alcohols is commonly attributed to the occurrence of structures such as

in the anion, whereby the negative charge is partly distributed over the ring. On the other hand, in the five acids on the right-hand side of Table I1 there is no formal possibility of structural change on ionization, and they might therefore be described as having “less pseudo-character than the carboxylic acids and the phenols. This falls into line with their enhanced catalytic effect. In the paper by Bell and Higginson (62) a number of other suggestions are made for correlating individual deviations from the Bronsted relation with the structures of the catalysts concerned. Although several of these suggestions are speculative, it seems possible that further kinetic data will at least provide a new approach for investigating structural problems concerning acid-base pairs. It may be suggested that there is no sharp distinction between pseudo-acids and other acids, but rather a whole range of types of acid, varying in the magnitude of the electronic shift which accompanies ionization. The kinetic approach offers a possibility of obtaining a quantitative measure of the extent t o which different acids behave as pseudo-acids, and hence of testing views as to their electronic structure.

REFERENCES Wilhelmy, L. F., Pogg. Ann. 81, 413, 499 (1850). Oatwald, W., J . prakt. Chem. SO, 93 (1884). Bell, R. P., Acid-Base Catalysis. Clarendon Press, Oxford, 1941. Daviee, C. W., J . Chem. Soc. 349 (1939). Owen, B. B., and Gurry, R. W., J . Am. Chem. Soc. 60,3074 (1938). Bell, R. P., and Prue, J. E., J . Chem. Soc. 362 (1949). 7. Bell, R. P., and Waind, G. M., J . Chem. SOC.1979 (1960). 8. Bransted, J. N., and Teeter, C. E., J . Phys. Chem. 28, 579 (1924). 8a. Guggenheim, E. A., Phil. Mag. (7) 19, 588 (1935). 9. Bransted, J. N., 2. phyaik. Chem. 102, 169 (1922). 10. Hammett, L. P., and Deyrup, A. J., J . Am. Chem. Soc. 64, 2721 (1932). 11. Bell, R. P., Quart. Rev. Chem. Soc. 1, 113 (1947). 12. Lowry, T. M., and Smith, G . F., J . Chem. Soc. 2539 (1937). 13. BrGmted, J. N., and Guggenheim, E. A., J . Am. Chem. Soc. 49, 2554 (1927). 14. Dawson, H. M., and Powis, F., J . Chem. SOC.2135 (1913) and later papers. 15. Bransted, J. N., and Pedersen, K. J., 2.physik. Chem. 108, 185 (1923). 16. Bell, R. P., and Baughan, E. C., J . Chem. Soc. 1947 (lQS7). 1. 2. 3. 4. 5. 6.

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17. Bell, R. P., and Burnett, R. le G., Trans. Faraday SOC.33, 355 (1937). 18. Bell, R. P., and Burnett, R. le G., Trans. Faraday SOC.36, 474 (1939). 19. Wassermann, A., J . Chem. SOC.618 (1942). 20. Wassermann, A., J . Chem. SOC.3046 (1949). 21. Bell, R.P., Ann. Rept. Chem. SOC.36, 82 (1940). 22. Haskell, V. C., and Hammett, L. P., J . Am. Chem. SOC.71, 1284 (1949). 22a. Levesque, C. L., and Craig, A. M., Znd. Eng. Chem. 40, 96 (1948). 22b. Walling, C., J . Am. Chem. SOC.72, 1164 (1950). 23. Bell, R. P., and Brown, J. F., J . Chem. SOC.1520 (1936). 24. Lapworth, A., J . Chem. SOC.30 (1904). 25. Kohler, E. P., and Thompson, D., J . Am. Chem. SOC.66,3,822 (1933). 26. Arndt, F., Ber. 74, 438 (1941). 27. Schwareenbach, G., and Wittwer, C., Helv. Chim. Acta 30, 669 (1947). 28. Pedersen, K. J., J . Phys. Chem. 37, 751 (1933). 29. Pedersen, K. J., J . Phys. Chem. 38,601 (1934). 30. Bell, R. P., and Lidwell, 0. M., Proc. Roy. SOC.(London) A.176, 88 (1940). 31. Bell, R. P., and Tantram, A. D. S., J . Chem. SOC.370 (1948). 32. Hammett, L. P., Chem. Revs. 13, 61 (1933). 33. Flexser, L. A., Hammett, L. P., and Dingwall, A., J . Am. Chem. SOC.67, 2103 (1935). 33a. Conant, J. B., and Wheland, G. W., J . Am. Chem. SOC.64, 1212 (1932). 33b. Hsu, S. K., Ingold, C. K., and Wilson, C. L., J . Chem. Soc. 1778 (1936). 34. Hsu, 8. K., and Wilson, C. L., J . Chem. SOC.623 (1936). 35. Ingold, C. K., and Wilson, C. L., J . Chem. SOC.773 (1934). 36. Bartlett, P. D., and Stauffer, C. H., J . Am. Chem. SOC.67,2580 (1935). 37. Reitz, O., Z. physik. Chem. 179, 119 (1937). 38. Hammett, L. P., Physical Organic Chemistry, p. 67. New York, 1940. 39. Leuchs, H., and Wutke, J., Ber. 46, 2425 (1913). 39a. Nelson, W. E., and Butler, J. A. V., J . Chem. SOC.957 (1938). 40. Walters, W.D., and Bonhoeffer, K. F., Z.physik. Chem. b182, 265 (1938). 41. Junell, R., Z. physik. Chem. A141, 71 (1929). 42. Pedersen, K. J., Kgl. Danske Videnskab. Selskab, Mat.-fys. Medd. 12, 1 (1932). 43. Reits, O., Z. physik. Chem. A179, 119 (1937). 44. Turnbull, D., and Maron, S. H., J . Am. Chem. Sac. 66, 212 (1943). 45. Kuhn, R., and Albrecht, O., Ber. 60, 1297 (1927). 46. Shriner, R. L., and Young, J. H., J . Am. Chem. SOC.62,3:332 (1930). 47. Kornblum, N., Lichtin, N. N., Patton, J. T., and Iffland, D. C., J . Am. Chem. SOC.69,307 (1947). 48. Kornblum, N., Patton, J. T., and Nordmann, J. B., J . Am. Chem. SOC.70, 746 (1948). 49. Lowry, T. M. et al., J . Chem. SOC.1898-1915. 50. Bell, R. P., and Sherred, J. A., J . Chem. Soc. 1202 (1940). 51. Hantssch, A., and Veit, A., Ber. 32, 615 (1899). 52. Thiele, J., and Lachmann, A., Ann. 288, 267 (1895). 53. Hantesch, A., Ann. 292, 340 (1896). 54. Hantssch, A., Ber. 63, 1270 (1930). 55. Cambi, L., and Szego, L.,Ber. 61, 2081 (1928). 56. Kortiim, G., and Finck, B., Z.physik. Chem. B48, 32 (1940). 57. Bell, R. P., and Trotman-Dickenson, A. F., J . Chem. SOC.1288 (1949). 58. Hunter, E. C. E., and Partington, J. R., J . Chem. SOC.30!J (1933). 59. Pedersen, K. J., J . Phys. Chem. 38, 581 (1934).

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CATALYSIS AND MOLECULAR STRUCTURE

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Booth, V. H., and Roughton, F. J. W., J . Physiol. 92, 36 (1938). Bell, R. P., and Darwent, B. de B., Trans. Faraday SOC.46,34 (1950). Bell, R. P., and Higginson, W. C. E., Proc. Roy. SOC.(London),A197, 141 (1949). Cohn, M., and Urey, H. C., J . Am. Chem. SOC.60, 679 (1938). Dieckmann, W., Ber. 49, 2213 (1916). Dieckmann, W., Ber. 60, 1375 (1917). Bell, R. P., and Baughan, E. C., J . Chem. SOC.1947 (1937). Bell, R. P., and Hirst, J. P. H., J . Chem. SOC.1777 (1939). Bergmann, M., Ann. 462, 123 (1927). Bergmann, M., and Miekeley, A.,Ber. 62,2297 (1929). Schaefgen, J. R., Verhoek, F. H., and Newman, M. S., J . Am. Chem. SOC.67,253 (1945). 71. McKenzie, A., and Mitchell, A. G., Biochem. 2. 208, 456 (1929). 72. McKeneie, A., and Ritchie, P. D., Biochem. 2. 231, 412 (1931); 237, 1 (1931); 260, 376 (1932). 73. Dawson, H. M., and Lowson, W., J . Chem. SOC.63,3217 (1931). 74. Rolfe, A. C., and Hinshelwood, C. N., Trans. Faraday SOC.30, 935 (1934). 75. Hinshelwood, C. N., and Legard, A. R., J . Chem. SOC.587 (1936). 76. Day, J. N. E., and Ingold, C. K., Trans. Faraday SOC.37, 686 (1941). 77. Pedersen, K. J., J . Phys. Chem. 38, 581 (1934). 78. Rakowski, A. V., 2. physik. Chem. 67, 321 (1907). 79. Bartlett, P. D., J . Am. Chem. SOC.66, 967 (1934). 80. Zucker, L., and Hammett, L. P., J . Am. Chem. SOC.61,2791 (1939). 81. von Euler, H., and Olander, A., 2. physik. Chem. 131, 107 (1927). 82. Taylor, T. W. J., J . Chem. SOC.62,2741 (1930). 83. King, C. V., J . Am. Chern. SOC.62, 379 (1940). 84. King, C. V., and Bolinger, E. D., J . A m . Chem. SOC.68, 1533 (1936). 85. Gross, P., Steiner, H., and Suess, H., Trans. Faraday SOC.32, 883 (1936). 86. Orr, W. J. C., and Butler, J. A. V., J . Chem. Soc. 330 (1937). 87. Nelson, W.E., and Butler, J. A. V., J . Chem. Soc. 957 (1938). 88. Brescia, F.,and La Mer, V. K., J . Am. Chem. SOC.60, 1962 (1938). 89. Brescia, F., J. Chem. Phys. 7, 307 (1939). 90. Wynne-Jones, W. F. K., Trans. Faraday SOC.34, 245 (1938). 91. Gross, P., Trans. Faraday SOC.34, 261 (1938). 92. Reite, O., 2. physik. Chem. A179, 119 (1937). 93. Hamill, W. H., and La Mer, V. K., J . Chem. I’hys. 4, 395 (1936). 94. Wilson, C. L., J . Chem. SOC.1550 (1936). 94a. Reite, O.,Z. physik. Chem. A179, 119 (1936). 95. Bell, R. P., and Caldin, E. F., J . Chem. SOC.382 (1938). 96. Bell, R. P., Lidwell, 0. M., and Wright, J., J . Chem. SOC.1861 (1938). 97. Flexser, L. A,, Hammett, 1,. P., and Dingwall, A., J . Am. Chem. SOC.67, 2103 (1935). 98. Flexser, L. A., and Hammett, L. P., J . Am. Chem. SOC.60, 885 (1938). 99. Zucker, L., and Hammett, L. P., J . Am. Chem. SOC.61,2791 (1939). 100. Pedersen, K. J., J . Phys. Chem. 38, 601 (1934). 101. Kimball, R. H., J . Am. Chem. SOC.68, 1963 (1936). 102. Bell, R. P., J . Chem. SOC.1637 (1937). 103. Westheimer, F. H., and Cohen, H., J . Am. Chem. SOC.60, 90 (1938). 104. Westheimer, F. H., Ann. N . Y . Acad. Sci. 39, 401 (1940). 105. Westheimer, F. H., and Jones, W. A., J . Am. Chem. SOC.63,3283 (1941). 106. Pacsu, J., J. Am. Chem. SOC.66, 5066 (1933); 66, 745 (1934). 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70.

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107. MoeIwyn-Hughes, E. A., Klar, R., and Bonhoeffer, K. F., 2. physik. Chem. A169, 113 (1934). 108. Moelwyn-Hughes, E. A., 2.physik. Chem. B26, 272 (1934:~ 109. Hamill, W. H., and La Mer, V. K., J . Chem. Phys. 4, 395 (1936). 110. Wynne-Jones, W. F. K., Trans. Faraday SOC.34,245 (1936). 111. Reitz, O., 2.Elektrochem. 44, 693 (1938). 112. Reitz, O., 2.physik. Chem. A183, 371 (1939). 113. Lowry, T. M., and Faulkner, I. J., J . Chem. SOC.2883 (1925). 114. Lowry, T. M., J . Chem. SOC.2554 (1927). 115. Dawson, H. M., and Spivey, E., J . Chem. SOC.2180 (1930). 116. Pedersen, K. J., J . Phys. Chem. 38, 581 (1934). 117. Swain, C. G., J . Am. Chem. SOC.72, 4578 (1950). 118. Ingold, C. K., and Wilson, C. L., J . Chem. SOC.93 (1934). 119. Hsti, S. K., Ingold, C. K., and Wilson, C. L., J . Chem. SOC.1778 (1935). 120. Richards, E. M., andLowry, P. M., J . Chem. SOC.1385 (11925). 120a. Bell, R. P., and Jones, P., unpublished work. 121. Bell, R. P., and Ryhicka, S. M., J . Chem. SOC.24 (1947). 121a. Bell, R. P., and Clunie, J. C., Nature 167, 362 (1951). 122. Hinshelwood, C. N., J . Chem. SOC.694 (1947). 123. Hantzsch, A., Ber. 32, 575 (1899). 124. Lewis, G. N., and Seahorg, G. T., J . Am. Chem. SOC.61, 1894 (1939). 125. Kilpatrick, M., J . Am. Chem. SOC.62, 1094 (1940). 126. Hantesch, A., 2. Elektrochem. 29, 244 (1923). 127. Hantzsch, A., 2. Elektrochem. SO, 202 (1924). 128. Hantesch, A., Ber. 68, 923 (1925). 129. Fajans, K., Naturwissenschuften 11, 179 (1923). 130. von Halban, H., 2. Elektrochem. 29, 443 (1923). 131. Ley, H., and Hlinecke, H., Ber. 69,510 (1926). 132. Bonhoeffer, K. F., Trans. Faraday SOC.34,252 (1938). 133. Ives, D. J. G., and Wilks, G. C., J . Chem. SOC.1456 (1938). 134. Hantzsch, A., and Kalb, M., Ber. 32, 3116 (1899). 135. Aston, J. G., J . Am. Chem. SOC.62, 5254 (1930). 136. Aston, J. G., J . Am. Chem. SOC.63, 1448 (1931). 137. Reyerson, L. H., J . Am. Chem. SOC.67, 779 (1935). 138. Ogston, A. G., J . Chem. SOC.1023 (1936). 139. De La Mare, P. B. D., Hughes, E. D., and Vernon, J. L., Research 3, 192, 242 (1950). 140. Bell, R. P., and Gelles, E., J . Chem. SOC.2734 (1961). 141. Bell, R. P., Trans. Faraday SOC.39, 253 (1943). 142. McEwen, W. K., J . Am. Chem. SOC.68, 1124 (1936). 143. Westheimer, F. H., J. Org. Chem. 2, 431 (1938). 144. Bronsted, J. N., Chem. Revs. 6, 322 (1928). 145. Bell, R. P., and Goldsmith, H. L., Proc. Roy. SOC.(London) in the press. 146. Bell, R. P., Smith, R. D., and Woodward, L. A., Proc. Roy. SOC.(London) A192, 479 (1948). 147. Bell, R. P., Gelles, E., and Moller, E., Proc. Roy. SOC.(London) A198, 310 (1949). 148. Lidwell, 0. M., and Bell, R. P., Proc. Roy. Soc. (London) A176, 114 (1940). 149. Turnhull, D., and Maron, S. H., J. Am. Chem. SOC.66, 212 (1943). 150. Wheland, G. W., and Farr, J., J . Am. Chem. SOC.66, 1433 (1943). 151. Bronsted, J. N., and Bell, R. P., J . Am. Chem. Soe. 63, 2,478 (1931). 152. Bell, R. P., Proc. Roy. SOC.(London) A143, 377 (1934).

Theory of Physical Adsorption TERRELL L. HILL Naval Medical Research Institute, Bethesda, Maryland

CONTENTS ..................................................

Page 212

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212

a. Uniform Surface with Nearest Neighbor Interactions. . . . . . . . . . . . . . . 222 b. Uniform Surface with Higher Neighbor Interactions. . . . . . . . . . . . . . . . 223 . . . . . . . . . . . 224 1. Introduction.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

. . . . . . . . . . 225

. . . . . . . . . . . . . . 230

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

242

2. Solution Thermodynamics. ..

. . . . . . . . . 246

References. . . . . . . . . . . 211

212

TERRELL L. HILL

I. INTRODUCTION This article consists of a review of those recent aspects of the statistical thermodynamical theory of physical adsorption which the writer has been particularly interested in. The discussion therefore is in no sense a really comprehensive review of the entire field, and this choice of subject matter accounts also for the disproportionate number of references to the writer’s own papers. A relatively nonmathematical discussion (omitting many derivations) is given, primarily for the benefit of experimentalists. In Secs. 11 and I11 we shall summarize the present state of the theory of physical adsorption, starting with the simplest case of essentially independent molecules on a surface (very low equilibrium gas pressure) and then considering monolayer adsorption with interactions, and finally multilayer adsorption. In Sec. IV, the thermodynamics of adsorption will be discussed. There has been a certain amount of confusion in recent years concerning the application of thermodynamics to adsorpt,ion data, and it is hoped that the present article will help to clarify this situation. While chemisorption will not be discussed explicitly, a good deal of what is said about monolayer adsorption and thermodynamics will of course have implications for this field.

XI. MONOLAYER ADSORPTION I . Adsorption of Independent Molecubes If we consider a gas in equilibrium with adsorbed molecules on a solid surface, at sufficiently low gas pressures the adsorbed :molecules will form a dilute two-dimensional system with negligible interactions. This is analogous to the perfect gas state in three dimensions. Under these conditions we are interested in the behavior of essentially independent adsorbed molecules, the only forces of interest being those between adsorbed molecules and the adsorbent. The interaction of an adsorbed molecule with the adsorbent is (in the usual approximation of additive forces) the sum of the separate interactions of the adsorbed molecule with the atoms or molecules of the adsorbent in the immediate neighborhood of the adsorbed molecule. a. Energy of Adsorption. The case with perhaps the least complications (no relatively long-range electrostatic forces, no polarization, etc.) is the adsorption of a monatomic gas on a nonpolar solid, so that only ordinary van der Waals forces need be considered. For simplicity, we shall discuss this case only, it being understood th.at the appropriate

THEORY OF PHYSICAL ADSORPTION

213

law of force must be substituted in more complicated cases (see, for example, Brunauer, 1). As a hypothetical example suppose the monatomic gas A is adsorbed on the 100 plane of the simple cubic lattice of solid B. If only van der Waals forces are operative, the potential energy of interaction between an atom of A and an atom of B as a function of the distance T between the atoms is (Lennard-Jones, 2) * I2 U(T)

=

6

(5)

- 2€

(;)?

to a good approximation. The form of u ( r ) is shown in Fig. 1. The associated force between A and B is of course given by f = -du/dr. A and B attract each other for T > T * and repel for r < r * .

FIG.1. Schematic form of intermolecular potential energy according to LennardJones.

The adsorbing force holding A to the surface of B, as mentioned above, is the sum of a number of such interactions. In Fig. 2 we show an adsorbed atom A at a distance z from the surface of B, and a few of the distances ri which should be substituted into Eq. (1) to give the various contributions ui to the total energy of interaction of A with the solid, U = c u i , for this particular location of A. i

Incidentally, it is generally considered an excellent approximation in physical adsorption on solids to assume that adsorbed molecules do not perturb or alter the properties of the solid adsorbent appreciably.

214

TERRELL L. HILL

This is because of the relatively strong forces, in most cases, holding the solid together. However, Cook, Pack and Oblad (2a) question the validity of this assumption, and their arguments would affect the treatment of molecules adsorbed in the first layer. A sufficiently good approximation t o U for som’e purposes (1) is obtained by assuming the atoms of B to be smeared out in a continuum, replacing summation by integration. Thus, suppose the surface of the solid is the XY-plane in a Cartesian coordinate system, with z increasing away from the surface (Fig. 2). Let p be the number of adsorbent atoms

0

0

0

0

0

0

FIG.2. Interactions of an adsorhed atom A with atoms of the adsorbent B.

per cubic centimeter. The energy of interaction of A at x, y, z with the atoms of B in dx’ dy’ dz’ a t x’,y‘, z’ is u(r)dx‘ dy‘ dz’

with T2

= (z

- x’)2

p

+ (y - y’)2 +

(2

-

2‘)2

A straightforward integration (London, 3; Hill, 4:) over - ~i < x’ < *, --co < y ’ < 4 - 0 0 , - QI < z‘ < 0 then gives for the total energy U ,

+

u = -cr-*1%p -

45z9

cr *%p

39

(2)

The function U ( z ) has the same general form as U ( T ) in Fig. 1. The depth of the potential well, U O= 0.497er*%p, is the energy of adsorption (at O”K., neglecting zero-point energy, which is of’ the order of 100 cal./mole). U o is of the order of magnitude of several kilocalories per mole, as expected from experimental heats of adsorption. The location of the minimum is a t z = 0.765r*, which is the equil.ibrium position for motion normal to the surface.

T H E O R Y O F PHYSICAL ADSORPTION

215

The free translational motion of a gas atom in the x direction is replaced on adsorption by vibration in the potential well of Eq. 2 (see also Fig. 1). The frequency of vibration is given approximately by the curvature of U ( z ) near the minimum. One can show (4) that v, = 2.383 ( e r * p / m ~ ) ~(where ' m is the mass of the adsorbed atom), which is of the 4 x 1013 set.-' for the internal order of 10l2 sec.-l, compared t o v vibration in most diatomic molecules. The above approximation of replacing the crystalline structure of the adsorbent by a continuum gives useful information about the z motion of an adsorbed atom, but may be inadequate for a discussion of the xy 0

0

0

0

0

0

0

0

& y

L X

o

0

Energy of

0

Potential Barrier VO

FIG.3. Periodic variation of energy of adsorption on a simple cubic crystal.

motion. If we take into account the actual crystalline structure (noncontinuous) of the adsorbent and use summation instead of integration to obtain U (Orr, 5) along a line normal to the surface, it is obvious that the U ( z ) obtained will be different for different locations (x,y) of the line. Thus, U ( z ) directly above a B atom will be different from V(z) above the center of a square of B atoms. In particular, the depth U oof the potential well in U ( z ) will vary periodically in both x and 1~ directions. Figure 3 illustrates this variation in a hypothetical case, where U o is plotted against x along one particular line in the surface (y constant). It is clear from Fig. 3 t h a t motion parallel t o the surface involves passing over periodic potential barriers ( V Oin Fig. 3). These barriers are of the order of 0.3 t o 1 kcal./mole (5). At low temperatures, where the thermal energy kT of the adsorbed atoms is small compared with the height of the barrier, these atoms will be trapped (localized) in the neighborhood of

216

TERRELL L. HILL

potential minima [in Uo(z,y)], except for occasional passages over the barrier or evaporation and recondensation (necessary for thermodynamic equilibrium). When kT is large compared with the height of the barrier, the periodic variation in U o becomes unimportant and the adsorbent, in effect, approaches the continuum already mentioned above. b. Degrees of Freedom for Adsorbed Atoms. We now summarize qualitatively the types of motion possible for adsorbed atoms, as compared to the gas phase. It is, however, possible to pul, this discussion in mathematical language, and in fact this is necessary if one wants to make quantitative calculations of thermodynamic properties such as the entropy of adsorbed atoms (Kemball, 6). In the gas phase each atom has three translational degrees of freedom. At low temperatures the “trapped” (localized) adsorbed atoms will have three vibrational degrees of freedom in place of the translational degrees. We have already discussed v, above. The frequencies v, and vY are determined by the shape of Uo(z,y)near its minima. These have been estimated as being of the order of v z = vV = 3 X 10” sec.-I (4). At higher temperatures, where the barrier in Uo is not important, the adsorption is said to be “mobile,” since the adsorbed atoms lhave translational freedom in the x and y directions, though the z motion is of course still vibrational. There is an interesting transition region in which the adsorption is neither mobile nor localized. Hill (7) has pointed out that the problem of hindered rotation in ethane (and other related problems), as discussed especially by Pitzer and Gwinn (8), is essentially equivalent to the localized-mobile transition problem for a dilute monolayer, and has put some of the results of Pitzer and Gwinn in adsorption language. As an illustration, Hill calculated the heat capacity of dilute adsorbed argon on potassium chloride through the transition range in temperature (25-100°K.). A maximum in the heat capacity curve is predicted by the theory. Recent measurements by Morrison, Los, and Drain (9), of a new order of refinement for this kind of work, are the first available for comparison with the theory. The system studied was argon adsorbed on rutile. At the lowest surface coverage used (about 0.17 of a monolayer) the heat capacity rises steadily with increasing temperature to a value of 11 cal./mole deg. at 130”K., and shows no definite maximum in the temperature range studied (14” to 130°K.). The transition from localized to mobile adsorption is presumably responsible to some extent for ithis large value of the heat capacity. However, as Drain and Morrison have shown (9), the main reason for the high heat capacity is probably the presence of a “configurational” term arising from the heterogeneous nature of the solid surface.

THEORY OF PHYSICAL ADSORPTION

217

c. Degrees of Freedom for Adsorbed Molecules. The motion of isolated molecules near a surface is naturally more complicated than in the atomic case. One has to consider the interaction energy of the adsorbed molecule with the atoms of the adsorbent, taking into account the additional coordinates necessary to specify the configuration of the molecule in space. Thus, for a diatomic molecule (4), the three translational degrees of freedom (of the center of mass) can be discussed as above for the monatomic case, but the two rotational and one (internal) vibrational degrees must also be considered. Actually, the internal vibrational degree of freedom should be relatively unaffected on adsorption since the forces introduced by the surface are rather weak compared to the restoring force in the internal vibration. However, the rotational motion will be seriously altered by the surface, the surface essentially introducing a potential barrier restricting the turning over of the molecule. In the extreme case of a high barrier or low temperature the two rotational degrees will become one degree of planar rotation (rotation of the molecule about an axis perpendicular to the surface) and one degree of vibration (rocking of the molecular axis out of a plane parallel to the surface). More complicated cases have been discussed by Kemball (10) and by Drenan and Hill (11). d. Heterogeneous Surface. We have discussed above the ideal case of a perfect crystalline surface. Actual heterogeneous surfaces will introduce further complications (9). At very low pressures, when N molecules are adsorbed (without dissociation) on a homogeneous surface, Henry’s law requires that N = ap, where a is a constant determined by the theoretical considerations discussed above. On a heterogeneous surface, suppose Nd molecules are adsorbed on a patch of surface with constant ai and occupying the fraction ji of the total surface area. Then N I = flalp N Z = fzazp

N = N I + N 2 . . . =ap a = flu1

+ fra2 + . - .

Henry’s law is of course still obeyed. Regardless of the nature of the adsorption on the surface (it is convenient though not necessary to consider the dissociation case separately), N = ap 6 = kT

I”

( N / a ) d In p

(3)

(Tconst.)

(4)

218

TERRELL L. HILL

and 4@.= N k T

(5)

where is the surface area and 4 the surface pressure, defined, if one likes, by Eq. (4) [see also Eq. (71)]. The two-dimensional equation of state, Eq. ( 5 ) , thus applies as p -+ 0, regardless of whether the adsorption is mobile or localized or whether the surface is heterogeneous or homogeneous. 2. Mobile Adsorption with Interactions We have so far examined the limiting .case p 0. At somewhat higher pressures, the surface concentration, r = N / ' @ becomes large enough so that interactions between adsorbed molecules themselves become appreciable, in addition to the interactions between adsorbed molecules and adsorbent discussed above. It is customary to assume, in the present state of the theory, except in special cases, that the (using the notation above but where A and B may be molecules) A-A and A-B interactions do not perturb each other. That is, the interaction energy curve u(r) for two adsorbed A molecules is the same as for two A molecules in the gas phase, and the A-B interaction energy is the same for an isolated adsorbed A as for an A with neighboring adso'rbed A's. This is an excellent approximation when only van der Waals forces between atoms need be considered, but is clearly less satisfactory with polyatomic or polar adsorbates, for example. I n the present section we discuss the mobile case making the above assumption, and consider pressures low enough so that multilayer adsorption is not important. (There are of course some systems which never form a monolayer in the usual sense but start on multilayers practically from the beginning, as in the adsorption of water vapor on graphite or other systems with B E T constant c very small.) Al.so, we assume for simplicity a uniform surface. a. Two-Dimensional van der Waals Equation. A simple and very approximate approach to this problem is to assume that the adsorbed molecules obey a two-dimensional van der Waals equation on the surface (Cassel, 12; Hill, 7). According to the assumptions discussed above, the two-dimensional van der Waals constants a' and b' in ---f

(4

+$ ')

( a - Nb')

=

NkT

must be related t o the three-dimensional constants a and b for the same molecules in the gas phase,

THEORY O F PHYSICAL ADSORPTION

219

One finds (7), in fact,

An interesting consequence of Eq. (9) is that it provides (4) an alternative method (Brunauer and Emmett, 13) of assigning the surface area per molecule in a filled monolayer, a quantity necessary t o deduce surface areas by the B E T method. The constant b’ is in fact just the surface area per molecule in a filled monolayer, assuming Eq. (6) t o be valid, and hence. since b = kT,/8p,, b’ = 6.354

(%)’s,

(in A.2)

(10)

where p , is in atmospheres. Surprisingly, the constant 6.354 from the van der Waals equations gives practically the sanle surface area values as the usual liquid density method. The general law of corresponding states determines the functional relation in Eq. (10) but one might expect this particular constant to be unsatisfactory. Livingston (14) has applied Eq. (10) in some detail. The use of b’ is actually slightly inconsistent since this requires a completely filled liquid “lattice”; in the BET method the actual liquid density is used, and real liquids do not have a completely filled “lattice.” Also, there must be connections between the two- and three-dimensional critical constants, and these turn out to be (7,12)

where 1’Zeis the two-dimensional critical temperature, etc. The two-dimensional 4-a isotherms will of course have the familiar (in p-V isotherms) van der Waals equation loops, indicating a condensation from a two-dimensional gas to a two-dimensional liquid if T < Tzc. Correspondingly, the adsorption isotherm, assuming the three-dimensional gas phase t o be perfect,

will also have loops (7) showing that for T < Tzca sudden jump in the amount adsorbed is to be expected a t a certain gas pressure for each temperature. I n Eq. (14), j , is the partition function for the vibration of the adsorbed molecule normal t o the surface, - el is the energy of adsorption a t 0°K. and 0 = rb’ (0 is the fraction of the surface covered).

220

TERRELL L. HILL

Earlier discussions of the van der Waals equation of state for adsorbed monolayers have been given, for example, by Volmer (15),Langmuir (16), and Semenoff (17). Volmer and Langmuir considered the special case a‘ = 0, and detailed thermodynamic functions for this case have been given by Hill (18), allowing b’ to be a function of temperature. b. Theory of Devonshire. Lennard-Jones and Devonshire (19) have developed a relatively successful approximate theory of liquids based on the following “cell” model. In a close-packed lattice, any given molecule is surrounded by twelve nearest neighbors which lie, when in their equilibrium lattice positions, on the surface of a sphere; in the liquid state it is assumed that the twelve molecules are, on a time average, smeared with uniform (surface) density over the surface of this s,phere; the motion of the molecule of interest inside the “cell” is then determined by the potential energy of interaction of the interior molecule with its smeared shell of neighbors, the interaction energy being a function of the distance of the interior molecule from the center of the cell. The fact that a molecule can in due course escape from any particular cell and thus wander over the entire volume V of the container is taken care of rather artificially by the insertion of an appropriate combinatorial factor “ I N ! in the partition function for the system of N molecules. This is the factor which gives the liquid “communal’1 entropy, TABLE I Values of T2,/Tsofor n-Heptane (from Jura et at., 21)

Adsorbent Ag Fen03 Graphite

Theory van der Waals Lennard-Jones and Devonshire

TZ,/Ta, 0 . 5 3 -0 . 5 5 0.56 0.56

0.50 0.53

Lennard-Jones and Devonshire arrive at theoretical critical temperatures which are in good agreement with experimental values for argon, neon, nitrogen, etc. However, the theoretical critical volumes and pressures are off by factors of about 1.6 and 3.9, respectively. Devonshire (20) has worked out the corresponding two-dimensional equation of state. On comparing critical constants,

THEORY OF PHYSICAL ADSORPTION

221

These equations may be compared with Eqs. (11-13). It is seen that the values of Tsc/T8,nearly agree but this is not true for the other critical constants. In view of what was said above about the comparison of Tsc, p,, and V c with experiment, we can in fact conclude that Tzc/Tse = 0.5 is probably about correct, but the other ratios are uncertain.

FIG.4. Low-pressure isotherms of n-heptane on ferric oxide. The curves at 15, 22, 25, 27, 28.5"C. exhibit a finite discontinuity of v as a function of p.

On the experimental side, first order phase changes have been investigated especially by Jura el al. (21), Jura (22), and Gregg (23). Thus, in Fig. 4 experimental low-pressure isotherms of n-heptane on ferric oxide are shown (21). By use of Eq. (4) these data can be put in the form r$ versus l/r, as shown in Fig. 5. The analogy with the more familiar first order phase change gas + liquid is striking. From measurements of Jura et al.,the ratios T2,/TsCin Table I may be obtained. The agreement with theory is better than can really be expected in view of the complicated shape of the heptane molecule (both

222

TERRELL L. HILL

theories assume spherical molecules). The experimental values of a, and 4,, however, are in surprisingly poor agreement with theory (even considering the approximate nature of the theories), as discussed by Jura et al. (21) and Hill (7). c. More Rigorous Theories. Two-dimensional versions of the more formal and rigorous theories of the liquid state due to Mayer (24) and Kirkwood (25) can easily be written down, but are of not much practical value for the present owing to mathematical difficulties.

TORCC.ARCA RCLATIONS FOR f l R S 1 OROCR CHANGE. N-HCPTANL ON fr203

Q,

i

59 ANGSTROYS F f R YOLECULL

FIG. 5. The pressure-area curves for the films of n-heptane on ferric oxide at various temperatures as calculated from the adsorption values of Fig. 4.

3. Localized Adsorption with Interactions

a. Uniform Surface with Nearest Neighbor Interactions. Suppose t h a t there are N adsorbed molecules on B surface sites, t h a t each site has z nearest neighbor sites, t h a t the energy of adsorption is the same on all sites for isolated molecules, t h a t the energy of interaction of two adsorbed molecules on nearest neighbor sites is w, and that on1.y nearest neighbor interactions are appreciable (w > 0 means repulsion). The problem is t o find the thermodynamic properties of this system of N molecules, in particular the chemical potential p since the adsorption isotherm follows from A N ) = pllas(p). I n the special case w = 0, this problem was solved kinetically by Langmuir (26) and statistically by Fowler (27), the adsorption isotherm being just the well-known Langmuir isotherm.

THEORY O F PHYSICAL ADSORPTION

223

With w # 0 the problem is not simple. It can in fact be shown t o be essentially equivalent to certain problems in the theory of the orderdisorder transition in alloys, regular solutions, ferromagnetism, etc. Perhaps the two best known approximate solutions when w # 0 are the Bragg-Williams and quasi-chemical approximations. These are discussed a t length by Fowler and Guggenheim (28) and by Rushbrooke (29) (see also Miller, 30). In the Bragg-Williams approximation one simply assumes that the N molecules are distributed statistically over the B sites as i f w = 0 (that is, the distribution is random), and then one counts up the number of nearest neighbor interactions (with w # 0) using the assumed random distribution. I n the more refined quasichemical approximation one considers that the Bz/2 pairs of nearest neighbor sites are of three types: both sites occupied by adsorbed molecules; one occupied; and neither occupied. One then assumes a quasichemical equilibrium between the three different types of pairs of sites, treating them as “diatomic molecules”:

88

+ 00 e 2 0 8

The equilibrium constant for the “reaction” as written is then 4ewIkT, where 4 is a statistical factor (symmetry number of 2 for 8 8 and 00). Using the “equilibrium constant,” it is then easy t o find the equilibrium number of pairs of both occupied nearest neighbor sites as a function of temperature-which essentially solves the problem. I n these approximations, as well as in higher ones, one finds t h a t when w < 0 (attraction) there exists a critical temperature below which a first order phase change will be observed-a sudden condensation, as the equilibrium gas pressure is increased, from a dilute localized monolayer to a relatively condensed localized monolayer. For a plane square surface lattice of sites, the Bragg-Williams approximation gives - w / k T , = 1 and the quasi-chemical approximation - w / k T , = 1.386. Recently the equivalent two-dimensional ferromagnetic problem for certain geometries has been solved exactly by Kramers and Wannier (31), Onsager (32), and Kaufmann and Onsager (33), using very elegant and powerful mathematical methods. Translating t o our notation, the exact critical temperature for this (idealized) problem is - w / k T , = 1.763. b. Uniform Surface with Higher Neighbor Interactions. I n actual adsorption systems it is probably as important t o take into account second and third, etc., neighbor interactions as it is to use an exact treatment of the nearest neighbor problem. It turns out in fact t o be possible t o include higher neighbor interactions approximately by generalizing the quasi-chemical approximation t o larger groups of sites than pairs (Hill, 34; Yang, 35; Li, 36). For example, using squares of sites in the

224

TERRELL L. HILL

plane square case above and considering the various possible quasichemical equilibria between squares with different types of occupation by adsorbed molecules, one finds the critical temperature to be given by

where w’ is the second neighbor (diagonal of a square) interaction energy, higher neighbor interactions being neglected. It will be noted that when w’ = 0, Eq. (18) gives a slightly better value of T , than the quasi-chemical approximation using pairs of sites. In fact, it is obvious that if one applies the quasi-chemical principle to larger and larger groups of sites, one must converge to the exact solution. This procedure is not practical, however. In summary, the recent work of Kramers, Wannier, Onsager, and Kaufmann provides an exact solution of two-dimensional nearest neighbor problems of the present idealized type; but approximate approaches, for example the quasi-chemical method, must still be used in taking into account second and higher neighbor interactions, heterogeneous surfaces, etc . c. Nonuniform Surface. The treatment of localized adsorption on a uniform surface, above, is a necessary prerequisite t o theoretical consideration of the actual problem of adsorption on a nonuniform or heterogeneous surface. Unfortunately one must expect that the nature of the heterogeneity will in most cases be different for each solid surface considered so that a single general theory is hardly to be anticipated. However, it is definitely worth while to examine some of the simpler possibilities, and some work has been done along these lines. Thus, Halsey and Taylor (37) and Sips (38) have investigated the relation between the adsorption isotherm and the frequency distribution of the energy of adsorption over the different adsorption sites, assuming localized adsorption without interactions and neglecting changes in the vibrational partition function of adsorbed molecules on sites of different energy. Hill (39) has derived expressions for the thermodynamic functions, including the configurational entropy, for localized adsorption without interactions on a nonuniform surface and also has discussed statistically one case in which interactions are taken into account, using the quasi-chemical pair approximation. In this particular case (a random distribution of sites of different energy assumed) it was found that at some temperatures two successive first order phase changes can occur. Hill (39) points out that in order to take into account approximately the dependence of the vibrational partition function on the energy of a

THEORY OF PHYSICAL ADSORPTION

225

site, one should take v,, v y and v, proportional to UO3', a suggestion originally due to Beebe and Kington (39a). Some of the above questions and others have also been considered recently by Tompkins (40). Drain and Morrison (9) have analyzed experimental data for argon adsorbed on rutile (see Sec. 1I.l.b) in terms of surface heterogeneity. They find that the observed zero point entropy of the adsorbed argon is about 0.4 cal./mole deg., which is very much less than that required on the basis of a uniform surface. A satisfactory model for the system, valid for volumes adsorbed less than 0.6 of a monolayer, was found to be that of localized adsorption on a heterogeneous surface without interactions between adsorbed atoms. The energy frequency distribution function was obtained from consideration of the heat of adsorption at 0°K. This model accounts satisfactorily for the magnitude of the experimentally deduced zero point entropy and for the variation of the entropy and heat capacity with surface coverage. Above 0.6 of a monolayer deviations occur as might be expected from the fact that interactions between adsorbed atoms and possible formation of multilayers were neglected in setting up the model.

111. MULTILAYER ADSORPTION 1. Introduction

Below the three-dimensional critical temperature of the adsorbate, multilayer adsorption and eventually bulk condensation (on a nonporous adsorbent) will occur as the equilibrium gas pressure p approaches the vapor pressure po of liquid adsorbate. From a theoretical point of view this problem is extremely complicated. It is well known that even for the simplest substances such as argon or krypton there is no satisfactory theory of the liquid state at the present time. The theories of Mayer (24), Kirkwood ( 2 5 ) , and Born and Green (41) may for practical purposes be considered rigorous and would presumably give excellent agreement with experimentally observed thermodynamic properties of classical (i.e., nondegenerate) liquids with spherically or effectively spherically symmetrical molecules-but the equations which can be written down are so complicated that they cannot be solved for useful numerical results. The best that, can be done along these lines at present is to use Kirkwood's superposition approximation. There are also a number of approximate theories of liquids, but none of these is really very adequate. Now even in the completely idealized case of an adsorbent which is mathematically plane and uniform the theory of multilayer adsorption is

226

TERRELL L. HILL

much more difficult than the ordinary theory of liquids. That is, the adsorption problem is essentially the problem of a fluid in an external field, which is a nonisotropic situation in contrast to the liquid. When an actual surface is considered, the completely general problem is still worse. The simplest “real” surface (“uniform” or “ homogeneous”) is not mathematically uniform but has a periodic nature owing to the crystalline structure, as already discussed above. This means that the external potential field presented by the surface is a function not only of z, the distance of a molecule from the surface, but is also a periodic function of x and y. Finally, we know that surfaces actually encountered experimentally are by no means uniform even in this periodic sense. Thus actual surfaces must present potential fields with a wide variety of wild and disorganized fluctuations along x, y, and z directions. Any general theory of physical adsorption, over the entire presslure range, of the Mayer-Kirkwood-Born-Green type is therefore quite out of the question. A formulation of the uniform surface (potential periodic in x and y ) problem ought to be possible, and preliminary work on the mathematically uniform surface (potential a function of z only) along these lines has already been done by Wheeler (42) and Ono (43). However, we have to recall that even for a mathematically uniform surface the equations must be more complicated than the (at present) unsolvable equations of the Mayer-Kirkwood-Born-Green theory of liquids. Wlneeler is using a “superposition approximation” to avoid this difficulty. More explicitly, the mathematically uniform surfa,ce problem is the following. Given a homogeneous gas at pressure p , we introduce a solid with a mathematically uniform surface having the property that the potential energy of interaction of a gas molecule with the solid as a function of the distance z of the gas molecule from the surface of the solid is U ( z ) [see, for example, Eq. (2)]. Interactions between gas molecules in all possible configurations are taken into account in the configuration integral, the total potential energy of any configuration being the sum of all pair interactions between gas molecules plus the interaction of all gas molecules with the solid. Sufficiently far from the surface the gas density will be unchanged in the presence of the solid but near the surface, if the surface attracts gas molecules, there will be an excess of gas molecules over the number which would obtain if the bulk gas density remained constant right up to the mathematical surface. This surface excess is the amount adsorbed, and in principle this quantity per unit area can be found from the treatments of Wheeler and Ono as a function of p , thus giving the adsorption isotherm. This approach should automatically include, at low p , first order phase changes in the mobile first adsorbed “layer.” I n general, of course, any complete theory of

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227

multilayer adsorption will include everything we have said about monolayer adsorption as a special case (small p ) . An interesting question (see also below) is whether, at low enough temperatures, multilayer adsorption will occur in steps (i.e,, by layers) on such an idealized surface, according t o this rigorous treatment. It should be mentioned that the recent discussion of surface tension by Kirkwood and Buff (44) is rather closely related to the approach outlined above. The Wheeler-Ono point of view is the only really correct way to approach the theory of physical adsorption. However, it is unfortunately true, in view of the status of the theory of the liquid state, that this method is likely to yield useful results in the near future only with great difficulty, after the introduction of mathematical approximations. Approximate theories must therefore be resorted to, and we shall discuss some of these below. Here one makes sufficiently simple assumptions so that the mathematics can be carried through. 2. The BET Theory and Modifications The Brunauer-Emmett-Teller (45) theory of physical adsorption was the first and is still probably the best and most useful theory of multilayer adsorption covering the complete range in p . This is not to say, however, that it is a really satisfactory theory. As will be made clear below, the assumptions of the theory are extremely crude, but they are still sufficiently good to contain a number of the important qualitative features actually observed experimentally. The role of the BET theory in physical adsorption is in fact rather analogous to that of the van der Waals equation in the theory of liquids: it was the first important theory in the field; it has stimulated virtually all the work that followed it; it is still extremely useful as a qualitative guide; but it is not quantitatively correct. The BET theory, of course, is still the basis for the most widely used method of determining surface areas. Especially after the confirmatory work of Harkins and Jura (46), it now seems clear that the determination of surface areas from the BET theory is fortunately not very sensitive to the simplifying assumptions in the BET model, and that BET surface areas are the best that can be had a t the present time. a. The BET Model. Brunauer, Emmett, and Teller (45) gave a kinetic derivation of the BET equation which may be considered a generalization of Langmuir’s kinetic derivation (26) of the Langmuir isotherm. By analogy with Fowler’s later statistical derivation (27) of the Langmuir isotherm, Cassie (47) and Hill (48) provided the statistical deduction of the BET equation. The usefulness of a statistical mechani-

228

TERRELL L. HILL

cal treatment lies in its explicit provision of all constants, parameters, etc., in terms of molecular properties. Also, all questions of particular kinetic mechanisms can be avoided; one considers only the various possible accessible states of the system without regard to transitions between them, etc. The precise nature of the model on which the BET theory is based is thus most clearly exposed by a statistical derivation. The following are the assurdptions one must make: (1) There are B equivalent sites for localized adsorption in the first layer, interactions between molecules in the first layer being neglected. This is just Fowler’s model for “Langmuir adsorption” on a uniform surface without interactions. (2) Each molecule in the first layer is a possible “site” for adsorption of a molecule in the second layer; each molecule in the second layer is a possible “site” for the adsorption of a molecule in the third layer, etc. (We are considering the case of a free surface, with no restriction on the number of layers.) (3) All molecules in the second and higher layers are assumed to have the same partition function (including the energy) as in the liquid state (different than the first layer partition function, in general). Interactions between molecules in the same layer are not talken into account. The adsorbed phase may thus be pictured as a collection of piles of molecules built up on the molecules in the first layer, the piles being energetically noninteracting but related statistically in that the distribution in pile heights will be such as to minimize the free energy of the system. Further, these assumptions in effect assign to the liquid state the completely unrealistic property of each liquid molecule having only two nearest neighbors (i.e., above and below in a pile) whereas a real liquid molecule has something like ten to twelve nearest neighbors. With the above assumptions one deduces (48)the BET isotherm e = - N= B

cx (1

- x)(l - x

+ cx)

(19)

where c is the familiar BET constant and z = p / p o . Equation (19) shows multilayer adsorption and other properties in qualitative agreement with experiment. But this is so far from a rigorous approach along the lines of Mayer-Kirkwood-Born-Green, aB discussed above, that, if anything, one is surprised that the equation deduced from these assumptions is as satisfactory as it is. b. Modijications of the BET Model. MacMillan (49) and Walker and Zettlemoyer (50) have retained the assumption of localized adsorption without interactions in the first layer, but have allowed for a hetero-

THEORY OF PHYSICAL ADSORPTION

229

geneous surface (see Sec. II.3.c). This is unquestionably a step in the right direction as is obvious from experimental heats of adsorption and the failure of the BET equation to follow experimental isotherms below xi Z0.05. Interesting and useful semi-empirical modifications of the BET assumptions have been introduced by Anderson (51), Cook (52), and Dole (53), among others, in some cases leading to a considerable improvement in agreement with experimental isotherms. These modifications have to do primarily with recognition of the fact that all second and higher layer molecules should not be treated as completely equivalent; there will certainly be a more or less gradual transition toward liquid-like properties as the number of the layer increases. Perhaps the outstanding fault of the BET model is its neglect of horizontal interactions in all layers (and in the liquid). As is clear from Eqs. (18-21) of reference 48 and from the original kinetic derivation, the BET model is actually equivalent to a Langmuir treatment of each layer, the nonconfigurational partition function of an adsorbed molecule being different in the first than in other layers, and the distribution of molecules among different layers being found by minimizing the free energy. Hill (54) has taken into account horizontal interactions in all layers, approximately, by retaining this point of view except for the use of the BraggWilliams (see Sec. II.3.a) approximation in place of the Fowler-LangmuirBET neglect of horizontal interactions. Numerical calculations were carried out for a simple cubic arrangement of sites, which corresponds to assigning a liquid molecule six nearest neighbors, a considerable improvement over the BET theory. However, the model is of course still very approximate, retaining the ideas of a uniform surface, localized adsorption in the first layer and a lattice type of structure for the liquid state. Complete calculations were not carried out, but the results obtained corrected the BET theory in the desired direction of reducing the amount of adsorption (above a monolayer) for a given relative pressure, p / p o . In addition, these calculations indicated that below a critical temperature (or temperatures) multilayer adsorption would occur in steps. Whether or not this result is “correct” for a hypothetical uniform surface or merely a property of this particular approximation cannot of course be decided with certainty until rigorous calculations of the Wheeler-Ono type can be carried out. The result may well be different above and below the melting point of the adsorbate. Halsey (55), however, taking into account horizontal interactions of the above type in an equivalent but somewhat more qualitative way, has concluded that multilayer adsorption in steps would indeed occur on-an ideally uniform surface, and that the reason steps are practical1y:never

230

T E R R E L L L. H I L L

observed experimentally is that actual surfaces are heterogeneousspoiling the exact geometrical arrangements necessary for step adsorption, and smearing such effects into a continuous isotherm. It has further been pointed out by Teller (56) that even in the case of adsorption on a perfect crystalline surface, unless the adsorbate and a,bsorbent crystal (or liquid) structures just happened to fit perfectly (e.g., water vapor on silver iodide) sufficient dislocations could probably occur to spoil step adsorption. I n this connection, Halsey (57) has emphasized that since successive steps in higher layers would be separated by very small differences in p / p o , only very minor dislocations would in any case be necessary to smooth out an adsorption isotherm above, say, two or three absorbed layers. As shown in the work of Hill (54) and Halsey (55) mentioned above, if one recognizes that liquid and adsorbed molecules actually have horizontal as well as vertical interactions, then adsorbed molecules, especially in second and higher layers, will have very different properties than in the BET theory. Halsey discusses these properties but unfortunately in addition makes the misleading remark that “on the basis of the BET hypotheses the adsorption for any value of c is actually confined to the amount accommodated in the first layer.” What is olsviously meant is that this result would obtain if a liquid with both vertical and horizontal interactions is incorrectly allowed only vertical interactions on adsorption. Actually, the BET hypotheses are self-consistent and lead to multimolecular adsorption for any value of c ; the hypotheses include the assumption that both liquid and adsorbed molecules have only vertical interactions. Halsey’s remark incorrectly implies that the BET theory uses the inconsistent hybrid assumptions mentioned above. c. Surface Tension and the BET Theory. Cassel (58) has emphasized a failure of the BET theory of a very fundamental nature. It will be recalled (“Gibbs equation”) that the integral (Eq. 4) kT

r d In p

where r is the surface concentration obtained from an adsorption isotherm, gives the surface or spreading pressure ~ $ ( p ) , +(PI = Ye - Y ( P ) ,

where y a is the surface tension of the clean surface and y ( p ) is the sum of the surface tension of the adsorbate and the solid-adsorbed film interfacial tension. A t p = pa, +(PO)

=

YE

- Y L - YLS,

THEORY O F PHYSICAL ADSORPTION

23 1

where y L is the surface tension of the bulk liquid and y L s is the bulk solidis a finite quantity. But liquid interfacial tension. Clearly &),I Cassel pointed out that substitution of the BET isotherm into the Gibbs equation gives @ ( p ) / a k T = In [(l - x

and hence @(Po) =

+

-+ cx)/(l

- z)]

(20)

O0,

where z = p / p o , c is the familiar BET constant, and a is the number of first layer adsorption sites per unit area. Cassel attributed this infinity catastrophe of the BET theory to the neglect of lateral cohesion.

FIG.6. Comparison of (1) the BET isotherm, (2) Eq. (38), and (3) experimental data for nitrogen on anatase.

The fact that the Gibbs integral for 4 ( p o )is infinite for the B E T theory agrees with the well-known fact that above about z = 0.35 the BET theory in general predicts a greater adsorption than is observed experimentally; the experimental isotherms approach the line x = 1 for large 0 much more closely and in such a way that although 0 - 03 a t x = 1, the Gibbs integral remains finite at z = 1 (see Fig. 6). Another way of putting this is to say that the BET theory assigns too low a free energy (and hence too low a vapor pressure p ) to the adsorbed phase for a given 0, relative to the liquid state. Thus, in the statistical

232

TERRELL L. HILL

BET theory the chemical potential of the bulk liquid is (48,54) pL/kT = - I n

j L

- (sL/kT)

(21)

while the chemical potential of the adsorbed film is &=ln--------lnjL-m N - X N

CL

where j, is a partition function and eL an energy, and X is the number of adsorbed molecules in the first layer. Both terms in p r / k T occur also in p / k T because liquid-like properties are assigned to adsorbed molecules in second and higher layers (or rather, in effect, the liquid is assigned the properties of a one-dimensional pile of molecules : each liquid molecule has only two nearest neighbors). The remaining term in p / k T , In ( N - X ) / N , has a configurational or entropic origin (i.e., it is connected with the number of ways of shuffling adsorbed molecules among different piles and layers; however X is found by minimizing the free energy so the separation between energy and entropy is not clean-cut) and accounts entirely for the difference between p and p L . After the first layer is filled, X = B , 8-1 kT = 1 n x =ins = I n rp- N E P

PL

or E = 1 / ( 1 - 2)

The function 1/(1 - x ) is responsible for the infinite integral mentioned above; its primarily configurational origin is made clear in the present argument. The more refined but still approximate treatment of horizontal interactions (Bragg-Williams approximation) discussed in Sec. III.2.b complicates the configurational term in Eq. (22) and also the term - (eL/kT), the net effect of the two alterations being to increase the free energy (lower the isotherm) of the adsorbed molecules. This is an indirect confirmation of Camel’s remarks about the importance of the BET neglect of horizontal interactions (though it must be remembered that neither the liquid nor the adsorbed film has horizontal interactions in the BET theory). Qualitatively, the introduction of horizontal interactions tends, despite the configurational entropy, to flatten the BET piles of molecules, since this allows a greater number of nearest neighbor interactions (see Sec. 111.4, below). d. Adsorption from a Mixture of Gases. The kinetic derivation of the BET equation has been extended to include adsorption from a mixture of gases (Hill, 59). As might be expected, recent experimental attempts

THEORY O F PHYSICAL ADSORPTION

233

to verify this development-as a further test of the BET postulateshave been fairly successful but not completely so (Arnold, 60; White and Schneider, 61). Ultimately, on the theoretical side, what will be needed here is an extension of the theory of liquid solutions analogous to the extensions of the theory of one-component liquids already discussed in the case of adsorption from a single gas. 3. The Huttig Equation

Hiittig (62) has derived the adsorption isotherm

where c is the BET constant, by a very fundamental modification of the kinetic BET argument (45). Namely, he assumes that, for example, the rate a t which the number of molecules in the second layer increases is proportional to the number of “uncovered” first layer molecules (in agreement with BET), and the rate at which this number decreases is proportional to the total number of molecules in the second layer (in disagreement with the BET theory which uses instead the number of uncovered second layer molecules). Otherwise the assumptions of the two theories are identical. Ross (63) and Fergusson and Barrer (64) have examined the Huttig theory in some detail, the latter authors giving a statistical mechanical derivation. Although no objection can be raised to the Huttig equation as an empirical equation, in the writer’s opinion the derivations (both kinetic and statistical) on which it is based are fallacious (65). We shall mention three qualitative reasons for believing this, and then give a more formal criticism: (1) Since liquid-like properties are assigned t o molecules in second and higher layers, as in the BET theory, one expects that on a plane surface e --+ a as 2 --+ 1. But the Huttig equation gives e I 2 for 2 = 1. (2) The Huttig and BET models are the same except for a change in the kinetic mechanism of attaining equilibrium. Now the statistical derivation (48) of the BET equation is independent of any choice of kinetic mechanism, and since otherwise the two models are identical, one is forced to conclude that the Hiittig kinetic mechanism is impossible [see (3)l. (3) The Huttig mechanism is thermodynamically unsound because it violates the principle of microscopic reversibility. Consider the second layer equilibrium again. The Huttig equations assume implicity (with BET) that there are no vacant positions in a “pile” of adsorbed mole-

234

TERRELL L. HILL

cules. Hence when a covered molecule escapes from the second layer (as allowed by Huttig) the molecules above it in the pile must all move down one position to fill the vacancy. In effect, then, such a molecule escapes from the top of the pile. The “principle of equal frequency for reverse molecular processes at equilibrium ” (66) is violated here because we are equating the frequencies of two processes which are not the reverse of each other: (a) adding molecules to uncovered first layer molecules; but ( b ) removing molecules from the top of any pile containing two or more molecules. (4) We show here the fallacy in the statistical derivation published by Fergusson and Barrer. Let Qi be the partition function of the molecules in the ith layer. Then Fergusson and Barrer write (see Eqs. 18-21 of reference 48)

where Xiis the number of molecules in the ith layer and pi( T ) is a molecular partition function (48,64). Now for the entire adsorbed phase 6? = &1&2&3 . . . N=X1+X*+Xa+

’.

where Q is the complete partition function of the N a,dsorbed molecules. To find the equilibrium distribution of the N molecules in the different layers, the correct procedure is to minimize the Helmholtz free energy (i.e., maximize In Q) at constant B , T, and N . That is, using the undetermined multiplier a, a [ a ~(In, Q - a N )

]

Xk.B.T

=o

i = l , 2 , .

or

where we understand X O= B. We define

so that the equilibrium condition is pi

+

/li.i+l

=

-akT

i = 1 , 2,

...

”

THEORY OF PHYSICAL ADSORPTION

Now d(1n Q

-

]

crN) =

Xk.5.T

=o

dX,

(B,Tconstant)

235

(30)

i

using Eq. (26). That is,

(%kQLT =

O1

But (8 In Q / c ~ N is) ~of , course ~ also just -p/lcT, where p is the chemical potential of the adsorbed phase (at equilibrium, p = p J . Therefore pi

+

pi.i+1 = p =

is the equilibrium condition. incorrect condition

pgss

i = l ,2,

...

(32)

Fergusson and Barrer used instead the

p i = p = pgm

i = 1,2,

...

(33)

which leads to the Huttig equation. Since Eq. (32) is consistent with the lowest possible Helmholtz free energy (and leads, incidentally, to the BET equation), Eq. (33) must give a higher free energy. This is the essential reason why, for the same value of c, the Huttig isotherm is always below the BET isotherm and hence accounts for the accident,al better agreement of the Huttig isotherm with experiment for x < 0.8 (i.e., for other reasons already discussed, the BET model has too low a free energy relative to the liquid state). Huttig (67) has recently proposed a “compromise” equation with an adjustable parameter y allowing a transition from the Huttig equation (y = 0) t o the BET equation (y = 1). It is obvious from the above discussion that this feature cannot improve the logical position of Huttig’s equation. The minimum free energy for given N , B, and T is obtained only for y = 1. All the discussions of Huttig (62,67), Ross (63), Fergusson and Barrer (64), and Hill (65) are based explicitly or implicitly [see, for example, Eqs. (25)] on the BET picture of molecular piles. We have shown above that the Huttig equation cannot be obtained f r o m this model without violation of statistical thermodynamical principles. This does not mean, of course, that the Huttig equation is not valid as a strictly empirical equation. In fact, one can go farther than this and assert that the actual Huttig distribution of N adsorbed molecules among different layers is conceivable on some entirely different grounds from BET molecular piles. Barrer (68) has investigated just this point, introducing an appropriate combined potential field owing both to the solid adsorbent and to the adsorbed molecules themselves. However, it is the present writer’s

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TERRELL L. HILL

opinion that such an attempt to “rescue” the Huttig equation, though very interesting, is quite artificial. Indeed a potential field can be found to “account” for a n y empirical distribution of molecules among layers.

4. The Adsorption Isotherm for 8 Large An approach completely different from any discussed above has been used by Frenkel (69), Halsey (55), Hill (70), and MacIMillan and Teller (71) in considering the restricted problem of adsorption of a spherically symmetrical molecule on a nonporous, nonpolar adsorbent for, say, 0 > 2. It now appears that a very simple and rather successful discussion of this problem can be given. This is possible because, as far as molecules being adsorbed in, say, the third and higher layers are concerned, (a) the effects of the detailed structure of the adsorbent surface are effectively smoothed out and ( b ) the molecular environment of these molecules is not very different from that of corresponding molecules in the bulk liquid. The remaining effect, which is predominant in determining the adsorption behavior, is the potential energy field in which a molecule is adsorbed compared with the field in the bulk liquid. It will of course be noticed that this resembles the point of view of Polanyi’s (72) well-known potential theory (see also Sec. 111.6). Frenkel discussed this problem first, but, independently of Frenkel’s work, it was also treated by Halsey and by Hill. Halsey deduced the isotherm lnx

=

-a/I”

(34)

by a semiquantitative argument which did not specify a and s in terms of molecular properties. Frenkel and Hill obtained Eq. (34) also, but related a and s explicitly to intermolecular forces. We shall give a simple derivation of Eq. (34) below, using the notation of Hill., Actually, Hill’s original treatment is more general than the derivation t o be given, taking into account the difference in molecular environments mentioned in (b) above. We assume that we have a solid plane adsorbent of density p (molecules per cubic centimeter) in the semi-infinite space z < 0, and an adsorbed film in the form of a slab with parallel walls in the region 0 5 z 2 h. The adsorbed film is assumed to have the bulk liquid density uniformly, p L = r / h . We consider only potential energy effects, as explained above. Suppose the potential energy of interaction of a liquid molecule with an adsorbent molecule is given by Eq. (1). Then the interaction with the entire adsorbent of a liquid molecule a distance E from the surface

THEORY OF PHYSICAL ADSORPTION

237

of the adsorbent is given by Eq. (2). The potential energy of interaction of two liquid molecules is assumed to be of the form of Eq. (1) with parameters €1 and n*. Now (Eq. 23) k T In x =

- pL = a(A - At)/aN

where A is the Helmholtz free energy of the film. Let U Sand UL be the potential energies of the N molecules in the adsorbed slab and in the bulk liquid, respectively. Then (see above) we are to assume that other effects cancel and A - A , = US - U L ,so that kT In x

=

aUs/aN

- (UL/N)

(35)

Since we are assuming the slab to have a uniform density independent of h, the change d U s in US owing to the adsorption of dN additional molecules comes entirely from the d N additional molecules. That is, a Us/aN is just the potential energy of interaction of a liquid molecule at z = h with (a) the liquid slab, 0 <: z 5 h, and ( b ) the adsorbent, z < 0. Now the energy (a) is equal to the interaction of a molecule a t z = h with (i) the semi-infinite liquid z < h less the interaction with (ii) the semiinfinite liquid z < 0. But the energy (i) is just U L / N ,the semi-infinite liquid being appropriate since in the infinite liquid every interaction would be counted twice on this basis. Equation (35) becomes then kT In x = potential energy ( b )

- potential energy (ii)

(36)

the other terms cancelling. The potential energy expressions in Eq. (36) are given by Eq. (2). Neglecting the very short-range repulsive term in Eq. (2) (for 6 > 2), Eq. (36) becomes - er *%p elrl*%pL kTInx = 7 +T 3h

Using h

= II/pL,

Near z

=

(37)

we have finally

1, using In x = z

r

- 1,

=

&/(I

- s)n

(39)

which is to be compared with Eq. (23),

r

= const./(l

- s)

for the BET theory. Equation (39), unlike Eq. (23), leads to convergence of the Gibbs integral for the surface pressure 4 ( p o ) . It will be

238

TERRELL L. HILL

recalled that Eq. (23) had a configurational origin while, in contrast, Eqs. (38) and (39) represent pure potential energy effects (entropy neglected). It appears that the latter point of view is more nearly correct for, say, 0 > 2, entropy considerations and other complications introducing corrections (which may, however, be important-see Sec. 111.6). In Fig. 6, adapted from MacMillan and Teller, the data of Jura and Harkins (73) for nitrogen on anatase are compared with the BET equation (Eq. 23) 0 = 1/(1 - x) and Eq. (38) in the form In x = - K / @ , for different values of K . It is clear that the choice of the exponent s = 3 and the coefficient K 4 gives excellent agreement with the experimental points, in distinct contrast to the BET theory. MacMillan and Teller have shown by an approximate argument that K = 4 is of the right order of magnitude, according to the expres,sions for 6 and 6 L in Eq. (38). Halsey (55) has in fact plotted experimental datab for a number of examples in the form In In x against In f3 t o obtain the exponent s. In the case of nitrogen (the only nonpolar adsorbate) adsorbed on anatase, using the data mentioned above, Halsey found the best fit to be given by s = 2.67, agreement with experiment extending (fortuitously for small x) from x = 0.0026 to 0.9936! It will be of great interest to examine in this way even simpler systems such as argon on nonpolar adsorbents, etc. Thus, Drain and Morrison (9) have found recently that s = 2 for argon on rutile (9), up to 0 = 4.5. Incidentally, Harkins and Jura (46) have used Eq. (34) as an empirical equation with s = 2. The case of most interest above is clearly a > 0, that is 6 > 6L. Roughly speaking, this means that the attraction between adsorbent and adsorbate molecules is greater than the attraction between adsorbate molecules themselves (the liquid “wets” the solid). The amount of adsorption decreases as a 4 0 (Fig. 6) ; when a = 0 there is no adsorption at all up to x = 1, at which point there is sudden condensation. When a < 0, questions of supersaturation and film instability arise which are discussed in detail by Frenkel (69) ; the concept of surface excess must be used and the adsorption is negative for p < po (see Sec. IV.5.a). In all of the above discussion it has been assumed without formal justification that if the adsorbent has a plane surface, the adsorbed film will have a parallel plane surface. MacMillan and Teller (71) have investigated the validity of this assumption. Using a Fourier analysis, they have considered the various surface waves possible in the adsorbed film and in the bulk liquid, calculating a potential (surface) energy in each case for use in a Boltzmann factor. Thus, large waves with considerable extra surface have a relatively high energy and a small Boltz-

THEORY OF PHYSICAL ADSORPTION

239

mann factor. The film and the bulk liquid differ in that the amplitude of waves in the film is limited by its finite depth. In the absence of a surface energy d, there is no energetic disadvantage to waves with large surface area. This in fact resembles the situation in the BET theory where lateral interactions are not taken into account and strong fluctuations in pile heights occur; indeed, MacMillan and Teller show that the BET result, Eq. (23), is obtained from their theory on putting e l = 0 [this is, incidentally, not inconsistent (74) with the discussion of Eq. (20) which implies that the surface tension yL = - w for the BET theory]. Putting E’ = co gives the plane film surface assumed by Frenkel, Halsey, and Hill. The actual situation is intermediate, with d finite and positive. The equations of MacMillan and Teller are not designed to take into account the effect considered by Frenkel, Halsey, and Hill, but rather to provide the correction term necessary because of the unjustified assumption of a plane surface. Intuitively, it is not obvious whether this correction term is going to be small or large, perhaps larger than the term it is supposed to correct. Actually, it turns out in the detailed analysis of MacMillan and Teller that the correction is relatively unimportant under usual conditions, and hence that the plane surface (c’ = m ) assumption is rather good. The correction term may, however, be important when a is very small or negative. Rather curiously, the correction term turns out to have the same form as Eq. (38). However, the coefficient is much smaller in general and is proportional to k T / d . MacMillan and Teller show, incidentally, that the two terms [Eq. (38) and correction] are additive, to a sufficient degree of approximation. It should be emphasized that the reason the MacMillan and Teller theory can reduce to Eq. (23) as a special case but does not include Eq. (38) is that the BET and MacMillan and Teller treatments both omit the potential fields (b) and (ii) of Eq. (36) (the solid interacts only with the first adsorbed layer in the BET theory). Included in the paper of MacMillan and Teller is an excellent analysis and review of the BET theory and related topics.

+

5. Capillary Condensation and Hysteresis

From a theoretical point of view there are of course no new physical principles involved in adsorption on a porous adsorbent. It should therefore be possible to retain the methods discussed in detail above and include capillary condensation and hysteresis within the framework already provided. Thus, it is easy to see that even with hypothetical

240

TERRELL L. HILL

pores consisting of two plane parallel walls both capillary condensation and hysteresis can occur. For example, suppose we have parallel adsorbent wdls several molecular (adsorbate) layers apart. Extending (Hill, 75) the argument leading to Eq. (38), as x increases there will first build up independent adsorbed films on the two walls; the predominant effect being, as in Eq. (38), the drop off in the attractive field of the adsorbent as the depth of the films increases. That is, the chemical potential (and hence the vapor pressure) of the adsorbed phase increases with 0, as usual. But as the two films approach each other through further building up on both walls, the van der Waals attraction between the two films becomes important and can cause the chemical potential (and vapor pressure) to decrease with increasing 0. When the capillary is filled or nearby filled the chemical potential must again increase, owing to the limit on the number of molecules that can be accommodated. This gives an isotherm in the form of a loop. The thermodynamically stable path is to cut across the loop by a vertical jump (capillary condensation) a t that value of x which has the same chemical potential at the two terminals of the jump. Experimentally, strictly vertical jumps are rare!; they are usually smoothed out by the pore size distribution. Hysteresis loops are possible (54) by following metastable but accessible parts of the loop in adsorption or desorptiom, or both. This predicts that the desorption branch will always lir! above the adsorption branch, in agreement with experiment. An alternative but completely equivalent discussion of this subject can be given (54) using the lattice and Bragg-Williams method discussed in Sec. III.2.b. Everett (75a) has adopted this same point of view as a starting point in a recent and more detailed analysis of hysteresis. 6. Some Recent Advances

It has been pointed out above that the Wheeler-Cho approach (see Sec. 111.1)to the idealized mathematically plane surface problem is the rigorous approach, though actual numerical calculations based on the general equations are not practical. On the other hand, the FrenkelHalsey-Hill method (see Sec. 111.4) is essentially a very approximate solution of this same problem resulting in a simple and surprisingly successful isotherm equation, Eq. (38), for B not too small. This method can be applied to capillary condensation (see Sec. 111.5) and is capable of accounting for isotherm types I1 to V (1,55,75). The Frenkel-Halsey-Hill theory considers the adsorbed film to have the bulk liquid density out from the surface to a point at which there is a

THEORY O F PHYSICAL ADSORPTION

24 1

discontinuous drop down to the equilibrium gas density. The difference in potential energy between this film, next to a solid, and the same amount of adsorbate in bulk liquid leads to Eq. (38). In a recent paper Barrer and Robins (76) have introduced an important refinement (in the direction of the Wheeler-Ono treatment) in the above picture by use of the van der Waals equation as equation of state not only for the gas and liquid but also for molecules in the film. The chemical potential must have the same value at all distances z from the surface, and this value is determined by the temperature T and equilibrium gas pressure p far from the surface (z = m). This chemical potential a t any point z’, according to Barrer and Robins, is made up of two terms: (1) the chemical potential of the usual van der Waals bulk fluid, appropriate to T and the local density p ( z ’ ) ; and (2) the potential field of the solid, V(z’),falling off as r3(Eq. 2). Since 1.1 ( = constant) and the potential field of the solid are known, one can solve for the density p at each 2. This gives a film density p ( z ) which is not constant going out from the surface and not the density of bulk liquid. The adsorption isotherm is found by computing the surface excess (amount adsorbed) from p ( z ) for each p . Isotherm types I to V (1,76) may be computed. A further refinement of this approach, again a step in the Wheeler-Ono direction, was suggested by Hill (77,78), removing the following difficulties contained in the theory of Barrer and Robins: (1) The contribution of the film t o the chemical potential at z’ is determined not by the local density at z‘ but actually by the entire curve p ( z ) and in particular the fact that p ( z ) = 0 for z < 0, responsible for the term in pL in Eq. (37), should not be ignored. (2) There is a physically unsatisfactory discontinuous jump in density* remaining, owing to the use of the local density p ( z ’ ) . (3) The theory does not reduce approximately to Eq. (37) for large e (the term in pL is missing, as mentioned above). (4) The limiting case z + 1 gives a discontinuous transition from the bulk liquid density (at po) to the bulk gas density (at PO). (5) The van der Waals equation of state can be replaced by better models of the liquid state, for example, the gas of hard spheres with intermolecular attractions superimposed (78), or the Lennard-Jones and Devonshire (19) theory of liquids. *The surface waves of MacMillan and Teller (71) would give a mean density transition curve p ( z ) that is continuous, although each individual wave has the same type of discontinuity as in the Frenkel-Halsey-Hill theory. The MacMillan and Teller approach is essentially one way of classifying a certain restricted set of configurations consistent with a continuous mean density transition.

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TERRELL L. HILL

The disadvantage of the treatment of Hill (77,78) is of course t h a t the computations, though certainly feasible, are lengthy; a nonlinear integral equation must be solved numerically by successive approximations t o get p ( x ) for each p . Calculations of this type are in progress. The method can also be applied t o the region of the critical point, capillary condensation, and adsorption from a mixture of gases, for example. A completely different type of extension of the Frenkel-Halsey-Hill theory has been made recently by Halsey (79), who has considered a heterogeneous surface of special type, consisting of a number of regions, all sites in a given region having the same energy of adsorption. These regions or patches must be large enough so that a given patch by itself would follow Eq. (38) and so that edge effects (between patches) are negligible. An empirical exponential distribution of population of sites of different energy is assumed. From this foundation and for this special kind of surface heterogeneity, Halsey derives an equa,tion which proves to be a very convenient tool for generating mathematically a wide variety of isotherms. Using some of these isotherms, Halsey shows how the BET surface area method can succeed or fail depending on the shape of the adsorption isotherm. All the refinements considered in this section can be expected t o alter somewhat the third power law of Eq. (38) to some other “effective” power. Halsey has discussed this point in conneclion with his new isotherm equation (79). It is certainly clear from an examination of the experimental differential entropy curve for argon on rutile (9) that the entropy terms dropped from Eq. (35) are not negligible in this case, at least up t o 0 = 4.5. I n fact, in the notation of Eq. (35), aA’,/dN - ( S L / N )is about one-half of R In x a t 0 = 2.4 and these two quantities are about. equal at 0 = 4.5. One may hope that the refinements discussed above will make some contribution t o understanding such entropy effects at8 well as deviations from the third power law (the two problems are of course related). The Frenkel-Halsey-Hill model has also been extended t o adsorption on small spherical particles, an important consideration in nucleation of condensation by foreign nuclei (75).

IV. THERMODYNAMICS OF ADSORPTION I . Introduction For some time there has been a certain amount of confusion about the application of thermodynamics t o adsorption data. However, it seems possible now, after the work of Cassel (80),Coolidge (81), Rowley and Innes (821, Hill (18,83), Gorter and Frederikse (84), Hansen (85),

THEORY O F PHYSICAL ADSORPTION

243

Everett (86), Kington and Aston (87), and Guggenheim (87a), to say that general agreement has been reached on all questions of thermodynamic correctness or incorrectness. There is still some difference of opinion as to relative utility and completeness of different approaches, etc., but at the present time these are to a large extent matters of personal taste. Future work will perhaps lead to general agreement even on some of these questions. We shall summarize below the present status of the more important aspects of this subject. This seems worth while since the papers listed above overlap each other considerably and contain more detail than is necessary for applications to experimental data. We consider a one-component gas and a one-component sorbent throughout for simplicity. In accordance with common thermodynamic practice we use in this section the gas constant R and “moles” instead of the Boltzmann constant k and “molecules.” 2. Solution Thermodynamics

If we consider a list of sorption systems such as argon-graphite, hydrogen-tungsten, hydrogen-charcoal, hydrogen-palladium, benzenerubber, water-sodium chloride, and water-sulfuric acid, it is clear that, from the point of view of pure thermodynamics, no really significant division of these systems into separate classes can be made, and therefore that the same thermodynamic treatment should apply to the entire list. This is a well-known point of view and has been especially emphasized by Coolidge (84). In particular, since ordinary solution thermodynamics obviously applies, say, to water-sulfuric acid, this approach can be extended to strictly adsorption systems as well, such as argon-graphite. Consider a two-component (sorbent “ A ” and sorbed gas, “ 1”)condensed phase in equilibrium with gas, “A” being nonvolatile. The condensed phase has energy El entropy S,etc., and contains nAand n1 moles of the two types. We include only those cases in which E is completely determined by S, V , n A , and nl. This requires no further comment if A is a liquid (e.g., sulfuric acid-water) or a solid (e.g., palladiumhydrogen) such that surface effects are negligible compared with bulk effects. However, if A is, for example, a finely divided solid with appreciable surface effects (the case of primary interest here), it is to be understood that in all of the thermodynamic processes contemplated a change in nAIdn,, refers to an addition of pure A in the same state of subdivision and specific surface, etc., as the pure A used in preparing the original sample. Thus, in cases where the term “surface area” has meaning, the area is proportional to nA and is not an additional independent variable.

244

TERRELL L. HILL

Using the standard methods of solution thermodynamics, we can then write dE dF dpl

= = =

+

+

TdS -PdV pldnl prdnr - S d T f V d P p l d n l 4-pAdnA -&dT elaP ( a p l / a r ) T . p d r

+

+

+

where I’ = nl/nZA, B1 = (aS/anl),A,p,T, etc., P is the hydrostatic pressure, exerted by a hypothetical inert piston or (in part) by a hypothetical inert additional gas, V is the volume of the condensed phnse (we ignore the fact that in certain cases it is not easy to define V precisely-see Sec. IV.5), and the p’s are chemical potentials. For the gas, we have the well-known expression dpo

=

-8adT

+ vodp

(43)

where so = So/nQ,etc., and p is the gas pressure. At equilibrium between gas and the condensed phase, d p l = dpa, so that -&dT

BldP -k (apl/ar)T.pdTI = -&dT

+ VQdp

(44)

Under ordinary experimental conditions P = p and we have a t constant I’ (ap/aT)r =

(SQ

- & ) / ( v u - iil) (P = P:)

(45)

I n the usual approximation of a perfect gas and vQ >> el,

=

(HQ

- 81)/RTB = q.t/RTZ

where H=E+PV HQ = EQ p V a

+

and, of course, HQ

- fIl

=

T(SQ -

&I)

a t equilibrium. Equations (45), (46), and (47) give the so-called isosteric (constant I’) heat of adsorption, q s t . It should be emphasized that since we have used the amount of sorbent nAand not the surface area a, Eqs. (45), (46), and (47) apply regardless of whether the sorbent swells on sorption, or the area or number of sorption sites changes with temperature or pressure, or even whether an “area” exists. Thus the above “textbook” derivation of Eqs. (45), (46), and (47) is quite sufficient, as Halsey (55) has remarked, to dispose of the proposed modification of the isosteric equation by Wilkins (88). The discussions of Armbruster and Austin (89) and Pierce and Smith (90) are also incorrect, in the writer’s opinion. Suppose we consider also a sample of pure A with the same state of

THEORY OF PHYSICAL ADSORPTION

245

subdivision and specific surface, etc., as used in preparing the LLsolution ” of Eq. (40), and with the same nA,P, and T :

since Eo,, S O A , and V O A are independent of nl. The ‘(approximate” equations ( G ) are in general very accurate because of the “solution” being a condensed phase with small pressure dependence. In fact, H , E E,. Eqs. (46) and (47) become then

Equations (52) and (53) hold regardless of any perturbations of the sorbent, etc., as discussed above. However, there is really no advantage in using H , and AS,,as formally defined above, over H and S [Eqs. (46) and (47)] except in the important special case of an inert adsorbent, by which we mean a hypothetical adsorbent whose own thermodynamic properties are unaffected by the presence of adsorbed molecules and whose surface area is independent of temperature and pressure. We can then replace nx by in Eqs. (52) and (53) and S, and H , become just the entropy and heat content of the one-component system of nl moles of adsorbed gas. In effect, the adsorbent merely plays the role here of an external potential field. At the present early stage of our understanding of physical adsorption, this approximation certainly seems justified in most cases and indeed is made implicitly by almost all workers in the field. We shall make this simplification below except where otherwise noted. S, and H , are the total or integral entropy and heat content of the onecomponent system of adsorbed molecules. An isosteric calculation [Eqs. (52) and (53)] from neighboring adsorption isotherms gives the diflerential entropy and heat content, (aS,/an,)a,, and (aH,/anl)a,,. The integral entropy S , is the quantity of direct statistical mechanical significance, being related to the number of possible quantum states of *We adopt the H . and Xa notation of Everett (86), for clarity.

246

TERRELL L. HILL

the system Q(E8,a,nl) by 8, = k In Q, where k is the Boltzmann constant. Arguments concerning order-disorder, randomness, etc., of the adsorbed molecules must be referred to S, (or conventionally and more conveniently to the molar integral entropy, s8 = S,/nl) rather than to (dS,/an,)a,T. It is unfortunately not uncommon in the literature t
where qat is the isosteric heat of adsorption [Eqs. (47) and 53)]. (2) If the heat of adsorption is measured isothermally and reversibly (including a PdV work term) there is in addition to q d an extra “heat of compression,” and one finds the “isothermal” heat of adsorption (18)

where Vois the volume of the gas phase. The derivative can be obtained from the slope of the adsorption isotherm. (3) If the heat of adsorption is measured adiabatically and reversibly (including a PdV work term) there is an analogous heat of compression with the result (87)

where aT/anl is a t constant total entropy (adiabatic), and q,, is the adiabatic heat of adsorption.

THEORY O F PHYSICAL ADSORPTION

247

There has been considerable experimental uncertainty for some time about the connections between calorimetric and isosteric heats. Hill (18) showed that in the reversible isothermal process, Eq. (56) must be replaced by Eq. (57). Kington and Aston (87) derived the analogous Eq. 58 for the reversible adiabatic process and then proceeded to show experimentally that their adiabatic calorimeter actually behaves reversibly. Thus for six values of 0 from 1.16 to 1.33, q. is larger than pat by successive values of 84, 107, 143, 128, 154, and 111 cal./mole. But q. is larger than the right-hand side of Eq. (58), i.e., qst properly corrected, only by the successive values -48, -25, +lo, -9, +14 and -34 cal./mole, which is excellent agreement in view of the estimated maximum error of & 15 cal./mole in q. and & 15 cal./mole in qRt. Hence the work of Kington and Aston completely clarifies for the first time the relations between isosteric and calorimetric heats of adsorption. Ward (91) had made use of the heat of compression in an empirical way in a paper which has been generally overlooked until referred to by Kington and Aston. Incidentally, Kington and Aston (87) derive an equation for qat and state that “ . . . Hill’s expression for qat explicitly contains a term involving 4.” To avoid any possible confusion, it may be remarked that this reference is to Hill’s paper V (18) which treated the adsorbed molecules as a one-component system. Paper IX of the series (83) gives the same derivation of qst as Kington and Aston (87) [this is the derivation given above, Eqs. (40) t o (47)l.

4. Adsorption Thermodynamics We have mentioned above the tendency to confuse ss and (a,S,/an,)a,,. The differential entropy can be obtained experimentally from pat, or calorimetrically. We now wish to discuss how s, may be found since, from the point of view of theoretical interpretation using statistical mechanics, this is the entropy of most interest (83). (1) In principle one can find the differential entropy down to very low coverages, extrapolate (aS,/an,>a,. or, better, qst to nl = 0 and integrate. However, this is a very inaccurate way of finding s,. (2) From the same data (adsorption isotherm measurements to low coverages) one can calculate = RT

[r d In p

(T const.)

(59)

using a much more accurate extrapolation and then employ

(y),=

to find s,.

(SG

s.)/RT

(60)

248

TERRELL L. HILL

(3) Integral calorimetric heats of adsorption or heats of immersion, together with a calculation of 9 at one temperature may be used, as shown by Jura and Hill (92) and Drain and Morrison (9). (4) In some cases (e.g., heptane on liquid mercury) i t is possible to measure 9 instead of I? directly, making the use of Eq. (60) especially easy (Cassel, 80; Kemball and Rideal, 93; Hill and Kemball, 94). We now derive Eqs. (59) and (60) from “adsorption thermodynamics” (83). This can be done by using the completely general approach of solution thermodynamics as a starting point, the value of which would be to emphasize that adsorption and solution thermodynamics are completely equivalent, are derivable from each other, have the same starting point, and apply to the same systems (regardless of adsorbent perturbations, swelling, etc.). However, this point of view has been stressed elsewhere (83) and we confine ourselves here, except for a few further remarks later, to the special case of an inert adsorbent, this being the case for which adsorption thermodynamics is particularly useful and natural. With an inert adsorbent it is both possible and desirable for purposes of understanding adsorption data to consider the adsorbed molecules as a one-component system (in the external field of the adsorbent). Eventually adsorbent perturbations will have to be taken care of, but this is certainly a second-order effect in almost all physical (not chemical) adsorption systems. [Cook, Pack and Oblad @a) would except the first adsorbed layer.] The system of n lmoles of adsorbed gas has an energy E,, entropy S,, volume V , and area a, as already discussed in connection with solution thermodynamics. There are thus two external variables, V , and a, although for practical purposes V , may be ignored. I[n ordinary threedimensional thermodynamics dE = TdS

- PdV

+ 4dn

where E is considered a function of S , V , and n.

We may also write

so that T, P , and p are defined, if one likes, by derivatives, for example

I n the present system there are two external variables and we have dE. = TdS.

- PdV. - +da 4-pldnl

(61)

249

THEORY OF PHYSICAL ADSORPTION

where # is deJined as

4 is the “surface pressure” and plays the same role for a one-component adsorbed phase ( V , being negligible) as P does for a one-component three-dimensional phase. Now since this is a one-component system the Gibbs free energy F , should have the property F , = n l p 1 , and we use this as our definition df F.. Integrating Eq. (61) keeping the intensive quantities T , P , 4, and

where the Helmholtz free energy A , is defined as usual by A , = E, - TS,. Then dF, = nldfil

+

(64)

pldnl

and, from Eq. (63), dF. = dE.

- TdS, - SadT

+ PdVe + VsdP 4-

+ ad4

(65)

Combining Eqs. (611, (64), and (651,

At equilibrium with gas

Putting P

=

p , we can write (warn = (sG- s , ) / ( v G - u,) (P = p ) ( a + / a p ) ~= (VQ - v,)r ( P = p )

or, approximately,

(%)+ (

%$=

(SG

- sJ/RT

RTr

=

p

On integration of Eq. (71) one obtains Eq. (59). Equation (70) is the same as Eq. (60). The entropy change SO - s8 is completely analogous to the entropy of fusion or vaporization (per mole) of a one-component solid or liquid. The question then arises as to what function for the adsorbed phase leads

250

TERRELL L. HILL

to a heat of adsorption AH related to the entropy of adsorption by A S = S. - S, = A H I T . From

we have so that the desired function is, using Eq. (63), TS.

+ F, = E. + P V , + +a =

That is, a t equilibrium, X.

- Ho

H,

+ +a c X.

= T ( S , - SQ)

In X, (Eq. 72), PV, is negligible and $6 is found from Eq. (59). By combining Eqs. (59), (60), (72), and (73) one can then find the molar energy of the adsorbed phase, E,. The quaqtities E,, s8) and p1 (from the equilibrium gas pressure and pl = p,) or E ~ s,, , and $I give an essentially complete thermodynamic description of the adsorbed phase. A comparison of the value of molar and differential entropies and energies for interpretive purposes has been given elsewhere in some detail and will be omitted here (83). It might be added, though, that in the theories of Wheeler-Ono, Barrer-Robins, and Hill, discussed in Seo. 111.6, the molar entropy and energy will drop out naturally for each p c.dculated, but the differential quantities must, be obtained indirectly. More extensive and accurate data and additional calculations are necessary to obtain s8, E,, and 4 from isotherm data over what is required to get the differential energy and entropy from the iaosteric equation. The first complete calculation of ss, E,, and 4, as well as the differential quantities, has recently been made by Hill, Emmett, and Joyner (95). This paper shows in detail how the methods of this sect ion can be applied in practice. Using heats of immersion, Harkins and Jura (96) made earlier equivalent calculations, but the relationship of their calculated quantities to the thermodynamic functions of the adsorbed molecules was not pointed out until recently by Jura and Hill (92). Jura and Hill (92) have suggested that in cases of strong binding to a surface, the location of the minimum in s, (or the associated loop in aS,/dnl) can be used as a surface area method not dependent on any particular theory. It is very interesting that the minima in s8 for nitrogen on graphon (95) and for argon on rutile (9) (using calorimetric and isotherm measurements) both give excellent agreement with the BET surface area method. Finally, we wish merely to indicate the relation between adsorption thermodynamics and solution thermodynamics [it is hardly clear from Eq. (Sl)] and incidentally to show that Eqs. (61) to (73) are valid regard-

THEORY O F PHYSICAL ADSORPTION

251

less of adsorbent perturbations, etc. For any kind of sorbent, it will be recalled that functions E,,S,, and V , were defined by Eq. (49), and that these reduce properly to the desired thermodynamic functions for a onecomponent adsorbed phase in the special case of an inert adsorbent. I n the general case let us define further @ = pod - p A and subtract Eq. (48) from Eq. (40). This gives dE. = TdS,

- PdV, - W n n + p l d n ,

174)

This equation is completely analogous to Eq. (61) so that all equations subsequent to Eq. (61) still hold, regardless of the nature of the sorbent, if 4 is replaced by 4, and a by nA. In the inert adsorbent special case, poA and p d for the adsorbent differ only by virtue of surface contributions (the bulk properties are assumed the same with and without adsorbed molecules). Hence 4J=Yb-Y

(76)

where ya is the surface tension of the pure adsorbent and y the surface tension of the adsorbent with an adsorbed film on it. 6 . Alternative Thermodynamic Treatments

For virtually all practical purposes the thermodynamic treatment above and in IX (83) is as general as necessary. For the sake of completeness, however, we make a few additional remarks in this section. For systems of the type water-sulfuric acid, benzene-rubber and hydrogen-palladium, in which interfacial effects (gas-condensed phase) are negligible compared to bulk effects, the analysis in IX is exact. But in systems where there are large interfacial effects (i.e., appreciable gas-condensed phase surface areas) there are two uncertainties, both connected with a precise definition of the volume of the condensed phase: (1) the exact location of the dividing surface between the gas phase and the condensed phase (e.g., in argon-graphite) ; and ( 2 ) the meaning and experimental measurement of the “volume” of the pure adsorbent in some cases (e.g., a highly porous charcoal). As pointed out in Appendix I of IX (83), these difficulties can be avoided by use of Gibbs’ surface thermodynamics or by considering gas plus condensed phase as the thermodynamic system. The Gibbs method can be applied to uncertainty (l),at least in some instances, while the second method is completely general and takes care of both (1) and ( 2 ) . a. The Gibbs Method. Let us apply the Gibbs method in the idealized case of an adsorbent with a precisely defined volume (when no gas is present) and, hence, surface area [i.e., to define the volume precisely a

252

TERRELL L. HILL

dividing surface (with an area) between vacuum and adsorbent must be visualized]. The surface need not be planar. The adsorbent is further assumed to be inert; the geometry and thermodynamic properties of the adsorbent are independent of pressure, temperature, and the amount of gas adsorbed. In our discussion of the inert adsorbent case in previous sections and in V (18) we implicitly chose, when gas was adsorbed, a dividing surface in the transition region between adsorbed film and gas (i.e., where the relatively high density of the film falls off rather suddenly to the gas density). This surface was not defined precisely. All molecules between the dividing surface and the surface of the adsorbent were considered as belonging to the condensed phase. Everywhere outside of the dividing surface the gas density was assumed to obtain. The treatment of this idealized system can be made ]precise by choosing a dividing surface in the Gibbs sense at the surface of the adsorbent. This is the approach used by Hansen (85). Also this method should be adopted in applying thermodynamics rigorously to statistical theories (e.g., the theories of Frenkel-Halsey-Hill, Wheeler-Ono, Barrer-Robins, and Hill-see Sec. 111.6) employing this same inert adsorbent model. Furthermore, the usual experimental procedure for obtaining the "amount of gas adsorbed" makes implicit use of this particular dividing surface and the related surface excess (85). We can ignore the inert adsorbent in all our equations here. Suppose the volume (aside from that occupied by the adsorbent) is V a ,that there are, at equilibrium, na moles in this volume, the energy is Ea, the entropy is Sa,the equilibrium pressure is p , the chemical potential is p and the surface area is a. Then we can write where The area can be changed by adding more adsorbent. Now consider also the hypothetical situation in which the gas density remains constant right up to the dividing (adsorbent) surface, retaining the same values of p , T , Va and I.L. In this case let the energy, entropy, and number of moles be Eo, So and no, respectively. Then

Subtracting Eq. (78) from Eq. (77),

THEORY OF PHYSICAL ADSORPTION

253

where

These are the usual Gibbs surface excesses. Note that there is no volume term in Eq. (79). Integrating Eq. (79) we have E' = T S d - cpaa

We define

+ pn8

An = Ea - TSa

P = pn' from which we find, for example,

+

dA' = - S a d T - Qada pdn. d p = -sadT (1/ra)dcpa

+

where s' = sd/n*,

=

na/a

Using d p = dfict and Eqs. (43) and (85), ( S O - s*)/VQ (acpa/ap)T = vctr.

(ap/dT)Va =

Eqs. (86) and (87) are exact and may be compared with Eqs. (68) and (69). Also, using azAs/dTanaand Eq. (84), we get ( a p / a T ) r 8 = - (aP/ana)T.a

Combined with dfi

=

(88)

dpa and Eq. (43), this gives (ap/aT)r"

= [SO

- (aS8/ana)a.~l/v~

(89)

which may be compared with Eys. (45) and (50). There are certain differences in detail between these results and the earlier ones, especially the definitions of ns,'pa, s8 and P. However these distinctions are in general of no practical significance in experimental applications. It should be recalled that we have restricted ourselves t o an inert adsorbent. In so doing we have been able to handle the volume of the adsorbed film rigorously, but with a considerable loss of generality. That is, we cannot include adsorbent perturbation, swelling, the transition from adsorption to solution, etc., as is done in IX (83). b. Entire Container as the Thermodynamic System. As implied above, there is one way of avoiding the various imperfections in the thermodynamic treatments discussed so far and that is to consider gas plus condensed phase as a single thermodynamic system. As a matter of fact, Guggenheim (87a) has recently analyzed this problem. The

254

TERRELL L. HILL

general approach is to employ only those thermodynamic variables that are actually measurable, regardless of the nature of the adsorption system inside the container. For example, we can measure the amount of adsorbent put in the container, the total amount of gas (“adsorbed” or not) in the container, the equilibrium temperature and gas pressure and the total volume of the container. The exchange of work (pressurevolume) and heat between the container and the outside world can also be measured. But one has to give up any attempt to divide the volume of the container into two parts: volume of gas and volume of condensed phase. This makes unnecessary, for example, the measurement of “dead space” by helium adsorption. The concepts of “amount of gas adsorbed,” “ adsorption isotherm,’’ “ surface pressure,” “ thermodynamic functions of adsorbed molecules,” etc., must, as a consequence, also be abandoned. This is the program that must be adopted to be absolutely rigorous thermodynamically, and it is certainly important that workers in the field realize it. However, in the present writer’s opinion, if this program were actually used by experimentalists, the severe price paid, in loss of contact with molecular reality inside the container, would far exceed the value of the last ounce of exactness gained. 6. Summary

We have given an incomplete discussion since details are readily available in the original papers. However, it should be clear that there is no longer any uncertainty about the relations between the two-component (adsorbent plus adsorbed molecules) point of view of solution thermodynamics, leading naturally to differential quantities, and the one-component (adsorbed molecules) point of view of adsorption thermodynamics, leading naturally to the molar quantities of more direct statistical mechanical interest. Also, the connectione between calorimetric heats and entropies of adsorption now seem to be straightened out. The “solution” or equivalent “adsorption” thermodynamics of IX (83) has complete generality* as :far as the system&included is concerned, but in many cases is not quite rigorous in the sense discussed in Sec. IV.5. Despite this slight inexactness, the writer feels that this approach is the single most useful method of treating the general problem. I n the special case of an “inert” adsorbent, the Gibbs method (Sec. IV.5.a) is rigorous and therefore more satisfactory than V(18), though from the experimental point of view the distinction between the two methods is in general undetectable. The completely rigorous approach of Sec. IV.5.6 is necessarily too general to be of much practical value, * The extension to an elastic adsorbent has been made (Hill, 97).

THEORY OF PHYSICAL ADSORPTION

255

except t o serve as a reminder of the slight imperfections in the other treatments. With these purely thermodynamic questions disposed of, one can hope that the future will bring a number of really complete experimental studies (both isotherm and calorimetric measurements, including heats of immersion and heat capacities in some cases) of the simplest possible systems-for example, argon or krypton adsorbed on nonpolar, nonporous adsorbents. Work along these lines is already in progress in the laboratories of J. A. Morrison and G. Jura. Aside from intrinsic interest, a backlog of detailed thermodynamic data of this type should prove invaluable for future theoretical attempts to understand the nature of adsorbed films.

REFERENCES 1. Brunauer, S., Adsorption of Gases and Vapors. Princeton University Press, Princeton, 1945. 2. Lennard-Jones, J. E., Proc. Roy. SOC.(London) 106A, 463 (1924); Physica 4, 941 (1937). 2a.'Cook, M. A., Pack, D. H., and Oblad, A. G., J . Chem. Phys. 19, 367 (1951). 3. London, F., 2.physik. Chem. B11, 222 (1930). 4. Hill, T. L., J . Chem. Phys. 16, 181 (1948). 5. Orr, W. J. C., Trans. Faraday SOC.35, 1247 (1939). 6. Kemball, C., Advances in Catalysis 2, 233 (1950). 7. Hill, T. L., J . Chem. Phys. 14, 441 (1946). 8. Pitzer, K . S., and Gwinn, W. D., J . Chem. Phys. 10, 428 (1942). 9. Morrison, J. A., Los, J. M., and Drain, L. E., Trans. Faraday Soe. (in press); Morrison, J. A., and Los, J. M., Faraday SOC.Discussion No. 8, 321 (1950); Morrison, J. A,, and Drain, L. E., J . Chem. Phys. 19, 1063 (1951); Drain, L. E., and Morrison, J. A., American Chemical Society Meeting, New York, LSeptember, 1951; Drain, L. E., and Morrison, J. A., to be published. 10. Kemball, C., Proc. Roy. Soc. (London) 187A,73 (1946); 190A, 117 (1947). 11. Drenan, J. W., and Hill, T. L., J . Chem. Phys. 17, 775 (1949). 12. Cassel, H. M., J . Phys. Chem. 48, 195 (1944). 13. Emmett, P. H., and Brunauer, S., J . Am. Chem. SOC.69, 1553 (1937). 14. Livingston, H. K., J . Colloid Sci. 4, 447 (1949). 15. Volmer, M., 2. physik. Chem. 116, 253 (1925). 16. Langmuir, I., Colloid Symposium Monograph 3, 72 (1925). 17. Semenoff, N., 2. physik. Chem. B7, 471 (1930). 18. Hill, T. L., J . Chem. Phys. 17, 520 (1949). 19. Lennard-Jones, J. E., and Devonshire, A. F., Proc. Roy. SOC.(London) 163A, 53 (1937); 166A, 1 (1938). 20. Devonshire, A. F., Proc. Roy. Soc. (London) 163A,132 (1937). 21. Jura, G., Loeser, E. H., Basford, P. R., and Harkins, W. D., J . Chem. Phys. 14,117 (1946). 22. Jura, G., and Criddle, D., J . Phys. & Colloid Chem. 66, 163 (1951). 23. Gregg, 5. J., in Surface Chemistry (Supplement t o Research). Butterworth's Scientific Publications, London, 1949. 24. Mayer, J. E., J . Chem. Phys. 6, 67, 74 (1937).

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Kirkwood, J. G., J . Chem. Phys. 3, 300 (1935). Langmuir, I., J . Am. Chem. SOC.40, 1361 (1918). Fowler, R. H., Proc. Cambridge Phil. SOC.31, 260 (1935). Fowler, R. H., and Guggenheim , E. A., Statistical Thermodynamics. Cambridge University Press, Cambridge, 1939. 29. Rushbrooke, G. S., Introduction to Statistical Mechanics. Oxford University Press, Oxford, 1949. 30. Miller, A. R., The Adsorption of Gases on Solids. Cambridge University Press, Cambridge, 1949. 31. Kramers, H. A., and Wannier, G. H., Phys. Rev. 60, 252, 263 (1941). 32. Onsager, Id., Phys. Rev. 66, 117 (1943). 33. Kaufmann, B., and Onsager, L., Phys. Rev. 76, 1232, 1244 (1'949). 34. Hill, T. L., J . Chem. Phys. 18,988 (1950). 35. Yang, C. N., J . Chem. Phys. 13, 66 (1945). 36. Li, Y. Y., Phys. Rev. 76,972 (1949). 37. Halsey, G., and Taylor, H. S., J..Chem. Phys. 16, 624 (1947). 38. Sips, R., J. Chem. Phys. 16, 490 (1948); 18, 1024 (1950). 39. Hill, T. L., J . Chem. Phys. 17, 762 (1949). 39a. Beebe, R. A., and Kington, G. L., personal communication (1948). 40. Tompkins, F. C., Trans. Faraday SOC.46, 569, 580 (1950). 41. Born, M., and Green, H. S., Proc. Roy. SOC.(London) 188A, 1.0 (1946); 189A, 103 (1947). 42. Wheeler, A., paper given at American Chemical Society meeting, Atlantic City, September 1949. 43. Ono, S., J. Chem. Phys. 18, 397 (1950); J. Phys. SOC.Japan 6, 232 (1950); 6, 10 (1951). 44. Kirkwood, J. G., and Buff, F. P., J . Chem. Phys. 17, 338 (1949). 45. Brunauer, S., Emmett, P. H., and Teller, E., J . Am. Chem. SOC.60, 309 (1938). 46. Harkins, W. D., and Jura, G., J . Am. Chem. SOC.66, 1362, 1.366 (1944). 47. Cassie, A. B. D., Trans. Faraday SOC.41,450 (1945). 48. Hill, T. L., J . Chem. Phys. 14,263 (1946); 17, 772 (1949). 49. MacMillan, W. G., J . Chem. Phys. 16,390 (1947). 50. Walker, W. C., and Zettlemoyer, A. C., J. Phys. & Colloid Chem. 62, 47, 58 (1948). 51. Anderson, R. B., J . Am. Chem. SOC.68, 686 (1946). 52. Cook, M. A., J . Am. Chem. SOC.70,2925 (1948); (with D. H. Pack) 71,791 (1949). 53. Dole, M., J . Chern. Phys. 16, 25 (1948). 54. Hill, T. L., J . Chem. Phys. 16,767 (1947). 55. Halsey, G., J. Chem. Phys. 16,931 (1948). 56. Teller, E., National Colloid Symposium, St. Louis, June 1950. 57. Halsey, G., National Colloid Symposium, St. Louis, June 1950. 58. Cassel, H. M., J . Chem. Phys. 12, 115 (1944); J. Phys. Cheni. 48, 195 (1944). 59. Hill, T. L., J . Chem. Phys. 14,268 (1946). 60. Arnold, J. R., J . Am. Chem. SOC.71, 104 (1949). 61. White, L., Jr., and Schneider, C. H., J . Am. Chem. SOC.71, 2593, 2945 (1949). 62. Huttig, G. F., Monatsh. 78, 177 (1948). 63. Ross, S., J . Phys. & Colloid Chem. 63,383 (1949). 64. Fergusson, R. R., and Barrer, R. M., Trans. Faraday SOC.46, 400 (1950). 65. Hill, T. L., J . Am. Chem. SOC.72, 5347 (1950). 66. Tolman, R. C., The Principles of Statietical Mechanics. Oxford university Press, Oxford, 1938, 25. 26. 27. 28.

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Huttig, G. F., and Theimer, O., Kolloid-2. 119,69, 157 (1950); 121,50, 54 (1951). Barrer, R. M., J . Chem. Soc. 1874 (1951). Frenkel, J., Kinetic Theory of Liquids. Oxford University Press, Oxford, 1946. Hill, T. L., J . Chem. Phys. 17, 590, 668 (1949). MacMillan, W. G., and Teller, E., J . Chem. Phys. 19, 25 (1951); J . Phys. & Colloid Chem. 66,17 (1951). 72. Polanyi, M., Verhandl. deut. physik. Ges. 16, 55 (1916). 73. Jura, G., and Harkins, W. D., J . Am. Chem. SOC.66, 1356 (1944). 74. MacMillan, W. G., personal communication. 75. Hill, T. L., J . Phys. Ce. Colloid Chem. 64, 1186 (1950). 75a. Everett, D. H., American Chemical Society Meeting, New York, September, 67. 68. 69. 70. 71.

1951.

Barrer, R. M., and Robins, A. B., Trans. Faraday SOC.47, 773 (1951). Hill, T. L., J . Chem. Phys. 19, 261, 1203 (1951). Hill, T. L., J . Chem. Phys. (in press); J . Phys. & Colloid Chem. (in press). Halsey, G., J. Am. Chem. SOC.73, 2693 (1951). Cassel, H. M., Physik. Z.26, 862 (1925); 28, 152 (1927); 2. Elektrochem. 37, 642 (1931); (with K. Neugebauer) J . Phys. Chem. 40, 523 (1936). 81. Coolidge, A. S., J . Am. Chem. SOC.48, 1795 (1926). 82. Rowley, H. H., and Innes, u’.B., J . Phys. Chem. 46, 537, 548, 694 (1942); 49,411

76. 77. 78. 79. 80.

(1945). 83. Hill, T. L., J . Chem. Phys. 17, 507 (1949); 18, 246 (1950); Trans. Faraday SOC. 47, 376 (1951). 84. Gorter, C. J., and Frederikse, H. P. R.,Physicu 16, 891 (1949). 85. Hansen, R. S., J . Phys. & Colloid Chem. 64, 411 (1950); 66, 1195 (1951). 86. Everett, D. H., Trans. F C I T U ~SOC. U ~ 46, 453, 942, 95’7 (1950). 87. Kington, G. L., and Aston, J. G., J. Am. Chem. SOC.73, 1929 (1951). 87a. Guggenheim, E. A., Boston University Conference on Nucleation, August, 1951. 88. Wilkins, F. J., Proc. Roy. SOC.(London) 164A, 496, 510 (1938). 89. Armbruster, M. H., and Austin, J. B., J. Am. Chem. Soc. 66, 159 (1944). 90. Pierce, C., and Smith, R. N., J . Phys. & Colloid Chem. 64, 795 (1950). 91. Ward, A. F. H., PTOC. Roy. SOC.(London) 133A, 506 (1931). 92. Jura, G., and Hill, T. L., to be published. 93. Kemball, C., and Rideal, E. K., PTOC. Roy. SOC.(London) 187A,53 (1946).

94. Hill, T. L., and Kemball, C., to be published. 95. Hill, T. L., Emmett, P. H., and Joyner, L. G., J . A m . Chem. SOC.73,5102, 5933 (1951). 96. Harkins, W. D., and Jura, G., J . Am. Chem. SOC.66, 919 (1944). 97. Hill, T. L., J. Chem. Phys. 18,791 (1950).

References of a Review Nature a. Brunauer (1). b. Miller (30).

c. Taylor, H. S., in Frontiers in Colloid Chemistry, edited by R. E. Burk and 0. Grummitt. Interscience Publishers, Inc., New York, 1950. d. Emmett, P. H., in Advances in Catalysis, Vol. I, edited by W. G. Frankenburg, V. I. Komarewsky, and E. K. Rideal. Academic Press Inc., New York, 1948. e. Seitz, F., in Advances in Catalysis, VoL 11, edited by W. G. Frankenburg, V. I. Komarewsky, and E. K. Rideal. Academic Press Inc., New York, 1950. f. DeBoer, J. H., in Advances in Colloid Science, Vol. 111, edited by H. Mark and E. J. W. Verwey. Interscience Publishers, Inc., New York, 1950.

258

TERRELL L. HILL

Glossary of More Important Symbols in Order of First Occurrence

rs

Potential of intermolecular force, as a function of intermolecular distance Potential energy of interaction of an adsorbed molecule, at a distance z from t h e adsorbent surface, with the adsorbent Boltzmann constant Vibrational frequency Number of molecules adsorbed Equilibrium gas pressure Surface area Surface concentration, N/a;or nI/nA in Sec. IV Surface or spreading pressure B E T constant N divided by the number of molecules necessary to fdl the first adsorbed layer. 6 may be greater (multilayer adsorption) or less than unity Number of localized adsorption sites in area a Chemical potential Equilibrium vapor pressure of bulk liquid Relative pressure, p l p o Statistical mechanical partition function Density in molecules/cubic centimeter Gas constant per mole Number of moles of adsorbed gas Number of moles of adsorbent Tsosterie heat of adsorption Differential heat of adsorption Isothermal heat of adsorption Adiabatic heat of adsorption Integral entropy of n1moles of adsorbate on surface S./nl, integral entropy per mole of adsorbate on surface Differential entropy of adsorbate on surface. Surface or spreading pressure using the Gibbs method. Number of moles of adsorbed gas (surface excess) using the Gibbs method. Surface concentration, n4/&,using the Gibbs method.

The Role of Surface Heterogeneity in Adsorption GEORGE D. HALSEY* Mallinckrodt Chemical Laboratory, Harvard University, Cambridge, Massachusetts

CONTENTS Page

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 259 . . . . . . . . . 259 11. Chemisorption. .... 1. The Freundlich Isotherm.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 259 2. Frankenburg’s Data.. . . . . . 3. The Isotherm at High Co 111. Physical Adsorption. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1. Criticism of the BET theory., . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 263 a. Localized Adsorption.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 263 b. Liquid-like Adsorption. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 264 2. Cooperative Adsorption on a Nonuniform Surface.. . . . . . . . . . . . . . . . . . . 265 3. Multilayer Adsorption. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 265 4. Origin of Heterogeneity.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 268 References.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 269

I. INTRODUCTION The role of surface heterogeneity in chemisorption is still a controversial one; all the data so far analyzed can be explained equally well on the basis of interaction between adsorbed molecules. Proponents of heterogeneity on the one hand assert that the interaction energies needed for this explanation are “unreasonably” large. On the other hand, the proponents of interaction think that it is “unreasonable” to expect heterogeneity on clean surfaces. Physical adsorption, however, represents the scene of a mariage de convenance because both attractive interaction and heterogeneity are needed to explain the usual isotherm.

11. CREMISORPTION 1. The Freundlich Isotherm

For approximately thirty years, the Langmuir equation has been used to explain and interpret the phenomena of adsorption. Its vogue is due more to its simplicity and easily understandable theoretical basis

* Present address: University of

Washington, Seattle. 259

260

GEORGE D. HALSEY

than t o its quantitative success. The Freundlich isotherm which preceded it is in general more successful from the empirical standpoint, but owing t o its lack of a direct derivation from simple principles it is employed rather reluctantly, if a t all. I n addition its lack of an adsorption limit a t a monolayer of adsorbed molecules makes it seem even more artificial, now t h a t it is possible t o assign definite areas t o solid adsorbents. Perhaps the simplest objection t o the unmodified Langmuir equation is the fact that a t apparent saturation of the surface, often only a small percentage of the area of the adsorbent is covered. This observation immediately suggests that the surface is a t least abruptly heterogeneous, t h a t is, there are certain sites t h a t can accommodate adsorbed gas, but t h a t these are in the minority and that most sites are less active or inactive. A second and more difficultly resolvable objection is encountered when a series of isotherms a t varying temperatures is considered. When the best Langmuir fit is made t o each of them it is often found t h a t the adsorption maxima a t “saturation” increase with decreasing temperature. I n discussing these deviations Brunauer (1) points out t h a t from the time of Langmuir’s original work, it had been realized that such deviations could be explained by a heterogeneous surface; however, it has often been forgotten t h a t such heterogeneity makes the simple Langmuir equation inapplicable. The treatment for noncooperative adsorption on a heterogeneous surface was formulated. by Fowler (2), who showed that if there are N , sites of energy x covered t o an extent Bx the total coverage is 8 =

$exNxdx

The coverage on sites of energy ex = 1/[1

+

(1)

x is given by the Langmuir equation

((l/p)eAS~/ke-~/LT)]

(2)

AS, is the entropy of adsorption on sites of energy x. I n general, ASx will be a function of x, because the frequency of vibration of a molecule bound t o the surface, which determines t o a large extenr; AS,, is presumably influenced by the strength of the bond with the surface x. A linear variation of ASx with x can be introduced without complicating the derivation which follows. Therefore we write

where ASo and r are constants. Substituting (3) and (2) in (1) yields

THE ROLE OF SURFACE HETEROGENEITY IN ADSORPTION

261

We shall integrate this expression (Halsey and Taylor, 3) first, using the exponential distribution N x -- C e - X / X m

(6)

when c and x,,, are constants, between the limits x = = to - CQ . There have been a number of criticisms (e.g., Sips, 4) of these limits, especially the lower one, which seems to indicate belief in a large number of sites with repulsive adsorption energies. Actually, the limits are a mathematical device only to facilitate integration. The physical assumption involved in using this device is that the actual distribution function on the surface is exponential somewhat beyond the weakest sites that are ever filled and the strongest sites that are ever emptied, to any measurable extent. Thus if one assumed the real limits and integrated (4) numerically, the results would be the same as if the infinite limits were assumed, within the accuracy of the experiment. Substituting (6) in (4),and writing

P"

4

e-(l-rT)x/kT

P

which for small values of the argument of the cosecant becomes =

(;)kT/xm(l-rT)

CXm

Thus, the distribution function (6) leads to the classical Freundlich isotherm, with a slope on a log-log plot, equal to d In p

=e=

xm(l - rT) ICT

and isosteric heat

where terms from differentiating log T have been omitted. If we define the units of 0 so that 0 = 1 when p / p " = 1 equation (9) becomes

-q

= xrnlog e

262

GEORGE D. HALSEY

If rT is small compared with unity, the new unit of 8 will be independent of 0, and all the isotherms, when extrapolated intersect at p = po and e = 1, where q also extrapolates to zero. It has been stated (Hill, 5) that the disadvantage of using a treatment based on equation (1) was the neglect of entropy variations over the surface. We have here remedied that defect by generalizing the treatment to include variation of the special type (3). Hill has pointed out that the frequency of vibration normal to the surface should vary with the square of the adsorption energy, t o the first approximation. The entropy of the adsorbed gas varies with the logarithm. of this frequency, so we may expect the approximate relationship &a. = const.

- 2R log x

(14)

The entropy of adsorption A S is equal t o Sgas- Sad.so A S = const.' + 2R log x 2R dAS = - d x X

This expression is not of the form (3), but in the neighborhood of a heat of adsorption x we should expect r to have the value r z -2R

(17)

X

2. Frankenburg's Data

The isotherms of Frankenburg ( 6 ) for the adsorption of hydrogen on tungsten have been criticized because the degree of cleanliness possible on the outgassed tungsten powder he used does not equal that attained with evaporated films. However, his results remain the only data extensive enough for a complete analysis. Here we revise some of the numericalidata of an earlier treatment (Halsey and Taylor, 3). The preceding argument fits his data from 200 to 800°K. with the constants xo = 5.70 kcal., r = 0.75 X per degree and p o == 3150 mm., with l/i substituted for p , to account for the atomic nature of the adsorption. If we take the values p = 100 mm. and T = 400"K., roughly in the middle of the range of values, r from Eq. (17) is per degree showing substantial agreement between Hill's argument, and the empirical value.

-

3. The Isotherm at High Coverage The Freundlich isotherm has no definite maximum adsorption or saturation value; therefore, when 0 approaches a monolayer, the Freundlich equation necessarily breaks down. For high coverages the exponential distribution function is conveniently replaced by a linear distribution

THE ROLE OF SURFACE HETEROGENEITY IN ADSORPTION

263

+

with the energy of adsorption descending from xo A x for 0 = 0 to xo for 0 = 1, with a constant number of sites in equal energy ranges between. For the case where the highest energy sites are always covered we have derived elsewhere that (Halsey and Taylor, 3) the isotherm is 1

-e

=

kT -In AX

[l

+(p~/p~)$$e-xa/~~]

(18)

the square root of pressure indicating that the isotherm is written for dissociation upon adsorption. The data of Rideal and Trapnell (7) for the adsorption of hydrogen on an evaporated tungsten film are fitted by (18), or rather by a limiting form of (18). When (1 - 0 ) A x / k T is large, equation (18) becomes 1 - e = -log kT

2Ax

polpn

-xo AX

The constants are A x = 24 kcal. and p o = 5 x lo5 mm. The isotherms do not go to high enough coverage to evaluate xo but they state it is 2 kcal. (Halsey, 8).

-

111. PHYSICAL ADSORPTION 1. Criticism of the BET Theory

a. Localized Adsorption. The multilayer adsorption isotherm of Brunauer, Emmett, and Teller (BET) (Brunauer, 1) was based on the following assumptions : 1. 2. 3. 4.

The surface possesses uniform, localized sites. The adsorption is of the noncooperative (Langmuir) type. The energy of adsorption in the first layer is El, a constant. The energy of adsorption in succeeding layers is EL, the energy of liquefaction. 5. Surface area available for the nth layer is equal to the coverage of the (n - l ) t h layer.

Assumptions 3 and 5 taken together mean that the whole energy of liquefaction is available when one isolated molecule is adsorbed on another isolated molecule, and that this energy is unchanged when the approximately eleven more contacts of the liquid phase are established. Such an assumption is obviously unreasonable. These assumptions have been replaced (Halsey, 9) by the more realistic ones that follow: 1. Unchanged. 2. Adsorption is cooperative with an attractive interaction energy of %EL (6 neighbors on the surface, 12 in the liquid; 9i.i~= %).

264

GEORQE D. HALSEY

3. The energy of adsorption of a completed layer is E l in the first layer. 4. The energy of adsorption in complete succeeding layers is EL. 5 . A site for adsorption must have a triangular array of adsorbed molecules beneath it. On the basis of these assumptions it has been shown that there is effectively no adsorption in any but the first layer, below p / p ~equal unity, and that adsorption in the first layer takes place cooperatively a t a constant pressure p l p o = exp ( E l - ELIRTI (20)

If we discard (4) and postulate adsorption energy A& greater than EL in the nth layer with El > Ez > E , > EL, a series of steps results, in contradiction to the experimental fact that the isotherm is smooth. We shall see below that discarding (1) leads to agreement with experiment. b. Liquid-like Adsorption. Substantially the same conclusions have been reached by McMillan and Teller (10). Instead of assuming localized adsorption, they approached the problem from the standpoint of the continuum. The BET theory postulates that all but the first layer of adsorbate has the energy of the liquid. Since the adsorbed molecules are piled at random on the various sites, that is, for example, five deep on one site and two deep on the next, there is a combinatorial entropy term associated with exchanging these piles, which is the physical explanation for the stability of the BET adsorption a t a pressure below the bulk condensation pressure. Such random piling disregards, however, the force of surface tension, which tends to pull down the hills and fill in the valleys, thereby lowering the combinatorid entropy and the adsorption. These hills and valleys can be thought of as surface tension waves in a continuous layer of liquid adsorbed on the surface. If the surface tension is negligible, these waves will be deep and choppy, leading to many equivalent configurations, and thus to a large entropy, approaching the BET entropy. On the other hand, if the surface tension is effectively infinite, the surface of the liquid will be stretched flat, and then will be only one configuration, and thus no combinatorial entropy. The 10 monolayers in authors have considered adsorption in the range 8 terms of the Fourier components making up the shape of the surface. Consider first the component waves of short wavelength. For an intermediate value of the surface tension these waves will possess only a small amplitude, much less than the thickness of the layer, because of the large surface area of short high waves, and the concomitant high energy. If we consider equilibrium between the bulk liquid and the (thick) surface -J

THE ROLE OF SURFACE HETEROGENEITY IN ADSORPTION

265

layer, the short wavelengths will exist in both to the same extent, and thus will not influence adsorption. Waves of longer wavelength, to which the continuum theory is quite applicable, can suffer larger amplitudes without exposing such a great surface. In a bulk liquid they may possess any amplitude but on the surface they are limited by the thickness of the layer. The increased number of configurations made possible by deepening the adsorbed layer is responsible for adsorption below condensation pressure. McMillan and Teller have solved this problem in terms of the surface tension and find that for zero surface energy the BET equation is predicted. However, when the actual surface tension is estimated the adsorption realized is negligible, corresponding closely to the flat surface of one configuration, resulting from infinite surface tension. 2. Cooperative Adsorption on a Nonuniform Surface

It is clear that on a uniform surface one would expect stepwise condensation. That this is not the case leads one to discard the assumption of uniformity, and assume that there is a distribution function of adsorption energies over the surface. If the surface is characterized by a distribution function NAE, and the sites of equal energy are collected in patches, each patch will fill at a pressure p/po = e-AE/RT. The total coverage will be NAEdAE

i.e., the sum of all sites at greater than the filling pressure. c exp { -AE/AE,} , an exponential distribution function 8 = const. X ( p / p o ) R T / A E m

If NnE= (22)

identical in form with the Freundlich isotherm (8) derived for noncooperative adsorption. The only difference is that in the noncooperative case AE,/RT > 1, whereas here, it is possible to have the positive second derivative a 2 0 / d p 2 characteristic of cooperative adsorption. 3. Multilayer Adsorption Equation (22), derived for cooperative adsorption, concerns only the monolayer. As we have seen, if only the heat of liquefaction is available beyond the first layer, no further adsorption is to be expected. We must assume therefore that a van der Waals field transmits energy to the second and higher layers. If we assume a uniform surface, with the van der Waals energy falling off with distance according to a power 1 / ~ an , isotherm equation can be derived. Hill (11) has law AE

-

266

GEORGE D. HALSEY

pointed out that for simple London forces n equals 3, and evaluating the constant of proportionality, has derived an isotherm of the form Fad..

- F~ic,.= -RT

log ( p / p o ) = const./e:'

(23)

when the adsorbed phase has the liquid entropy. The critical assumption in deriving this equation ie that the distance from the surface x: in the London law can be set proportional to 8. Clearly 8

7 6 5

1

9

P/P, FIG. 1.

below one layer, x will be constant at the closest distance of approach as e varies. Thus x is in reality a step function of 61. The continuous function (23) and the step function are plotted ih Fig. 1. The continuous and dashed lines represent two different step functions, corresponding t o two different values of the molecular diameter of the adsorbent. It is clear that the continuous approximation is far from being valid below four or five layers and has no relation to the volume of the monolayer. However, the isotherm Eq. (23) is strikingly successful in fitting adsorption data, if the power three is replaced by an adjustable constant r, with a value frequently between 2 and 3. At first it was thought that

THE ROLE OF SURFACE HETEROGENEITY IN ADSORPTION

267

nonuniformities in the surface would “smooth out ” the steps causing the empirical equation to be valid. Recent investigations have shown this not to be the case (Halsey, 12). Equation (22) was derived for cooperative adsorption in the first layer. If the source of the van der Waals energy A E is located one molecular diameter below the first layer, a weaker van der Waals field exists in the second layer. It has the same distribution function except that all energies are divided by the factor 23, because the second layer is twice as far away. Similarly, in the nth

VPo FIG. 2.

layer the energy is divided by n3. Isotherms can be derived for all the n layers and summed to give the total coverage. Some of the isotherms so calculated are presented in Fig. 2. Curve a shows the result of assuming great heterogeneity of the surface; there is no indication of a “point B ” except a pseudo “point B ” a t -0.1 monolayer, and no possibility of estimating the surface. Curve b is for a less heterogeneous surface and has the characteristic true point B; it gives a successful BET surface area. Curve c is for an almost uniform surface, and shows the steps at each layer characteristic of such a surface. The surface of intermediate heterogeneity b gives an area measurement without having

268

GEORGE D. HALSEY

pronounced steps and corresponds closely t o the usual BET nitrogen isotherm. The reference t o surface area is t o be contrasted with the lack of such a reference in the isotherm Eq. (23) plotted in Fig. I. It is important t o note that the three diverse isotherms of Fig. 2 were derived from the same third power law that led to (23). The latter part of the three isotherms are all fitted by (23) with the power 2.5 substituted for 3. Only a t coverages greater than 10 layers does the power approach three. The value 2.5 is also closer t o the experimental results than 3.0. We may conclude, then, that the third power law of van der Waals forces in conjunction with a heterogeneous surface is flexible enough t o explain any isotherm, with the possible addition of short range forces operative only in the first layer.

4. Origin of Heterogeneity Rhodin has recently presented (Rhodin, 13) some elegant experimental data that give an insight into the origin of surface heterogeneity. Thin plates were prepared from large single crystals of copper in such a manner as t o expote only a chosen crystal face. Although the plates had a surface area of only a few square centimeters, when suspended from an extremely accurate torsion balance, adsorption isotherms of nitrogen could be determined with remarkable accuracy. It was of course impossible t o maintain the standard of cleanliness possible with evaporated films; however, the plates were reducefd in hydrogen and baked out in such a manner that the results on thle three faces [llO], [lOO], and 1111) appear t o be evidence for an almost uniform surface. When isosteric heats were computed, they did not show the usual monotonic decline from high t o initial heat to AHl,q. Instead the heat curve a t low coverage was flat; increased t o a large maximu rn a t the monolayer, and then declined slowly t o AHli,. These heats have their explanation in initial noncooperative adsorption on the periodic lattice determined by the crystal face, followed by transition t o cooperative adsorption determined by the packing diameter of the nitrogen as 0 reaches unity. For the purpose of the present discussion, it is the heat curve determined on polycrystalline copper that is of particular interest, however. It shows the usual monotonic decline of heat with 0, and is by no means a simple average of the three curves on single faces, since a t low coverage it is higher than any of the three. It would appear t h a t although the same type of adsorption is taking place on the mixture of faces, the intercrystalline boundories exert enough influence t o spread out the onset of cooperative adsorption over a range of pressure and t o provide some sites of higher energy than present on any of the single faces.

THE ROLE OF SURFACE HETEROGENEITY IN ADSORPTION

269

When we recall that practically all adsorbing systems studied are polycrystalline it is no surprise that Rhodin’s results are almost unique, being most nearly approximated by isotherms on salt crystals. Furthermore, the entirely different results for the same absorbate on two crystals of the same substance purified in the same way is a strong argument for heterogeneity operating in at least one case, presumably the polycrystalline preparation. We can only hope that studies will be made on other single crystal faces, to test this conclusion. The author is indebted to Dr. T. N. Rhodin, Dr. T. L. Hill, and Dr. E. Teller for valuable discussions. REFERENCES 1. Brunauer, S., The Adsorption of Gases and Vapors. University Press, Princeton, 1943. 2. Fowler, R. H., Statistical Mechanics. University Press, Cambridge, England, 1936. 3. Halsey, G., and Taylor, H. S., J . Chem. Phys. 16, 624 (1947). 4. Sips, R., J . Chem. Phys. 16, 490 (1948). 5. Hill, T. L., J . Chem. Phys. 16, 181 (1948). 6. Frankenburg, W. G., J . Am. Chem. SOC.66, 1827 (1944). 7. Rideal, E. K., and Trapnell, B. M. W., Discussions of the Faraday SOC.8, 114 (1950). 8. Halsey, G., Trans. Faraday SOC.47,649 (1951). 9. Halsey, G., J. Chem. Phys. 16, 931 (1948). 10. McMillan, W. G., and Teller, E., J . Phys. & CoZZ. Chem. 66, 17 (1951). 11. Hill, T. L., J. Chem. Phys. 17, 590, 668 (1949). 12. Halsey, G., J. Am. Chem. SOC.73, 2693 (1951). 13. Rhodin, Jr., T. N. J . Am. Chem. SOC.72,5692 (1950).

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Twenty-five Years of Synthesis of Gasoline by Catalytic Conversion of Carbon Monoxide and Hydrogen HELMUT PICHLER Hydrocarbon Research, Inc., Trenton, N . J .

CONTENTS Page I. Introduction. ..... ..................................... 11. Historic Review of earch Work Concerning the Catalytic Con Carbon Monoxide and Hydrogen to Higher Hydrocarbons.. . . . . 1. The Work of Frans Fischer and His Co-workers.. . . . . . . . . . . . a. Fischer-Tropsch (1922-1928). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 274 (1) Work Preceding the Fischer-Tropsch Synthesis. . . . . . . . . . . . . . . . . 274 (2) The Discovery of the Normal-Pressure Synthesis.. . . . . . . . . . . . . . 274 b. Fischer-Meycr-Koch-Roelen (1928-1934). . . . . . . . . . . .

(4) Pilot Plant Experiments.

.......

c. Fischer-Pichler (1935-1937). . . . . . . . . . . . . . . . (1) Synthesis in Several Steps with Intermed

. . . . . . . . . . . . . . 281 (1) Research and Development Work on Medium-Pressure Synthesis 286 with Iron Catalysts.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . (2) High-pressure Synthesis of High-Melting Paraffins with Ruthe289 nium Catalysts. ............................................ (3) High-pressure Synthesis of Branched Hydrocarbons (Iso-Synthesis) in the Presence of Thorium Oxide and Other Oxide Catalysts.. . . . . . . . . . . . . . . . . . 2. Development Work Outside o a. France.. . . . . . . . . . . . . . . . . b. Germany.. . . . . . . . . . . . . . . (1) Ruhrchemie-Modificat Technical Purposes. ... (2) Ruhrchemie-Lurgi-Re (3) Schwarsheide-Experiments with Iron Catalysts (Comparison of Different Precipitation Catalysts with a N H a Type Fused Iron Catalyst) ....................... 271

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HELMUT PICHLER

Page (4) I.G. Farbenindustrie. Rheinpreussen. and Others.- Modification of

Medium-Pressure Synthesis with Iron Catalysts (hot gas recycle. oil slurry. oil recycle) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . (5) Synthesis of Alcohols and Other Oxygen-Containhg Organic Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . (6) Basic Research Work . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . c Great Britain . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

.

................................................ .............................................. Early Research Work., . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Agreement with Ruhrchemie . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Synthesis with Fluidized Catalyst. . . . . . . . . . . . . . . . . . . . . . . . . . . .

(1) (2) (3) (4) Synthesis with Oil Recycle in the Presence of Iroin Catalysts of the Synthetic Ammonia T y p e . , . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . (5) Basic Research Work . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . f . Other Countries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111. Solved and Unsolved Problems of Hydrocarbon Synthesis . . . . . . . . . . . . . . . . 1 Summary of Relations between Catalyst, Synthesis Con.ditions and Reaction Products . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2. Poisoning of Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 . Hydrocarbon Synthesis and Water Gas Equilibrium . . . . . . . . . . . . . . . . . . 4. Pro and Con Carbide Theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . a . Chemical Investigations. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . h . Ortho-Para Hydrogen Equilibrium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . c . C ~ as C Auxiliary Substance . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . d . Thermomagnetic Studies . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . e. X-ray Diffraction Studies . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . f . Thermodynamics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

.

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

h . General Considerations. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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

305 307 308 309 310 310 310 310 310 314 316 318 319 319 321 327 331 332 332 333 333 334 335 335 335 337

I . INTRODUCTION A quarter of a century has passed since Franz Fischer and Hans Tropsch of the Kaiser Wilhelm Institut fur Kohlenforschung a t Miilheim Ruhr made the important discovery that carbon monoxide and hydrogen react in the presence of active iron and cobalt cata1;ysts a t atmospheric pressure. producing higher gaseous. liquid. and solid aliphatic hydrocarbons (July. 1925). The first publication was issued a t the beginning of 1926 (1) . The title of this paper was. “Synthesirci of Petroleum from Gasification Products of Coal a t Normal Pressure.” The new reaction very rapidly became an object of greatest interest to scientists engaged in catalytic problems as well as t o experts concerned with a possible shortage of natural petroleum sources. The magazine Fuel in Science and Practice stated (2). “Nothing is more surprising. considering the enormous amount of work that has been done on the chemistry of carbon com-

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

273

pounds, than that there should still be so much to learn about as simple a compound as carbon monoxide.” Between 1925 and 1950 an amazing amount of diligence and money were devoted toward the chemical and technical development of the Fischer-Tropsch process. Many countries participated in this work. The first technical scale plant was erected by Ruhrchemie A.G. (Oberhausen, Germany) in 1935. The erection of eight other plants in Germany followed at short intervals. Factories were also built in France, Japan, and Manchuria. The United States, Great Britain, and other countries began the erection of pilot plants. The largest commercial plant for the production of liquid fuels from carbon monoxide and hydrogen, designed by Hydrocarbon Research, Inc., for Carthage Hydrocol, Inc., is now in operation in Brownsville, Texas. A great many papers were published during the past years on ‘(the synthesis.” These papers consist partly of excellent reviews on the basis of other publications and reports and partly of valuable contributions t o the new technical development of the process and the explanation of the reaction mechanism (3). It is not the object of this paper to summarize again all the literature. It would also be beyond the scope to discuss in detail industrial and engineering features of hydrocarbon synthesis, as well as questions which are only indirectly connected with the synthesis such as synthesis gas manufacture or conversion of primary synthesis products to other compounds. The scope of this paper inevitably limits the review t o important steps of research work on the reactions of carbon monoxide and hydrogen, with special consideration regarding the behavior of the catalysts necessary to carry out the reactions.

11. HISTORIC REVIEWOF RESEARCH WORKCONCERNING THE CATALYTIC CONVERSION OF CARBON MONOXIDE AND HYDROGEN TO HIGHER HYDROCARBONS 1 . The Work of Franz Fischer and H i s Co- Workers The successful development of the synthesis of liquid fuels was the result of the cooperation of many scientists. Franz Fischer, however, was the spiritual center of this work. He was of the opinion that this synthesis would be of greatest importance not only t o Germany, with only a few natural oil sources, but also to the whole world (4). Publications issued by people unfamiliar with the work carried out at Miilheim were often in error concerning the authors of the different syntheses of Fischer and his co-workers. Some papers presented new names for one and the same process merely on the basis of unimportant

274

HELMUT PICHLER

changes, while other papers mentioned the same inventor for all processes converting carbon monoxide and hydrogen to hydrocarbons. Fischer did not agree with such interpretations. He confirmed this many times, the last time, a short time before his death in 1947, in a statement which he sent to the author of this article. This document clarifies the authorship of the different processes, as did Fischer’s farewell address delivered in 1943 at the Mulheim Institute for Coal Research ( 5 ) . a. Fischer-Tropsch (2926-2968). (1) Work preceding the FischerTropsch synthesis. Four important publications precede the discovery of the normal-pressure synthesis: A. The publication of the classic reaction of carbon monoxide chemistry, the synthesis of methane in the presence of nickel and cobalt catalysts by P. Sabatier and J. B. Senderens in 1902. B. Patent applications issued by Badische AniIin und Sodafabrik in 1913, which claim the production of a mixture of hydrocarbons and oxygenated organic compounds from carbon monoxide and hydrogen at high pressures (100 atm. and more). As possible catalysts B.A.S.F. mentioned a wide variety of different metals such as :nickel, cobalt, iron, manganese, chromium, titanium, osmium, cerium, :molybdenum, zinc, palladium. Thermodynamic calculations were the baais for this work (6). Many years later, some of the metals tested in these patent applications turned out to be useful as catalysts, others not. C. A paper concerning the first important work by Fischer and Tropsch on catalytic hydrogenation of carbon monoxide. It was the synthesis of Synthol (1922-1923). Alkalized iron filings were used as catalysts for the conversion of water gas at pressures above 100 atm. and temperatures of about 400°C. The reaction product was a mixture of different oxygenated compounds (alcohols, aldehydes, acids, etc.) . D. Patents of Pier and Winkler (7) (1923) for controlled catalytic conversion of carbon monoxide and hydrogen to methanol exclusively (ZnO-Cr20s catalyst). (2) T h e discovery of the normal-pressure synthesis. I n 1923, Fischer and Tropsch produced the first measurable quantities of higher hydrocarbons by catalytic reduction of carbon monoxide. By recycling carbon monoxide and hydrogen over alkalized iron filings at 410°C. they were able to vary the ratio of oxygen-containing molecules t o pure hydrocarbons by varying the pressure. I n the range of 115 atm. the products were almost entirely oxygen-containing compounds. However, as the pressure was decreased to 7 atm., hydrocarbons appeared in increasing, but still very small amounts. In July 1925, Fischer, Tropsch, and Ter-Nedden, were successful in producing measurable quantities of higher hydrocarbons by the

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

275

catalytic reduction of carbon monoxide at atmospheric pressure and 250-300°C. in the presence of Fe-ZnO and Co-Crz03 catalysts. The first publication of Fischer and Tropsch (spring, 1926) had already contained a great many facts which were of importance for the later development: A. Iron, cobalt, and nickel are the catalysts for the hydrogenation of carbon monoxide to hydrocarbons. Cobalt has the greatest tendency to produce aliphatic hydrocarbons with more than one carbon atom per molecule. Iron was less active and nickel showed too high a hydrogenation activity favoring the formation of methane. B. Oxides difficult to reduce, such as ZnO and Cr203,promote the conversions and lower the tendency of the catalysts to sinter. C. The presence of small amounts of alkali favors the formation of larger hydrocarbon molecules. D. Addition of certain amounts of copper are useful in the case of iron catalysts (better reduction at low temperatures), but dangerous in the case of nickel catalysts (formation of alloys). E. The synthesis gas must be free of sulfur. This can be achieved by catalytic conversion of the sulfur of organic sulfur compounds to hydrogen sulfide which can be removed by known methods. ( 3 ) Basic research work. I n connection with a n exhaustive research program, Fischer, Tropsch, and Ter Nedden (8) also carried out some experiments with iron catalysts at a pressure of 10-15 atm. (250-280°C.). They obtained aqueous reaction products and nonaqueous oil in the ratio 3 :1 to 1:1, and paraffins with melting points up to 117°C. However, the catalysts used at that time were not suitable for the production of hydrocarbons at increased pressures. The comparatively low activity of the catalysts declined very rapidly as the result of the adsorption of high melting waxes and possibly also of oxygenated products. Tropsch (9) stated: “If the process is to be carried out at increased pressures, it is necessary to increase the temperature in order to make it possible that the reaction products leave the catalyst. In this case the reaction products are oxygenated products, probably formed by a secondary reaction.” Before Tropsch left Miilheim (1928) in order to become director of a coal research institute in Czechoslovakia, Fischer and Tropsch summarized their findings on hydrocarbon synthesis (10) : The yields of liquid products in one throughput amounted t o 20-25 C C . / M . ~ carbon monoxide-hydrogen mixture with iron (Fe:Cu = 4 : l 0.2% KzC03)and cobalt (Co:Cu = 9: 1, no alkali) as catalysts. Thirty grams of oxides were used for the conversion of 4 liters of synthesis gas per hour. Figure 1 presents results of experiments on the influence of alkali on

+

276

HELMUT PICHLER

iron catalysts (gas contraction as measure of conversion-about 35 % contraction corresponds to 100% CO conversion). The curves show the great influence of alkali on the behavior of the Fe catalysts. 0.5-1.0% K&Os yield the best results. Smaller amounts of alkali produce less active catalysts, and higher amounts yield catalysts with a shorter lifetime. Sodium, lithium, ammonia, and alkali earth salts have less pronounced influence. Rubidium acts similarly to potassium. Even today, these results are, in principle, still valid for all kinds of iron catalysts. The presence of alkali is not favorable in the case of Co catalyst. An Cncrease of the molecular weight of the reaction products in this case is connected with a rapid decline of catalyst activity.

o oez

HRS

K ~ C O ~

I -

OPERATION

FIQ.1. Influence of alkali on behavior of iron catalyst.

The best catalysts were produced by thermal decomposition of nitrates in the presence of porous carriers. Fischer and Tropsch carried out experiments with precipitated catalysts too. However, the precipitation of cobalt was performed with hydroxides (NaOH, KOH, etc.) which produce catalysts of very low activity. Unknown up to that time was the outstanding behavior of kieselguhr as carrier of the catalyst, and the fact that only carbonates can be used for the production of active precipitation catalysts. Fischer and Tropsch assumed an intermediate formation of carbides (carbide theory) as mechanism of the reaction of carbon monoxide and hydrogen to higher hydrocarbons. Methane was assumed to be formed via an intermediate formation of hydrides. The competition of carbon monoxide and hydrogen in connection with the formation of carbides and hydrides was considered to be responsible for the tendency of different, catalysts to form preferentially either higher hydrocarbons or methane. With this theory, Fischer and Tropsch explained why iron presented art

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

277

unfavorable catalyst for methane synthesis, while nickel could not be used for synthesis of higher hydrocarbons. In this connection, cobalt ranked between iron and nickel. The conversion of carbon monoxide and hydrogen proceeds in two different ways, according t o the equations:

The first conversion was preferred in the reaction on cobalt catalysts, the second in the reaction on iron catalysts. b. Fischer-Meyer-Koch-Roelen (1928-195.4). The work performed by Fischer and Tropsch resulted in the discovery of the reaction known today as “normal-pressure synthesis,” or “ Fischer-Tropsch reaction.” Experiments by Fischer with Meyer, Koch, and Roelen turned the scientifically interesting conversion into a technically important process (1928-1934). (1) Nickel ~ r e c i ~ i t u ~ icatalysts. an Fischer and Meyer (11) carried out exhaustive research work trying to develop active Ni catalysts for the synthesis of higher hydrocarbons. In spite of the failure of all previous attempts, Fischer and Meyer found two types of nickel catalysts with a behavior similar to that of the best cobalt catalysts. The first type was a nickel-thorium-kieselguhr precipitation catalyst and the second type a nickel-manganese-aluminum-kieselguhr precipitation catalyst. Hundreds of experiments were necessary to find the optimum conditions for manufacturing these catalysts. The first encouraging results were obtained with a catalyst produced by thermal decomposition of nickel manganese nitrates in the presence of kieselguhr. Fifteen cubic centimeters of oil per cubic meter of synthesis gas were obtained. The amount of reaction products could be increased to 43 cc. of liquid hydrocarbons by precipitation of nickelmanganese mixtures with sodium carbonate in the presence of kieselguhr. Using potassium carbonate instead of sodium carbonate as agent of the precipitation improved the yields to 93 cc. liquid hydrocarbons per cubic meter of gas. After finding the optimum conditions for preparation, washing, and reduction, the yields went up to 120 CG. Addition of small amounts of aluminum nitrate to the nitrates of nickel and manganese finally resulted in a further improvement of the yields to 130 cc. of liquid hydrocarbons (C, hydrocarbons). The best nickel-thorium oxide catalyst (100 Ni: 18Th02: 100 kieselguhr) was also precipitated with potassium carbonate and showed

+

+

278

HELMUT PICHLER

similar behavior as the nickel-manganese-aluminum-kieselguhrcatalyst. The optimum reduction was obtained with a great surplus of dry hydrogen a t a temperature of 450°C. The stepwise increase of catalyst activity made it possible to lower the synthesis temperature from 250°C. t o 175°C. Thirty per cent of the synthetic hydrocarbons consisted of methane a t 175"C., compared with 90% a t 250°C. (2) Cobalt precipitation catalysts. The development of cobalt catalysts (Fischer and Koch, 12) was similar to the development of the corresponding nickel catalysts. In the case of cobalt,, however, it was easier t o prevent extensive methane formation. The 1OOCo:18Th02: 100 kieselguhr catalyst became the so-called standard cobalt catalyst. The catalyst was prepared as follows: Two solutions were made up: Solution I: 1000 g. of cobalt and 180 g. of thorium oxide a s the nitrates were dissolved in 20 liters of distilled water. Solution 2: 2100 g. of sodium carbonate were dissolved in 30 liters of distilled water. Shortly before using, 1000 g. of kieselguhr were added t o the solution. (On a large scale, precipitation was carried out with 103 g. of sodium carbonate per liter of solution, and 1,he kieselguhr was added with stirring immediately after precipitation.) Both solutions were heated to boiling, and Solution 2 (containing the kieselguhr) was poured as rapidly as possible into Solution 1. This was then heated t o boiling for one minute and filtered. The catalyst cake was washed six to eight times with a total of 20 liters of boiling distilled water (until alkali free), dried a t 11O"C., and granulated. The reduction of the catalyst was effected in the laboratory by passing 20 liters of dry hydrogen per hour per 10 g. of cobalt over catalyst beds 30 cm. deep. At a temperature of 365°C. it was found that reduction could be stopped after 4% to 5 hours. Particular1.y active catalysts were obtained (high degree of reduction) when the reduction was continued for 20 hours. Minimum reduction temperatures were necessary in order t o prevent sintering. Increasing contents of kieselguhr required increasing reduction temperatures. Addition of small amounts of copper lower th e necessary reduction temperature, but lower th e lifetime of th e catalyst as well. Increasing amounts of thorium increase the average molecular weight of the reaction products. Optimum yields (153 cc. of liquid hydrocarbons per cubic meter of synthesis gas) were obtained a t a synthesis temperature of 190°C. I n order t o maintain the activity of th e catalyst it is necessary t o increase t he temperature slowly (10-15°C. in 6 weeks). One of the reasons for the decline of the catalyst activity is the adsorption of high molecular products. One of the disadvantages of th e normal pressure

GASOLINE SYNTHESIS FROM CARBON MONOXIDE A N D HYDROGEN

279

synthesis is the fact t h a t it is necessary to “clean” the catalyst surface from time t o time by extraction with a solvent or by treatment with hydrogen. I n the first case it is possible t o obtain the high melting paraffins, and in the second case hydrocracking occurs with methane formation. (3) Raney type skeleton catalysts (13). I n the case of the nickelkieselguhr and cobalt-kieselguhr precipitation catalysts, the accepted theory was t h a t the high activity of these catalysts was connected with special surface conditions on the voluminous kieselguhr. No other carrier could replace the kieselguhr in any case of precipitation catalysts. It seemed t o be very interesting therefore, t h a t a Raney type nickelcobalt catalyst produced only 10-20 % less liquid hydrocarbons than the best cobalt-kieselguhr catalysts. I n the case of the skeleton catalyst the volume per weight of metal was only one-tenth of that of the precipitation catalysts. The best skeleton catalyst was produced from a nickel-cobaltsilicon alloy (1 : 1 :2) by dissolving the silicon with sodium hydroxide. (4) Pilot plant experiments. Pilot plant operations (Roelen, 14) were carried out with a gas throughput of 5-10 m.3/hr. (compared with 4 l./hr. in the laboratory scale). Two important problems had t o be solved in particular in this operation: (a) the purification of the synthesis gas from sulfur compounds and (b) the removal of the exothermic heat. Extensive removal of sulfur from the reacting gases was necessary in order t o obtain economic lifetime of the catalysts. With increasing catalyst activity it was necessary t o increase the purity of the synthesis gas. The limit for industrial purposes was 0.2 g. sulfur per 100 m.3. A smaller amount of sulfur was more favorable. The removal of HzS was performed in the normal manner with iron oxide (Luxmasse). The organic sulfur was converted in a second step in the presence of alkaliaed iron oxide (200-300°C.) t o hydrogen sulfide, which reacts with iron forming iron sulfide (15). Quantitative removal of the reaction heat which is necessary t o keep the catalyst temperature constant, was for a long time one of the most important problems. About 20% of the combustion heat of the carbon monoxide-hydrogen mixture is released during the reaction. A temperature increase of only a few degrees lowers the yields of useful products. Greater temperature increases result in extensive formation of methane and also of carbon. Various procedures were developed t o meet this problem. Although this paper is not intended t o discuss thoroughly the technical details or the advantages of different types of reactors, it seems of importance t o mention that heat removal on a semitechnical scale was decisive in limiting the amount of gas conversion per reactor space and time (space-time yield). The first solution satisfactory for the purposes

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HELMUT PICHLER

of a pilot plant was the use of small catalyst layers (8-.12 mm.) between metal surfaces cooled with circulating oil of constant temperature. (5) Composition of synthetic products. In 1935 Fischer published (16) data showing the average composition of the synthetic lxoducts obtained with cobalt catalysts (Table I). The composition of the reaction prodTABLE I Reaction Products o n Cobalt Catalysts “C. VOl. % Boiling Range Amount in Wt. % Olefins

+

Gasol (CS C,) Gasoline Oil Paraffin in the oil Paraffin, obtained by extraction of the catalyst

<30 30-200 >200 -

23 7

50 30 10 -

4

-

4 62

ucts changes with many factors. For example: A. Increasing catalyst age and increasing temperatures give lower molecular products. B. Increasing amounts of thorium oxide, increasing amounts of unreduced cobalt, and traces of alkali increase the percentage of higher hydrocarbons. C. Increasing amounts of carbon monoxide and decreasing amounts of hydrogen increase the percentage of olefins (mono-olefins). D. The percentage of olefins decreases with increasing molecular weight of the synthetic hydrocarbons. E. The reaction products consist mainly of straight chain hydrocarbons and smaller amounts of oxygenated products. The amount of “oxygenates ” (mainly alcohols) changes with the conditions of the reaction and the catalyst composition. I n the case of cobalt catalysts, the amount of oxygenated reaction products is very low (less than 1 %). The raw gasoline has a comparatively low octane number, which depends on the boiling range and the olefin content (see Fig. 2). The synthetic oil shows, with its cetene number of 105, an excellent behavior in the diesel engine. c. Fischer-Pichler (1935-1937). (1) Synthesis in several steps with intermediate removal of liquid reaction products. I n 1934 Fischer and Pichler (17) began investigations directed towards the increase of yields of useful synthesis products by changing the physical performance of the process. They found it unfavorable, for yields of higher hydrocarbons, to force too high carbon monoxide conversions in one step. The carbon

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+

monoxide hydrogen partial pressure changes, in a single pass operation, from the gas inlet to the outlet from almost 1 atm. to a very small fraction of this pressure. In view of the fact that specific optimum conditions correspond to any given partial pressure, of CO/H2, a one-step operation (which must operate at average conditions) cannot yield optimum results. A separation of the process into two, three, or more steps, with removal of the reaction products or at least of the reaction products condensable a t room temperature (hydrocarbons and water) resulted in a 10-20% increase in yields of liquid products and decreased the formation of methane (without increasing the amount of catalyst). The intermediate removal of water (in the case of iron catalysts and carbon dioxide) also lessens the danger of catalyst oxidation at high

-

CC

Pb (‘22 H5)4 /LIT. GASOLINE

FIQ.2. Octane numbers of synthetic gasoline (cobalt catalyst).

carbon monoxide-hydrogen conversions (increase of catalyst life). Experiments with gas recycle and intermediate removal of condensable reaction products were carried out too. The results of these recycle experiments, however, were not favorable enough (in comparison with the two- or three-stage operations) to justify the expensive recycle operation of the 1-atm. gas. It was of great advantage t o use less active catalyst (“old catalyst”) in the first stage and highly active catalyst (“fresh reduced catalyst”) in the following stages. ( 2 ) Medium-pressure synthesis with cobalt catalysts. Earlier publications (18) indicated that the synthesis of hydrocarbons ought to be carried out at atmospheric pressure. The formation of high molecular products seemed to be the reason for rapid inactivation of the catalysts in the case of superatmospheric experiments. The decline of catalyst activity could not be compensated by an increase of the temperature. At higher temperatures, the conversion of carbon monoxide and hydrogen

282

HELMUT PICHLER

yielded preferably oxygenated products. These results were obtained with catalysts (prepared by decomposition of different salts) used in fixed beds as well as with catalysts suspended in oil (oil slurry system). Fischer and Pichler resumed such experiments at varied pressures during the winter of 1935-1936. They used cobalt-thorium-kieselguhr precipitation catalysts for the synthesis at increased pressures arid found a very interesting behavior of the catalysts at pressures of 5-20 atm. (medium pressure) : A. The yields of solid paraffins (obtained with the liquid products) increased tenfold per cubic meter of carbon monoxide-!hydrogen mixture

40r 'V 6o

50

t

\

I

-5 30

10

J

O1

2

3

4

5

678910

20

d 40

60

80 100

ATMOSPHERES

FIG.3. Products of hydrocarbon synthesis a t different pressures (Co catalysts).

(compared to the yield in reactors with the same heat transfer surfaces). B. The total yield of Ca+ hydrocarbons increased by 20% (same reactor conditions for both systems). C. The periodic treatment of the cobalt catalyst, necessary in the case of the normal-pressure synthesis, was not necessary at the conditions of medium-pressure synthesis in spite of the higher molecular weight of the reaction products. The life of the catalyst (without regeneration) was increased many times. D. The reaction products (using 2Ht: 1CO as synthesis gas) are more saturated than in the case of the normal-pressure synthesis. This is an

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

283

advantage for the production of paraffins and diesel oil, but a disadvantage for the use of primary gasoline as liquid fuel. Figure 3 shows yields (g./m.3) of paraffin, diesel oil, gasoline and gaseous hydrocarbons (CO-ThO,-Kg = 100 : 18 : 100, 1 stage, average of H z / l g. Co/hr., 4-week periods, no intermediate regenerations, 1 1. CO reactors with the same heat transfer surfaces) as a function of the pressure of the synthesis gas. The yields of solid paraffins show a maximum a t pressures of 10-20 atm. The amounts of gas and gasoline decrease with

+

!

I

I

0

I

100

200

300

I0

VOL. SYNTHESIS GASIVOL. CATALYST I H R .

FIG.4. Yields of C I S hydrocarbons (at optimum temperatures).

increasing pressures. The diesel oil yields show only comparatively small changes. A t pressures above 50 atm. the formation of volatile carbonyls, accompanied by a decrease of catalyst activity was observed. Figure 4 present yields of C3+ hydrocarbons in grams per cubic meter of synthesis gas and gram per hour (using the same amount of catalyst). The space velocity (volume of gas throughput per volume of catalyst) varied between 25 and 500. I n every case a temperature was maintained a t which the largest production of C3+ hydroc,arbons took place. This optimum temperature was 175°C. a t a space velocity of 25,

284

HELMUT PICHLER

and 210°C. a t a space velocity of 200. The two curves of the figure show a decrease of the yields per cubic meter of gas with increasing space velocity and a n increase of the yields per hour per catalyst volume. These two curves are similar for all types of carbon monoxide hydrogenation catalysts. They are important for explaining the different views concerning the optimum operation conditions tor the synthesis. If the price of the synthesis gas determines the suitability of the process, working a t comparatively low space velocities will be inevitable. If, however, the capital costs of reactor and catalyst are the more important items, it will be more favorable to work a t high space velocities, if the problem of t he removal of th e exothermic reaction heat can be solved. (3) Medium-pressure synthesis with iron catalgsts. Up to January, 1935, the maximum yields of Cs+ hydrocarbons obtained with iron catalysts a t atmospheric pressure were 30-40 g./m.3 synthesis gas. The decline of catalyst activity amounted to 20% within 8 days (19). Fischer and Meyer (20) improved the yields of the normal-pressure synthesis with iron catalysts (in 1934-1936) t o 50-60 g./m.3 synthesis gas and the lifetime of the catalyst from 8 days to about 30 days. 'These results were obtained with iron-copper precipitation catalysts (1 a,tm., 230-240°C.). The decline of catalyst activity was closely connected with changes of the composition of the reaction products. The color of the synthetic products changed from white t o yellow and formation of fatty acids and organic iron salts was detected. Increased carbon monoxide content of the synthesis gas and increased alkali content of the catalyst accelerated this phenomenon. Meyer and Bahr (21) carried out similar experimenlcs with iron-kieselguhr catalysts. The type of catalyst support, which was very important in the case of cobalt and nickel catalysts, showed no pronounced influence in the case of iron catalysts. Iron catalysts produced from ferro salts and kieselguhr were not active, while iron catalysts produced from ferri salts and kieselguhr yielded comparatively good results. In connection with their work on medium-pressure synthesis with cobalt catalysts, Fischer and Pichler began (1936-1937) some research work with iron catalysts too. The first positive results, comparable with results obtained with cobalt catalysts were achieved when a n iron precipitation catalyst that had been in use a t atmospheric pressure for several weeks was switched to operation at a synthesis gas pressure of 15 atm. The initial pretreatment of the catalyst a t 1 atm. proved t o be necessary for the successful use of the catalyst a t higher pressure. The combination of proper pretreatment (reduction and carbonization) followed by synthesis a t elevated pressures increased the yields of Cs+ hydrocarbons t o more than double and the lifetime of the catalyst was

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

285

increased many fold (2la). Unreduced catalysts which were used immediately under pressure never reached the same activity as catalysts which were pretreated several hours at atmospheric pressure with synthesis gas. Fischer reported these results in the fall of 1937 to the German industry, and Pichler gave detailed data in a talk on “Medium-Pressure Synthesis with Iron Catalysts” on September 10, 1940 (22). The alkali content of the iron catalyst is particularly important as will be seen from Table 11. The amounts of alkali indicated represent TABLE I1 Yields Obtained with Iron Catalyst Containing Different Amounts of Alkali (Single Stage”)

Precipitant NH3 Na2C03 NazC03 NazCOa NazCOa NanCOa NanCOs Na2C03 Na2COa NazCOs NazC03 a

Alkali Addition %

Solid, Liquid Gasol Hydrocarb. Liquid g./NM3 of Paraffins Hydrocarbons Gasol Ideal Gas % % % 141 141 148 157 155 158 163 154 143 161 155

12 13 26 42 41 43 46 38 43 45 46

67 67 56 47 45 41 42 52 44 43 44

21 20 18 11 14 16 12 10 13 12 10

Table 6 , H. Pichler. Lecture, MBlheim/Ruhr, Sept. 10, 1940 (Reel 101. Doc. 21574-NID).

the weight per cent of potassium carbonate based on metallic iron. Where other alkali salts (for example, permanganate and fluoride) are indicated, the notation “1% KMnOr” means the amount equivalent t o 1 % of K2COX. The table shows the nature and quantity of the products formed in the presence of various alkalized precipitation catalysts during operation at 235°C. and 15 atm. with a synthesis gas of three parts of carbon monoxide and two parts of hydrogen. During the first month, the yields in grams per normal cubic meter of ideal gas fluctuated between 140 and 160 g. It was apparent that the addition of alkali was not necessary for the preparation of an active catalyst (in contrast t o the normal-pressure synthesis). The yields obtained over a long period of time with a catalyst prepared by precipitation with ammonia and containing no

286

HELMUT PICHLER

alkali were not significantly lower than those from a n iron catalyst containing %% of K2C03. On the other hand, the alkali content plays a n important part in controlling the type of products obtained. In the absence of alkali, the solid paraffin fraction, determined by the butanone method and based on the total yield of solid, liquid, ansd gasol hydrocarbons, amounted to 12%. Addition of 1, 2, and 5 % of potassium carbonate, respectively, increased the yields obtained t o 26, 42, 43, and 45-46%. As the alkali content increases, a corresponding decrease is observed in the amounts of liquid and gasol hydrocarbon produced. Table I1 also shows that during the first week of operation the amount and type of products are not influenced by the nature of the potassium salt added t o the catalyst. Experiments with 1% potassium carbonate and those using catalysts containing equivalent amounts of potassium permanganate, silicate, fluoride, and phosphate, all gave the same results, within the limits of experimental error. Addition of copper seems t o have an important effect on the behavior of iron catalysts. Just as in the case of cobalt, the a,ddition of copper lowers the reduction temperature for iron. However, th e effect of copper on the life of the iron catalyst is not as adverse as on cobalt. d. Fischer-Pichler (1938-19.45). (1) Research and development work o n medium-pressure synthesis with iron catalysts. The work on mediumpressure synthesis with iron catalysts was continued with W. Dienst, H. Merkel, K. Ruckensteimer, and F. Weinrotter. For shorter periods H. Buffleb, W. Lohmar, A. Meusel, and K. Ziesecke also participated in this work. The experiments were carried out on a laboratory scale and in a small pilot plant (maximum gas throughput 10 m.3/hr.). Many types of catalyst preparations and promoters were tested a t different synthesis conditions. There is a wide range of possible synthesis conditions (in contra.st t o the Co catalyst). The iron catalysts can be divided into the four following groups. A. Precipitated catalysts can be used with advantage when highly active iron catalysts, are desired for satisfactory yieldis of hydrocarbons a t low temperatures (200-240°C.). These catalystc, are usually prepared from a dilute solution of the iron salt by precipitation with a solution of sodium carbonate. Ferric solutions ordinariliy give hard glassy catalysts while ferrous solutions give less compact catdysts of a n earthy consistency. Ferrous-ferric mixtures yielded particularly satisfactory results. At the Mulheim Institute for Coal Research, iron precipitation catalysts were subjected t o a conditioning treatment with a carbon

x,

GASOLINE

SYNTHESIS

FROM CARBON MONOXIDE

AND HYDROGEN

287

monoxide containing gas, such as water gas. Active iron catalysts, containing little (1% or less) or no copper were also obtained by pretreatment with carbon monoxide at ?$o atm. and 325°C. With catalysts containing more copper, best results were obtained by pretreatment with synthesis gas at 220-240°C. and atmospheric pressure. The following description refers t o an iron precipitation catalyst of outstanding activity as developed by Pichler and Weinrotter, for example, of the composition: 75Fe+2 - 25Fe+3 - 20Cu - 1K2C03

+

+

Solution 1 contained 134 g. FeC1z.4Hz0 61 g. FeCI3.6HzO 26 g. CuC12.2H20 in 2 liters of H2O. Solution 2 contained 200 g. Na2C03dissolved in 2 liters HzO. TABLE 111 Experiments with Active Iron Precipitation Catalysts Synthesis Pressure Temperature, "C. Space velocity Yields g./m.3 (1 stage, no recycle) Product distribution Wt. % CH4 CrC4 Liquid Products 180°C. Liquid Products and paraffins 180°C. Extraction necessary a t intervals of

1 Atm.

21 1 50 135 2.4 10.1 13.5 74.0 2-3 days

10 Atm. 200 100 137 3.2 9.0 10.0 77.8 N o extraction necessary '

Solution 1 was heated to 70°C. and Solution 2 to boiling. Precipitation was effected by adding Solution 2 to Solution 1 as rapidly as possible until a slightly alkaline reaction persists. The precipitate was washed free of alkali, then alkalized with 0.5 g. K2C03dissolved in 100 cc. H20. The mixture was evaporated, dried at 105"C., and granulated. The catalyst was pretreated at atmospheric pressure with synthesis gas (2H2:1CO) at a rate of 500 liters gas per liter of catalyst and a temperature of 225°C. After 48 hours it was necessary to extract paraffin collected in the catalyst. Pretreatment was continued, under the same conditions for another 24 hours. After a second extraction, the catalyst was ready for use in normal-pressure and medium-pressure synthesis. A corresponding iron catalyst with lower copper and alkali content (50Fe+2-50Fe+3;1% Cu, %% KzCOa), prepared under the same conditions as the catalyst with 20% copper and 1% KzC03, produced a lower paraffin yield for the same degree of conversion and the same total yield

288

HELMUT PICHLER

a t atmospheric pressure. This catalyst was extracted once a week. At medium pressure, no extraction was necessary. B. Decomposition catalysts may be used when it is not essential that the synthesis be carried out at the lowest possible temperature. These catalysts, prepared by thermal decomposition of various iron salts mixed with appropriate promoters, have a slightly lower activity than corresponding precipitated catalysts. Decomposition catalycrtswere obtained, in the absence of special carriers, as a powder, suitable for liquid phase suspension units (oil slurry system). C. Fused NH3 synthesis type iron catalysts were used for hydrocarbon synthesis for the first time by I.G. Farbenindustrie [see Sec. II.2.b(4)] and by American investigators (U.S. Patents and U.S.Bureau of Mines publications [Sec. II.2.e(4)]. In some respects the fused catalysts are reminiscent of the Raney type skeleton catalysts. In the latter the porosity and lattice rearrangement necessary for catalytic activity is obtained by removing aluminum or silicon, whereas in the case of fused catalysts this rearrangement is brought about by “extracting” the oxygen during the reduction of Fe304. I n connection with the fused iron catalysts developed by the I.G. Farbenindustrie, a Russian patent (Troitzkii, 23) is worth mentioning. According to this patent, fused catalysts not only of iron but also of nickel and cobalt should be suitable as hydrogenation catalysts. D. Pretreated iron ores were successfully used for hydrocarbon synthesis at higher temperatures during the American development of the process [see Sec. II.2.e(3)]. Pichler a n d Merkel (24) investigated the composition of iron catalysts at various stages of pretreatment and synthesis by chemical and thermomagnetic analysis. Copper-free iron catalysts, carburized at 325°C. before medium-pressure synthesis, were virtually completely transformed to a ferromagnetic higher iron carbide with a Curie point of 265”C., whose formula corresponded to approximately Fe&. Carburization of copper promoted (20% Cu) iron catalysts at 220230°C. yielded not only the 265°C. Curie point carbide, but also a second ferromagnetic iron carbide whose Curie point was 380°C. At 205°C. iron was carburized largely to the 380°C. Curie point carbide. Equilibrium conditions were established between carbide and oxide. The higher carbides formed during carburization were very stable against hydrogen at synthesis conditions. At low temperatures it was possible t o produce carbides and to carry out the synthesis without formation of free carbon. Copper and alkali increase the velocity of carbide formation.

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

289

(2) High-pressure synthesis of high-melting parafins with ruthenium catalysts (25). The synthesis of high-melting paraffins in the presence of ruthenium catalysts was not the result of systematic experiments on hydrocarbon synthesis. This synthesis was developed in connection with attempts to synthesize carbohydrates from carbon monoxide and hydrogen. In order to have a chance of success the experiments were carried out at the lowest possible temperatures, highest possible pressures, and long contact times (batch operation). Nickel and cobalt could not be used because of the formation of carbonyls, iron was not active enough. The probability of success seemed t o be small because ruthenium was not active a t atmospheric pressures and temperatures below 250°C. Working in a n autoclave, however, rapid gas conversion (pressure drop) was observed at a temperature of 140°C. and pressures above 100 atm. The autoclave was cooled down (after the pressure had reached a constant level and refilled several times with a fresh carbon monoxide-hydrogen mixture. The same pressure drop was observed every time. After opening the pressure vessel, a large amount of a solid white product was observed, partly soluble in boiling benzene. It showed a melting range of 120-132°C. Analyses showed that the product was pure paraffin, with higher molecular weights and higher melting points than known before. Systematic experiments with streaming gases demonstrated (Fig. 5) that for the formation of solid paraffin (below 220°C.) high pressures are indispensable. At 300°C. high CO conversion occurs at atmospheric pressure, too, but only methane is formed. The lowest possible temperature (at high pressures) is very close to the melting point of the highest melting parts of the synthetic paraffin (134°C.). * Table IV shows the influence of the pressure at a temperature of 180°C. on CO conversion and on the distribution of the converted carbon TABLE IV Synthesis of Hydrocarbons on Ruthenium Catalysts at 180°C. and Different Pressures % Converted CO t o

CO Pressure Atm. 1

50 100 1000

Conversion, %

Paraffin

Liquid Hydrocarbons

Gaseous Hydrocarbons

0 48 68 92

46 53 59

33 31 26

21 16 15

290

HELMUT PICHLER

'

80% co Conversiorl

cn

W

[L

w

1

a cn

50% CO Conversion

0

I

H I-

a

20% CO Conversion 10

---

c-increasing amounts of high-melting hydrocarbons

I

I

I

140

160

increasing umounts of low-boiling hydrocarbons

200

180

220

240

TEMPERATURE, "C. FIG.5. Synthesis of paraffin on Ru catalysts. 1 liter product gas per hour.)

(Experiments with 3 g. Itri and

TABLE V Average Composition of Oil-Free Parafin (Extraction at Different Temperatures) Fraction 100 Atm.

1000 Atm.

Solvent

Extraction Temp., "C.

1 2 3 4

Benzene Benzene Residue Toluene

110

1 2 3 4 5

n-Pentane n-Hexane Gasoline fraction n-Heptane Gasoline fraction

34 68 90 98 121

20 80

% Dissolved 34 53 13 Traces on catalyrit 30-33 14-17 14-1 6 20-:!5 12-15

Melting Point, "C. 50-60

113 126-128 132 51-57 {33-95 121-123 130 132-134

GASOLINE

SYNTHESIS

FROM CARBON MONOXIDE

AND HYDROGEN

291

monoxide t o solid, liquid, and gaseous hydrocarbons. The oil-free paraffin was separated into fractions by extraction a t different temperatures (Table V). Paraffins with melting points of 120°C. and higher were not previously known. The paraffin obtained a t a synthesis pressure of 100 atm. contained l0-20% of those products, th e “1000 atm. paraffin” 50-55%. The average molecular weight of the paraffin of fraction 5 of the 1000 atm. experiment was 23,000 and the melting range of the products 132-134°C. The highest-melting paraffins show high viscosity above the melting point (35,600 centistokes a t 150°C.). A particularly satisfactory ruthenium catalyst was prepared as follows. Commercial ruthenium powder was fused with a mixture of potassium hydroxide and potassium nitrate (1 part ruthenium, 10 parts potassium hydroxide, 1 part potassium nitrate) preferably in a silver crucible and stirred with a silver spatula. Fusion was complete after 1 t o 2 hours. After cooling, the fused mass was dissolved in water; a deep red solution of potassium ruthenate resulted, which was heated t o boiling. Methyl alcohol was added dropwise t o th e boiling solution. The reduction of potassium ruthenate t o ruthenium dioxide began with the addition of the first drops and went rapidly to completion. The precipitate settled after a few hours. It was washed on a fritted glass plate, first with water acidified with nitric acid and then with distilled water. Finally the catalyst was dried at 110°C. The reduction t o metal proceeds just as smoothly under synthesis conditions as b y a hydrogen treatment, which latter is usually required with catalysts of the iron group. Ruthenium catalysts show a very particular behavior in many connections. No promoters and no carriers could be found which improve the behavior of the ruthenium catalyst. Working with pure carbon monoxide-hydrogen mixtures, the catalyst activity does not change for exceptionally long periods (an experiment a t 195°C. and 100 atm. during 6 months of operation showed no change in amount and distribution of reaction products). This outstanding constancy of catalyst activity was not t o be expected because the carbon monoxide pressures and temperatures used in these experiments are favorable for the formation of the volatile Ru-carbonyl (RU(CO)~). The type of products was very similar whether the catalyst was used in a dry fixed bed, or suspended in neutral liquids, or in concentrated alkali solution or in dilute acids (26). Traces of sulfur compounds however caused a rapid decline of catalyst activity. Ruthenium catalyzes th e conversion of carbon dioxide and hydrogen (to methane) at much lower temperatures (below 100°C.) than the conversion of carbon monoxide and hydrogen. The presence of carbon monoxide, however, prevents a conversion of carbon dioxide (27).

292

HELMUT PICHLER

On the basis of the results obtained with ruthenium, other platinum metals were also tested as catalysts for high pressure carbon monoxide hydrogenation. In contrast to the early experiments a t atmospheric pressure, rhodium and osmium gave appreciable amounts of liquid and solid products when operated under pressure, although the paraffin yields were smaller than for ruthenium. The extent of carbon monoxide conversion with rhodium was comparable to that of ruthenium. However, in addition to liquid and gaseous hydrocarbons, rhodium produced greater amounts of oxygenated compounds. Osmium showed catalytic activity only above 220°C. In this case, also, liquid and solid hydrocarbons were obtained but as was to be expected at the higher reaction temperature, a considerable quantity of gaseous hydrocarbons was produced. Platinum is much less suitable as a catalyst for synthesis of higher hydrocarbons. Palladium and iridium gave oiily traces of reaction products. (3) High-pressure synthesis of branched hydrocarbons (iso-synthesis) i n the presence of thorium oxide and other oxide catalysts (28). Experiments a t high pressures and comparatively high temperatures were initiated in order to develop a method to synthesize benzene from carbon monoxide and hydrogen. The evaluation of these experiments included in each case a low-temperature distillation if gaseous hydrocarbons were obtained. When water gas was passed through an empty steel tube at a pressure of 30 atm. and a temperature of 450"C., carbon monoxide and hydrogen were converted t o carbon, carbon dioxide, and low molecular weight hydrocarbons, especially methane. The results of a low-temperature distillation of the C&4 fraction are represented in the distillation curve I in Fig. 6. No characteristic break is to be seen at the boiling point of isobutane (-11.8"C.). In addition to n-butane and n-butene, small amounts of isobutene were identified in the fraction between - 10 and 0°C. Catalysts prepared from metals of the iron and platinum group of the periodic system convert carbon monoxide and hydrogen a t temperatures above 400°C. to methane, carbon, and other undesirable products. Similar negative results were obtained with molybdenum and tungsten. When carbon monoxide and hydrogen were treated at 400-450°C. in a copper-lined reaction tube in the presence of a methanol synthesis catalyst, small amounts of hydrocarbons were formed in addition to oxygenated products. Curve I1 (Fig. 6) shows the distillation curve for the C3-C, fraction obtained with a zinc-copper catalyst. The Cd fraction contained very small amounts of isobutane (0.2 g. per cubic meter of synthesis gas). Curve I11 represents the distillation curve for the G-C4 fraction of a

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

293

hydrocarbon mixture obtained with a precipitated aluminum oxide catalyst (water gas as synthesis gas, 150 atm. 450°C.). The portion of the curve corresponding to the fraction that distills between -15" and -5"C., indicates the presence of isobutane. However, the quantity was small. Completely different results were obtained when thorium oxide was used as catalyst. Curve IV shows that with this catalyst, isobutane *I0

,

I

0 VOLUME %

FIG. 6. Low-temperature distillation of CS-G fractions from different catalysts.

represents one of the principal products. This new observation was the basis of the iso-synthesis. In addition to isobutane, liquid branchedchain hydrocarbons were produced. Table VI presents the influence of pressure on the synthesis with pure thorium oxide as catalyst (450°C. 10 1. product gas per 28 g. ThOn). Increasing pressure increased the carbon monoxide conversion from about zero at atmospheric pressure t o almost 100% at 1000 atm. The influence of pressure was not very great on the distribution of the products between 30-300 atm. At 600 atm. dimethyl ether, in addition to hydrocarbons was formed. The formation of dimethyl ether increased with the pressure. If the temperature was lowered to 400°C. or below, increasing amounts of the liquid products consisted of isobutanol, and other oxygenated products. Increasing temperatures resulted in increasing amounts of

294

HELMUT PICHLER

TABLE VI Iso-Synthesis with Tho2 as Catalyst (Influence of Pressure at a Temperature of l+5OoC.) Pressure Total Yield -~ Atm. g./m.3 Liquid product

Traces 30. I 79.2 109.1 153.7 200.0

1 30 150 300 600 1000

52.5 56.0 56.5 34.5 19.5

Wt. % of Total Products i-C, C3 nC4 CI -tC1 Dimethyl ether

+

17.3 20.7 20.6 29.3 16.5

16.9 10.4 12.4 14.7 6.0

~

__

-

13.3 12.9 101.5 18.2 33’.5

0 0 3.3 24.5

naphthenes. At temperatures of 475°C. and higher, aromatics could be obtained as well. The “standard thorium catalyst ’’ was prepared by precipitation with sodium carbonate. The catalyst was pretreated by sintering the precipitate in a stream of air at 300°C. The catalyst had a very long life (particularly when regenerated from time to time with air). It is not affected by sulfur compounds. A large number of unpromoted oxides were studied t o determine their activity as catalysts for the reaction between carbon monoxide and hydrogen. The oxides of thorium, aluminum, chromium, titanium, beryllium, zirconium, uranium, zinc, manganese, magnesium, cerium, and lanthanum were investigated. Of these, thorium oxide, zirconium oxide, and to a small extent cerium oxide were active a s catalysts for the iso-synthesis. Aluminum oxide is not satisfactory as a single component catalyst for the synthesis of iso-hydrocarbons, but when added t o thorium oxide its effect on t he activity of that compound is striking. A considerable increase in iso-C4 yields was obtained with Th02-A1203catalysts. Experiments in which 20% of aluminum oxide was added to thorium oxide produced the best results. The yields were particularly good, especially of iso-C4 hydrocarbons, when the thorium and aluminum oxides were precipitated separately and thoroughly mixed after washing (61 g. iso-C4 34 g. liquid hydrocarbons per cubic meter of synthesis gas). However, when thorium and aluminum oxides were used in successive layers no improvement in yields was observed over those obtained with thorium oxide alone. When 3% of potassium carbonate (based on Al203) was added to the aluminum precipitate before mixing with the thorium precipitate, a further appreciable increase in iso-C4 formation (85 g./m.3) was observed. If, on the other hand, the same amount of alkali (0.6% K&O3 based on ThOz) was added to the thorium precipitate

+

100-

90-

80

-

E 70-

8 a a

60-

*

(L

0 I

2

XI-

c)

aJ b 40 -

9

s? i

9

50-

20

-

10

m 4

U

0t l = ~ o t m dhroffinr Singla banCh.6 P a o f h s m = D o u b h b r m c h e d POmNms S i b on$ mulli-tumched Pomffinr DIIIID=NooMhenes

eza-

m:

FIG.7. Composition of liquid hydrocarbons obtained with different catalysts of iso-synthesis.

296

HELMUT PICHLER

before mixing with the aluminum compound, the promoting action of aluminum oxide diminished. The presence of zinc oxide in thorium catalysts increased the amount of liquid products. The search for a possible substitute for thorium oxide led to the discovery that a two-component catdyst, composed of aluminum and zinc oxides, was capable of catalyzing the formation of higher hydrocarbons as well as that of isobutane. However, the spacetime yields obtained with the catalysts of this type are lower than those obtained with the corresponding Tho2 catalysts. Three-component catalysts, thorium, zinc, and aluminum oxides, gave results similar to those obtained with Thoz-ZnO catalysts (in the absence of A1203). Aluminum and chromium oxides used as two-component catalysts showed little activity. However, it is interesting to note that catalysts of this type, operated at 500"C.,produced liquid hydrocarbons composed almost exclusively of aromatics. Pichler, Ziesecke, and Titzenthaler (29) carried out exhaustive investigations on the influence of the catalyst composition on the composition of liquid hydrocarbon produced by the iso-synthesis. Figure 7 shows a comparison of the composition of liquid products obtained with four typical catalysts at optimum conditions (fractionation of the liquid hydrocarbons after hydrogenation). 2. Development Work Outside of the Research Institute at Mulheim

a. France. Fischer's interest in catalytic reactions goes back to 1900-1901 when he was a co-worker of Moissan in Faris. During this same period Sabatier carried out basic research work. on catalytic conversions such as the catalytic hydrogenation and polymerization of unsaturated hydrocarbons, followed a year later by the classic synthesis of methane from carbon monoxide and hydrogen in the presence of nickel catalysts. Sabatier and Senderens (30) found that carbon monoxide and hydrogen can be converted at temperatures between 180' and 250°C.according to the equation: CO

+ 3Hz = CH, + HzO

At higher temperatures the conversion should proceed according t o the above authors as follows: 2co =

CO1

c + COe

+ 4Hz = CHI + 2Hz0

The velocity of the carbon monoxide decomposition increased rapidly with increasing temperatures.

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

297

Sabatier proposed to use his reaction for the production of town gas with high heating value from carbon monoxide and hydrogen (31). It is difficult to understand why Sabatier and Senderens were unable to observe a synthesis for higher hydrocarbons from carbon monoxide and hydrogen. Sabatier and Senderens, in 1902 [as well as FischerTropsch (1925-1928) and Fischer and Meyer (193l)] used temperatures of 180°C. and higher, and atmospheric pressure for the conversion of carbon monoxide and hydrogen on nickel catalysts. As discussed in Sec. II.l.b(l), however, the method of catalyst preparation is decisive for the course of the synthesis. In most of his experiments Sabatier (32) used nickel catalysts prepared by reduction of pure nickel oxide at a temperature of 700°C. Such a catalyst has high hydrogenation activity and is able to catalyze many kinds of hydrogenations, e.g., the conversion of benzene to cyclohexane or the conversions of carbon monoxide and carbon dioxide to methane. However, it is not suitable for building up a carbon-carbon linkage as is necessary for the synthesis of higher hydrocarbons. Patart published (1921-1926) some papers on synthesis of methanol and mixtures of other oxygenated compounds and hydrocarbons from carbon monoxide and hydrogen at high pressures. The conditions are similar to the synthesis of synthol (Fischer-Tropsch) and of methanol (Badische Anilin and Sodafabrik, 33). Audibert and Raineau (34) carried out experiments on the FischerTropsch reaction with iron catalysts. Similar studies were also performed by Decarrihre and Antheaume (35). Lefebvre and LeClerc (36) carried out thermodynamic studies on catalysts of the Fischer-Tropsch synthesis. They drew attention t o the significance of the specific Curie points, of various compounds, for their activity or inactivity as catalysts. They assumed that the active catalysts were cubic iron oxide ” and “hexagonal nickel.” Pichler and Merkel* (37) found that the Curie point attributed by Lefebvre and LeClerc to “cubic iron oxide” is actually the Curie point of one special form of Fe&. The hexagonal nickel seems to be actually a nickel carbide. During the past few years Prettre (38) reported on thermodynamic studies on hydrocarbon synthesis. He also reported on the observation that methane could enter into the conversion. Trambouze (39) et al. carried out investigations on the activity of nickel catalysts. Teichner and Pernoux (40)found that the efficiency of kieselguhr important for Fischer precipitation catalysts is proportional to the quantity of reactive silica. Examinations with electron microscope

* See Sec. III.(4).

298

KELMUT PICHLER

show that untreated kieselguhr contains pores of about 0.5 p diameter. Acid treatment fills the smaller pores with silicic acid gel. Sodium carbonate solutions remove the gel and enlarge the pores,. The Societ6 Courrihres-Kuhlmann erected a Fisoher-Tropsch plant in 1937. This plant was operated under license from1 Ruhrchemie A.G. b. Germany. I n 1934 the pilot plant experiments of the Coal Research Institute in Mulheim yielded results encouraging enough t o justify the erection of a bigger plant. The Ruhrchemie A.G. , a company owned by a number of mining interests of the Ruhr district, purchased general licensing rights from Studien und Verwertungsgesellschaft (owner of the patents of Kaiser Wilhelm Institut fur Kohlenforschung, MulheimRuhr) and erected a semitechnical plant and shortly thereafter a fullscale plant (41). The contract between Studien und Verwertungsgesellschaft and Ruhrchemie required that all further improvements in the field of normalpressure synthesis, invented a t the Kaiser Wilhelm Institut, had t o be made known t o Ruhrchemie. A separate agreement later extended these conditions t o the medium-pressure synthesis. Eight other German companies received operating rights between 1935 and 1939: Gewerkschaft Victor (Castrop-Rauxel) Braunkohlen-Benzin A.G. (Schwarzheide) Gewerkschaft Rheinpreussen (Mom) Friedrich Krupp A.G. (Wanne-Eickel)

Essener Steinkohlen A.G. (Bergkamen) Hoesch A.G. (Dortmund) Wintershall A.G. (Luetzkendorf) Schaffgotsch G.rr1.b.H. (Odertal)

The total capacity of these plants (normal- and medium-pressure synthesis) reached about 600,000 tons synthetic products per year. Ruhrchem<e also made agreements on synthesis with Lurgi A.G. and T.G. Farbenindustrie A.G. and several companies outside of Germany. The German synthesis plants used the following main operations : A. Water gas manufacture (from solid fuels or solid fuels plus coke oven gas). B. Gas purification (removal of water, dust, and hydrogen sulfide, conversion of organic sulfur in the presence of highly alkaliaed iron catalysts). C. Partial carbon monoxide conversion t o hydrogen and carbon dioxide (production of synthesis gas with 2H2: 1CO by partial conversion of carbon monoxide with steam in the presence of Fe.-Crz03catalysts). D. Synthesis of hydrocarbons in the presence of cobalt catalysts a t 1 or 7 atm. (thin layers of “fixed bed” catalysts between heat transfer surfaces, cooled with circulating water boiling a t synthesis temperature).

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

299

E. Condensation of liquid products and recovery of gasoline and Cs- C4-hydrocarbons from product gas with active charcoal. F. Fractionation and refining of synthesis products. The synthesis factories founded research laboratories and pilot plants in order to investigate special questions concerning the process, the testing and development of catalysts, and the use of the synthetic hydrocarbons as raw materials for other processes (production of lubricating oils from olefins, fatty acids and fats from paraffins, high octane gasoline by cracking, polymerization, isomerization, aromatization, and alkylation, etc.). These refining and related processes were also the subject of many thorough investigations by Fischer and Koch (42) a t the institute at Mulheim. Koch and his co-workers have been continuing their work to the present time. (1) Ruhrchemie-modification of the cobalt standard catalyst for technical purposes. After many investigations Ruhrchemie (Roelen) changed the cobalt standard catalyst lOOCo : 18Th02: 100 kieselguhr t o lOOCo :5Th02:7.5MgO :200 kieselguhr. This catalyst did not increase the laboratory yields. However, it seemed advantageous to save a part of the thorium without serious decline of catalyst activity. The magnesium oxide increased the mechanical strength of the catalyst. Th‘e lower the content of thorium oxide, the greater the tendency of the catalyst t o form lower-boiling products. The higher kieselguhr content caused a greater “dilution” of the cobalt and therewith a lower heat production per volume catalyst per hour. This made it easier to carry out the process on a technical scale without extensive formation of methane and its gaseous homologues. Figure 8 presents yields of hydrocarbons obtained in two-stage technical operations with a (100Co :5Th02:7.5Mg0 :200Kg) catalyst a t 1 atm. and a t 7 atm. A comparison of these figures with the curves of Fig. 3 (lOOCo : 18Th02 : 1OOKg) shows a lower tendency of the MgO containing catalyst t o form solid paraffins a t the conditions of the mediumpressure synthesis. This comparatively lower paraffin formation, however, is partly due to the fact that the medium-pressure converters of Ruhrchemie contain less heat transfer surface than the normal-pressure converters. The fact that it is possible to carry out the medium-pressure synthesis with less heat transfer surface without increasing the methane formation presents another advantage of the medium-pressure operation. The quality of the kieselguhr (diatomaceous earth) is of importance for the behavior of the cobalt catalyst. Kieselguhr of different origin resulted in catalysts of different activity. Ruhrchemie therefore investigated the reasons for the different behavior and issued standards for kieselguhr. The preliminary thermal treatment should be carried out

300

HELMUT PICHLER

at lowest possible temperatures (400-700°C.), in order to prevent extensive sintering. The density in grams per liter of kieselguhr was 90-120. Reduction of the catalysts on a large scale is governed by the same considerations as in the laboratory. I n order t o avoid sintering, which would impair catalytic activity, the reduction should take place a t as low g/m3

200.

180.

C I -c2

160.

140.

120

cn L3

4 100 * 80

6 0.

40.

20

I ATM. :

The C,-$ond

PARAFFINS

63;

7 ATM. OLEFINS

>GIB fractions include small amounts of cilefins

FIQ.8. Typical yields with Co-Th02-Mg0-kieselguhr catalyst at 1 atm. and 7 atm.

a temperature and as rapidly as possible. The reduction process is facilitated by a rapid flow of hydrogen, containing as small an amount of oxygen compounds as possible (exhaustive removal of water formed, short catalyst layers). In industrial synthesis apparatus, 100% reduction gives too great an initial activity and shortens the life of the catalyst. Hence, for technical purposes, the degree of reduction was kept a t about

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

301

60%. The remaining oxide is useful in preventing further sintering. The reduction temperature, set a t 360-370°C. for the 1OOCo: 100Kg catalyst, must be increased for the 1OOCo :200Kg catalyst to about 390-400°C. (2) Ruhrchemie-Lurgi-recycle experiments. The production of synthetic hydrocarbons with a high content of olefins (mono-olefins) is desirable in many cases. The olefins increase the octane number of the gasoline and can be used as valuable raw material for many processes (polymer gasoline, lubricating oils, oxo-synthesis, etc.). It is possible to increase the olefin content of the synthetic hydrocarbons by increasing the carbon monoxide content of the synthesis gas. This step cannot be taken in the case of the normal-pressure synthesis, because of its deleterious influence on the activity of the cobalt catalyst. High carbon monoxide content of the synthesis gas, however, can be used with advantage for the medium-pressure synthesis. I n the presence of cobalt catalysts at 7 atm., synthesis gas with a hydrogen-carbon monoxide ratio of 2 :1 yields a synthetic gasoline with 20% olefins, while water gas with 1.2Hz:lCO will be converted to gasoline with 40% olefins. A gas with 1H2:2CO produces a gasoline with almost 70% olefins. Such a gas could be fed continuously to the catalyst by mixing (recycling) two t o three parts of product gas to one part of fresh water gas. In order to keep the amount of fresh synthesis gas per unit of catalyst constant, the velocity of the gas passing the catalyst would have to be increased with increasing " recycle-fresh feed ratio." Figure 9 however, shows that the increase in olefin production depends neither on the recycle process nor on the increased gas rate, but on the ratio carbon monoxide to hydrogen of the synthesis gas entering the catalyst bed. The different points of the figure correspond t o experiments with single pass and different recycle ratios, carried out at different temperatures and different gas throughputs. The curve presents the olefin content as a function of the Hz:CO ratio. The recycle process (separation of the conversion in several steps) contributes to the uniformity of the operations and, consequently to an increase in the life of the catalyst. Recycle operation is also advantageous for the medium-pressure synthesis with iron catalysts. The advantages of the gas recycle are not the same as with cobalt catalysts due t o the fact that iron catalysts have a greater tendency to form carbon dioxide in addition to hydrocarbons (instead of water) and to produce highly unsaturated hydrocarbons also in single pass operation. I n the case of iron catalysts the gas recycle increases the tendency t o produce water and changes the consumption ratio of hydrogen-carbon monoxide, and has a favorable influence on the

302

HELMUT PICHLER

stability of the catalyst activity and the yields of synthetic products per cubic meter of synthesis gas. with iron catalysts (comparison of (3) Schwarzheide-experiments different precipitation catalysts with a NHI type fused iron catalyst). In view of the favorable results obtained by the Coal Reeiearch Institute in Miilheim (FischeraPichler, 1937), industrial compani'es started experiments on medium-pressure synthesis with iron catalysts.

0.2

I

I

10

20

x)

40

50

60

713

80

% OLEFINS

FIG.9. Influence of HrCO ratio on olefin content of synthetic hydrocarbons.

In 1943 it seemed preferable t o replace the cobalt catalysts of the medium-pressure synthesis with iron catalysts without changing the converters. Comparative experiments were carried out by different groups in order to find out the status of the development work. The experiments were performed on the basis of the following rules: A. Standard water gas to be used as synthesis gas. B. Synthesis temperature must not exceed 225°C. C. Synthesis pressure-10 atm. D. Length of experiments-3 months, without regeneration of catalysts. E. The synthesis products must be as similar :as possible t o those obtained with cobalt catalysts (demands of the market).

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

303

F. The experiments were to be carried out in a single stage without recycling of the gas. The converters consisted of water-cooled tubes, 4.5 m. long, with a catalyst layer thickness of 10 mm. Catalyst volumes of 4800 cc. were used for each experiment. Converter 1 was operated by Kaiser Wilhelm Institut fur Kohlenforschung, Mulheim-Ruhr, Converter 2 by Lurgi, GeseYschaft fur Warmetechnik, Frankfurt-Main, Converter 3 by Braunkohlenbenzin A.G. Werk, Schwarzheide, Converter 4 by I.G. Farbenindustrie A.G., Ludwigshafen, Converter 5 by Ruhrchemie A.G., Oberhausen-Holten, and Converter 6 by Treibstoffwerk, Rheinpreussen, Homberg-Niederrhein. The following Fe catalysts were used for the tests: Converter 1 100Fe: 1Cu:0.75KzC08 (no carrier)

The catalyst was precipitated from a warm nitrate solution (60-70°), containing ferrous and ferric ions in a ratio of approximately 1: 1 and 1 % copper, by a boiling solution of sodium carbonate. After precipitation, the slurry was heated briefly to lOO"C., filtered, and washed free of alkali. The catalyst was then reslurried in water and an aqueous solution of potassium carbonate (0.75%, based on iron) was carefully added with stirring. The mixture was evaporated to a paste on the water bath, and dried at 110°C. for 24 hours. The dry catalyst was crushed to granules of 2-4 mm. in size and pretreated with water gas at 325°C. and atm. (4 1. gas at standard conditions per 10 g. of iron per hour).

x.i~

Converter 2 100Fe: lOCu: 30K2Si03 (silicic acid as carrier)

The catalyst was prepared by precipitation and reduced with hydrogen at a temperature of 250-300°C. The metallic iron content amounted t o about 30%.

Converter 3 1OOFe: lOCu :lOZn:0.5K&03

(no carrier)

Precipitation catalyst with synthesis gas pretreatment at atmospheric pressure and 230°C.

Converter 4

+

100Fe:2(Alz03 CaO):1.OK&03

(no carrier)

The catalyst was prepared by fusion in a current of oxygen. Careful reduction is of particular importance in the case of fused catalysts (high velocity of dry hydrogen, 450-500°C.).

304

HELMUT PICHLER

Converter 5 lOOFe :5Cu :50 kieselguhr :0.5-2.0 % K2COa (kieselguhr as carrier)

Usual precipitation with sodium carbonate. gen.

Reduction with hydro-

Converter 6 1OOFe:5Cu: 0.5-1 % K&Os

(dolomite as carrier)

Precipitation catalyst, pretreated a t atmospheric. pressure, first with hydrogen a t 300-400°C. and then with water gas a t 245°C. 2

I

6 1081 5 224

1

39 1069 220

i

5

69 1107

18 9 1150

25 1038

225

4

222

6 21 I040 224

-3 0 48 -2 6

.

Rmlfh.

.Ol.fl".

1r n m e n m

FIG.10. Comparison of the behavior of different Fe catalysts.

All six converters were operated for 3 months a t the required conditions. Figure 10 shows the results in metric tons per 10 .m.3catalyst per day (technical units). It is of little interest t o discuss small differences of yields. Proper conditions could bring the CI+ hydrocarbons in all cases t o satisfactory yields. I n some cases comparatively small change,s in the alkali content could change the ratio of lower-boiling produclts t o higher-boiling products. The following conclusions from these tests seem to be of general significance : A. The total yields per volume of catalyst (not per weight of catalyst) were of the same magnitude in spite of the use of very different types of

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

305

iron catalysts (fused iron catalyst with the apparent density 2.3 against precipitated catalyst with and without carriers with apparent densities between 0.44 and 1.37). B. The synthetic products were composed of paraffins, olefins, and oxygenated products. The amount of these components varied over wide limits. In all cases the olefin content of the product decreased with increasing molecular weight of the hydrocarbons, while the oxygenates seemed to have a maximum in the low range (C,) and a second maximum in the diesel oil fraction. C. No promoters are known (except alkali) which are indispensable for the activities of iron catalysts. In contrast to cobalt, supports are not decisive for the results obtained. D. Iron catalysts for medium-pressure require pretreatment with reducing gases. A very cautious reduction of fused iron catalysts with pure dry hydrogen was necessary while active catalysts could be pretreated with synthesis gas only. (4) I.G. Farbenindustrie, Rheinpreussen, and others-modification of medium-pressure synthesis with iron catalysts (hot-gas recycle, oil slurry, oil recycle). Hot gas recycle. As already discussed, recycle operations (with removal of liquid reaction products after each step) with a ratio of recycle gas to fresh feed of 1: 1 td 5 :1 change the consumption ratio of hydrogen to carbon monoxide, increase in some cases the olefin content of the products and the total yields, and result in a smoother operation which can be important for the lifetime of the catalysts. The purpose of the so-called hot gas recycle process (without intermediate removal of products) with a ratio of recycle gas to fresh feed of about 100 : 1 was to remove the exothermic reaction heat (700 kcal. per cubic meter of synthesis gas) from the catalyst bed to heat exchangers outside of the reactor. The temperature increase within the catalyst was limited to 10°C. Michael (I.G. Farbenindustrie) carried out these experiments. The hot gas recycle process made it necessary t o use iron catalysts of adequate mechanical strength. Sintered catalysts showed better resistance against the erosive influence of fast moving gases than highly active precipitation catalysts. Table VII presents some results of Michael’s experiments. The method used for heat removal made it difficult and expensive t o prevent extensive carbon and methane formation. Oil slurry. Numerous attempts have been made to use a catalyst suspended in liquid media for synthesis (43). Such a process should offer ideal conditions in regard t o temperature uniformity and result therefore in large yields of higher hydrocarbons and in low methane

306

HELMUT PICHLER

formation. Experiments of a different scale were carried out by the Coal Research Institute a t Mulheim, by Ruhrchemie, Rheinpreussen, I.G. Farbenindustrie, and others, and during the past few years also b y the Bureau of Mines. The efficiency of the suspended catalyst is determined essentially by the degree of uniform dispersion of catalyst and synthesis gas in the oil. Table VII shows an experiment with liquid phase iron catalyst. TABLE VII

I. G.ExperimentsMichael Hot Gas Recycle Oil Slurry

Process Pressure, atm. CO :HZratio Temperature, "C. Catalyst

20 1:1.2 325 Sintered Fe

Number of stages Number of times CO, removal Conversion, % Yield g./m.3 (CO H,) Liquid products C3 Ca (% olefins)

+

+ ci + cz

Space velocity (m.3 gas/ m.3 catalyst/hr.)

10

1:o.a

Duftschmitt Oil Recycle

25 1:o.a 240-200

2

240-250 Fe oxide powder reduced 3

Fused iron (reduced) 2

1 9 1-92

2 90

87

112 35(80) 48

I70 5 6

128 22 20-30

250

100

0

Report by Dr. K. Peters.

An interesting observation made by Kolbel and Ackerman (44) in connection with experiments with iron catalysts suspended in highboiling synthetic oil fractions was that the oil participates in the synthesis of high molecular hydrocarbons. Oil recycle. Another method for removal of the exothermic reaction heat of the synthesis consists in flushing a fixed catalyst bed with oil, which absorbs the heat, transporting it in a recycle stream t o a heat exchanger outside the reactor. Such experiments were carried out by Duftschmidt A.G. Oppau (see Table VII) and by Storch and Crowell and co-workers (Bureau of Mines) * (45). Some experiments were performed with oil boiling a t synthesis temperature and some with higher-boiling oil fractions. Increasing amounts of recirculating oil lower the tempera-

* See Sec. ILe(4).

GASOLINE SYNTHESIS

FROM CARBON MONOXIDE

A N D HYDROGEN

307

ture difference between the gas inlet and gas outlet of the catalyst bed. Duftschmidt allowed a temperature differential between gas inlet and outlet of 50°C. (5) Synthesis of alcohols and other oxygen-containing organic compounds. I n 1923 Fischer and Tropsch published the “Synthesis of Synthol ” (which is essentially a mixture of oxygenated compounds). Alkalized iron filings were used as catalyst for th e conversion of water gas a t a pressure above 100 atm. and a temperature of about 400°C. Later it was established that decreasing pressures (in connection with decreasing temperatures and increasing catalyst activity) favor the production of hydrocarbons. However, it was never possible t o synthesize hydrocarbons with Fe catalysts without simultaneously making some oxygenated products. The Schwarzheide experiments (Fig. 10) showed t ha t a t 10 atm. and 220°C. iron catalysts yield 10 to 30% of the total liquid synthesis products as alcohols and esters. An increase in pressure results in a considerable increase in alcohol yields. Under the title “Synol Process ” Wenzel (46) published experiments in the field of the medium-pressure synthesis with iron catalysts. The following conditions were mentioned as promoting the formation of alcohol : ( a ) low synthesis temperature (preferably below 2OO0C.), ( b ) high pressure (e.g., 20 atm.), ( c ) operation in several stages, preferably with recycling of gas, minimum conversion in each stage, and removal of carbon dioxide between the stages. Figure 11 shows the product distribution of Synol synthesis. A comparison with the results of the Schwarzheide experiments (Fig. 10) shows that the amount of alcohols was on the average somewhat lower in the Schwarzheide experiments than in the Synol experiments. In all cases the percentage of alcohols reaches a maximum in the diesel oil fraction (Cll-cld. An indirect method of producing aldehydes and alcohols from carbon monoxide and hydrogen is the 0x0-synthesis (originated by Roelen, Ruhrchemie, A.G.), in which aldehydes are produced by addition of carbon monoxide and hydrogen to olefins a t relatively low temperatures and high pressures (cobalt as catalyst), according t o the equations:

R--CH=CH*

+ C 0 + H, 7

R-CHz-CHZ-C

// H‘

I

R-CH-CH,

A=o I

H

0

308

HELMUT PICHLER

The aldehydes reacting with further hydrogen form the corresponding alcohols. The constitution of the alcohols formed as direct reaction products of the medium-pressure synthesis is closely related to the corresponding synthetic hydrocarbons. Thus, for the greatest part they have straight chains. The alcohols produced by the oxo-synthesis, however, show a high amount of branched alcohols; this is also true if the raw materials consist of primary olefins. The formation.of branchecl-chain compounds depends on double bond isomerization during the oxo-reaction. r o l w wlwbl. b*.r olwblr

DiUl

011

180-300%

W3oO.C

hydrocurbom

69 FIG. 11. Synol synthesis. GI and C) hydrocarbons.)

oxygenates

(Product distribution without consideration of

(6) Basic research work. Different German research laboratories were concerned independently with basic questions of the behavior of hydrocarbon synthesis catalysts. Scheuermann (I.G.) (47) reported about x-ray investigations on iron catalysts. Carefully reduced fused iron catalysts, used for medium-pressure synthesis, contained hexagonal iron carbide (Fe2C), previously identified by Halle (I.G. Oppau). The structure of this carbide shows great similarities t o the iron nitride (Fe2N). The hexagonal iron carbide is unstable. A t higher temperatures it can, according to Herbst, I.G., be converted to Hagg’s FezC (48). Both carbides are identical to the Fe2C carbides identified by Pichler and

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

309

Merkel (49) in iron precipitation catalysts by magnetochemical investigations. The conversion of the hexagonal FezC t o Hagg’s Fe&% irreversible. Small amounts of copper stabilize the hexagonal iron carbide and raise the conversion temperature 50-100°C. (50). With magnesium oxide, activated fused iron catalysts produce hydrocarbons with higher olefin content during the synthesis. The ratio of potassium t o magnesium seems t o be of importance. H. Kolbel (Chemische Werke Rheinpreussen) and co-workers carried out x-ray investigations as well as magnetochemical investigations on iron catalysts (51). The results of this work are in agreement with the work of Halle and Herbst, and of Pichler and Merkel in so far as it also describes the formation FezC (Hagg) as a product of low-temperature carbiding of iron catalysts. Kolbel and co-workers report (in agreement with Pichler and Merkel) that copper and alkali accelerate the carbide formation. c. Great Britain. England’s great interest in the synthesis of gasoline from carbon monoxide and hydrogen has been evident since Fischer and Tropsch issued their first publication. The British Fuel Research Board carried out thorough investigations on theory and practice of the process (52) and many papers were published by different authors. England honored Franz Fischer for his work by awarding him the Lord Melchior Medal. Elvin and Nash (53) discussed (1926-1928) the reaction mechanisms and assumed that the formation of hydrocarbons occurred by way of oxygenated products. Several years later Craxford and Rideal (54) published some interesting results regarding the question of formation of carbides as intermediate products of the synthesis. Higher hydrocarbons should be formed by the reaction of carbides with molecular hydrogen, while methane should be the result of a reaction of carbide with chemisorbed hydrogen (see Sec. 111). The Fuel Research Board operated a small pilot plant ( 5 5 ) . Experiments on a larger scale (150 gallons per day) were reported by Myddleton (56). The Committee of Imperial Defense, subcommittee on oil from coal (Falmouth) and the Committee of the British Labor Party (Hall) made thorough investigations on the economics of production of synthetic fuels for England (57). After World War 11, Hall published further papers on synthesis of hydrocarbons from carbon monoxide and hydrogen ( 5 8 ) . Interesting views on the reaction mechanism of the synthesis with iron catalysts offer x-ray investigations on the iron-carbon and ironnitrogen system by Jack (59). FezoC9 is proposed as formula of iron-

310

HELMUT PICHLER

percarbide. Decomposition of iron percarbide yields cementite and carbon while decomposition of cementite gives x-Fe and carbon. The rate of decomposition of both carbides is negligible below 350°C. d. Japan. Many papers on the Fischer process were published between 1930 and 1940 by K. Fujimura, G. Kita, S. Kodama, S. Matsumura, Y. Murata, s. Tsuneoka, and other Japanese scientists [ J . SOC. Chem. Ind. Japan and Sci. Papers Inst. Phys. Chem. Research (Tokyo)]. Cobalt, nickel, and iron catalysts combined with all types of promoters were tested. The investigations, which are valuable contributions t o the research work on hydrocarbon synthesis, confirmed the observations of Fischer and his co-workers. Two industrial plants were completed in 1938, and the erection of other plants followed. e. United States. (1) Early research work. I n 1928-1930, Smith and co-workers (60) published experiments on the Fischer-Tropsch process in the presence of cobalt-copper-manganese catalysts a s well as in th e presence of iron-copper catalysts. The composition of the reaction products (paraffins, olefins, and oxygenated products) was examined. Of particular interest was the discovery th at when ethylene is added t o the carbon monoxide-hydrogen mixture in the presence of cobalt catalysts, but not in the presence of iron catalysts, it enters into the reaction of carbon monoxide and hydrogen, increasing the production of oxygenated compounds. The ethylene did not react when passed over the catalyst alone or when in mixture with either carbon monoxide, or hydrogen. (2) Agreement with Ruhrchemie. Because of the discovery of sufficient sources of natural petroleum, there was no acute interest in the production of synthetic fuels in the United States between 1930 and 1938. I n 1938 however, the Fischer process achieved increased international importance through a n agreement made by a group of United States petroleum interests (Standard Oil Development Co., Standard Oil Co. of Indiana, Shell Development Co., The Kellogg Co.), I.G. Farbenindustrie, A.G., and Ruhrchemie A.G. The agreement brought into being an international cooperation on processes for manufacturing synthetic liquid fuels. The first result was the foundation of “Hydrocarbon Synthesis, Inc.” This cooperation was interrupted by World War 11. (3) Synthesis with fluidized catalysts. Investigations made in the United States during the last decade (61) showed that the Fischer process could economically produce gasoline in this country if the necessary carbon monoxide-hydrogen mixture was produced from natural gas. New methods, however, seemed necessary for the commercial application

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

311

of the process in the United States. Encouraged by the successful use of the so-called fluidized catalyst system for the catalytic cracking of oils, a new development also was started in the United States for the synthesis of hydrocarbons from carbon monoxide and hydrogen, P. C. Keith, president of Hydrocarbon Research, Inc., is a pioneer in this development work. He called the new type of process the Hydrocol process. It was developed by Hydrocarbon Research, Inc., under the sponsorship of the Texas Co., Standard Oil Development Co., Socony-Vacuum Oil Co., LaGloria Corp., I.S. Abercrombie Co., and Hydrocarbon Research, Inc. Experimenhl work, including pilot plant operations of all the major steps of the process, was carried out by P. C. Keith and his staff a t Hydrocarbon Research, Inc., laboratories at Olean, New York, and Trenton, New Jersey. The first commercial Hydrocol plant was erected for Carthage Hydrocol, Inc., at Brownsville, Texas. It was designed by Hydrocarbon Research, Inc., and was erected by Arthur G. McKee & Co. Important contributions to the research, upon which the design was based, were made by the Texas Co., Stanolind Oil and Gas Co., and the Standard Oil Development Co. The basic principles of the process have been described in the literature by Keith (62), Latta and Walker (63), Pichler (64), and others. The main steps of the Hydrocol process, converting natural gas to gasoline, are : A. Recovery of C4 and higher hydrocarbons from natural gas, B. Separation of high purity oxygen from air, C. Partial combustion of natural gas with oxygen to carbon monoxide and hydrogen without catalysts, D. Conversion of the carbon monoxide-hydrogen mixture to gasoline and other synthetic products in the presence of fluidized iron catalysts, E. Separation of synthetic products, F. Treatment of gasoline to remove oxygenated compounds, polymerization of propylene and butylene, blending, etc., G. Manufacture of catalysts. The first plant was built for the conversion of 64,000,000 cu. ft. natural gas per day with 40,000,000 cu. f t . oxygen per day at a pressure of 20 atm. The synthesis was carried out in large vertical cylindrical converters in which the iron catalyst, of comparatively small particle size, was in continuous movement around cooling tubes containing water at elevated pressure. In order t o maintain the desired conditions of fluidization, it was necessary to introduce the synthesis gas at the bottom of the reactor and to maintain a certain gas velocity and certain flow conditions (65). The catalyst was made by a comparatively simple treatment of cheap iron ore. Comparatively high synthesis tempera-

312

HELMUT PICHLER

tures (300-350°C.) eliminated expensive manufacture of highly active catalysts. I n spite of an 8-40 times greater throughput of synthesis gas per weight of iron catalyst than in fixed bed experiments with iron catalysts (see Schwarsheide experiments), it is possible t,o achieve yields per volume of synthesis gas equal to those obtained with cobalt catalysts a t 180-200". A one-stage operation with recycling of a part of th e effluent gas resulted in 95% synthesis gas (Hz CO) conversion. I n contrast t o other experiences with iron catalysts, the oxygen of the carbon monoxide was converted mainly to water. The Hydrocol process produces predominantly gasoline with a high content of mono-olefins and small amounts of higher-boiling products. The characteristics of the gasoline are described by Bruner (66). The C3-C4fraction is valuable as raw product for manufacturing polymer gasoline. Between 10 and 20% of the total product are oxygenated products, mostly alcohols (67).

+

TABLE VIII Composition of Hydrocol Fractions Weight, %

Component C3 Fraction Propane Propene Ca Fraction Isobutane n-Butane Isohutene 1-Rutene 2-Butene Cs Fraction Isopentane n-Pentane I-l'entane 2-l'entane 2-Methyl-I-butenc 3-Methyl-1-butene 2-Methyl-2-hutene Cyclopentane Cyclopentenc

20.2 79.8 1 .'J

13.6 8.7 64. :3 11.5 3.5 7.0 67.2 5.8 3.5 11.1 0.7 0.1 0.2

Normal Monomethyl Dimethyl Cyclics

75.9 20.0 0.4 3.7

C7 60.2 29.3 1.7 8.8

55.4 36.6 2.4 5.6

IJnsaturates, wt. % total

85.2

89.2

89.8

C6

CS

GASOLINE

SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

313

Table VIII presents results of analytical investigations made by Bruner (Texas Co., Beacon, New York). The hydrocarbons of the Hydrocol process are predominantly olefinic, with straight-chain or single-branched structure and with double bonds mostly in the l-position and not adjacent to the branching. Cyclic compounds are present as minor constituents and contain cyclic aliphatic homologs as well as benzene homologs. The refined (isomerized) gasoline mixed in the ratio of production with the polymerized C3-C4fraction and brought to a vapor pressure of 9.2 lb. (Reid) shows an octane number of 80.2 according to the motor method, and 91.4 according to the research method. Addition of 3 ml. of tetraethyl lead increases these numbers to 84.1 and 97.2, respectively. Ten per cent of the product boils below 56”C., 50% below 105”C., and 90% below 175°C. The end point is 200°C. Table IX presents the typical composition of the water-soluble chemicals (oxygenated compounds) and the oil-soluble chemicals. The fundamental advantage of the use of “fluidized catalysts” for the highly exothermic hydrocarbon synthesis consists of a radical solution of the crucial heat transfer problem, which limited the yield per space and time in the case of fixed catalyst beds.” The fluidized system presents the possibility of going to higher synthesis temperatures which means higher conversions with cheaper catalysts and more efficient heat recovery. This can be done without producing excessive amounts of carbon or methane. The yields of valuable olefinic hydrocarbons are very high in comparison with other hydrocarbon synthesis processes. The fluidized system also makes possible a continuous operation by regeneration of the catalyst in a side stream or by continuous addition of fresh catalyst and removal of spent catalyst. A great many patents describe preparation, use, and regeneration of the fluidized catalyst. Fused iron of the synthetic ammonia type is mentioned for catalyst manufacture and such ores as hematite, limonite, magnetite, pyrite, as well as mill scale, bloom scale, and other materials. I n all cases there is preliminary alkalization of the iron with about 0.5-1.5 % alkali necessary. Potassium carbonate is generally used. Some patents, however, mention special treatment of the ores, for example with potassium fluoride (68). Some patents describe the use of mixtures of different catalysts, for instance: Mixtures of iron ores (particle size 50-300 p ) and aluminum oxide (1-50 p ) lower the yields of oxygenated products (69). Mixtures of old and new catalysts lower the carbon formation. Mixtures of synthesis catalyst (e.g., Fe, 1-2% KzO, 2-3% A1203) and water gas shift catalyst

314

HELMUT PICHLER

TABLE IX Typical Composition of Water-Dissolved Chemicals and Oil-Llissolved Chemicals (Water-dissolved chemicals :oil-dissolved chemicals = 2.64: 1) Water-Dissolved Chemicals Alcohols Methanol 0 . 3 mrt. % Ethyl 36.9 Isopropyl 0.8 n-Propyl 8.7 4.0 Butyl Aniyl and higher 1.2 Aldehydes 6.0 Acetaldehyde 2.2 Propionaldehyde Butyraldehy de 2.1 Ketones 7.5 Acetone Methyl ethyl ketone 2.2 Methyl propyl ketone 0.9 Methyl butyl ketone 0.2 Acids Acetic 18.1 Propionic 4.7 Butyric 3.4 Valeric and higher 0.8

100.0

Oil-Dissolved Chemicals Alcohols Aldehydes and ketones Acids

3 2 . 9 Wt. % 33.8 33.3 100.0

(e.g., 55FezO3, 8Cr203, 26Ca0, 4AI20s, 7Mg0) (70) in order to change the consumption ratio of hydrogen and carbon monoxide. The preliminary reduction of the catalysts is carried out in a “fluid bed” at 350-500’C. with dry hydrogen of sufficient rate to keep the steam content very low during the reduction process and to maintain satisfactory conditions for the fluidization of the iron. (4) Synthesis with oil recycle in the presence of i r o n catalysts of the synthetic ammonia type. Storch and his co-workers (U.S. Bureau of Mines, Bruceton, Pennsylvania) have carried out exhaustive research and development work on hydrocarbon synthesis since the end of World War 11. Two variations of the process were particularly investigated, both using oil for direct heat exchange (71). One method employs the

GASOLINE SYNTHESIS

FROM CARBON MONOXIDE

AND HYDROGEN

315

use of a fixed bed of granular catalyst (or slightly moving catalyst) with recycling oil, while the other employs suspensions of finely powdered catalyst in oil. The first process will be used for a demonstration plant in Louisiana, Missouri. The reactor used by the Bureau of Mines consists of a vertical cylindrical vessel into which the synthesis gas and recycle oil enter at the bottom. The catalyst is either in a fixed bed, or is kept in slight motion by sufficiently increasing the gas velocity. The recycling oil transports the exothermic reaction heat to a heat exchanger outside of the reactor. Synthetic diesel oil with a boiling range of 300-450°C. may be used as recycle oil since it meets the requirement of a sulfur-free oil. After several days the oil and the production of new synthetic hydrocarbons result in an "equilibrium11 composition of the recycle oil so that 65% boils below 45OoC.* (72). At a pressure of 22 atm., temperature of 244-249"C., gas throughput of 310 liters per kilogram of iron, synthesis gas with 1.3Hz:lCO and recycle ratio (tail gas:fresh gas) of 1:1, a conversion of 79.8% carbon monoxide and 63% hydrogen was reported. Liquid product yields (including C3and C,) of 117 g. per cubic meter of feed gas were produced. Assuming a two-stage operation, 95% conversion of the synthesis gas will be possible. Liquid reaction products consist of 52% gasoline (C, 400"F.), 10% diesel oil (400-600°F.), 1691, heavy distillates (600-842"F.), 11% wax (842"F.), and 11% oxygenated compounds. The Bureau of Mines reported the use of the iron catalyst of the synthetic ammonia type given in Table X.

+

TABLE X Raw and Reduced Iron Catalyst, % Before Reduction

Fep04 MgO Cr203 MnaOa

KzO SiOl

93.51 4.61 .65 .03 .56 .64

After Reduction Total iron Reduction of t h e iron

89.2 95.5

Loss of weight on reduction 24.2

It is known that the conditions of the reduction are important t o the behavior of ammonia type iron catalysts during the synthesis of hydro* Kolbel and Ackermann report on the basis of synthesis experiments with iron catalysts suspended in hydrogenated synthetic oil, not only a catalytic cracking of the oil, but also participation of the oil in the conversion of carbon monoxide and hydrogen, resulting higher molecular hydrocarbons (44).

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HELMUT PICHLER

carbons. The reduction was carried out a t 450°C. and atmospheric pressure for 48-72 hours. The space velocity of the substantially dry hydrogen amounted to 2000. The particle sizes of catalyst used in fixed bed operations were 2-4 and 4-6 mesh, whereas in the moving bed particle sizes of 8-16 and 20-42 mesh were satisfactory. The fact t ha t fused iron catalysts of the synthetic ammonia type were successively used in many investigations of hydrocarbon synthesis for both fluidized and fixed catalyst bed operations is of interest in different respects. Due to this fact it is possible to make use of the valuable experience obtained during development work of the ammonia synthesis (73). This applies to the reduction, the tendency to oxidize, and the effect of promoters and poisons, and to a certain extent also to questions regarding the reaction mechanism. In the case of the ammonia synthesis, iron nitrides instead of carbides were identified .in spent catalysts (Fe4N, FesN, and Fe2N). I t is of particular interest that spent ammonia catalysts, containing iron-nitrogen compounds were superior in many ways t o fresh “ammonia catalysts ” for hydrocarbon synthesis. Anderson, Schults, Seligman, Hall, and Storch (74) studied the behavior of iron nitrides as catalysts for hydrocarbon synthesis. T h e r-phase nitrides which have the same crystal structure as the hexagonal FezC have a similar favorable influence on catalyst activity. The nitrides are gradually converted to the corresponding carbon compounds. Nitrided catalysts are more resistant to oxidation and the formation of free carbon. These factors may be important for a longer life of these catalysts (75). ( 5 ) Basic research work. Emmett, Storch, Taylor, and co-workers, and other scientists carried out basic research work which directly or indirectly influenced the development work on catalysts for hydrocarbon synthesis. The studies of Emmett and Brunauer on adsorption of gases by different types of catalysts (75a) and the publications of Taylor and co-workers on activated adsorption (75b) are of particular importance. Emmett also published investigations on the reaction mechanism of hydrocarbon synthesis (Sec. 111.4). Storch described our knowledge concerning the rnechanism of the synthesis in Volume I of Advances in Catalysis (19413). He based his review on experiments carried out with Weller, Hofer, and Anderson (76), and Hofer and Peebles (77) and came t o the conclusion th a t bulk carbide cannot be an intermediate product of the synthesis. Anderson, Krieg, Seligman, and O’Neill (78) and Weller (79) report activation energies of the synthesis as 24-27 kcal./mole and 26 kcal./ mole, respectively.

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

317

Basic research work on hydrocarbon synthesis was continued by the Bureau of Mines during 1949 and 1950, and the following publications were issued. Anderson, Hall, Krieg, and Seligman (80) : Activities and surface areas of reduced and carburized cobalt catalysts. Anderson, Krieg, and Friedel (81): Differential reaction rate studies on FischerTropsch catalysts. Leva, Weintraub, and Grummer (82) : Heat transmission through fluidized beds of fine particles. Weller and Friedel (83): Isomere distribution in hydrocarbons from FischerTropsch process. Clark, Andrews, Fleming (84): Composition of synthetic gasoline (identification of different aromatic compounds by infrared analysis). Hofer, Cohn, and W. C. Peebles (85): The modification of the carbide Fe&; their properties and identification. Hofer and Cohn (86) : Synthesis of cementite. Hofer and Cohn (87): Thermomagnetic determination of Hagg carbide in used Fischer-Tropsch catalysts. Hofer, Peebles, and Bean (88): X-ray diffraction studies of the action of carbon monoxide on Co-ThOz-kieselguhr catalysts. Cohn and Hofer (89): Mode of transition from Hhgg iron carbide to cementite.

The work of the Bureau of Mines is devoted to the behavior of cobalt and iron catalysts, and particularly to questions concerning the reaction mechanism and the relationship between catalyst composition and synthesis reaction. In the case of cobalt, unstable cubic cobalt was identified as the product of the reduction of standard cobalt catalysts, while hexagonal cobalt was found as a product of the hydrogenation of cobalt carbide. Used cobalt catalysts show no carbide by x-ray examination. Bulk phase carbide decreases the activity of cobalt catalysts. Surface area measurements show no appreciable change when the cobalt of cobalt catalysts was converted to cobalt carbide. Carburization a t conditions where free carbon is formed increases the area considerably. In the case of iron catalysts, x-ray and thermomagnetic investigations confirm the work of Pichler and Merkel and show that the FezC with the Curie point 265°C. of Pichler and Merkel is identical with Hagg’s carbide. Hofer, Cohn, and Peebles found the inflection point of the thermomagnetic curve at 247°C. The FezC with the Curie point 380°C. of Pichler and Merkel seems identical to a hexagonal carbide identified independently in the research laboratories of I. G. Farbenindustrie by work of Halle and Herbst (90). Outside of the Bureau of Mines, Eckstrom and Adcock (91) (Standard Oil and Gas Co., Tulsa, Oklahoma) report the discovery of a new iron

318

HELMUT PICHLER

carbide with the formula FeC. They obtained this carbide from reduced, promoted mill scale (97% Fe with minor amounts of IUn, Cu, Ni, All S, and P), used for synthesis in a fluidized catalyst bed (27 atm., 360OC.). Ninety per cent of the catalyst was reported to consist of carbide. Chemical analysis of the product involved some difficulties. Separation of the carbide from condensed hydrocarbons and free carbon was performed by magnetic separation in a solution of xylene. The Curie point of the described carbide is identical to the Curie point of Hagg’s carbide (250°C.). f . Other Countries. Research laboratories in nearly all the countries of the world have been concerned with questions regarding synthesis of hydrocarbons and oxygenated products from carbon monoxide and hydrogen. The great flexibility of the process made it applicable to many different needs, not only in the field of liquid fuells but also to the field of organic compounds. I n Italy, Natta and co-workers (92) published papers on variations of the synthesis of methanol. Alkalization of the zinc oxide catalysts of the methanol synthesis resulted in changes in the conversion. It is possible to produce a methanol-isobutanol mixture with 20-30 % isobutanol, instead of pure methanol. Italy was also interested in the manufacture of liquid fuels. Ruhrchemie made proposals for a medium-pressure plant with iron catalysts, but the termination of World War I1 prevented the realization of this project. The U.S.S.R. has been interested in the synthesis for the past ten to fifteen years. However, in spite of exhaustive discussions at the beginning of World War I1 an agreement with the German industry was not reached. After 1945 Russia continued the development work in Eastern Germany and also continued operation of a part of the Brabag-Plant at Schwarzheide. Russian scientists have also published papers, particularly Eidus and co-workers (93), on the reaction mechanism of the synthesis. Kinetic studies showed that the rate of the synthesis on cobalt, nickel, and iron is greater than the rate of carbiding. The Shell Development Co. is carrying out some research work on synthesis in Holland. At present, South Africa and other countries are investigating the technical and economical possibilities for the erection of large plants on the basis of medium-pressure synthesis. The economic situation is connected with local conditions of the costs, necessasy raw materials, and the value of the end products. Some time in the future, when the petroleum sources of the world will be insufficient to se,tisfy the demands of the market, it will be necessary in all countries to produce more and

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

319

more chemicals and liquid fuels by synthesis. The synthesis from carbon monoxide and hydrogen will be one of the most interesting possibilities. 111. SOLVED A N D UNSOLVED PROBLEMS OF HYDROCARBON SYNTHESIS 1. Summary of Relations between Catalyst, Synthesis Conditions, and

Reaction Products The pressure-temperature diagram of Fig. 12 shows the optimum regions of operation for the various processes used in the hydrogenation

m

FIG. 12. Pressure-temperature regions for synthesis processee on a carbon monoxide and hydrogen basis.

of carbon monoxide. The arrows show the sequence in which the development work has proceeded. Each of the catalysts in Fig. 12 has a specific action toward hydrogen and carbon monoxide as shown by the tendency to form methane or higher hydrocarbons, or alcohols, not to mention the ability t o hydrogenate, t o dehydrate alcohols, to polymerize olefins and to promote reactions such as isomerization, alkylation, and cyclization. Moreover, by suitable choice of synthesis conditions these characteristic properties of the catalyst may be guided into a particular direction. For instance, the less the hydrogenating power of a given catalyst, the higher must be

320

HELMUT PICHLER

the temperature and the pressure, in order t o obtain a given reaction. Addition of promoters also may serve t o carry the conversions in a certain direction. Under the conditions of the Fischer-Tropsch synthesis, i.e., a t 200” and atmospheric pressure, cobalt, and t o a lesser extent iron, were suitable catalysts for the synthesis of higher hydrocarbons. Nickel, which has a much greater hydrogenating action than cobalt or iron promotes the formation of methane and could be used for synthesis of higher hydrocarbons only by “diluting ” with kieselguhr. Fischer and Tropsch used cheap catalysts (alkalized iron turnings) and high temperature and pressures in their initial development work. The products of the reaction were oxygenated compounds (Synthol). Later they used catalysts of high activity, low temperature, and low pressure. As a result of these experiments the opinion was established that the optimum conditions for production of hydrocarbons are atmospheric pressure and lowest possible temperatures (normal-pressuresynthesis). Later development work demonstrated that each specific catalyst should be used within certain limits of pressure in orlder t o obtain optimum yields of hydrocarbons. Nickel cannot be used a t pressures above atmospheric because of its tendency t o form volatile carbonyls. The optimum pressure range of cobalt catalysts which form carbonyls a t higher pressures is 5-20 atm., of iron catalysts 10-30 atm., and of ruthenium 100-500 atm. Working at high pressures I.G. came first t o a Synthesis of mixtures of different organic compounds and later t o the specific synthesis of methanol. Oxide catalysts, in general, show a smaller degree of activity toward carbon monoxide and hydrogen than the metal catalysts. High pressures and temperatures are required for conversion which is the result of surface reactions. Whereas the high hydrogenating power of cobalt, nickel, and ruthenium orient the hydrogenation of carbon monoxide almost entirely toward hydrocarbons, and the less active iron also produces some alcohols; oxide catalysts favor the formation of alcohols. Just as conditions can be adjusted so that nickel produces methane as the sole hydrocarbon, so, in the case of zinc oxide, methanol may be produced as the only alcohol. Similarly, just as methane is the sole hydrocarbon that can be formed by direct synthesis from carbon dioxide and hydrogen, so methanol is the only alcohol of which this is true. Addition of alkali t o the catalysts of the iron group shifts the reaction toward the formation of higher-boiling compounds. I n a similar manner, addition of alkali t o zinc oxide favors the formation of higher alcohols.

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

321

When added to thorium oxide, it also results in the formation of higherboiling hydrocarbons. The fact that hydrocarbons constitute the chief product of isosynthesis results from the intense dehydrating action of thorium oxide and related catalysts. In the region of operation for this synthesis, low temperatures produce high yields of alcohols. A rise in temperature at first increases the iso-olefin yield; then, as the hydrogenating power of the catalyst increases, isoparaffins are formed. A further rise in temperature produces naphthenes and finally, because of their high thermal stability, aromatics. 2. Poisoning of Catalysts

Decline of catalyst activity during the synthesis of hydrocarbons may be caused by two basically different reasons: A. Poisoning by impurities brought into the system with the synthesis gas, for example: dust, tar, resin formers, hydrogen sulfide, organic sulfur compounds, and chlorine and other elements. B. Decline of catalyst activity for reasons connected with the synthesis itself, for example : sintering, adsorption of high molecular synthetic products (waxes, etc.), excessive carbon formation, salts, produced on the catalyst surface by conversion with synthetic acids, and formation of volatile carbonyls and carbonyl hydrogen compounds. The unfavorable influence of gas impurities (group A) can be eliminated (within certain economic limits) by proper purification of the synthesis gas. Unfavorable results of the synthesis reaction itself and of undesirable by-products (group B) must be eliminated or reduced t o a minimum by suitable synthesis conditions. Figure 13 presents an example for changes in the behavior of a cobalt catalyst used for synthesis of hydrocarbons from carbon monoxide and hydrogen, and analytical investigations of a cobalt catalyst at different distances from the gas inlet (94). Figure 13A shows the progress of carbon monoxide conversion (standard cobalt catalyst, fixed catalyst bed, 100 vol. synthesis gas throughput per volume of catalyst per hour) at different points of the gas inlet on the first day of operation. The freshly reduced Co catalyst is very active. Carbon monoxide conversion was 75% in the first 10% of the catalyst bed. Figure 13B presents a corresponding picture for the same catalyst after 30 days of operation. During this time it was necessary t o increase the temperature of the reactor some degrees in order to obtain a CO conversion similar t o that at the beginning of the operation. The first 15%

322

HELMUT PICHLER

of the catalyst bed were “dead” at the end of the 30 clays run. Up to this time 75% CO conversion was reached in 55% of :the catalyst bed. The failure of the catalyst at the gas inlet must be attributed to two facts: first, excessive production of reaction heat at the gas inlet may cause some sintering of the catalyst and formation of carbon and other undesirable products (oxide) at the beginning of the operations. SecDIRECTION OF GAS FLOW-

301h Day of Synthesis 50

(XCO-conversion)

25

0

-

10 20 30 40 50 60 70 80 90 100

2 ,

1

t

X Pomffin m the Cotolyst I

2or5zd 0

10 20 X , 40 50 60 70 80 90 100

bCork + r891n-lihe suastonces

10

0

10 20 50 40 50 60 70 Bo 90 100

’

LENGTH OF CATALYST BED (INU

’

FIQ.13. Poisoning of cobalt catalyst.

ondly, and the chief reason for the decline of the catalyst activity, this part of the catalyst is a screen for impurities, particu1a:rly of sulfur. Figure 13C, D, and E show the percentages of some impurities detected in the catalyst after one year of operation. Figure 13E shows sulfide- and sulfate-sulfur. Both are concentrated within the first part of the catalyst bed. The formation of sulfate is the result of the presence of 0.1% O2 in addition to traces of sulfur compounds. The conditions concerning the promotion of carbon and resin-like substances are very similar. The amount of high-melting synthetic paraffins is higher a t the side of the gas outlet of the catalyst bed, due to lower catalyst temperatures.

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

323

The waxes can be removed from the catalyst by extraction with solvents (synthetic oil). Resins are usually difficult to dissolve. However, it is possible to remove resins and a great part of the carbon by treatment with hydrogen a t increased temperatures (about 400°C.). It is also theoretically possible to remove the sulfide-sulfur, particularly if the hydrogen used for the decomposition of the sulfides passes the catalyst bed countercurrently to the synthesis gas. In this paper it is not possible to discuss in detail all the reasons responsible for activity decline of synthesis catalysts. Sintering can be prevented by introduction of high-melting oxides in the catalyst structure. Carbon formation is mostly connected with the disability of the reactor to remove pompletely the exothermic reaction heat. High hydrogen content of the synthesis gas and low temperatures can lower or prevent the formation of carbon. Cobalt and iron salts were detected in many cases during the regeneration of the catalysts by dissolving with nitric acid. Formation of volatile carbonyls, connected with decline of catalyst activity, can sometimes be observed particularly in the case of nickel catalysts. The presence of carbon monoxide prevents the conversion of carbon dioxide with hydrogen to methane (95). Of great interest are two questions. A. To what degree does steam (and carbon dioxide) spoil the activity of catalysts during reduction and synthesis? The degree of drying of hydrogen used for reduction of certain carbon monoxide hydrogenation catalysts (particularly fused iron catalysts, to a smaller extent standard cobalt catalysts) is important to the behavior of the catalyst during the synthesis. It is best to use very dry hydrogen and high gas velocities, to keep the partial pressure of the reduction water low and to continue the “reduction” for some time after “the last traces” of reduction water disappears in the off gas. It is difficult to approach this problem by thermodynamical calculations. Almquist and Black (96) discussed this many years ago in connection with a similar problem, the reduction of iron catalysts used for ammonia synthesis: At a reduction temperature of 444°C. the stable iron oxide phase is FesOl. The equilibrium of the reaction fFe

+ HzO = $FesO, + Ha

was represented by K = H2/H20 = 5. “This is equivalent to a volume percentage of water vapor of about 16%. That is, iron at this temperature should not be oxidized by hydrogen-steam mixtures containing less than 16% of steam. I n our experiments some oxidation of catalytic iron occurred at the lowest

324

HELMUT PICHLER

oxygen concentration used, namely 0.008 %, which is equivalent t o 0.016% of water vapor. At 0.02% of oxygen or 0.04% of water vapor, the oxidation corresponded t o the retention of 4.8 mg. of oxygen at th e steady state, that is, after the water in the gas leaving the catalyst was equal t o that entering. The ratio of H2/H20 was then 75/0.04, for which K = 1875. While the ratio of these equilibrium constants may be taken as an approximate measure of the greater ease of oxidation of a certain range of unsaturated iron atoms, as compared t o crystalline iron it must be recognized t ha t the two constants are not strictly connparable in t h a t we do not know what the solid oxide phase is in the latter case. Indeed, since the small number of unsaturated atoms undergoing oxidation is widely distributed among a large number of normal iron atoms, it would not be expected that a definite oxide phase such as Fe304 would be formed. On this basis, the variation in the equilibrium constants is a function of both solid phases, iron and iron oxide. If we make the assumption, however, that the solid oxide phases are the same, and th a t all of the change is attributed t o a difference in the iron phase, it is possible t o determine a maximum energy difference between crystalline iron and the average of the unsaturated atoms over the range covered, by combination of the equations :

+

+

PFe H 2 0 = $Fe20, HZ; K1 = 5 $Fe (active) H10 = $Fe304 HI; K2 = 1875

+

+

(1 1 (2)

By subtracting (1) from (2), :Fe(,, = QFe(K2/K1) = 375

for which, a t 444”C., AF = -RT In ( K 2 / K 1 )= -8600 cal., or -11,500 cal. per gram atom of iron, This value then represents an upper limit for the average free energy of a range of unsaturated iron1 atoms which are responsible for a greater part of the ammonia-synthesizing activity of the catalyst. A similar calculation employing the lowest concentration of vapor used, namely 0.016%, gives an energy value of 13:200 cal., which is of the same order of magnitude.” Schutza (97) discussed similar problems in connection with the possible oxidation of active cobalt catalysts according to: CO

+ HzO+

COO

+ Hz

The surface conditions of the cobalt determines the maximum water contained in hydrogen-steam mixtures, without oxidizing the cobalt catalyst. It is interesting to note that traces of steam present during the reduction lower the catalyst activity considerably, while the same

GASOLINE

SYNTHESIS FROM CARBON MONOXIDE

AND HYDROGEN

325

catalyst does not lose its activity during the following synthesis in the presence of concentrations of steam (and carbon dioxide) of greater magnitudes formed with the synthesis according to:

+ 22.Hz = (CH& + z.HzO + 2.Hz = (CHz)s + PCOZ

z.CO 22CO

During the first days of synthesis a partial oxidation of the catalyst occurs, acaompanied by a slight decrease of activity. The synthesis temperatures however, are lower than the reduction temperature and the “re ”-oxidation of the catalyst is less dangerous a t synthesis conditions than “insufficient preliminary reduction.” After some time, the reoxidation process reaches a kind of equilibrium. The influence of a certain oxygen-content at these “equilibria conditions ” on catalyst activity, is not identical with the influence of a similar oxygen content remaining in the catalyst at the end of the reduction process. The conditions of the reduction are not of the same importance for different kinds of catalysts. If the partial pressure of steam and carbon dioxide exceed a certain level during the synthesis (depending on the kind of catalyst, the composition of the gases present, and the temperature and pressure), oxidation begins with a dangerous decline of catalyst activity. B. What is the maximum amount of sulfur tolerable in gases passing a synthesis catalyst* (98)? It is a well-known fact that in order to prevent poisoning of catalysts the use of iron, cobalt, nickel, or ruthenium demands strict observance of regulations concerning the sulfur content of the reacting gases. An equilibrium exists between metal, sulfur compounds, hydrogen, steam, and other gases; e.g., in the case of nickel:

+ + + + + + + + + +

Ni HzS NiS Hz NiS H 2 0 F? NiO HzS NiO Hz d Ni Ha0 CSz 4Hz CHc 2HzS

Organic sulfur compounds will be converted to hydrocarbons and hydrogen sulfide. The nickel/nickel-sulfide hydrogen/hydrogen-sulfide equilibrium favors the formation of nickel sulfide at low temperatures and favors its decomposition at high temperatures (98a). An increase in the steam concentration decreases the extent of conversion of nickel to inactive nickel sulfide similar to an increase of the hydrogen partial pressure. It would appear valuable to calculate the maximum amount of sulfur which might be present at different temperatures, without converting the nickel to nickel sulfide. The energy differences of differ-

* Contrary to some publications synthesis gas with the lowest sulfur content gives the best yields.

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HELMUT PICHLER

ent atoms of the surface of an active catalyst however, make the value of such calculations doubtful, similar t o the case of the oxidation of active iron by hydrqgen-steam mixtures. In order to obtain a positive picture of the limits a t which the sulfur content of gases is not dangerous for reactions over certain catalysts, it is 100.0

10 0

v)

3 “E

8 -’

D

0

I .o

0.1 I

I

I

200

400

.

I

600

800

TEMPERATURE :C

FIG. 14. Poisoning of Ni-catalysts used for adjustment of equilibrium: CO 3Hz i=? CH, H10.

+

necessary to evaluate experiments or commercial operations. example may illustrate these facts. The equilibrium CO

+ 3Hz

CH4

+

An

+ HzO

can be approached from both sides in the presence of nickel catalysts. At low temperatures (below 3OOOC.) methane and water are stable and at high temperatures (above SOO0C.) carbon monoxide and hydrogen. The production of methane (Sabatier’s synthesis) as well as the “re-forming” of methane to carbon monoxide and hydrogen are technically important conversions.

GASOLINE

SYNTHESIS

FROM CARBON MONOXIDE

AND HYDROGEN

327

Figure 14 shows maximum sulfur contents for the equilibrium mentioned above a t different temperatures of operation. Point 1 corresponds to experiences of the Coal Research Institute in Mulheim, point 2 presents conditions of a technical scale operation for production of synthesis gas for ammonia synthesis. The points 3-8 are literature values for similar processes (99). With increasing temperatures it is possible to tolerate increasing amounts of sylfur. However, increasing temperatures decrease the demands in catalyst activity also. At temperatures above 1200°C. it is possible t o carry out the methane conversion in the absence of metallic catalysts. Fischer and co-workers mentioned 0.2 g./100 m.3as the “economic” limit of the sulfur content of synthesis gas for production of hydrocarbons in the presence of cobalt, nickel, and iron. Ruthenium shows at least the same sensitivity toward sulfur.in spite of its great stability in the presence of other compounds. The thorium oxide catalyst of the iso-synthesis cannot be poisoned by sulfur or sulfur compounds. The activity of this catalyst slowly declines only in cases when carbon covers its surface. A periodical treatment with oxygen containing gases restores the activity of the high melting oxide catalysts without activity decline by sintering. 3. Hydrocarbon Synthesis and Water Gas Equilibrium

The reactions taking part in hydrocarbon synthesis from carbon monoxide and hydrogen are very heterogeneous. Table XI and Table TABLE XI Some Typical Conversions, Important I Synthesis “on Co catalysts” I1 Synthesis “on Fe catalysts” 111 Water gas equilibrium (water gas shift reaction) IV Synthesis of methane V Synthesis of methane VI Methane from carbon dioxide (in the absence of carbon monoxide) VII Formation of carbon VIII Hydrogenation of carbon IX “Oxygenation” of carbon with steam

for Synthesis of Hydrocarbons z(C0) 22(H2) = (CHz), 2~(C0) = (CHz),

+ +

+~(Hz0) +~(C0z)

+ + Hz + + Hz0 + + COz COz + 4Hz = CH, + 2Hz0 2co = c + coz C + 2H2 = CHa C + 2H20 = COZ+ 2Hz CO HzO = COz CO 3Hz = CHa 2CO 2Hz = CH,

XI1 present a few typical conversions. Equilibrium conditions are only of secondary importance for the composition of the reaction products. Rather, the velocities at which certain conversions occur, control the process.

328

HELMUT PICHLER

TABLE XI1 Values of the Logarithm of the Equilibrium Constanf, loglo K a

Reactions

CO

+ + + +

+

HzO (gas) = COz Hz CO 3Hz = CH4 HzO (gas) 2H2 = CHa f COz 2CO COz 4Hz = CH4 2 H z 0 . (gas) 2co = c COZ

+

+ +

400°K.

500°K. 600°K. 700°K. 800°K.

3 . 1 7 0 2.100 1.433 15.611 10.059 6.296 18.782 12.179 i . 7 2 9 12.441 7.939 4.863 13.283 8.752 5.729

0.955 3.570 4.526 2.616 3.57'4

0.606 1.506 2.112 0.900 1.963

D. D. Wagman, J. E. Kilpatrick, W. J. Taylor. K. S. Pitzer, F. D. Rossini, J . Reaeareh Natl. Bureau o f Standards, 92 (February. 1945). (The paper presents logarithms with live numbers after the decimal point.)

Iron catalysts show a tendency to convert the oxygen of carbon monoxide to carbon dioxide, while under ,similar conditions cobalt produces water. In all cases long contact times of the reacting products on the catalyst surface increase the carbon dioxide formation and short contact times increase the formation of water. These results (confirmed by experiments with many different catalysts) were the basis for the hypothesis that water is a primary reaction product while carbon dioxide is a secondary product produced by the water gas shift reaction according to equilibrium conditions. Methane can be the product of one of the conversions mentioned in Table XI. However, methane can also be the product of hydrocracking reactions of higher synthetic hydrocarbons. Some catalysts have a tendency to produce oxygenated organic products, while some produce only hydrocarbons and water. The molecular weight and degree of saturation of the hydrocarbons can be changed by changing the catalyst composition or by changing the process conditions. The same is true for carbide and carbon formation and also for degree and kind of branching of the reaction products, etc. All these questions have been of great interest, since the start of research work on hydrocarbon synthesis. Smith (100) calculated equilibria conditions for the synthesis of methane, ethane, propane, and higher hydrocarbons with both water and carbon dioxide as by-products, and also the equilibria between carbon monoxide, hydrogen, and alcohols. Smith came to the conclusion that there is a greater tendency t o produce higher hydrocarbons and water than lower hydrocarbons and carbon dioxide. The statement that higher pressures would be favorable for synthesis of higher hydrocarbons, was opposed by Tropsch (101), who pointed out that carrying out the synthesis a t increa,sed pressures results in increased production of oxygenated compounds.

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

329

Thermodynamic questions often were discussed in connection with the reaction mechanism of the synthesis (102). Montgomery and Weinberger (103) report that the distribution of paraffins obtained a t normal-pressure synthesis on cobalt is similar to that predicted by thermodynamic equilibrium calculations. Fischer and Koch (104) carried out experiments on the formation of carbon dioxide and water as by-products of the normal-pressure synthesis. They came to the following conclusions: Cobalt and nickel catalysts convert the oxygen of the carbon monoxide preferentially t o water, iron catalysts to carbon dioxide. The different behavior of the iron can not be explained by conversion of primarily produced water with carbon monoxide. The water gas shift reaction can be carried out in the presence of cobalt catalysts as well as in the presence of iron catalysts. The amount of carbon dioxide increases with increasing synthesis temperatures, and also in the presence of cobalt catalysts. Craxford (105), who in several publications supported the carbide theory, explained the formation of carbon dioxide (in the presence of cobalt catalysts) as a secondary reaction, taking place on parts of the catalyst surface which aie covered with uncarbided reduced metal atoms. This part of the catalyst could also be responsible for hydrocracking of higher hydrocarbons to methane. Kolbel and Engelhardt (106) investigated the conversion of carbon monoxide with steam (water gas shift reaction) in the presence of carbided and uncarbided reduced cobalt and iron catalysts (24OoC., CO HzO = 1 :1). Both carbide and metal catalyze the water gas reaction. Uncarbided catalysts, however, were better catalysts of the water gas shift reaction. Cobalt catalysts which at normal conditions convert 99 % of the oxygen of the converted carbon monoxide to water (180°C., Hz :CO = 2 : 1, space velocity = loo), can reverse their behavior relative to water and carbon dioxide formation by changing the synthesis conditions. A conversion of the oxygen of CO t o COZ oxygen amounting to 99% was obtained at the same temperature by lowering the space velocity to 10 and by changing the Hz:CO ratio of the synthesis gas t o 2:3.55. This is due to a conversion of primarily produced water with excessive carbon monoxide. In the case of iron catalysts it was possible t o reverse the formation of carbon dioxide and water (from 81:19 t o 18:82) by carrying out the reaction in nine steps (removal of water in eight steps with corresponding changes of the conditions for the water gas equilibrium). Bashkirov, Kryokov, and Kagan (107) also report studies on the mechanism of the formation of carbon dioxide and water. Water forma-

+

330

HELMUT PICHLER

tion according to the equation CO

+ 2Hz+

CHz

+ HzO

is a primary reaction on nickel, cobalt, and iron catalysts, while carbon dioxide is a secondary product of the conversion. Anderson, Krieg, Friedel, and Mason (108) studied the progress of the synthesis reaction in a cobalt catalyst bed. The reaction rate was high in the initial part of the bed, but lower and nearly constant throughout a large portion of the catalyst bed, until the reactants were fairly completely consumed. Methane and carbon dioxide were formed by primary and secondary reactions, at least to some extent, throughout the entire catalyst bed. Most of the carbon dioxide, however, was formed by the water gas reaction. When the partial pressure of carbon monoxide went down to less than 0.01 atm., the carbon dioxide, formed in the reaction, reacted with excessive hydrogen to methane. At the beginning of the carbon dioxide consumption the quotient of the partial pressures (0) (H20) (C02) (Hz)

was 50 to 100 times greater than the equilibrium constant. The fact that the synthesis reaction and the water gas shift reaction occur in parallel, but with different velocities, is very important for the result of the overall conversion. Optimum synthesie catalysts are not necessarily optimum catalysts of the water gas shift equilibrium too. A patent by Watson (109) (to The Texas Co.) [see See. II.2.e(3)] describes the use of a mixture of a typical synthesis catalyst aind a typical water gas shift catalyst for carrying out the synthesis in a fluidized bed. Both iron and cobalt catalysts are able to convert the oxygen of the carbon monoxide to water and carbon dioxide during hydrocarbon synthesis. At the same synthesis conditions however (temperature and gas composition), the iron catalysts show a greater tendency to form carbon dioxide and the cobalt and nickel catalysts show a greater tendency to form water. These results, typical for iron .and cobalt, can be changed with the help of the water gas reaction by changing the partial pressures of the different reacting components. Achievement of equilibrium conditions according to the water gas reaction occurs very slowly a t low synthesis temperatures. Increasing temperatures increase the rate of the conversion, yielding normally increasing a mounts of carbon dioxide. However, it is also possible to convert all the oxygen of the carbon monoxide to water in the presence of iron catalysts at high temperatures (300°C.). After consumption of the greatest part of the carbon monoxide, carbon dioxide starts to react with hydrogen producing (over inter-

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

331

mediate formation of carbon monoxide) higher hydrocarbons in addition to methane. Figure 15 shows carbon balances for different carbon monoxide conversions. Increasing carbon monoxide conversion increases the production of higher hydrocarbons, C1 and Cz hydrocarbons and carbon dioxide. The percentage of carbon dioxide, however, decreases a t very high carbon IOC

90

80

70

60

i 3 m

5 50 z >

s

0 40 0

se 30

20

10

8

I

I0

20

I

30

40 CARBON

I

1

50 60 BALANCE

I

70

I

BO

90

0

FIG.15. Carbon balances at different CO conversions (iron catalyst).

monoxide conversions. Both carbon monoxide and carbon dioxide approach a value close t o zero, almost at the same time. (The synthesis gas contained a hydrogen-carbon monoxide ratio of 2:l. The operation was carried out in one step with a recycle ratio of 1:1.)

4. Pro and Con Carbide Theory In 1926 Fischer and Tropsch (110) discussed the mechanism of the reaction between carbon monoxide and hydrogen t o form hydrocarbons.

332

HELMUT PICHLER

They assumed that carbon monoxide and metal react first to form carbides (formation of higher carbides is thus to be expected) and that hydrogen then reacts with carbidic carbon to form (CF[z) radicals. These methylene groups polymerize t o form more or less saturated hydrocarbon chains. This theory of the formation of hydrocarbons over an intermediate formation of carbides is generally referred to as the “carbide theory.” In 1932 Fischer and Koch (111) formulated a mechanism of the synthesis as follows. Carbon monoxide and hydrogen are first adsorbed on the catalyst surface. The carbon monoxide then undergoes chemisorption at the active centers (112) with loosening of the carbon-oxygen bond. Hydrogen, or excess carbon monoxide on the iron, reacts with the oxygen to water, or carbon dioxide, respectively. The carbon bound as carbide is next transformed by further action with hydrogen t o form CH, CH2, and CH3 radicals, which then polymerize to a wide variety of hydrocarbons. As the length of the chain increases, the desorption of the compounds from the catalyst surface proceeds more and more slowly, thus increasing the interval of time available for complete hydrogenation by the hydrogen. This assumption would easily explain the steady decrease in unsaturation of the products as the boiling point rises. Under certain conditions, high molecular weight hydrocarbons may be split into smaller fractions by a cracking process, thus establishing an equilibrium between polymerization and cracking. A great deal of research work was performed in connection with the question of whether or not the carbide theory presents the correct explanation of the conversions occurring on the catalyst surface during the hydrocarbon synthesis. Many experiments seemed t o support the hypothesis of Fischer and co-workers, while results abf other more recent investigations makes intermediate formation of carb.ides more and more unlikely. a. ChemicaZ Investigations. Experiments were performed in which the conversion of carbon monoxide and hydrogen t o hydrocarbons was divided into two steps: (1) the conversion of carbon monoxide with the catalyst metal t o carbide and (2) the conversion of carbide with hydrogen to hydrocarbons. Carbided cobalt reacts with hydrogen a t synthesis conditions. However, the product of the conversion is methane only (113). Iron carbides are very stable against hydrogen a t synthesis conditions. The decomposition of iron carbide with acids yields higher hydrocarbons in addition to methane (114). b. Ortho-para Hydrogen Equilibrium. Craxford and Rideal (115) studied the question of the conversion of para to ortho hydrogen in the

GASOLINE SYNTHESIS FROM CARBON MONOXIDE A N D HYDROGEN

333

presence of cobalt catalysts. The para-ortho conversion could be observed at conditions which are favorable for methane formation. A large part of the conversion was prevented a t conditions which are favorable for the formation of higher hydrocarbons. The investigators concluded that methane is a product of a reaction which can occur in the absence of carbide by action of chemisorbed hydrogen. Chemisorbed hydrogen serves also in the breaking u p of hydrocarbon chains. T h e synthesis of higher hydrocarbons, however, should occur over a formation of CH2 groups, produced by a reaction of the carbide with molecular hydrogen. Emmett and Harkness used the ortho-para hydrogen conversion for investigations on surfaces of iron catalysts (115a). c. C14 as Auxiliary Substance. Iron and cobalt catalysts were carburized (Kummer, Dewitt, Emmett, 116) partly with normal carbon monoxide and partly with radioactive C L 4 0 . The rate of appearance of C14 among the hydrocarbons produced by the synthesis should clarify whether or not the carbides are intermediate products. The conclusion drawn from this work was that only a small part of the hydrocarbons could be produced by hydrogenation of bulk carbide, provided that a uniformly active catalyst is assumed. On the basis of the assumption of a nonuniformly reacting surface, however, it cannot be completely ruled out that a greater part of the hydrocarbons were formed via a n intermediate carbide formation. d. Thermomagnetic Studies. Exhaustive thermomagnetic studies on iron catalysts were carried out by Pichler and Merkel (117). Hofer, Cohn, and Peebles (118) continued this work a t th e Bureau of Mines. Similar studies were performed by Kolbel, Ackerman, Juza, and Tentschert (119). Pichler and Merkel found that the iron of iron catalysts is converted during the carburization with carbon monoxide a t temperatures below 400°C. t o two carbides with the approximate formula Fe&, one with a Curie point at 265°C. (according t o Hofer, Cohn, and Peebles, 247°C.) and the other with a Curie point at 380°C. Both carbides are unstable a t temperatures above 400°C. At temperatures between 250°C. and 400°C. the 265" Curie point carbide was observed and a t lower temperatures increasing amounts of the 380" Curie point carbide. This carbide predominates a t temperatures of 200°C. and lower (most active iron catalysts). Copper and alkali promote the carbide formation. Iron catalysts with the highest activity for the synthesis of hydrocarbons show the highest carbide content. It is possible to convert the iron almost completely to Fe2C. These facts seem to be in good agreement with the carbide theory. However, other results obtained b y the same authors, and also by Kolbel and co-workers, show that the bulk carbides cannot

334

HELMUT PICHLER

be intermediate products of the synthesis. The bulk carbides are very stable against hydrogen at synthesis temperatures. An almost completely carbided iron catalyst showed no appreciable change in its composition (no changes of thermomagnetic curves) after 121 hours treatment with hydrogen (3000 vol. Hz/vol. catalyst/hour at 220°C.). e. X-Ray Difraclion Studies. Carbiding of finely dispersed nickel with carbon monoxide at temperatures below 270”C., converts the nickel to Ni& (120) which has, according to Jacobsen and Westgren (121), a closely packed hexagonal structure. Carbiding of cobalt with carbon monoxide at temperatures below 230°C. yields CozC (122) which reacts with hydrogen at synthesis conditions to cobalt metal and methane. Freshly reduced cobalt catalysts (with hydrogen at 360-400°C.) and spent cobalt catalysts used for hydrocarbon synthesis, contain primarily cubic cobalt. Hydrogenation of cobalt carbide however, converts the cobalt to hexa,gonal cobalt (123). This result seems to be a strong argument against the carbide theory. X-ray diff ract,ion studies on iron samples carburized with carbon monoxide were carried out by Hagg (124), Halle and Rerbst (125), Hofer, Cohn, and Peebles (126), Kolbel, Ackerman, Juza, and Tentschert (127), and by Eckstrom and Adcock (128). Hagg prepared his carburized iron samples below 260”C., following a method described by Bahr and Jessen (129), and determined x-ray diffraction patterns, according to Hofer, Cohn, and Peebles, identical to the Fe&, with a Curie point at 265°C. of Pichler and Merkel. Halle and Herbst obtained a hexagonal carbide by carburization of iron-copper catalysts, and later also by carburizalion of copper-free catalysts (reduction and carburization at low temperatures). The x-ray pattern is not identical to that described by Hagg. On the basis of their x-ray investigations Hofer, Cohn, and Peebles believe that the carbide of Halle and Herbst is identical to the FezC carbide with a Curie point at 380°C. of Pichler and Merkel (see Sec. III.4.d). Halle and Herbst found that hexagonal FezC is not stable, as did Pichler and Merkel in their magnetic studies. The hexagonal Fe2C Curie point at 380°C. can be converted to FezC wii,h a Curie point at 265°C. in copper-free catalysts at temperatures of 290°C. and higher, and in copper containing catalysts a t 350°C. and higher. Herbst assumed that the new hexagonal carbide would be the “active carbide’’ while the carbide of Hagg would be “inactive.” This does not agree with results obtained by Pichler and Merkel, who show that active catalysts of medium-pressure synthesis are in many cases composed only of Fe2C (Curie point at 265”C.), although catalysts containing mixtures of both carbides are more active.

GASOLINE SYNTHESIS FROM CARBON MONOXIDE A N D HYDROGEN

335

f. Thermodynamics. Kummer, Browning, and Emmett (130) carried out thermodynamic calculations concerning the possibility of the formation of hydrocarbons by direct reduction of Fe2C. Equilibrium constants for the synthesis of n-paraffins (C,-C8) and of n-l-olefins (CZ-C,) were calculated for 227°C. and 327°C. When these equilibrium constants are compared with the yields actually obtained, it appears that these products can not be formed by the reduction of Fe2C. Schuman (131) calculated the equilibrium constants for the reactions:

+

+

+

Fe& (2n - 1)H2 (n - 1)CO = CnHzn -t 2Fe (n - 1)HzO FezC 2nH2 (n - 1)CO = C,HZ,+Z 2Fe (n - 1)H20

+

+

+

+

for n = 2 or 8 at 500°K. and 600°K. These calculations indicate that the formation of hydrocarbons from FezC may be thermodynamically possible. g. Kinetics. Eidus (132) carried out investigations on the velocities of synthesis and carbide formation on cobalt and iron catalysts. The rate of synthesis was greater than the rate of carbiding. The ratio, rate of synthesis to rate of carbiding increased with the duration of the experiment; in the case of cobalt and nickel, from 5 to 20-30 and in the case of iron from 2 to 8. Eidus says that the formation of carbides can he an intermediate step of the synthesis on iron catalysts, but not on cobalt catalysts . Weller (133) passed carbon monoxide over a cobalt catalyst at synthesis conditions and found high carbon monoxide consumption during the very first moment of the treatment. The carbon monoxide consumption diminished rapidly to a much lower but steady rate. Evidently a flash carbiding of the surface occurred, followed by a slower reaction of the bulk metal. The rate of synthesis is comparable with the initial carbiding rate. Hydrogenation of the cobalt carbide proceeds faster than the carbiding (in contrast to the experiences with iron catalysts). h. General Consideraiions. A review of the research work, carried out to explain the reaction mechanism of hydrocarbon synthesis does not yield entirely satisfactory results, in spite of the great efforts of many investigators and the use of different kinds of scientific tools. The assumption that carbides are intermediate products of the synthesis is not probable, at least in so far as the bulk carbide, which can be identified by thermomagnetic investigations, is concerned. The greatest part of the results, however, cannot be recognized as satisfactory proof against the carbide theory if only small parts of an irregular catalyst surface were involved in the reactions. A comparison of the behavior of the different catalysts seems to be useful.

336

HELMUT PICHLER

Active iron catalysts react with carbon monoxide t o carbides at synthesis conditions. These carbides accumulate in the catalyst because they are resistant against hydrogen. Active cobalt catalysts can be converted t o carbide too. Cobaltcarbide however, reacts easily with hydrogen. Spent cobalt catalysts, therefore, contain only small amounts of carbide. Carbides of ruthenium which might possibly be formed a t synthesis conditions are not known. The optimum conditions of pressure and tempera,turefor hydrocarbon synthesis on nickel, cobalt, iron, and ruthenium are close t o the conditions a t which formatlion of carbonyls and carbonyl hydrogen compounds can be detected. Pressure “requirements” for the catalytic synthesis and for carbonyl formation appear t o be closely parallel in the case of metallic catalysts. The physical structure, which can be changed by suitable methods of catalyst manufacturing, is .of decisive importance (promoters : highmelting oxides; supports : kieselguhr of cobalt and nickel catalysts; pretreatment : low-temperature reduction which limits the size of the crystals, or carbon monoxide treatment of iron catalysts which increases the surface by “breaking u p ” the structure with carbon). The oxide catalysts of the iso-synthesis produce hydrocarbons without a detectable intermediate formation of definite chemical compounds between catalyst and reactants. The iso-synthesis produces hydrocarbons, as has been proved, over a primary formation of oxygenated products and a secondary dehydration t o olefins. Some facts make primary formation of alcohols probable also for iron catalysts : Alcohols and hydrocarbons produced by the synthesis show similar structure of the branching, and process conditions favorable for recovery of intermediate products (short contact times, high recycle ratio, low temperatures, etc.) usually increase the yields of alcohols. Conversion of the bulk of the catalyst metal t o carbide during hydrocarbon synthesis is only observed in the case of iron catalysts. The carbide formation, which occurs parallel with the activity of iron catalysts, may have an important influence on the conditions of the catalyst structure and catalyst surface. The carbide formation, however, seems t o be insufficient for the catalyst activity. Treatment of carbided iron catalysts with sulfur does not change the carbide content but makes the catalyst inactive for hydrocarbon synthesis (134). A comparison of the results of the basic research work on hydrocarbon synthesis does not make it probable t o assume basically different reaction mechanisms for the different catalysts, a t least for the metallic catalysts.

GASOLINE SYNTHESIS FROM CARBON MONOXIDE A N D HYDROGEN

337

Gradual differences, which can be adjusted by process conditions show greater probability. The first step is a chemisorption of carbon monoxide, yielding a wide variety of compounds on the catalyst surface, ranging from very loosely bound surface compounds to well-defined carbonyls and also carbides. However, a very small part of the catalyst may be “active” for the synthesis. The second step will be, at least for a great part of the conversions, a formation of oxygenated compounds, probably of compounds containing HCOH groups which can be converted in the sequence, alcohols, olefins, and paraffins. Formation of methane and chain lengths of higher molecular hydrocarbons depend on the presence of active hydrogen and are the result of different types of conversions. Secondary polymerizations and depolymerizations (135), as well as hydrogenations and hydrocracking, are important for the final composition of the reaction products.

REFERENCES 1. Fischer, F., and Tropsch, H., Ber. 69, 830, 832, 923 (1926). 2. Fuel 6, 323 (1926). 3. Storch, H. H., Chemistry of Coal Utilization, Vol. 11, p. 1797. Edited by H. H. Lowry. John Wiley I% Sons, New York, 1945; Advances in Catalysis, Vol. I, p. 115. Academic Press Inc., New York, 1948. 4. Fischer, F., Proc. 1st Intern. Conj. Bituminous Coal, pp. 238-246 (1926). 5. Fischer, F., OeZ u. KohZe, No. 21, 22 (1943). 6. Mittasch, A., Advances in Catalysis, Vol. 11, p. 97. Academic Press Inc., New York, 1950. 7. French Patent 571,356: 580,905 (B.A.S.F. 1923). 8. Fischer, F., and Tropsch, H., Brennstoff-Chem. 8, 165 (1927). 9. Tropsch, H., Brennstoff-Chem. 8, 376 (1927). 10. Fischer, F., and Tropsch, H., Ges. Abhanal. Kenntnis KohZe 10,313-501 (1928). Fischer, F., Brennstoff-Chem. 11, 489 (1930). 11. Fischer, F., and Meyer, K., Brennstofl-Chem. 12, 225 (1931). 12. Fischer, F., and Koch, H., Brennstog-Chem. 13, 61 (1932). 13. Raney, M., J . Am. Chem. SOC.64, 4116 (1932). Fischer, F., and Meyer, K., Brennstof-Chem. 16,84, 107 (1934). 14. Fischer, F., Roelen, O., Feisst, W., Brennstog-Chem. 13, 461 (1932). Fischer, F., Brennstofl-Chem. 11, 489 (1930); 16, 1 (1935). 15. Roelen, O., Brennstoff-Chem. 12, 305 (1931). Roelen, O., and Feisst, W., U.S. Patent 2,110,240-1 (1938). Griffith, R. H., Gas World 107, 381 (1937). Gollmar, H. A., Chemistry of Coal Utilization, Vol. 11, p. 947. Edited b y H. H. Lowry. John Wiley & Sons, New York, 1945. Sands, A. E., and Schmidt, L. E., Ind. Eng. Chem. 42,2277 (1950). 16. Fischer, F., Brennstoff-Chem. 16, 6 (1935). Tropsch, H., and Koch, H., Brennstof-Chem. 10, 337 (1929). Koch, H., and Horn, O., Brennstoff-Chem. 13, 164 (1932). Hartner-Seberich, R., and Koch, H., Brennstog-Chem. 13,308 (1932). 17. Fischer, F., and Pichler, H., Brennstoff-Chem. 17, 24 (1936). 18. Tropsch, H., Brennstoff-Chem. 8, 376 (1927). Fischer, F., and Pichler, H., Brennstof-Chem. 12, 365 (1931). Fischer, F., and Kiister, H., Brennsto$-Chem. 14, 3 (1933).

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19. Fischer, F., Brennstoff-Chem. 16,2 (1935). 20. Fischer, F., and Meyer, K., unpublished. 21. Meyer, K., and Bahr, Th., unpublished experiments by Kaiser Wilhelm Institut in Miilheim. 21a. German Patent Appl. ST56470, July 30, 1937. 22. TOM Reel 101, Doc. Pg 21559 NID; Pg 21574 NID. 23. Troitekii, K. V., Russian Patent 54,392 (1939); see OeZ u. Kohle, 36, 73 (1940). 24. Pichler, H., and Merkel, H., U.S. Bureau Mines Special Rept., 1947; U.S. Bureau Mines T’ech. Paper 718, 1949; Brennstoff-Chem. 31, I, 33 (1950). 25. Pichler, H., Brennstog-Chem. 19, 226 (1938). Pichler, H., and Buffleb, H., Brennstoff-Chem. 21, 247, 273, 285 (1940). 26. Fischer, F., Pichler, H., and Lohmar, W., Brennstoff-Cheni. 20, 247 (1939). 27. Fischer, F., and Pichler, H., Brennstoff-Chern. 14, 306 (1933). 28. Pichler, H., Special Rept. for the U. S. Bureau of Mines (11947). Storch, H. H., Advances in Catalysis, Vol. I, p. 136. Academic Press lnc., New York, 1948. Pichler, H., and Ziesecke, K. H., Brennstoff-Chem. 30, 13, 60, 81 (1949). 29. Pichler, H., Ziesecke, K. H., and Titzenthaler, E., Brennstofl-Chem. 30, 333 (1949). 30. Sabatier, P., and Senderens, J. B., Compt. rend. 134, 514 (1902). 31. Sabatier, P., 6th Intern. Congress for Chemistry, Rome, 1906. 32. Sabatier, P., Die Katalyse, Akadem. Verlagsges., p. 104, Leipzig, 1927. 33. Patart, Ch. 2. 49,564; C 1925 I1 1562. 34. Audibert, C., and Raineau, A., Ind. Eng. Chem. 21, 880 (1929). 35. Decarriire, E., and Antheaume, A., Compt. rend. 196, 1889 (1933). 36. Lefebvre, H., and LeClerc, G., Compt. rend. 203, 1378 (1936); 207, 1099 (1938); 208, 1583, 1650 (1939). 37. Pichler, H., and Merkel, H., U.S. Bureau of Mines Special Rept. 1947; U.S. Bureau of Mines Tech. Paper ‘718, 1949; Brennstoff-Chem. 31, 1, 33 (1950). 38. Prettre, M., Rev. Znst. franc. pktrole 2, 131, 195 (1947); .J. Chem. Phys. 16,424 (1948); Compt. rend. 224, 278 (1948). 39. Trambouee, Y., et al., Compt. rend. 327, 971; 228, 837, 1015, 1432 (1949). 40. Teichner, S., and Pernoux, E., Compt. rend. 228, 1644, 1646 (1949). 41. Martin, F., Ind. Chem. 13,320 (1937); Oel u. Kohle 13,693 (1937); Chern. Fabrik 12,233 (1939). Paul, H., and Tramm, H., Erdol u. Koh.le 2, 229 (1949). 42. Fischer, F., and Koch, H., Brennstoff-Chern. 1930-1950. 43. Fischer, F., and Peters, K., Brennstoff-Chern., 12, 288 (1931). Fischer, F., Roelen, O., and Feisst, W., Brennstoff-Chem. 13, 461 (1!332). Fischer, F., and Kiister, H., Brennstoff-Chern.14,3 (1933). Fischer, F., and Pichler, H., Brennstoff-Chem. 20, 247 (1939). 44. Kolbel, H., and Ackermann, P., Brennstoff-Chem. 31, 10 (1950). 45. Crowell, I. H., Benson, H. E., Field, I. H., Storch, H. H., Ind. Eng. Chem. 42, 2376 (1950). 46. Weneel, W., Angew. Chem. B20, 225 (1948). 47. Scheuermann, A., Angew. Chem. 60, 211 (1948). 48. Hagg, G., 2. Krist. 89, 92 (1934). 49. Pichler, H., and Merkel, H., U.S. Bureau Mines Special Rept., 1947; U.8. Bureau Mines Tech. Paper 718, 1949; Brennstoff-Chem. 81, 1, 33 (1950). 50. TOM Reel 26, Bag 2463; Reel 134, Item 11/10. 51. Kolbel, H., Ackermann, P., Juza, R., and Tentschert, H., Erdol u. Kohle 2, 278 (1949). Kolbel, H., and Langheim, R., Erdol u. Kohle 2, 544 (1949). 52. Sinnatt, F. S., Gas J . 212, 711 (1935).

GASOLINE SYNTHESIS FROM CARBON MONOXIDE AND HYDROGEN

339

53. Elvin, 0. C., andzNash, A. W., J . Sac. Chem. Ind.'46, 876 (1926); 46, 473T (1927); Fuel 6,263 (1926). 54. Craxford, S. R., and Rideal, E. K., Trans. Faraday Sac. 36, 946 (1939); J. Chem. Sac. Pt. 11, 1604 (1939); Trans. Faraday SOC.42,580 (1946); Fuel 26, 119 (1947). 55. Dept. Sci. Ind. Research, Fuel Research Dept., March 1937, p. 142. 56. Myddleton, W. W., J. Inst. Fuel 11, 477 (1938). Aicher, A., Myddleton, W. W., and Walker, J. J., J . Sac. Chem. Ind. 64, 313T (1935); 66, 121T (1936). 57. Egloff, G., Ind. Eng. Chem. 30, 1091 (1938). 58. Hall, C. C., Gas World 126, 58 (1947); Engineering 163, 30, 54, 79 (1947); Trans. Inst. Chem. Engrs. (London) Advance copy, January 7, 1947; J . Junior Inst. Engrs. (London) 67, 309 (1947); J . Inst. Petroleum 33, 439 (1947). 59. Jack, H., Proc. Roy. SOC.(London) A196, 34, 41, 56 (1948). 60. Smith, D. F., Davis, I. D., and Reynolds, D. A., Ind. Eng. Chem. 20, 462 (1928). Smith, D. F., Hawk, C. O., and Reynolds, D. A., Ind. Eng. Chem. 20,1341 (1928). Smith, D. F., Hawk, C. O., and Golden, P. L., J . Am. Chem. Sac. 62,3221 (1930). 61. Alden, R. C., Oil Gas J . 46, No. 27, 79, 97 (1946). Crutchfield, J. W., and Purett, H. T., Petroleum Bng. 13, 208 (1946). Keith, P. C., Gas 22, No. 6, 21 (1946); Am. Gas. J. 164, No. 6, 11 (1946). Fieldner, A. C., Science Counselor 10, No. 1, 8-24 (1947). Gunness, It. C., Chem. Eng. News 26, 2123 (1947). Lee, J. A., Chem. Engr. p. 105, October, 1947. Roberts, G., Phinney, J. A., Oil Gas. J . 46, No. 45, 72 (1947). Murphree, E. V., Oil Gas J . 46, No. 49, 66 (1948). Schroeder, W. C., Chem. Ind. 62, 574, 682 (1948). Storch, H. H., Chem. Eng. Progress 44, No. 6, 649 (1948) ; Advances in Catalysis, Vol. I, p. 115. Academic Press, Inc., New York, 1948. 62. Keith, P. C., Oil Gas J . 46, No. 6, 102 (1946). 63. Latta, J. E., and Walker, H. W., Chem. Eng. Progress 44, 173 (1948). 64. Pichler, H., Brennstof-Chem. 30, 105 (1949). 65. Symposium on Dynamics of Fluid Solid Systems, Ind. Eng. Chem. 41, 1099 (1949). 66. Bruner, F. H., Ind. Eng. Chem. 41, 2511 (1949). 67. Sullivan, F. W., Chem. Eng. Progress 43, 13 (1947). Eliot, T. Q., Goddin, C. S., Pace, B. S., Chem. Eng. Progress 46, 532 (1949). 68. British Patent 621,902 (1949). 69. U.S. Patents 2,479,439 and 2,453,874. 70. U.S. Patent 2,486,895. 71. Crowell, I. H., Benson, H. E., Field, I. H., and Storch, H. H., Ind. Eng. Chem. 42, 2376 (1950). 72. Kolbel, H., and Ackermann, P., Brennsto$. Chem. 31, 10 (1950). 73. Mittasch, A., Advances in Catalysis, Vol. 11, pp. 80-104. Academic Press, Inc., New York, 1950. Frankenburg, W. G., Synthetic Ammonia, in F. TJllman Eneyklopadie der Technischen Chemie 2nd. ed. Berlin, Wien, 1928, Vol. 1, pp. 383-400, Vol. 6, p. 436. Emmett, P. H., Synthetic Ammonia in Fixed Nitrogen, ed. by H. Curtis, pp. 150-240. The Chemical Catalog Co., New York, 1932. Group Leader Conference, Ludwisgshafen (I.G.). TOM Reel 26, Bag 2463; TOM Reel 134, Item 11/10. 74. Storch, H. H., J . Am. Chem. Sac. 72, 3502 (1950). 75. Frankenburg, W., (to Hydrocarbon Research, Inc.) U.S. Patent 2,531,420 (1950). 75a. Emmett, P. H., and Brunauer, S., J. Am. Chem. Sac. 69, 310, 1553 (1937). Brunauer, S., and Emmett, P. H., J . Am. Chem. Sac. 69,2682 (1937). Brunauer, S., Emmett, P. H., and Teller, E., J. Am. Chem. SOC.60, 309 (1938). Brunauer, S., and Emmett, P. H., J . Am. Chem. Sac. 62, 1732 (1940). Emmett, P. H.,

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HELMUT PICHLER

J . Am. Chern. SOC.68, 1784 (1946). Emmett, P. H., Advances in Catalysis, Vol. I, p. 65. Academic Press, Inc., New York, 1948. 75h. Taylor, H. S., Eleventh Report of the Committee on Contact Catalysis, p. 52, John Wiley bi Sons, New York, 1935. Morikawa, K., Benedict, W. S., Taylor, H. S., J. Am. Chem. SOC.68, 1445, 1795 (1936). Morikawvri, K., Trenner, N. R., Taylor, H. S., J. Am. Chem. SOC.69, 1103 (1937). Taylor, H. 8., Advances in Catalysis, Vol. 1, p. 1 . Academic Press, Inc., New York, 1948. 76. Weller, S., Hofer, I,. J. E., and Anderson, It. B., J. Am.. Chetn. Soc. 70, 799 (1948). 77. Hofer, L. J. E., and Peebles, W. C., J. Am. Chem. SOC.69, 893 (1947). 78. Anderson, R. B., Krieg, A., Seligman, B., and O’Neill, W. E., Ind. Eng. Cheni. 39, 1548 (1947). 79, Weller, S., J. Am. Chem. SOC.69, 2432 (1947). 80. Anderson, R. B., Hall, W. K., Krieg, A., and Seligman, H., J . Am. Chem. SOC. 71, 183 (1949). 81. Anderson, R. B., Krieg, A., Friedel, R. A., Ind. Eng. Chem. 41, 2189 (1949). 82. Leva, M., Weintraub, M., Grummer, M., Chem. Eng. Progress 46, 563 (1949). 83. Weller, S., and Friedel, R. A., J. Chem. Phys. 17, 801 (1949); 18, 157 (1950). 84. Clark, A,, Andrews, A., Fleming, H. W., Znd. Eng. Chem. 41, 1527 (1949). 85. Hofer, L. J. E., Cohn, E . M., and Peebles, W. C., J. Am. Chem. SOC.71, 189 (1949). 86. Hofer, L. J. E., and Cohn, E. M., J. Chem. Phys. 18, 766 (1950). 87. Hofer, L. J. E., and Cohn, E. M., Anal. Chenz. 22, 907 (1950). 88. Hofer, L. J. E., Peebles, W. C., and Bean, E. H., J . Am. Chem. SOC.72, 2698 (1950). 89. Cohn, E. M., and Hofer, L. J. E., J. Am. Chem. SOC.72, 4662 (1950). 90. Halle, F., and Herbst, M., TOM Reel 26, Bag 2463; TOM Reel 134, Item 11/10. 91. Eckstrom, H. C., and Adcock, W. A., J . Am. Chem. SOC.72, 1042 (1950). 92. Natta, G., et ul., Giorn. ehim. ind. applicuta 14, 217 (1932); Publication Akad. Luftfahrtforschung, Vol. 9, p. 79 (1939). 93. Eidus, Ya. T., et al., Bull. Acad. Sci. U.R.S.S. Classe sci. chim. 190, 1942; 145, 303, 1943; 255, 349, 1944; 62, 1945; 447, 1946; J. Gen. Chem. (U.S.S.R.) 16, 869, 875 (1946). 94. Alberts, L., [I. S. Bureau Mines Special Rept. 1947. 95. Fischer, F., and Pichler, H., Brennstog-Chem. 14, 306 (1933). 96. Almquist, J. A., and Black, C. A., J. Am. Chem. SOC.48, 2814 (1926). Curtis, H. A., Fixed Nitrogen, Chapter 8; Emmett, P. H., A.Ch.S. Monograph series, 1932, p.jl86. Emmett, P. H., and Brunauer, S., J . Am. Chem. SOC. 62,2682 (1930). 97. Schutza, H., Angew. Chem. 60, 211 (1948). 98. Herington, E. F. G., and Woodward, L. A., Brennstof-Chem. 20, 319 (1939). 98a. Kirkpatrick, W. J., Advances in Catalysis, Vol. 3, p. 329,,Academic Press, Inc., New York, 1951. 99. Bosch, C., Chem. Fabrik 7, 1 (1934). Gluud, G., Ber. Ges. Kohlentech. 9, 211 (1930). Karzhavin, W. A., Trans. Chem. Engr. Congr. World Power Conf. London, 1936, Vol. 2, p. 383. Sachsse, H., Chem. Fabriit 21, 3 (1949). Osterrieth, J. W., and Dechamps, G., Gas World 98, No. 2531, Supp. 14 (1933). 100. Smith, D. F., Znd. Eng. Chem. 19, 801 (1927). 101. Tropsch, H., Brennstofl-Chem. 8, 376 (1927). 102. Francis, A. W., Ind. Eng. Chem. 20, 277, 283 (1928). Myddleton, W. W., J . Znst. Petroleum 30, 211 (1934). Prettre, M., Rev. inst.jraric. petroZe2, 131 (1947). 103. Montgomery, C. W., and Weinberger, E. B., J. Chem. Phys. 16, 424 (1928).

GASOLINE

SYNTHESIS FROM CARBON MONOXIDE

AND HYDROGEN

341

Fischer, F., and Koch, H., Brennstoff-Chem. 13, 67, 428 (1932). Craxford, S. R., Trans. Faraday SOC.42, 576 (1946). Kolbel, H., and Engelhardt, F., Erdol u. Kohle 2, 52 (1949); 3, 529 (1950). Bashkirov, A. N., Kryokov, Y. B., and Kagan, Y. B., Doklady Akad. Nauk S.S.S.R. 67, 1029 (1949); Chem. Ab. 43, 9415 (1949). 108. Anderson, R. B., Krieg, A., Friedel, R. A., and Mason, L. S., Znd. Eng. Chem. 41, 2189 (1949). 109. Watson, C. W., U. S. Patent 2,486,894 (1949). 110. Fischer, F., and Tropsrh, H., Brennstoff-Chem. 7, 97-104 (1926). 111. Fischer, F., and Koch, H., b'rennstoff-Chem. 13,428 (1932). Herington, E. F. G., Chemistry & Industry 66, 347 (1946). 112. Taylor, H. S., Proc. Roy. SOC.(London) A108, 105 (1925); J. Phys. Chem. 30, 145 (1926). 113. Fischer, F., and Koch, H., Brennstof-Chem. 13, 68 (1932). Tebboth, J. A., J. SOC.Chem. Znd. 67, 62 (1948). 114. Fischer F., and Bahr, H. A., Ges. Abhandl. Kenntnis Kohle 8, 269 (1928). 115. Craxford, S. R., and Rideal, E. K., J. Chem. SOC.,1604,1939; Trans. Faraday Soc. 42, 576, 580 (1946). 115a. Emmett, P. H., and Harlaess, R. W., J. Am. Chem. SOC.67, 1624 (1935). 116. Kummer, J. T., Dewitt, T. W., Emmett, P. H., J. Am. Chem. SOC.70, 3632 (1948). 117. Pichler, H., and Merkel, H., U.S. Bureau Mines Special Rept., 1947; U.S. Bureau Mines Tech. Paper '718,1949; Brennstoff-Chem. 31, 1, 33 (1950). 118. See references 85-89. 119. Kolbel, H., Ackermann, P., Juza, R., and Tentschert, H., ErdoL u. hohZe 2,278 (1949). Kolbel, H., and Langheim, R., Erdol u. Kohle 2, 544 (1949). 120. Bahr, H. A,, and Bahr Th., Ber. 61, 2177 (1928). 121. Jacobsen, B., and Westgren, A., 2.physik. Chem. B20,361 (1933). 122. Bahr, H. A,, and Jessen, V., Rer. 63, 2226 (1930). Hofer, L. J. E., and Peebles, W. G., J. Am. Chern. SOC.69, 893, 2497 (1947). 123. Weller, S., Hofer, L. J. E., Anderson, R. B., J. Am. Chem. SOC.70, 799 (1948). 124. Hagg, G., 2. Krist. 89, 92 (1934). 125. Scheuermann, A., Angew. Chem. 60,211 (1948). Wykoff, R. W. G., and Crittenden, E. D., J. Am. Chem. SOC.47,2866 (1925). Brill, R., and Mark, H., J.Phys. Chem. 133, 443 (1928). 126. Hofer, L. J. E., Cohn, E. M., Peebles, W. C., J. Am. Chem. SOC.71, 189 (1949). 127. Kolbel, H., Ackermann, P., Juaa, R., and Tentschert, H., Erdol u. Kohle 2, 278 (1949). Kolbel, H., and Langheim, R., Erdol u. Kohle 2, 544 (1949). 128. Eckstrom, H. C., and Adcock, W. A., J. Am. Chem. SOC.72, 1042 (1950). 129. Bahr, H. A., and Jessen, H., Ber. 66, 1238 (1933). 130. Kummer, J. T., Browning, L. C., and Emmett, P. H., J. Chem. Phys. 16, 739 (1948). Browning, L. C., DeWitt, T. W., and Emmett, P. H., J. Am. Chem. SOC.70, 3632 (1948); 72, 4211 (1950). 131. Schuman, S. C., J. Chem. Phys. 16, 1175 (1948). 132. Eidus, Ya. T., Bull. acad. sci. U.R.S.S. Classe sci. chim. 190, 1942; 255, 1944; Hall, C. C., and Smith, S. I,., J. SOC.Chem. Ind. 66, 128 (1946). 133. Weller, S., J. Am. Chem. SOC.69, 2432 (1947). 134. Pichler, H., and Merkel, H., Rrennsloff-Chem. 31, 40 (1950). 135. Koch, H., and Gilfert, W., BrennstoJ-Chem. 30, 213 (1940) 104. 105. 106. 107.

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The Free Radical Mechanism in the Reactions of Hydrogen Peroxide JOSEPH WEISS University of Durham, King’s College, Newcastle-upon-Tyne, England

CONTENTS Page 343 347 351 352 354 357

I. The Reaction between Hydrogen Peroxide and Ferrous Ions.. . . . . . . . . . . 11. The Reaction between Hydrogen Peroxide and Ferric Ions.. . . . . . . . . . . . . 111. The Reaction between Hydrogen Peroxide and Cupric Ions.. . . . . . . . . . . . IV. The Decomposition of Hydrogen Peroxide a t Different Metal Surfaces. . . V. The Photochemical Decomposition of Hydrogen Peroxide. . . . . . . . . . . . . . . VI. The Decomposition of Hydrogen Peroxide by Ionizing Radiations. . . . . . . VII. The Reaction between Ozone and Hydrogen Peroxide and the Decomposition of Ozone in Aqueous Solution.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . VIII. Detection of Free OH Radicals.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . IX. Some Thermodynamics Data Concerning the Radicals OH and H O Z . .. . . 1. Electron Affinities and Heats of Hydration.. . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

358 361 361 362 364

In 1931 Haber and Willstiitter (1) suggested that the radicals OH and HOz play an important role in the reactions of hydrogen peroxide and that these radicals could operate in a simple chain mechanism. Previously, Franck and Haber (2) had introduced these radicals into the mechanism of autoxidation processes in solution.

I. THEREACTION BETWEEN HYDROGEN PEROXIDE AND FERROUS IONS The reactions of hydrogen peroxide were taken up by Haber and Weiss (3a,b), who studied certain aspects of the reaction between hydrogen peroxide and ferrous salts and also outlined the importance of the free radical mechanism in many other reactions of hydrogen peroxide. After the early work of Schonbein and the preparative work of Fenton the H202-Fe11-saltreaction was investigated by Manchot and Lehmann (4),who claimed to have obtained the following results: 1. In the presence of excess of ferrous salt, e.g., if a given amount of dilute hydrogen peroxide solution be added with stirring t o a solution of excess of ferrous sulfate the total reaction was represented by the equation: 2FeSOd

+ HZOZ+ ~ H z O= 2Fe(OH)3 + ZHZSO~ 343

(ml)

344

JOSEPH WEISS

2. In the presence of excess of hydrogen peroxide e.g., if a given amount of a dilute solution of ferrous sulfate is added t o a dilute solution of hydrogen peroxide (the latter being in excess) the total reaction was represented by :

+

2FeS01+ 3HzOZ 2Hz0 = 2Fe(OH),

+ 2HzSO4 + OZ

(m2)

which was supposed to be valid over a very wide range of concentrations. In order to explain the two different stoichiometric relations (ml) and (m2) Manchot and Lehmann postulated the intermediate formation of a quinquevalent iron oxide (Fez06). Haber and Weiss (3) were able to show that this explanation could not be the correct one as, although Eq. (ml) was generally valid in the presence of excess of Fe"-salt, no real significance could be attached to Eq. (m2), as the mean consumption ratio fi, defined by: fi=

Number of moles of Ha02 decomposed __Number of moles of ferrous salt oxidiae'd

depended not only on the concentration of the reactants but also on the manner in which the component solutions were mixed. Furthermore, under suitable experimental conditions mean consumption ratios far in excess of 3/2 [corresponding to Eq. (m2)] were actually obtained. By studying the reaction with special mixing devices, e.g., a rapidly moving centrifugal jet or by mixing the solutions inLthe form of two turbulent jets, it was soon established that fi depends on the concentration of the reactants in the actual reaction space and that in order t o obtain reproducibles results well-defined initial concentrations of the reactants must be ensured which is possible only if the rate of mixing is rapid compared with the reaction rate. The results obtained with different mixing devices led Haber and Weiss (3) to the formulation of the reaction as a radical and chain reaction. If FeII-salt is at all times in excess the mechanism can be represented by the two simple electron transfer processes:

+ H ~ O-+kikaZ Fea+ + OH- + OH Fe*+ + OH -+ Fe3f + OH-

Fez+

The mean consumption ratio (t, = time required for the reaction i.e. until the stationary state has been reached) :

FREE RADICALS I N THE REACTIONS O F HYDROGEN PEROXIDE

345

in this case has the value fi = 0.5. For values fi > 0.5 it was originally suggested that Eqs. (1) and (4) had to be supplemented by the chain reactions (1): ka

+ + + + +

Hz02$- OH --+ H 0 2 HzO 28 kcal. HzOz H0z-t 0 2 HzO OH 20 kcal.

+

(2) (3’)

It is obvious that these four equations could account for these results. However, various difficulties still existed (dependence on pH, etc). I n 1935 Weiss ( 5 ) proposed a modification of reaction (3’) the HO, being replaced by its anion 02-with the dissociation equilibrium: HO2 a H+

+

02-

(Dis. constant, KHOJ

leading to the simple electron transfer process: Hz02

+ Oz-+ 0%+ OH- + OH -t kr

-

10 kcal.

(3)

It is well known that under the conditions of Manchot’s experiments the resulting ferric salt is always precipitated, not as Fe(OH)3 as Manchot’s equations suggest but presumably as a basic salt, and, therefore, originally, it was assumed that any reactions of the ferric ions could be completely neglected. However, recently (6,7) it has become clear that under the conditions of Manchot’s experiments the ferric salt is not eliminated sufficiently rapidly and that whenever ferric ions are present in the solution to any appreciable extent, practically all the oxygen is produced by the reaction:

+

Fea+

ks

02-

+ Fez+

+

0 2

(5)

This reaction is not new. I n 1935 it had been advanced in connection with the reaction between Fez+-ions and molecular oxygen ( 5 ) , of which it represents the “back” reaction, and it was also discussed by Haber and Weiss (3) in the form of: Fe*+ -tHO2+ Fez+

+ + H+ 0 2

a reaction recentIy suggested by Baxendale et al. (6). Recent work has shown (8) that in general processes of the type: Me+”

+

02- +

Meh-’

+

0 2

can play an important part whenever polyvalent metal ions of higher oxidation potential are present, e.g., in the case of cupric, cobaltic and eric ions. Thus in the reaction between hydrogen peroxide and ferrous salts Eqs. (11, (21, (31, (4), ( 5 ) , and also Fez+

+ HOS+ks Fe8+ + Hot-

(6)

346

JOSEPH WEISS

are necessary. One can, however, easily distinguish between two limiting cases. I . Ferric ions are suflciently rapidly eliminated. This case can be realized if suitable anions are present in the solution and in sufficient excess t o convert all the ferric ions rapidly into relatively inert complex ions. This is the case in the presence of excess jluoride or pyrophosphate ions. Such experiments have been carried out recently by Humphrey and Weiss (7,s)over a wide range of concentrations wing a flow apparatus on the lines suggested by Schmid (9). It could be shown that in the presence of fluoride or pyrophosphate reaction (5) can be practically neglected as an oxygen-producing reaction. However, even under these conditions mean consumption ratios far greater than 0.5 could be obtained : this clearly establishes reaction (3) as the oxygen-producing reaction under these conditions. This case which, in a sense, corresponds more closely to the one treated originally by Haber and Weiss (3) can be described by the five Eqs. (l), (2), (3), (4),and (6). For the stationary state of the radicals we obtain the equations:

at

=

0 = k1[Fe~+l[H2021 - k ~ [ H ~ 0 ~ 1 [ 0-Hk4[Fez+][OH] l t- k3[H~0~][0~-] (Ia)

=

0 = k2[H20~1[0Hl - ~~[HZO~I[O Zke[Fe2+l[H0~1 -I

(Ib)

and for the rate equations:

with

and

This gives for the mean consumption ratio:

2. Ferric ions remain in the solution. In this case one can in general neglect reaction (3) and for the stationary state corresponding to the reactions (I), (2), (4),(5), and (6) one obtains:

m) = 0 = kl[Fe*+][H202] - k2[H2OZ][0H] - k4[Fe2+][OH] at

(IVa)

d(Ho2) ~-

(IVb)

at

- 0 = k2[H202][0H]- k~,[FeS+][0~-] - k6[Fe2+][HO2]

F R E E RADICALS I N T H E REACTIONS OF HYDROGEN PEROXIDE

347

with the rate equation :

In general according to Eq. (V) the rate of oxygenevolutionincreases with increasing ferric ion concentration, but there are also ferric and ferrous terms in the denominator of Eq. (V) and if these are of importance a slower increase results. Under these conditions the mean consumption ratio assumes a somewhat different meaning as ferric ions are reduced to ferrous ions during the course of the reaction until some stationary state with regard to the (ferric/ferrous) ratio has been reached. This will become clearer in connection with the reaction between hydrogen peroxide and ferric-ions which is discussed below. In general, the individual rate constants in' the above expression cannot be measured directly. It is not difficult, however, to obtain the velocity constant of reaction (1) because if the process is carried out in the presence of excess of ferrous salt the rate equation jcorresponding to reactions (1) and (4)] assumes the simple form:

Such measurements were first carried out by Haber and Weiss (3b) in the concentration range of up to molar of the reactants. Under these conditions the reaction is fairly rapid and a special apparatus had to be employed. Their measurements (in the temperature range 6" t o 45") can be represented by: kl

=

4.0 X lo4 * exp.

(- g) moles-'

. liters . set.-'

Evans et al. (10) used a colorimetric method and worked in the concentration range of molar, where the reaction could be followed over a considerable period. They obtained (in the temperature range 12" to 35") : kl = 1.78 X lo9exp.

(- T)moles-'

*

liters * sec.-l

11. THEREACTION BETWEEN HYDROGEN PEROXIDE AND FERRIC IONS This reaction has been studied by a number of authors. rate law:

The simple

348

JOSEPH WEISS

where f denotes the total amount of the iron salt added as ferric salt, was first proposed by Bertalan (11) on the basis of experiments in sulfuric acid (10-3m to 10-lm). This was confirmed by variouri authors (3b,12) as being valid under certain conditions. The above rate law was interpreted by Haber and Weiss (3b) as resulting essentially from the electron transfer process : Fea+

+ HO2--+

ki

Fez+

+ HOe

(7)

which is the reverse process of reaction 6 and represents the reduction of ferric ions by the anion HOz-, connected with hydrogen peroxide by the equilibrium : HOz-

+ H+ @ HZOZ (Dis. constant, K H ~ o J

It is obvious th at the rate law (VII) does not permit an unambiguous decision with regard t o the oxygen-producing reaction which follows. This has also been pointed out recently by Kolthoff and Medalia (13). However, the recent more extensive work of Andersen (14) who investigated the reaction between ferric salts and hydrogen peroxide in nitric acid solutions (where the tendency for the formation of ferric complexes is greatly reduced) permits a more detailed treatment of the theory. Andersen’s empirical rate law is represented by the equation : Bt a = (H202)t-o) z

=

+

7

=log:

(H202)t-t, A

+ A

= propf,

(S -a> B

=

(H+)

(VIII) -r =

small constant

Andersen himself suggested that his rate law could be derived from the mechanism (14) : ( A l ) Fe3+

+ H02-

FeOH2+

+ 0,

(A2) 0

+ H02--+ OH- +

0 2

Kolthoff and Medalia (13) have already pointed out t h a t reaction (Al) involves a very high change in standard free energy (- 18 kcal.) which is very difficult to reconcile with the high velocity constant of the forward reaction. The general Haber-Weiss mechanism involves an intermediate reduction oxidation of the ferric-ferrous ion system. This has been accepted by a number of authors and Simon et al. (15) have shown t h a t catalytic decomposition of hydrogen peroxide by ferric salts only takes place under conditions where the formation of ferrous ions can be demonstrated. There is, therefore, no experimental basis for a mechanism without the intermediate reduction of the ferric ions. This brings out also the difference between the ferrous and the ferric reactions : starting from ferric ions and hydrogen peroxide there must be an induction period during which tke reduction process, ferric -+ fer-

FREE RADICALS IN THE REACTIONS OF HYDROGEN PEROXIDE

349

rous, takes place until the stationary equilibrium ([Fe3++I8/[Fe2+]J = const. has been established. This is normally almost entirely on the ferric side and, therefore, in this case the shift from the initial state is very small. Consequently the induction period is of very short duration and practically the whole reaction proceeds at the stationary (ferric/ferrous) level (see Fig. la). It is different if one starts with ferrous salt: as the stationary equilibrium is Iargely on the ferric side, ferrous ions have to be oxidized is reached until the very low stationary ferrous concentration ([Fe2++I8) which, under comparable conditions, is identical with the one obtained by starting from the ferric salt. So, here there is an appreciable conversion of Fez+to Fe3+before reaching the stationary state and the reaction during the conversion phase is the “ferrous reaction’’ proper (see Fig. lb). Thus, the difference in the initial behavior of ferrous and ferric salts is simply due to the fact that different magnitudes of chemical change are required t o reach the stationary equilibrium as this is situated almost completely on the ferric side. This has been recognized by most of the workers in this field and Wieland (16) has called the Fe1’-H202reaction (leading to the stationary state) “Primiirstoss.” Abel (17) has disputed this: his objections which were based on a misunderstanding have been dealt with by Kolthoff and Medalia (13) and need not be considered here in detail. Apart from the qualitative evidence mentioned above which is in favor of a ferric *ferrous change, it has been shown recently (8) that an expression which is, under certain conditions identical with Andersen’s empirical Eq. (VIII) can actually be obtained from the free radical mechanism discussed above by taking into account the six Eqs. (l), (2), (4), (5), (6), and (7). Treating these equations for the stationary state of OH, HOz, and of the ferrous/ferric system one obtains:

(1x4

=0 =

kl[Fea+][H2O2] - kz[H2O2][0H]- k4[Fea+][OH]

-0

=

~Z[HZO~][OH] - k~,[Fe3+][0~-] - ks[Fe2+][H02] kr[FeS+][H02-]

-~ d(Fe2+) = 0

=

(IXb) kl[Fez+][HzO~] k4[Fez+][OH]- k~[Fe3+1[02-] ke[Fe2+1[HO~l

dt

d(Hoz) -dt dl

+

+

+

- kr[FeS+][HO~-] (IXc)

and for the rate of decomposition:

- d(HzOz) -= 21cl[Fe2+],[HzOzl dt

From these equations we can obtain an expression which contains only the total concentration of the iron salt f added originally as ferric

W

'fI

I I

I\

/

I

1

I I

1 I

I

I+'

------

Ferrous ions 8

[Fez+],

\I

I

I

I

0 ti TIME

a

FIG. 1. Schematic representation of the change of the concentrations of the ferrous and ferric ions in the reaction with hydrogen peroxide. a. Starting with ferric ions (initial concentration [Fe3+],) b. Starting with ferrous ions (initial concentration [Fez+],,) [Fe3+].and [Fez+]*denote the stationary concentrations.

FREE RADICALS IN THE REACTIONS

OF HYDROGEN PEROXIDE

351

+

salt. Thus, f = ([Fe3++l. [Fe2++I8) represents the sum of ferrous and ferric ions in the stationary state. from the above equations is somewhat The calculation of [Fe2++I8 lengthy, but taking into account the fact that, [Fe2+],<< [Fe3+I8,higher terms in [Fe2+],can be neglected and the resulting equations assume a relatively simple form. In this way we obtain eventually the following differential equation for the rate of decomposition:

(x

. . . concentration

of the hydrogen peroxide a t time t )

with 01

= 2- 9 k 4

[H+l

6 = prop f,

fl and

y

-constant

Equation (XI) can be easily integrated and gives:

(a

. . . initial concentration of the hydrogen peroxide, F

= (6/y))

Under the conditions in question, for the principal part of the reaction, the relations ( F / a ) and ( F / x ) < 1 hold, with good approximation, and thus the square roots can be developed up to the linear terms, which gives:

{S 3 + log -:

0.4343 - - -

=

'$. t

0.43401

(XIII)

This is identical with Andersen's empirical Eq. (VIII) in which his constants B and A are expressed as:

$

B = 0 . 4 3 4 ~ ~ = prop . f

W+I' ~

A

=

0.434F = prop f

With regard to the small constant, r , in Andersen's empirical equation (VIII), Bray and Peterson (18) have shown previously that it is without significance for the actual mechanism of the reaction. 111. THEREACTION BETWEEN HYDROGEN PEROXIDE A N D CUPRICIONS According t o Haber and Weiss (3b) the mechanism of the action of cupric ions corresponds closely to that of ferric ions so that instead of, for example, eqs. ( 5 ) , (6), and (7), one has: Cu*+ 02-3 Cuf 0 2

+ + CU'+ + HOz- F? CU+ + HOz

352

JOSEPH WEISS

The promoting effect of cupric ions in decomposing hydrogen peroxide by ferric ions (19) was considered by Haber and Weiss (3b) to be mainly due to the reaction: Cu+

+ FeS+ 7rt Cu2+ + Fez+

Recent work of Baxendale et al. (6) has confirmed theae views.

Iv. THE DECOMPOSITION O F HYDROGEN PEROXIDIC

AT

DIFFERENT

METALSURFACES The catalytic decomposition of hydrogen peroxide on solid and colloidally dispersed metals has been studied particularly by Bredig et al. (20). These reactions have been treated by Weiss (21) from the point of view of the radical mechanism. On the basis of the electron theory the metal can be considered both as a source and as a sink of electrons. I n conformity with the electron transfer processes outlined above, the reactions a t the surface of the metal catalyst can be considered to be represented by processes of the type: HZ02

+

+ OH+ +

~ m e t -+ s ~ OH HOz-4 HO,

02--+

0 2

€metal

€metal

On this basis the strong poisoning effect of certain substances (e.g., of KCN, Na& CO in the case of platinum or palladium) can be easily understood: it is likely that the metal ions on the surface can form (surface) coordination compounds with these inhibitors, and it is well known that in many of these coordination compounds closed electron shells are present, which, on account of their high ionization potentials would greatly reduce the occurrence of electron transf'er processes on the surface which appear to be essential for the catalytic activity of the metal. More recently Dowden (22) has considered the catalytic activity of metals on the basis of the electron theory from a more general and unified point of view. *fThe actual mechanism of decomposition on a metal surface can be described by the following reactions (S and S+ stand for the uncharged and charged parts of the metal surface respectively, for a given surface area: (S) (S+) = const. = u.

-

+

+ H2022S+ + OH- + OH H,Oa + OH 2 H20 + HO2 s+ + s+ S + HOa+ S+ + H0zs+ + H 0 2 - 2 S + HOn S

(SO (52)

ksr

02--+

0 2

kns

635)

353

F R E E RADICALS I N T H E REACTIONS O F HYDROGEN PEROXIDE

For the stationary state of the radicals and of the ratio (S+)/(S) we obtain the equations:

do dt = 0 = k,l(S)[HzO~l- kB~[H2021[0Hl ~- 0 = k,z[HZOz][OH] - kea(S+)[Oz-]- k,e(S)[HOz] + k,?(S+)[HOz-] at dt

+

= 0 = k.l(S)[Ho021 - kB6(S+)[Oa-I k.&3)[HOz]

(XIVa) (XIVb)

- ka~(S+)[HOz-1 (XIVcl

The rate constants and all the concentrations refer here to the adsorbed phase on the surface of the catalyst. Adding Eqs. (XIVa), (XIVb), and (XIVc) we obtain the simple relation: k,&3+)[HOa-] = k,~(s)[HOs]= const.

(XV)

This relation, which holds for the stationary state, was actually found experimentally by Wiegel (23) in the catalytic decomposition of hydrogen peroxide by colloidal silver. From the above equations one derives the following expression for the velocity of decomposition of hydrogen peroxide : where

Introducing the total analytical concentration of hydrogen peroxide present in the soIution at any given time ( t ) Eq. (XVI) becomes:

(2)

(XVII)

The reaction is monomolecular with regard to the concentration of the hydrogen peroxide in the surface phase. With regard to the kinetics it must be borne in mind that all the concentrations in the above equations as well as the dissociation constants KHOa(l) and KH20Z(,) refer to the adsorbed phase. Unless a linear adsorption isotherm is valid throughout fractional exponents may appear in these equations. One of the most striking features of the catalytic decomposition of hydrogen peroxide by noble metals is the fact that an optimum pH is generally observed for the specific rate of the decomposition (24). This can be derived from Eq. (XVII). The constant in Eq. (XVII) is a function of the pH and the optimum hydrogen ion concentration ([H+JOp,.) is found from the Eq. (XVII) by differentiation, viz. : [H+lopt. = hBKa,o2cs)

(XVIII)

354

JOSEPH WEISS

V. THE PHOTOCHEMICAL DECOMPOSITION OF HYDROGEN PEROXIDE

It is well known that hydrogen peroxide which absorbs fairly strongly in the region below 3100 A. undergoes photochemical decomposition. Urey, Dawsey, and Rice (25) suggested in 1929 that the photochemical primary process in the quartz ultraviolet was represented by: Hz02

+ hu

-+

20H

(11)

If this process is combined with the chain reactions (2) and (3), one should obtain, in conjunction with a process for the disappearance of the free radicals (chain breaking), an adequate representation of the photochemical reaction. I n fact, it has been shown that the experimental material, although incomplete in many respects, is in good agreement with such a mechanism. Kornfeld (26) found that the quantum yield for the decomposition is far greater than unity (values up to 50 were recorded) which clearly indicates the operation of a chain mechanism. The interaction of two radicals such as:

-

HOZ

+ OH-,ko

0 2

+ H20 + 73 kcal.

(9)

was suggested as a suitable chain-breaking process aEi early as 1934 (3b). A t very low radical and hydrogen peroxide concentrations and in the presence of impurities a first order disappearance such as: OH

+ X + chain breaking kio

(10)

(where X represents, for example, a surface or impurities) has t o be taken into account. Considering, e.g., Eqs. (11) (2), (3), and (9), one obtains for the stationary state ( I , = absorbed light intensity) :

do = 0 = 21, dt d(Ho2) = 0 dt

=

+

- k2[H2O21[0H1 k3~Hz0z1[02-l - ke[HO~l[OHl (XIXa)

kz[H2O21[OH1 - k3[H2021[Oz-I- k~[HOz1[0H1

(XIXb)

and for the rate equation (as a good approximation, i.e., when [HOZ]<<

and for the (differential) quantum yield y :

FREE RADICALS IN THE REACTIONS O F HYDROGEN PEROXIDE

355

These expressions are in general agreement with the experimental work in dilute solutions, with regard to the dependence on the light intensity and hydrogen peroxide concentration. Kornfeld’s results (26) are also dependence of the quantum yield, but not incompatible with a [H+]-$’” there is no general agreement regarding the pH dependence. The effect of the [H+] as well as the magnitude of the minimum quantum yield depend both to a very considerable extent on the nature of the chain-breaking process. In the case discussed above, ymio.= 2. This actually corresponds to the non-chain mechanism comprising the reaction 11, 2 and 9. According to Volman (27) these three reactions may play an important part in the photochemical decomposition of hydrogen peroxide in the vapor phase where a quantum efficiency 2 has been found. In Table I several chain-breaking approaching y processes are listed together with the corresponding expressionsfor the rate, quantum yield, and values of ymin.. It is an experimental fact that even at the highest light intensities employed so far, the quantum yield never 1.5. This falls below unity and mostly remains in the region of y result excludes any appreciable recombination of OH radicals to give HzOz according to Eq. (13). This is in agreement with the experimental findings of Bonhoeffer and Pearson (28) in the gaseous phase and is supported by theoretical considerations (29). However, it cannot be 1 to 2 might be due to a combination of ruled out that the cases of y chain-breaking processes such as Eqs. (9) and (13), although the fact that ymin.never falls below unity makes this rather unlikely. With quantum yields approaching unity the main chain breaking should be due to reaction (12), which by itself would also give a quantum yield independent of pH (see Table I). It is clear that in the cases considered above a dependence on pH is only to be expected under conditions where the quantum yield is appreciably greater than unity. There is therefore no obvious discrepancy between the earlier findings of Kornfeld (26) who worked under conditions of y > 2 and those of Heidt (30) and Lea (31). It was pointed out in 1934 (3b) that the interaction of HOz radicals according to

-

-

-

2H02 + HzO

+ 0s + 34 k d .

(14

might also take place and could be the source of the traces of ozone which are often observed in the decomposition of hydrogen peroxide. This, however, could also be due to the well-known reaction: 02

+ 0 + 0 s + 24 kcal.

with the oxygen atoms produced by reaction 12.

356 V

T

O

N

JOSEPH WEISS

C i

0

x T

Pt

T

0

x

2T

a a a

FREE RADICALS I N THE REACTIONS O F HYDROGEN PEROXIDE

357

A number of photochemical experiments using the sector method have been carried out mainly by Allmand and Style (32) and more recently by Lea (31). These experiments give lifetimes of 0.5 sec. to 1 sec. for the active radicals. However, there are some unexplained features in this work and experiments under well-defined conditions are still required for a detailed discussion of the lifetimes. In neutral or acid solutions the hydrogen peroxide is present in solution almost entirely as H202, while in alkaline solutions (pH > 12) appreciable amounts of the anion HO2- must be present.* This is also shown clearly by the change in the absorption spectrum (33). In the latter case the photochemical primary process is presumably represented by the electron transfer process: HOz-eHOH

+ h ~ - +HOz +OH- + H

(16)

At present the experimental material on the photochemical decomposition in alkaline solutions is insufficient to enlarge upon this point.

VI. THE DECOMPOSITION OF HYDROGEN PEROXIDE BY IONIZING RADIATIONS

It is well known that hydrogen peroxide decomposes under the influence of ionizing radiations (x-rays, y-rays, etc.), but hardly any quantitative data are available at present except for the decomposition by x-rays in dilute aqueous solutions (34,35). According to the theory put forward in 1944 the action of ionizing radiations on water produces H atoms and OH radicals according to the net process (36): HzO

-- + H

OH

(17)

If dilute solutions of hydrogen peroxide are irradiated by, for example, x-rays, the primarily formed radicals will be able t o react according to:

+ + +

HZOZ OH-, H 2 0 HZOZ H -+ HzO HZOZ H -+ HOz

+ HOZ+ 28 kcal. + OH + 65 kcal. + Ht + 13 kcal.

and the radicals can then enter into the chain reactions (2) and (3) or interact with each other as discussed above. Reaction (19) may be responsible for the small amounts of hydrogen which have been observed in the decomposition of hydrogen peroxide solutions by x-rays (34). A quantitative treatment of the decomposition of hydrogen-peroxide by ionizing radiations is still lacking. A recent preliminary investigation has shown, however, that this case presents a number of interesting features (37).

* K H ~ =o 2.4 ~ X

lo-'* (25°C.).

358

JOSEPH WEISS

In the case of ionizing radiations one has to consider the reaction in the tracks of the ionizing particles where the primary radicals are formed. However, in the case of relatively long lived active radicals (as indicated here by the photochemical sector experiments) many of these radicals will diffuse away from the original tracks. This can give rise to a pseudo-homogeneous distribution of the radicals in the solution leading to an appreciable amount of interaction between the individual tracks, and t o a considerable amount of general “inter-track” reaction. On the basis of these considerations we find that with certain simplifying assumptions, the ionic yield could be approximately proportional t o [HzOzlH [Dose rate]-% as was found experimentally by Fricke (35). No published data exist on the action of ionizing radiations on concentrated solutions of hydrogen peroxide. The primary processes are possibly represented by the equations (37) :

--

HeOz

followed by :

+

HzOz+ electron

+

HzOz++ H+ HOZ HZOz+ 4 OH+ OH

+

and

+ electron -+H2OzHzOz- -+ OH + OH-

H202

followed by :

HzOa- H

+ HOz-

VII. THEREACTION BETWEEN OZONEAND HYDROGEN PEROXIDE AND THE DECOMPOSITION OF OZONEIN AQUEOUS SOLUTION As pointed out above there is a close connection between HO1 radicals and ozone represented by the equation: 2H02-+ 08

+ HzO + 34 kcal.

(14

According to Weiss the reverse reaction is strongly catalyzed by hydroxyl ions according to (38) : Oa

+ OH- + HOz + 02-

(01)

A reaction which has also been studied spectroscopically in KOH (1m. to 7 m.) at temperatures between -3” and -30°C. When ozone is passed into concentrated solution of KOH at these temperakures the solution turns orange-yellow (38). This is due to the appearance of Oz- ions which are also responsible for the color of, for example, solid potassium peroxide, which, from magnetic and crystallographic measurements, is an ionic compound of the type K+.OZ- (39). In dilute aqueous solutions

F R E E RADICALS I N THE REACTIONS OF HYDROGEN PEROXIDE

359

the formation of the radicals according to Eq. (01) is followed by the chain reactions (38) :

+ H O z + 2 0 2 + O H + 31 kcal. + oz-+ 202 + 0O3 + OH + Oz 4-HOz + 39 kcal.

(or possibly) and

Oa

(02)

0 3

(02’)

(03)

Chain breaking could be brought about by interaction of the radicals, e.g., by reaction (14), or more probably by the reaction: HOz

+ OH+

Oz

+ HzO + 73 kcal.

(9)

Taking into account Eqs. (Ol), (02), (03), and (9), one obtains for the stationary state : d(Hoz) = 0 = 2kol[Oa1[OH-] - koz[O3][HO2] dt

d(oH) = 0 dt

=

+ ko3[Oo][OH]- k ~ [ H 0 ~ 1 [ 0 H(XXIIa) ]

koz[O3][HOZ] - k03[0a][OH]- ks[HOz][OH]

(XXIIb)

and under some plausible assumptions one obtains for the rate equation the relatively simple expression : d(od dt

- kol[O,lIOH-]

+ (2 @?)

[03]%[OH-l%

(XXIII)

the first term (const. [03][OH-])is due to the initiation reaction thesecond terms (const. [03]’*[OH-I5‘)comes from the chain reaction. Although Rothmund and Burgstaller (40) and also Sennewald (41) claimed t o have found that the reaction was of second order with regard to [Oa], Weiss (38) showed that the 35 order constants gave a somewhat better constancy, although the results are not wholly satisfactory. On the other hand there is almost general agreement that the reaction rate is proportional to [OH-If’. Recently Alder and Hill (42) have published some experiments which support a first order reaction with respect to ozone and proportionality with [OH-]”. This is not compatible with the mechanism proposed above as can be seen from the rate Eq. (XXIII). However, the mechanism proposed by Alder and Hill (42) also does not account for this result. Their mechanism which is represented by the equations:

+ Hz0 OaH+ + OH01H+ + OH- F? 2H02 0 s + HOz+ 202 + OH H01 + OH-+ 02 + HzO Oa

kHI

--*

360

JOSEPH WEfGS

leads to the following rate equation for the stationary state:

i.e., in the case of this simple sequence of reactions (which do not represent a chain mechanism) the initiating process is rate determining. In view of the appearance of OH and HOz radicals in the mechanism of the decomposition of ozone it is to be expected that, ozone in aqueous solution will also interact with hydrogen peroxide. The reaction between ozone and hydrogen peroxide was studied by Brodie (43) who described it by the equation: HzOz O3+ 202 HzO. This corresponds to a mean consumption ratio: f i = A(Os)/A(HzOz) = 1. Rothmund and Burgstaller (40) who investigated this reaction in greater detail found that this was only true in the case of relatively high concentrations of 20, could hydrogen peroxide. A t low concentrations values up to fi easily be obtained. The reaction between HzOzand ozone can be adequately described if one adds to the Equations (Ol), (02), (03), and Eq. (9) the reaction:

+

+

-

HiOn

+ OH+

Hz0

+ HOz

(2)

which occurs in the hydrogen peroxide decomposition. Treating these five equations for the stationary strite one obtains: dt

=0

-

+

+

2k01[03l[OH-]- k0~[0~l[HOzlk~a[Oal[OH] I;:Z[HZOZ~[OH] -ke[HO~][OH] (XXIVa)

I

do = 0 = ~OZ[OX][HOZ] - ko,[Os][OH] - ~Z[HZOZ][OH] - k,[H021[0H] at (XXIVb)

and from this one gets for the mean consumption ratio the approximate expression :

-

from which follows that if [H202] is sufficiently high fi 1. This reaction was studied later also by Taube and Bray (44), who applied the chain reactions of the ozone decomposition but replaced reaction (01) by the reaction: HzOz

+ Ox + OH + H0z + - 14 kcal. 0 2

(TI)

The experimental evidence does not allow any c1ea:r decision between reactions (01) and (TI), but nothing is gained by substituting the latter reaction which, furthermore, is rather endothermic.

FREE RADICALS IN THE REACTIONS OF HYDROGEN PEROXIDE

361

VIII. DETECTION OF FREEOH RADICALS Evans et al. (10) have shown that OH radicals produced by the Fenton reagent can initiate the polymerization of vinyl compounds and that the OH radicals, which in the first instance attack the double bond of the monomer, are built into the polymer chain. Stein and Weiss (45) have used the hydroxylation of benzene and of other simple aromatic compounds (e.g., benzoic acid and nitrobenzene) to detect these radicals. In the action of OH radicals on benzene in aqueous systems the formation of phenol and of diphenyl indicated a free radical mechanism of the following type : CaHs CeHo.

+ OH + OH

-+

--t

2CaHs. -+

and also: CeHs 4- CsHs. -+

It was found that the action of OH radicals on nitrobenzene or benzoic acid produced all the three isomeric phenolic compounds. OH radicals produced, for example, by reaction (1) can be used to generate other free radicals or atoms from inorganic or organic compounds. These reactions have been studied more recently by Kolthoff and Medalia (13), Merz and Waters (46), Stein and Weiss (45), and others. Reaction (1) also proved useful in the study of the so-called “active oxalic acid” (47). The free radical mechanism of hydrogen peroxide has been discussed also in connection with the mechanism of the action of the enzymes catalase and peroxydase, the prosthetic groups of which are iron porphyrin complexes which presumably also undergo oxido-reduction processes in the course of their catalytic activity (48). IX. SOMETHERMODYNAMICS DATACONCERNING THE RADICALS

OH

HOz

AND

In the foregoing discussion some of the heats of reaction of the free radical processes have been given t o indicate that these reactions are compatible with the available thermodynamics data. Apart from some well-established earlier values the following data :

+ +

+ 118 koal. + 100 kcal.

H O H 4 HzO 0 H 3 OH

have been taken from the recent work of Dwyer and Oldenberg (49)and 20H + HzOz

+ 53 k d .

(25)

362

JOSEPH WEISS

is the mean of the value 50.9 kcal. given by Holt, McLane, and Oldenberg (50) and the more recent value of 55 kcal. of Baxendale, Evans, and Uri (51). Combining the equations : 2(0

20H = H202 + 53 kcal. + H) = 20H + 2 X 100 kcal. O2 = 20

one obtains: Oa

- 118 kcal.

+ 2H = HZOZ+ 135 kcal.

The dissociation energy DH-oIof the HOz radical according to: H & - + H -I-oz

- DH-o~

-

(27)

cannot be regarded as settled. Recently Walsh (52) has put forward the rather high value of 70 kcal. while Sokolov’s considerations (53) lead t o the value of 35 kcal. For the present discussion the value 45 kcal. will be adopted, which has been calculated on the basis of the theory of Rumer and Heitler (54). This value in combination with Eq. (26) leads to the equation:

-

H

+ HOZ= H202+ 90 kcal.

(28)

I . Electron Afinities and Heats of Hydration

The sum [electron affinity ( E ) + h e a t of hydration ( S ) ] of the radicals OH and HOz and of the O2 molecule can be estimated from the following cycles : OH radical - V ~ ~ o ( 0 . 4e.v.1 6

HzOliq.

I+

+

-DE-OH(5.2 e.v.)

HzO,----+

Q ~ ~ o ( 0e.v.) .6

(Has.+

(H

-

- 1 ~ ( 1 3 . 6e.v.)

+ S ~ + ( 1 2 . 3e.v.)

OH,,.-) ___---

+SOH-

1

5

+ OH),

(H+

+ OH-),

This cycle has been used by several authors (55,56,57). With the values given above it leads to:

-

+

(EOE SOH-) 6.4 electron volts HOz radical

--DH-Ho,(-

-S~,o,(0.5 e.v.)

HzOz,,.

HzOzg ___--

--A

4.0 e.v.) ---+

(H -t HO2)g

I

--la(13.6 e.v.) + E E o ,

FREE RADIOALS I N THE REACTIONS O F HYDROGEN PEROXIDE

363

This cycle has been used by Weiss (55) and recently by Evans and Uri (57). With the value for D H - H O n given above it leads to: ( E E o ~ SHO*-) 5.40 electron volts

+

N

As a first approximation, a value can also be derived from the absorption limit of the HOz- ion in the near ultraviolet, if this is interpreted as an electron affinity spectrum as expressed by Eq. (16). Taking as absorp3200 A., which corresponds to an energy of 3.8 electron tion limit X volts, one obtains from Eq. (16) : @EOZ SRO,-) (EOE SOE-)- DK-OH Nhv (29)

-

-

-

+

+

+

With the values given above this yields: 0 2

-

+

(EHOI SHO~-)5.0 electron volts

molecule

-SH02(-

-DH-o,(-

0.5 e.v.)

HOzq. +

I+

* (H + 0 2 1 ,

- I ~ ( 1 3 . 6e.v.)

QHO*(- 0.1 e.v.)

(Haq.+

2.0 e.v.)

HOzg +9a+(12.3 e.v.)

O~aq.-)

+So,-

I

+Eo,

1

(H+

+

02-)g

With the values given above this yields: (Eo, 4-Soa-)

-

3.7 electron volts

The original values given by Weiss (55) in 1935 were all too high; this was almost entirely due to the low value used for the heat of hydration of the proton (SH+).The electron affinity of 0 2 can also be estimated by a Haber-Born cycle on the (ionic) alkali peroxides (M+.02-) as follows:

(M).

+

0 2

--*

~

(M+.Oz-)a

-

Calculations have been carried out both by Kazarnovskii (58) who finds a value of 22 kcal. and by Evans and Uri (57) who give a value of 16 kcal. The uncertainty lies here in the calculation of the lattice energy (V,),for example, of the KOz crystals. It is possible to estimate (Eon So,-) also from the long wave absorption limit, of the 0 2 - ion in aqueous systems if this is interpreted as an electron affinity spectrum, i.e.,

-

+

02-.HOH

-

+ hv+

0 2

+ OH- + H

(30)

According to the spectrum determined by Weiss (54) the long wave 5200 A. which corresponds to an energy of absorption limit lies at X 2.3 volts.

-

364

JOSEPH WEISS

One can then obtain an estimate from the following equation: (Eel+ Sot-)

+ DH-OH- (EOH+ SOH-)+

(31)

if one neglects any effects due to the delayed orientation of the water dipoles and uses the full energy of hydration of OH- ion. This may not cause a very serious error here since we are only dealing with one singly charged (negative) ion both in the initial and in the final state. This leads t o the value: (Eol

+ Sot-)

-

3.5 electron volts

which is not very different from the value obtained above and by other authors.

REFERENCES 1. Haber, F., and Willstatter, It., Ber. 64,2844 (1931). 2. Franck, J., and Haber, F., Sitzber. preuss. Akad. Wiss. 250 (1931). 3a. Haber, F., and Weiss, J., iVaturwissenschaften 20, 948 (193'2). 3b. Haber, F., and Weiss, J., Proc. Roy. SOC.(London) A147, 332 (1934). 4. Manchot, W., and Lehmann, G., Ann. 460, 179 (1928). 5. Weiss, J., Nolurwissenschaflen 23, 64 (1935). 6. Barb, W. G., Baxendale, J. H., George, P., and Hargrave, EL R., Nature 163, 692 (1949). 7. Humphrey, C. W., and Weiss, J., Nature 163, 691 (1949). 8. Humphrey, C. W., and Weiss, J. In preparation. Weiss, J., h'xperientia. 7 , 135 (1951). 9. Schmid, H., 2. physik. Chem. A148, 321 (1930). 10. Baxendale, J. H., Evans, M. G., and Park, G. S., Trans. Faraday SOC.42, 155 (1946). 11. v. Bertalan, J., 2. physik. Chem. 96, 328 (1920). 12. cf. Bray, W. C., Chem. Revs. 10, 170 (1932). 13. Kolthoff, I. M., and Medalia, A. I., J . Am. Chem. SOC.71, 3777, 3784 (1949); J. Polymer Sci. 4, 377 (1949); 6, 391 (1950). 14. Andersen, V. S., Acta Chern. Scand. 2, 1 (1948). 15. Simon, A., Haufe, W., Reetz, Th., and Preissler, R., 2. anorg. u. allogem. Chem. 230, 129 (1936). 16. cf. Wieland, H., and Francke, W., Ann. 467, 1 (1927); 476, 1 (1929). 17. Abel, E., Monatsh. 80, 776 (1949). 18. Bray, W. C., and Peterson, S., J . A m . Chem. Soe. 7 2 , 1401 (1950). 19. Bohnson, V. L., and Robertson, A. C., J . A m . Chem. SOC.46, 2493 (1923). 20. Bredig, G., and von Berneck, Miiller, 2. physik. Chem. 31, 258 (1899). Bredig, G., and Ykeda, K., Z. physik. Chem. 37, 1 (1901). Bredig, G., and Teletov, J., 2. EZektrochem. 12, 581 (1906). 21. Weiss, J., Trans.Faraday SOC.31, 1547 (1935). 22. cf. Dowden, D. A., Discussions Faraday SOC.,1950. 23. Wiegel, B., 2. physik. Chem. 143A,81 (1929). 24. cf. Bredig, G., and Reinders, W., 2. physik. Chem., 37, 323 (1901). Bredig, G., and Fortner, M., Ber. 37, 798 (1904). Rius, A,, 2. Elektrochem. 36, 149 (1930).

FREE RADICALS IN THE REACTIONS

OF HYDROGEN PEROXIDE

365

Urey, H. C., Dawsey, L. H., and Rice, F. O., J . Am. Chem. SOC.61, 1371 (1929). cf. Kornfeld, G., 2. physik. Chem. B29, 205 (1935). Volman, D. H., J . Chem. Phys. 17,947 (1949). Bonhoeffer, K. F., and Pearson, T. G., 2. physik. Chem. B14, 1 (1931). Weiss, J., Trans. Faraday SOC.36,856 (1940). Heidt, L. J., J . Am. Chem. SOC.64,2840 (1932). Lea, D. E., Trans. Faraday SOC.46,81 (1949). Allmand, A. J., and Style, D. W. G., J . Chem. SOC.698,606 (1930). Bredig, G., Lehmann, H. L., and Kuhn, W., 2. anorg. u. allgem. Chem. 218, 16 (1934). 34. Risse, o., 2. physik. Chem. 140A, 133 (1928). 35. Fricke, H., J . Chem. Phys. 3, 364 (1935). 36. Weiss, J., Nature 163, 748 (1944); Trans. Faraday SOC.43, 314 (1947). 37. Weiss, J., in preparation. 38. Weiss, J., Trans. Faraday SOC.31, 668 (1935). 39. cf. Ann. Repts. Progress Chem. (Chem. SOC.London) 44, (1947)62 ff. 40. Rothmund, V., and Burgstaller, A., Monatsh. 34, 668 (1913); 38, 159 (1917). 41. Sennewald, K., 2. physik. Chem. A164, 305 (1933). 42. Alder, M. G., and Hill, G. R., J . Am. Chem. SOC.72, 1884 (1950). 43. Brodie, B. C., Phil. Trans. Roy. SOC.(London), 162, 454 (1872). 44. Taube, H., and Bray, W. C., J . Am. Chem. SOC.62, 3357 (1940). 45. Stein, G., and Weiss, J., J. Chem. SOC.3245, 3254 (1949); Nature 166,1104 (1950). Loebl, H., Stein, G., and Weiss, J., J. Chem. SOC.2074, 2704 (1949). 46. Merz, J. H., and Waters, W. A., J. Chem. SOC.2474 (1949). 47. Weiss, J., “The Labile Molecule,” Discussions Faraday SOC.2, 188 (1947). 48. Weiss, J., J . Phys. Chem. 41, 1107 (1937). 49. Dwyer, R. J., and Oldenberg, O., J . Chern. Phys. 12, 357 (1944). 50. Holt, R. B., McLane, C. K., and Oldenberg, O., J . Chem. Phys. 16,229 (1948). 51. Baxendale, J. H., Evans, M. G., and Uri, N., Trans. Faraday SOC.37, 236 (1949). 52. Walsh, A. D., J . Chem. SOC.331 (1948). 53. Sokolov, N., Acta Phgsicochim. U.R.S.S. 19,208 (1944). 54. cf. Weiss, J., Trans. Faraday SOC.31, 668 (1935). 55. Weiss, J., Trans. Faraday SOC.31, 966 (1935). 56. Baughan, E. C., Evans, M. G., and Polanyi, M., Trans. Faraday SOC.37, 377 (1941). 57. Evans, M. G., and Uri, N., Trans. Faraday SOC.46, 224 (1949). 58. Kazarnovskii, I. A., Doklady Akad. Nauk. S.S.S.R. 69,67 (1948); J. Phys. Chem. (U.S.S.R.) 14,320 (1940).

25. 26. 27. 28. 29. 30. 31. 32. 33.

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The Specific Reactions of Iron in Some Hemoproteins PHILIP GEORGE Department of Colloid Science, University of Cambridge, England

CONTENTS Page .............................................. 367 bin, Myoglobin, Peroxidase, and Catalase.. 369 ...................................... 369 1. Structure. . ematin . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 372 3. Compounds of Hemoglobin, Peroxidase, and Catal . . . . . . . . . . . 374 111. Hemoglobin and Myoglobin Autoxidation and Other 5 . . . . . . . . . . . 381 IV. Peroxidase Reactions. . . . . . ... V. Catalase Reactions.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 393 1. Catalytic Decomposition of Hydrogen Peroxide. . . . . . . . . . . . . . . . . . . . . . 393 2. Complexes of Catal and Alkyl Hydro. . . . . . . . . . . . . . . . . . 398 peroxides. . . . . . . . . . 3. Catalase in Coupled . . . . . . . . . . . . . . . . . . 400 4. Reaction Mechanisms with the Catalase-Peroxide Complexes. . . . . . . . . . 402 VI. The Mechanism of These Hemoprotein Reactions. . . . . . . . . . . . . . . . . . . . . . 404 1. Free Radical Mechanisms with Ferrous and Ferric Ions.. . 2. An Examination of the Possibility of Similar Free Radical Reactions with the Hemoproteins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 411 3. The Applicability of These Free Radical Mechanisms in the Hemoprotein Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 415 a. Reactions of Peroxidase and Catalase. . . . . . . . . . . . . . . . . 415 b. Oxidative Reactions of Hemoglobin and Myoglobin Involving Oxygen 420 VII. Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 424 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 425

I. Introduction. ....

I. INTRODUCTION I n the editorial preface to the first volume of Advances in Catalysis the decision was made known not to publish reviews of specialized topics in biocatalysis but from time to time to bring reports in which the relationship and parallelism between this special field and “normal ’’ catalysis are discussed. This is the first of these reports. Its purpose is to examine the reactions of four hemoproteins, hemoglobin, myoglobin, peroxidase, and catalase, which all contain the same coordination compound of iron-ferrous or ferric protoporphyrin attached to different protein molecules, with oxygen, hydrogen peroxide, and in a few cases additional reducing substances. Some of these reactions are specific; 267

368

PHILIP GEORGE

that is, they are either more highly developed, or developed to an extent that they can be regarded as unique reactions, when compared with the corresponding reactions of the other three hemoproteins or the reactions of ionic iron. These specific reactions may be listed as follows: Hemoglobin and Myoglobin. Ferrous compounds which combine reversibly with molecular oxygen, e.g., Hb

+ 02

HbOz

Both association and dissociation reactions are extremely rapid. The oxidation to the corresponding ferric compound is in comparison a very slow reaction. Perozidase. A ferric compound which utilizes hydrogen peroxide to oxidize reducing substances according to the general ;scheme:

+ +

Peroxidase HzOl -+ Peroxidase . [H&] Peroxidase [H202] AH1 -+ Peroxidase A -t2Hz0

+

which has been named “peroxidatic activity.” Peroxidase can be reduced by strong reducing agents to the ferrous compound. Catakse. A ferric compound whose biological function may be the oxidation of alcohols or other substrates by hydrogen peroxide in an analogous reaction to peroxidase and which is the most efficient catalyst known for the decomposition of hydrogen peroxide into water and molecular oxygen at room temperature. This has often been called “catalatic” activity. It is a most remarkable feature of catalase that its ferrous form cannot be prepared directly using strong reducing agents or electrolytic reduction. A ferrous compound is formed in the presence of hydrogen peroxide and azide ions, but its relation to the true ferrous form or whether in fact it is the ferrous form corresponding to the ferric form of the free enzyme is not known. Hemoglobin and myoglobin in their ferric forms show rudimentary peroxidatic and catalatic activity, but ferrous peroxidase does not combine reversibly with molecular oxygen. Ionic iron also gives the hydrogen peroxide reactions but not the combination with oxygen. Many experimental studies now support a free radical mechanism for these reactions of ionic iron, and it is the main theme of the report to examine the data on the hemoprotein reactions to see whether similar reaction mechanisms are first of all possible and if so whether there is evidence that the reactions do proceed in $his way. It is unavoidable that many topics of great biochemical importance and physico-chemical interest must be omitted in a, report of this kind. Two recent reviews by Wyman (1) and Theorell (2), articles in Respiratory Enzymes (3) and Haemoglobin (4) by various authors, and the

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

369

recent monograph Haematin Compounds and Bile Pigments by Lemberg and Legge ( 5 ) deal in great detail with these other themes. I n addition these sources review the very extensive biochemical investigations on the hemoprotein reactions discussed in this report and so reference is often made to these sources where parallel references may be found. A short section summarizing the relevant properties and reactions of iron protoporphyrins and the four hemoproteins precedes the detailed account of their reactions with hydrogen peroxide and molecular oxygen.

11. GENERALCHEMISTRY OF HEMOGLOBIN, MYOGLOBIN, PEROXIDASE, AND CATALASE 1. Structure

All four hemoproteins have the same reactive group, iron protoporphyrin :

where R = -CH8

;J-fR R

R Ri z

qN.

= -CH&HzCOOH -CH=CHz

7 \

R2

R2

The assignment of the double and single bonds in this figure is purely formal, and it is well established from X-ray studies of the structurally similar phthalocyanines that these molecules are resonance hybrids. Even in free phthalocyanine where two of the four pyrrole nitrogen atoms carry hydrogen atoms the interatomic distances in the inner 16-atom ring are almost identical, having values of 1.33 A. or 1.34 A. The porphyrin forms its iron compound by the replacement of the two pyrrole hydrogen atoms. Thus the ferrous derivative, called heme, has no ionic charge. The ferric derivative has a single positive charge; as the chloride it is called hemin and as the hydroxide hematin. It will appear later that the net charge on the iron atom plays an important part in determining what type of addition compound the molecule forms. These formal charges on the iron atom may not necessarily represent the charge on the molecule as a whole, for in the case of the free molecule the ionization of both propionic acid side chains could clearly result in heme carrying two negative charges and the ferric compound being no longer a chloride but a zwitterion carrying an additional negative charge. There is good evidence, however, particularly in the case of peroxidase, that in the hemoproteins these side chains may be concerned in the linkage of the iron protoporphyrin to the protein molecule, and the contribution

370

PHILIP GEORGE

of the iron protoporphyrin to the total charge in the molecule as a whole depends on the precise nature of this linkage. The four hemoproteins, hemoglobin, myoglobin, peroxidase, and catalase, differ only in the protein to which the iron protoporphyrin is attached. Molecular weight determinations give values 68,000, 17,000, 44,000, or 40,000 (Keilin and Hartree, 48) and 225,000 respectively. Analytical data show hemoglobin and catalase to have four heme or hematin groups per molecule whereas myoglobin and peroxidase have only one. The recent experimental data on the size and shape of these hemoproteins has been reviewed by Wyman (1). Not only do the four hemoproteins differ in molecular weight but also in the relative proportions of their constituent amino acids. Only for. hemoglobin and myoglobin are nearly complete analytical data available. Tristram (4) has recently summarized the results of many different workers. Both proteins are made up of about twenty different amino acids. In units of weight per cent, leucine and isoleucine 16.:3%, glutamic acid 16.48%) lysine 15.5%) histidine 8.50/,, aspartic acid 8.2% and alanine 7.95% predominate in myoglobin, six amino acids accounting for 73.4% of the molecule. Hemoglobin has a rather more varied structure; eight amino acids account for 75.75% of the molecule: leucine 15.4%) aspartic acid 10.6%) valine 9.1%, histidine 8.7%, lysine 8.5%) glutamic acid 8.15%, phenyl alanine 7.9%, and alanine 7.4%. In one respect there is a marked difference between these two proteins: myoglobin contains no cystine or cysteine, hemoglobin contains enough t o give a maximum of three disulfide links per molecule, but since -SH groups are known to be present there cannot actually be more than two disulfide links. The analytical data obtained by Theorell and Akeson (6) for peroxidase and catalase respectively enabled them to calculate the lysine, arginine, and histidine contents (see also Bonnichsen, 7). These results are summarized in Table I together with the relevant values for hemoglobin and myoglobin in terms of amino acid per molecule (R/M), and the TABLE I Typical Amino Acid Contents

Hemoprotein

Histidine R/M R/68,000

A rginine R/M

R/68,000

2 14 18 27

8 14 28 8

R/M ~

Myoglobin Hemoglobin Peroxidase Catalase

9 36 3 15-16

36 36 5 5

Lysine R/68,000

~~

18 38 12 32 ~~

Note. R / M Number of residues per molecule. R/68,000 Number of residues assuming molecular weights all to be 68,000.

72 38 19 10

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

371

equivalent number of residues if the molecular weight of all four hemoproteins was 68,000 to facilitate comparison (R/68,000). The total number of residues for a molecular weight of 68,000 is of the order of 580. These data illustrate very well the differences in composition. Histidine and lysine are present in hemoglobin and myoglobin in relatively high amounts but not in peroxidase and catalase. The arginine content of peroxidase is higher than in the other three proteins. A further distinguishing feature of peroxidase is that it contains a high percentage (18.4% of insoluble hydrolyzate) of substances which are not amino acids (Theorell and Akeson, 6). These have been tentatively identified with acid carbohydrate molecules in accord with the observation that peroxidase gives a positive result in the Molisch reaction, and that the isoelectric point is lower than would be expected from a predominance of basic amino acids. X-ray studies of hemoglobin and myoglobin have yielded very valuable information on the size and shape of these molecules. Reference may be made t o the review articles by Perutz (4) and Kendrew (4). The hemoglobin molecule can be pictured in an idealized form as a cylinder of 57 A. diameter and 34 A. in height containing four layers of scattering matter. Myoglobin is similar in shape and dimensions to one of the four layers of hemoglobin: it consists of a disk about 9 A. thick with other dimensions not greater than 57 A., the most probable values being 51.5 by 37 A. The Link between the Prosthetic Group and the Protein. I n all the four hemoproteins, hemoglobin, myoglobin, peroxidase, and catalase, the iron protoporphyrin group is joined to the protein by a coordinate bond to the iron atom. This may be regarded as occupying the fifth coordination position about the iron atom which then binds a water molecule or OH group to complete a stable octahedral coordination unit (Keilin and Hartree, 112). There is some evidence that the protoporphyrin side chains are also involved for not all iron porphyrins combine with the free protein when obtainable to give active compounds. Examples are given below. The observation that hemoglobin is split into heme and free globin in weakly acid solution and the fact that globin contains a high proportion of basic amino acids suggests that the hemes are bound to basic groups in the protein. This reaction is reversible and the reconstituted hemoglobin shows all the properties of the original hemoglobin except for slight spectroscopic differences (Jope, 4). Other ferrous porphyrins have been found to combine with free globin to give compounds which still give the reversible oxygenation reaction : These are mesoheme, hematoheme, and diacetyldeuteroheme containing ethyl, hydroxy ethyl

372

PHILIP OEORQE

and acetyl groups in place of the vinyl side chains of ferrous protoporphyrin. Further discussion of this topic is beyond the scope of this review for it involves the effect of pH and carbon dioxide on the hemoglobin oxygen reaction and the determinations of the ionizable groups on hemoglobin, oxyhemoglobin, and methemoglobin by various titration procedures (Wyman, 1). The results of these studies mainly favor the view that the heme is joined to a histidine group, but a linkage of this type cannot account for the effect of carbon dioxide (Roughton, 8). There is very little evidence on the nature of the linkage in myoglobin and catalase. All attempts to split the hematin group off catalase reversibly have failed. In the case of peroxidase the hematin can be removed by treatment with an acetone-hydrochloric acid mixture a t -5" t o -10°C. and the enzyme reformed by treating the colorless protein with alkaline hematin. Replacing the hematin by deuterohematin or mesohematin which have hydrogen atoms or ethyl groups instead of vinyl groups gave an enzyme with 63% and 57% of the original activity. Hematohematin with a-hydroxy ethyl groups gave no activity (Theorell, 2; see also Theorell and Maehly, 113, and Maehly, 114). Trials with other ferric porphyrins showed that the two propionyl side chains are essential and that these must occupy the positions they do in protoporphyrins. On the basis of differential titration data obtained for peroxidase itnd its free protein Theorell has discussed the possibility that the hematin group is joined to a carboxyl group of the protein (2,6). When heme or hematin are joined t o the various proteins in hemoglobin, peroxidase, and catalase, they combine with a larger range of compounds to give stable octahedral coordination complexes than they do by themselves. To illustrate this the chief derivittives of heme and hematin are listed below before proceeding to describe the hemprotein complexes. 2. Compounds of Heme and Hematin

Both heme and hematin show a marked tendency to aggregate in aqueous solution which has complicated many of the experimental investigations of their reactions. Four of the six valencies of the central iron atom are held in the planar porphyrin ring, and when heme and hematin are in true solution the fifth and sixth bonds completing an octahedral complex are assumed to be occupied by water molecules. These compounds, i.e., heme and hemin, can be represented as H20.Fe,.H20 and H20.Fe,+.H20, where Fe, and Fe,+ represent ferrous and ferric protoporphyrin respectively. Ionization of one of the water molecules attached to heme in solution

T H E SPECIFIC REACTIONS O F I R O N I N SOME HEMOPROTEINS

373

has not been observed, but in the case of hematin a pK of 7.6 has been measured for the ionization (Shack and Clark, 9) : H20,Fe,+Hz0

H20.Fe,.0H

Hemin

+ H+

Hematin

In strongly alkaline solution, however, there is good evidence that heme forms a compound with two hydroxyl groups [H0.Fep.0H]2-, (Keilin, 10). Heme but not hemin or hematin reacts with carbon monoxide giving a compound presumably H20.Fe;CO; it also combines with methyl isocyanide. Hematin but not! heme gives a complex with hydrogen peroxide which Haurowitz represents as H20.xFe;OH (11; see also Haurowitz, Brdicka, and Kraus, 12). A great variety of nitrogenous bases combine with heme including pyridine, nicotine, a-picoline, imidazole and its derivatives e.g. 4 methylimidazole, piperidine, methylamine, and ammonia. These compounds, which can be formulated as B.Fe,.B, are called hemochromogens. Many denatured proteins form hemochromogens particularly denatured globin. One of the molecules of the base can be replaced by carbon monoxide giving the compound B.Fe,CO. Cyanide ions react with heme forming the complex [CN.FeP.CNl2-,and hemin yields [CN.Fe,.CN]-. The same nitrogenous bases that react with heme also react similarly with hemin but have in general a much lower affinity. These compounds are called parahematins and with the corresponding hemochromogens form a well-known redox system. B.Fe,.B

[B.Fe,.B]+

Hemochromogen

+ electron

Parahematin

The potentials of these systems will be discussed later in connection with estimates of the ionization potential of the hemochromogen in aqueous solution. The effect of increasing alkalinity on these systems is somewhat complex for there is evidence in some cases that an equilibrium of the following type can occur. [B.Fe,.B]+

+ OH-

B.Fe,.OH

+B

These topics are discussed in greater detail by Lemberg and Legge (5) and Mansfield Clark (13). Like hemin the parahematins react with hydrogen peroxide. For pyridine parahematin Haurowitz, Brdicka, and Kraus (12) have suggested that the following replacement reactions occur : [Py.Fe,.Py]+Cl-

HzOz

HlOZ

[H202.FeP.Py]+C1--+ [H202~Fe,~Ha02]+C1-

374

PHILIP GEORGE

This behavior may be contrasted with t h a t of the hemochromogens where there is no evidence for similar replacement reactions, but considerable oxidative attack on the porphyrin ring occurs giving compounds named by Lemberg, Cortis-Jones, and Norrie oxyporphyrin hemochrome and verdohemochrome (14).

I n the former a methene group ‘CH

/

>COH and in the latter the methene group is replaced by

>

becomes

0 and an

adjacent carbon atom on a pyrrole ring carries an OH group.

3. Compounds of Hemoglobin, Peroxidase, and Catalase

One of the most characteristic features of these hemoproteins when compared with many other coordination complexes of iron is the very great stability of the protein-iron protoporphyrin unit. I n none of the normal reactions of these hemoproteins, where the protein remains in its native state, is there any evidence of t h e link being broken. The origin of this stability is unknown but would appear t o arise from t h e dissociation process being extremely slow. Indirect evidence of this is provided by the case of peroxidase where the combination of hematin with the free protein is itself a slow reaction yet the compound is very stable. The great stability of the protein-iron protoporphyrin unit enables these hemoproteins t o form a wide variety of coordination complexes in straightforward bimolecular reactions involving replacement of the water molecule or OH group attached t o the sixth iron valency. The evidence that a water molecule or OH group is attached in this way is largely indirect (Keilin and Hartree, 112). The best evidence is in the case of hemoglobin. Haurowitz has shown t h a t when reduced hemoglobin is dried there is a very marked change in the absorption spectrum and the single diffuse band a t 555 mp is replaced by two sharp bands typical of a hemochromogen (4). Addition of water restores the reduced hemoglobin unchanged. Oxyhemoglobin does not show this behavior, its characteristic color and absorption spectrum are identicad in solution and in the dry state. Furthermore, dry oxyhemoglobin does not dissociate into hemoglobin and oxygen a t low oxygen pressures, but does if water is added. The essential oxygenation reaction is thus seen t o be: Globin.Fe,.(HnO)

+ + Globin.Fe,.Oz + H 0 2

2 0

where globin.Fe, represents one of the four heme groups of the hemoglobin molecule. Haurowitz obtained similar results with myoglobin and reconstituted hemoglobins containing protoheme, mesoheme, and the dimethyl ester of mesoheme. The hemochromogen formed when hemoglobin is dried was shown t o be a n intermolecular compound in

THE SPECIFIC REACTIONS O F IRON I N SOME HEMOPROTEINS

375

which a basic group of one molecule is joined t o the heme group of another, i.e., globin.Fe;globin.Fe. * . . , by the absence of any hemochromogen formation when glucose is present in the solution. It was t o be expected that the multiple polar groups of the glucose molecule would adsorb on the surface of the hemoglobin molecule while in solution and thus prevent the mutual association on drying. There is no similar evidence as yet for a water molecule attached t o the iron atom in methemoglobin but the change in color from brown t o red which occurs as a solution is made more alkaline is best explained by the ionization of such a water molecule. [Globin.FeP+H2O]S Globin.Fe,.OH

+ H+

The alkaline methemoglobin formed has a magnetic susceptibility of 4.45 Bohr magnetons compared with 5.77 for methemoglobin in moderately acid solutions. This marked difference is indicative t h a t the group ionizing is associated with the iron atom. Similar p K values of 8.10 a t 30°C. are obtained for the ionization by both spectroscopic and magnetic determinations (see Lemberg and Legge, 5 ) . I n the case of peroxidase the reduced form is assumed to have a water molecule attached, by analogy with hemoglobin. The oxidized form, according t o titration data obtained for the free protein and the reconstituted enzyme and a study of the dissociation of the fluoride complex, undergoes ionization of the water molecule with a p K of 5.0 (Theorell and Paul, 15; and Theorell, 115). There is no marked spectroscopic change, however, and this leaves unexplained a n ionization with a p K 10-11 which is accompanied by a change in color and magnetic properties. Lemberg and Legge have discussed these results fully ( 5 ) , and it seems more probable that the ionization of the water molecule corresponds t o this second pK. As was mentioned in the introduction, catalase is unique among the hemoproteins because its reduced form is unknown; all attempts t o prepare it from the free enzyme including electrolytic reduction have failed. On the other hand Keilin and Hartree (16,17) have found that in the presence of azide ions reduction does occur when hydrogen peroxide is added. The nature of this reaction is unknown. It may be a consequence of catalase possessing four hematin groups per molecule and that azide combined with one group alters the redox properties of the others in a manner analogous t o heme-heme interaction in the hemoglobin-oxygen reaction. This is unlikely since azide catalase like free catalase is resistant t o other reducing agents, even sodium hydrosulfite. An alternative explanation based on the azide ion undergoing oxidation in the reaction and t h e stabilization of the ferrous form of catalase by a

376

PHILIP GEORGE

reaction product of the azide ion has no experimental support, for no destruction of azide has been detected (Foulkes and Lemberg, 18). The question whether a water molecule is bound in this reduced form of catalase has clearly little meaning until the detailed chemistry of this reaction is known. Catalase itself has been shown by Agner and Theorell (19) t o have a hydroxyl group attached t o the iron protoporphyrini group with a pK of 3.8. The evidence for this rests on changes in light absorption on addition of different anions such as phosphate, acetate, and formate and the inhibition of catalase activity by these anions increasing as the hydroxyl ion concentration is decreased. Hemoglobin and myoglobin, peroxidase, and catalase combine with a very wide range of compounds in simple bimolecular reactions involving the replacement of the water molecule or OH grou:p. In general the ferrous compounds show preferential reaction with neutral molecules and the ferric compounds with anions, as can be seen for the typical compounds listed in Tables I1 and 111. The most ch:zracteristic properTABLE I1 Ferrous Compounds

Reagent

None 0 2

co NO CNEtNC Ph.NO

Hemoglobin

m.i. s.d., m.c. s.d., m.c. s.d., m.c. s.d., m.c. s.d., m.c. s.d.

Reduced Peroxidase

m.i.

-

s.d., m.c.

Spectroscopically similar to parent compound. 8.8. s.d. Spectroscopically different to parent compound. m i . or m.c. Bonds to iron atom shown to be essentially ionic or ,covalent by magnetic ausceptibility measurements. A dash indicates that the compound is not formed, a blank space that it has not been investigated.

Note.

ties of these compounds are their color, type of absorpt8ionspectrum, and magnetic susceptibility. The six bonds around th.e iron atom are predominantly ionic in the parent compounds as shown by the susceptibility values which correspond very closely to th.e expected values for four and five unpaired electrons in the ferrous and ferric parent compounds respectively. Those compounds in which the bonding is still ionic have absorption spectra and colors similar t o the parent compounds. The formation of many compounds, on the other hand, results

THE SPECIFIC REACTIONS OF IRON I N SOME HEMOPROTEINS

377

TABLE I11 Ferric Compounds Reagent

Methemoglobin

Peroxidase

Catalase

None OHAnions FCNHSN 1NHzOH N2H4 NO NH 8 EtOH Imidarole Cyanate Thiocyanate

m.i. s.d., m i . &

m.i. s.d.

m.i. parent cpd.

s.s., m.1. s.d., m.c. s.d., m.c. s.d., m.c.

s.s., m.i. s.d., m.c. s.d., m.c.

s.s., m.i. s.d. s.d., m.c. s.s., m.1.

9.8.

S.S.

9.9.

8.5.

9.8.

s.d. m.c. m.i. s.d., m.c.

s.d.

s.d. 9.9.

S.8.

S.S.

Note. 8.8. Spectroscopically similar to parent compound. s.d. Spectroscopically different t o parent compound. m i . or m.c. Bonds to iron atom shown t o be essentially ionic or covalent by magnetic susceptibility measurements. a Alkaline methemoglobin has an anomalous susceptibility of 4.47 Bohr magnetons corresponding most closely to the value for three unpaired electrons (see Pauling, 4).

in a large decrease in magnetic susceptibility corresponding to no unpaired electrons in the ferrous compounds and one unpaired electron in the ferric compounds. These are the values expected for covalent octahedral iron complexes. This susceptibility change is usually accompanied by a profound alteration in color, type of absorption spectrum, and the intensity of the absorption bands which is usually enhanced. Tables I1 and 111indicate those compounds which are spectroscopically similar (s.s.) and those which are spectroscopically different (s.d.) from the parent compounds, and in addition whether magnetic measurements have shown the compounds to contain essentially ionic (m.i.) or covalent (m.c.) bonds around the iron atom. Table IV which is based on a general classification suggested by Theorell (20) and Hartree (21) shows the predominant correlation between color and spectrum type. I n addition t o the bands in the visible region of the spectrum all these compounds have a very intense band between 380 and 440 mp, in the so-called Soret region, where position and intensity is characteristic for each compound. I n some spectroscopic analyses it has proved more satisfactory t o work at these wavelengths. Hartree has recently reviewed the magnetic properties of hematin derivatives (21) and a detailed

378

PHILIP GEORGE

TABLE I V General Correlation between Bond T y p e , Color, and Absorption Spectrum (Weaker absorption bands are indicated by brackets; the wavelengths are given in millimicrons) Color

Spectrum Type

Examples with the Band Maxima

Ionic ferric compounds

Green-brown

Absorption band in MetHb; 637, (582), (548), red between 600 504 and 640 mp; strong Catalase: 639, (544), 506 band in blue some- Fluoride catalase: (622), 597, times faint bands (542) in the green Aside ca.talase: 624, (581), (567), (.536), 494 Peroxidme: 645, 583, 548, 498

Covalent ferric compounds

Bright red

Two diffuse bands in the green

Aside MetHb: 575, 542 HS. MetHb: 570, 545 Cyanide peroxidase: 581.5, 542

Ionic ferrous compounds

Carmine-red to purple

Diffuse band in the green

Hb: 555

Covalent ferrous compounds

Scarlet to pink

Two very sharp bands in the green

HbOz: 5‘75, 540 HbCO: 670, 540 HbNO: 568, 531 CO peroxidase: 578, 545.5

Note. Alkaline methemoglobin has the spectrum typical of a covalent ferric eornpoiind y e t its magnetic susceptibility value is anomalous (Table 111).

description of all the various compounds with references t o the original investigations is given by Lemberg and Legge ( 5 ) . Mention has been made above that the formation and properties of these compounds show them t o be octahedral coordination complexes in which the bonds t o the central iron atom are either essentially ionic or essentially covalent. There is, however, no correlaiion between the type of bond being formed in the complex and the speed of the reaction. The kinetic study of the reaction between hemoglobin and oxygen or carbon monoxide prompted the development of the rapid flow ” technique by Hartridge, Roughton, and Millikan for following fast reactions in solution (see Roughton, 4). These kinetic investigations showed both combination and dissociation reaction t o be very rapid. A comparison of velocity constants for hemoglobin and myoglobin reacting with O2 and CO is given in Table V based on data compiled by Millikan (22). ((

THE SPECIFIC REACTIONS OF IRON I N SOME HEMOPROTEINS

379

TABLE V Velocity Constants for Hemoglobin and Myoglobin Reactions at 2O"C., p H 7.4 Velocity Constants lteaction

O2 combination O p dissociation CO combination CO dissociation

1.9

Mb

Hb

x

106 M.-1 set.-' 40 sec.-l 1.3 X lo5 M.-1 set.-' 4 X sec.-l

107

37 3 4

x 106 x 10-2

4

x

The combination reactions involve a change in bond type from ionic to covalent and the dissociation reaction the reverse : Hb(H20) (Ionic)

+

0 2

HbOz

+ Hz0

(Covalent)

but it is clear from the data in Table V that both reactions are rapid. Eley (23) has examined the kinetic data on these reactions and has shown that the combination reactions are fast on account of low heats of activation and the dissociation reactions fast because of fairly high positive entropies of activation. It is thus not possible t o deduce from the speed of one of these reactions the change in bond type occurring. Apart from these reactions of hemoglobin and myoglobin, and the formation and subsequent reactions of the hydrogen peroxide complexes of peroxidase and catalase which have been intensively studied by Chance and will be described later, few have been examined kinetically. There are some kinetic studies for instance on the reduction of methemoglobin by hydrosulfite, the oxidation of hemoglobin by ferricyanide ions and the formation of methemoglobin cyanide. However, there are qualitative indications that the formation and dissociation of the complexes listed in Tables I1 and I11 are all fast reactions. Hence when a reaction is found t o be slow and the dissociation of some of the peroxide complexes of peroxidase and catalase are slow, it suggests that the complex may not be formed in a simple bimolecular displacement reaction like the formation of HbOz or HbCO. The rapidity of the hemoglobin reaction is a fair indication that the heme groups lie flat on the surface of the protein molecule. The observations of Perutz on the behavior of crystalline acid methaemoglobin offer further support (4). X-ray studies show that hemoglobin molecules are rigid and impenetrable t o liquid and crystalline oxy- or carbonmonoxyhemoglobin can be transformed into methemoglobin without any change in crystal structure. Acid methemoglobin crystals suspended

380

PHILIP GEORGE

in ammonium sulfate solution can be converted into alkaline methemoglobin by the addition of diammonium phosphate, and into azide methemoglobin by the addition of sodium azide without causing any change in unit cell dimensions. Thus these reactions between hemin groups and ions diffusing through the liquid of crystallization show the hemin groups to be located on the surface, and the constant unit cell dimensions show that these groups do not play a part in the bonding between neighboring hemoglobin molecules in the crystal lattice. In contrast t o the simplicity of these results it should be mentioned that in some way the crystal structure of reduced hemoglobin is different. Attempts to convert crystals of acid methemoglobin into reduced hemoglobin and crystals of reduced hemoglobin into oxyhemoglobin are always accompanied by a complete break up of the original crystal structure and recrystallization in a new form occurEi (Haurowitz, 24). It is not yet established whether these differences are only important in considering the structures of the crystals or whether they are reflected in the reactions of these compounds in solution. I n studying the reaction of the heme group in hemoproteins it is always necessary to ensure that no alteration in protein structure is affecting the results. Complete denaturation is easily recognized by precipitation and flocculation and the liberation of heme of hematin, but minor changes can also occur. I n the case of peroxidase two forms were isolated by Theorell (25) and named peroxidase I and 11. The first of these, now called paraperoxidase (Theorell, ZS), has been found t o have undergone some alteration possibly the removal of the carbohydrate portion of the protein molecule or possibly reversible denaturation t o a slight extent (Keilin and Hartree, 48). The paraperoxidase shows enzymatic activity comparable to that of the intact enzyme. Oxyhemoglobin can also undergo minor changes in structure which are apparent in a slight shift of the Soret band maximum absorption first from 414.5 mp t o about 410 mp, and then t o 406 mp (Jope, 4). It is an interesting feature that hemoglobin reconstituted from heme and free globin also shows maximum absorption at 410 mp. In addition to structural changes of this kind occurring, when the hemoproteins take part in oxidation-reduction react ions there is the further possibility of irreversible oxidative attack on the porphyrin ring itself. Lemberg, Legge, and Lockwood (27) have studied this process in the case of hemoglobin and have named the oxidation product choleglobin: the precise nature of this compound has not yet been established. The reactions described above and the specific reactions described below are to be regarded as characteristic of the four hemoproteins i n general. There are often differences in amino acid content and chemical

THE SPECIFIC REACTIONS

OF IRON IN SOME HEMOPROTEINS

381

behavior of the ‘ I same ” hemoprotein derived from different sources. For instance there are wide variations in the oxygen dissociation curves of mammalian and avian hemoglobins. Peroxidase obtained from horseradish has different properties from leucocyte peroxidase which contains a different porphyrin group (Lemberg and Legge, 5 ) . Bacterial catalase has a little greater catalatic activity than erythrocyte catalase although both hemoproteins contain four intact hematin groups; and both are more active than liver catalase in which one or more of the hematin groups have been converted into a catalytically inactive bile pigment molecule. These differences are described in detail by Lemberg and Legge (5). The reactions described in this paper refer mainly to horse hemoglobin and myoglobin, horseradish peroxidase, and catalase obtained from horse liver or erythrocytes. 111. HEMOGLOBIN AND MYOGLOBIN AUTOXIDATION AND OTHER REACTIONS The reversible combination of these hemoproteins with oxygen and carbon monoxide has been very extensively studied. Reviews of kinetics and equilibrium measurements are found in Haemoglobin (4), Lemberg and Legge ( 5 ) , and Wyman (1). The main themes of these studies have been t o explain the sigmoid shape of the oxygen equilibrium curve of hemoglobin which is in marked contrast with the normal hyperbolic curve found for myoglobin and t o explain the change in acidity which accompanies the oxygenation of hemoglobin, both of which are very important aspects of its biological function as an oxygen carrier. In hemoglobin it has been shown that interaction occurs between the heme groups such that when one has combined with oxygen the affinity of the remaining groups is altered. Several estimates have been made of the equilibrium constants for the formation of the successive intermediates Hb4(02),Hb4(02)2,Hbr(02)% and Hb4(0& (see Roughton, 4). Similar interaction occurs in the redox system hemoglobin-methemoglobin (Wyman, 1). There is a possibility of a similar sort of interaction occurring in catalase which has four hematin groups per molecule, and mention will be made of this later. In the case of hemoglobin it has been suggested that some oxidation-reduction reactions can proceed by an intramolecular mechanism on account of four reacting heme groups being attached to the same molecule. This mechanism has been developed particularly to account for the autoxidation of hemoglobin by Lemberg and Legge (5). The autoxidation reaction has received relatively little attention, but since it has considerable bearing on the problem of why oxyhemoglobin is stable, it will now be discussed in detail. It is often stated that the globin protects the heme from oxidation,

382

PHILIP GEORGE

and t h e problem is usually regarded as structural, the nature of the bond linking the ferrous protoporphyrin group to protein, conferring the highly specialized activity on the central iron ato:m. The way this happens may be considered as follows. The existence of oxyhemoglobin as a stable molecule is a particular case of the general problem of the stability of coordination compounds. As is well-known, this stability is very dependent on the nature of the coordinating groups and though combination with molecular oxygen is rare the ability of hemoglobin t o combine with it can reasonably be attributed to a favorable balance of bond energies, including as the crucial one that of the heme-protein link. The resistance to oxidation is also in part a structural problem for the nature of the coordinating group around II metal ion affect its oxidationreduction potential which is in part a function of the more fundamental property of the ion, the ionization potential of its reduced form in solution. The magnitude of this ionization potential immediately determines whether the first step in the oxidation is exothermic and so possibly very rapid, or how endothermic it is which would give an indication of how slow the step might be. But apart from this structural aspect of the resistance to oxidation there is also a kinetic aspect. Hemoglobin in the presence of oxygen does form methemoglobin, and so the mechanism of the oxidation, revealed by an analysis of the kinetics, should provide additional data for understanding the slowness of the oxidation when compared with other oxidation processes. Brooks (28,29) showed that the oxidation of hemoglobin at pH 5.69 is first order in unoxidized hemoglobin and the first order velocity conatant kob.. is dependent on the oxygen pressure according t o the complex function illustrated in Fig. 1A. Brooks summarized his results in the following rate equation:

where a is the fraction of uncombined hemoglobin, (a - t) the percentage of unoxidixed hemoglobin at time t , p the oxygen pressure, and k' and b numerical constants. In addition he made observations on the effect of varying the hydrogen ion concentration-the reaction is faster in more acid solution-the concentration of buffer salts, the effect of additions of sodium chloride, and the effect of temperature. Coryell, Stitt, and F'auling (30) later confirmed the first order dependence on unoxidized hemoglobin. Conant and Fieser (31) considered three possible paths the reaction could take: (a) the spontaneous decomposition of Hb02 into methemoglobin, ( b ) the oxidation of Hb by HbO2, ( c ) the direct oxidation of Hb

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

383

by 0 2 . In the light of his experimental results Brooks favored a mechanism based on ( c ) , for in (a) kObs.would increase t o a maximum value with increasing oxygen pressure, and in ( b ) the reaction would be second order in unoxidized hemoglobin. Direct reaction between Hb and 0 2 accounted bp) term for the term a(a - z ) p in his rate equation and left the b / ( l t o be explained. Brooks explored the possibility of certain intermediates in t,he oxygenation sequence Hb4(02)I, Hb4(02)2, Hb4(02)3 reacting preferentially with 02,and showed that no such selective reaction could account for the results. He put forward two possible explanations for the b / ( l b p ) term: (i) The presence of an oxidation catalyst R, in equilibrium with RO2 which oxidizes Hb (cf. Neil1 and Hastings, 32). (ii) Some type of chain mechanism for the oxidation in which O2 acts as an inhibitor. No further evidence has been found for an independent oxidation catalyst and recent work on autoxidation sheds no light on a chain mechanism which would explain the results. Lemberg and Legge ( 5 ) proposed that the oxidation catalyst is HbOz itself, but this point will be discussed later. I n 1942 Legge (33) suggested the spontaneous decomposition of one of the intermediates in the oxygenation sequence into methemoglobin as a possible mechanism. Using Pauling’s equation (34) for the fraction of hemoglobin in the form Hb4(02)1, Hb4(02)2, etc., as a function of oxygen pressure, he showed that the spontaneous breakage of Hb4(02)2 gave good agreement with Brooks’ data at low oxygen pressures. Lemberg and Legge ( 5 ) later amplified this mechanism by picturing the Fe2+02groups as oxidative catalysts and accounted for the selective action of Hb4(02)2by assuming the participation of two oxidizable XH2 groups on the protein t o make an intramolecular reaction possible fitting the stoichiometric relation:

+

+

+ 2XH2

Hbd(Oz)P

-+

Hb4(0H)r

+ 2X

in which the essential valency change occurring is: Fe2+(02).

. . XH2 . .

+

Fe2+--, Fe3+(0H) .

*

*

X

. . . Fea+(OH)

where Fez+, Fe2+02,and Fe3+(OH) represent the protoporphyrin iron atoms in hemoglobin, oxyhemoglobin, and methemoglobin respectively. Two particular assumptions underlie an intramolecular mechanism of this kind: (a) Fe2+and Fe2+02groups can react with each other (and with XH2 in the later mechanism) even though they are fixed on the globin molecule at considerable distances from each other. ( b ) I n spite of many collisions with other hemoglobin molecules, no significant inter-

384

PHILIP GEORGE

molecular reaction occurs. This intramolecular change is a necessary assumption in these cases to account for the first order dependence in unoxidized hemoglobin. The existence of this type of reaction is not yet established, and one difficulty in accepting it is the necessary assumption ( b ) that a normal reaction path by collision is blaicked in some way. On kinetic considerations alone the hypothesis of XH2 groups participating in the reaction is unnecessary. A mechanism of this kind has better claims if H 2 0 2is the reaction product and XH2 groups play no part at all, because a reaction path involving such groups would be an addition t o potential reaction paths similar to those responsible for the autoxidation of heme itself or hemochromogens which are rapid reactions. Bringing XH2 groups into the reaction mechanism makes it more difficult to understand why the reaction is slow. The validity of this mechanism rests primarily on its ability to explain the experimental results and over the low oxygen pressure range the agreement is good. However, if similar calculations are made for the higher pressure data serious discrepancies are found, especially at 723 mm. 0 2 where the reaction proceeds about forty times faster than it should according t o Hb4(02)2 decomposing (Brooks, 35). Even though the Pauling equation used in these calculations must be abandoned in the light of recent equilibrium measurements (Roughton, 4), there is no doubt that the fraction of hemoglobin present as Hb4(02)2 at high oxygen pressures will still be extremely small and this discrepancy in the oxidation rate still .appear. It seems doubtful therefore whether Hb4(02)2does play this unique role. An alternative approach to the problem was suggestNedby George (36) who showed that Brooks observed first order constant icob.. has an oxygen pressure variation very close t o that given by the product of the concentration of Hb and HbO2, i.e., a(1 - a). The maximum value of kob..should thus occur at the 0 2 pressure for half saturation in accord with experiment as in Fig. 1A. However a kinetic paradox appears for a velocity constant proportional t o [Hb] X [Hb02] should result from a reaction that is second order in unoxidized hemoglobin and not first order as was fully established by Brooks. This paradox may be resolved if the rate equation is written:

and a chemical mechanism looked for which would lead t o the concentration of unoxidized hemoglobin appearing in the denominator. Such a mechanism will be discussed later. George and Stratmann (37) investigating the kinetics of the oxidation of myoglobin t o metmyoglobin by molecular oxygen, under the aame

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

385

100

0 OXYGEN PRESSURE, mm.

ij i

A: HEMOGLOBIN

l I

I I

0

I

I

I

I

FIG. 1. Variation of the first order rate constant with oxygen pressure for: A, hemoglobin (Brooks, 29) ; B, myoglobin (George and Stratmann, 37). Temperature 30"C.,pH = 5.69. The broken lines indicate percentage saturation determined a t the various oxygen pressures. For rnyoglobin the two lines show the experimental variation.

conditions that Brooks used, found similar behavior. The reaction which proceeds about six times faster is first order in unoxidized myoglobin and the first order constant shows the same type of variation with oxygen pressure (see Fig. 1A and B). The maximum rate again occurs at the oxygen pressure required for half saturation indicating that both Fe022+ and free Fe2+ groups are involved in the reaction. Since myo-

386

PHILIP GEORGE

globin contains only one heme group on each protein imolecule an intramolecular mechanism such as t h a t suggested by Lemberg and Legge ( 5 ) is ruled out and for myoglobin the mechanism must be intermolecular. If it is accepted that similar oxidation kinetics result from an identical reaction mechanism, and the more complicated the kinetics the more likely this becomes, then it is very probable that with hemoglobin too t h e reaction is intermolecular. A system which shows some resemblance t o the autoxidation reaction is the oxidation of various substrates in the presence of hemoglobin and oxygen accompanied by the transformation of hemoglobin into choleglobin. The coupled oxidation of ascorbic acid in this system has been extensively investigated by Lemberg, Legge, and Lockwood (27), who compared its action with the production of verdohemochrome when hydrogen peroxide is added t o hemochromogens in the presence of hydrogen donors such as ascorbic acid (Lemberg, Cortis-Jones, and Norrie, 14). They showed t h a t the coupled oxidation is initiated by t,he direct reaction of oxyhemoglobin with ascorbic acid and suggested t h a t the reaction could be formulated:

+ tFez+ + A < Fez+

0 2

HrA

~

[Fez+02] HrA

A

+ [Fe2+H2021-+ FeS+

Cho:leglobin

*(

The ratio of ascorbic acid oxidized t o choleglobin formed is about 10: 1. B y using methemoglobin in place of oxyhemoglobin they found t h a t the rate of choleglobin formation was the same. Since Holden had established that the coupled oxidation could be poisoned by cyanide (38) this is evidence that methemoglobin acts as a true reaction intermediate. No detailed kinetic study of the system has been carried out but Lemberg, Legge, and Lockwood (39) showed that a t an oxygen pressure of 15 mm. the reaction proceeded about four times faster than at 150 mm. which indicates that the reaction mechanism is rather inore complicated than the above scheme where the reaction rate should rise t o a maximum with increasing oxygen pressure. The similarity with the autoxidation kinetics suggests that both H b and HbOz play a part in the reaction. A far more rapid oxidation of ascorbic acid occurs when it is mixed with a solution of oxyhemoglobin and acid then added t o denature the protein. I n the absence of ascorbic acid, oxyhemoglobin yields denatured protein and acid hematin, only 40% of the oxygen is evolved, hydrogen donor groups in the globin are oxidized, and a small part of the hematin also undergoes oxidation with liberation of its iron. If ascorbic acid is present it is oxidized instead of the globin. The oxidation occurs instantaneously if the ascorbic acid is added before acidification, but much

THE SPECIFIC

REACTIONS

OF IRON IN SOME HEMOPROTEINS

387

less is oxidized if it is added after. Lemberg and Legge ( 5 ) suggest that the ascorbic acid reacts directly with activated oxygen in an oxyhemoglobin-ascorbic acid complex, the oxygen being activated as a result of the acid altering the oxygen-heme link. There is a second way in which this system can oxidize substrates. With biliverdin and bilirubin (open-chain derivatives of porphyrins) there is no immediate decay of the oxidizing system after acidification. Lemberg (40) and Lemberg, Legge, and Lockwood (41) showed that this was due t o the hydrogen peroxide present produced by the action of acid on oxyhemoglobin, oxidizing the substances with the acid hematin also produced acting as a peroxidase. A possible free radical interpretation of these reactions will be discussed later. A different type of reaction is that between methemoglobin and hydrogen peroxide. Kobert (42) showed that a red-colored complex was formed and Keilin and Hartree (43) demonstrated that one peroxide molecule per iron atom was required. The complex which has absorption maxima at 585 mp (a band) and 545 mp (0band) slowly decomposes, liberating methemoglobin, but is rapidly destroyed by hydrogen donors such as ascorbic acid and hydroquinone. Haurowitz (44) claimed that addition of alkali t o the complex yielded alkaline methemoglobin and represented the reaction:

+ OH-

MetHb.+HzOz

4

MetHb.OH

+ Hz02

Keilin and Hartree showed too that ethyl hydrogen peroxide gives a complex with absorption maxima a t the same wavelengths. I n a more recent paper (45)) they have extended their observation on the hydrogen peroxide complex, confirming their results at low HzOz/MetHb concentration ratios. As this ratio is increased they found that the absorption spectrum is modified, the a band which previously showed a maximum at 585 mp splitting into one with two maxima at 578 and 592 mp. Increasing the HzOz/MetHb concentration ratio still further, resulted in the disappearance of the longer wavelength band and the spectrum with maxima now at 578 mp (aband) and 545 mp ( p band) was recognized as that of oxyhemoglobin. At the concentration of hydrogen peroxide required to obtain the oxyhemoglobin spectrum catalytic decomposition of the peroxide occurred. This establishes beyond any doubt that the catalytic decomposition of peroxide in this case proceeds by a mechanism involving a valency change of the iron atom. Keilin and Hartree (45) showed too that metmyoglobin forms a similar compound with hydrogen peroxide which undergoes the same reactions. George and Irvine (116) found that this complex is formed at all hydrogen ion concentrations between pH 5.5 and 12.0, but that only

388

PHILIP QEORQE

when the formation occurs in the pH range 8-9 can complete regeneration of the metmyoglobin be effected by adding mild reducing agents. The spectrum of the complex in this region shows a sharp band at 549 mp with emM. = 9.8 and a broad band appearing as a shoulder at 570-585 mp with cmy. = 8.4. In the Soret region there is a sharp band at 423 mp with emy. = 107. With solutions of pH < 8 some complex is formed, but an additional reaction involving the hematin also occurs, giving an equally intense spectrum but with the bands shifted. When the reaction is carried out in solutions of increasing alkalinity above pH 9.0 progressively less complex is formed and a hematin degradation product, colorless in both the visible and Soret regions of the spectrum, is produced. Titration of the complex using ferrocyanide as the reducing agent showed that one ferrocyanide ion reacts per molecule of the complex. The complex thus has only one oxidizing equivalent compared with the two oxidizing equivalents of hydrogen peroxide itself. On the other hand, when varying amounts of ferrocyanide and metmyoglobin were mixed and hydrogen peroxide then added, the full oxidizing capacity of the peroxide was obtained. This suggests very strongly that the formation of the complex proceeds in a reaction of the type MetMb

+ HzOz

(2 oxid. equiv.)

--*

+

Complex X (1 oxid. equiv.) (I oxid. equiv.)

and that a transient oxidizing entity X is produced along with the complex, which can react with added reducing agents if present initially, but which disappears if the complex formation is allowed t o proceed independently. These results indicate that the complex cannot be of the ion pair type [MetMb-OOH] for such a compound would have two oxidizing equivalents. A tentative mechanism which can account for these results involving the formation of a quadrivalent iron compound “ferrylmyoglobin ” will be discussed later.

REACTIONS IV. PEROXIDASE The catalytic activity of peroxidase is intimately connected with its ability to form complexes with hydrogen peroxide. Three such complexes are formed, depending on the experimental conditions and they are interrelated :

spontaneous

+

Peroxidase Ha09 + Per. (HnOi) I (brown) (1 mole) green

excem

Per. (Hz01) 11 Per. (HaOa) I11 H:Oi pale red deep red

decomposition

Per. (HzOz) I1 pale red

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

389

As the structure of these complexes is unknown they are denoted as I, 11, and 111. Keilin and Mann in 1935 described complexes I1 and I11 (46). Complex I1 is formed from equimolar amounts of peroxidase and hydrogen peroxide and with an excess of peroxide it changes into complex 111. Both these complexes react rapidly with hydrogen donors. Carbon monoxide does not affect their spectrum, which led Keilin and Mann (46) to suggest that the iron atom is in the trivalent state. Later in 1942 Theorell (20,47) observed that the pale red complex I1 was not formed immediately when peroxidase was added to hydrogen peroxide but that a transient green-colored complex was formed first which very rapidly changed into the pale red complex 11. The position of the absorption maxima of these compounds together with those of free peroxidase in the visible spectrum are listed in Table VI. TABLE VI Compound of Peroxidase Neutral peroxidase Peroxidase (H20z)I Peroxidase (H202)I1 Peroxidase (H2Oe) I11

.

Position of Absorption Maxima (mI*)

Reference

645, 583, 548, 498 658 - - - 561, 530.5 - 555," 527 - 583, 545.5 583, 546

46 20,47 46 48 46 48

a Asymmetric absorption band, 555 mp, refera to the observed maximum, not the "center" of the band.

Chance (49,50) has extended the spectroscopic examination of these complexes to the Soret region 370-450 mp and has examined too the complexes with methyl and ethyl hydroperoxide. The primary complexes are also green and have nearly identical spectra in the Soret region: these maxima are at 407 mp. The Soret band of the free enzyme is considerably diminished in intensity on forming the complex, and it is slightly shifted toward the visible. The alkyl hydroperoxides also yield pale red secondary complexes with spectra in the visible region very similar to the hydrogen peroxide complex 11. In the Soret region all have about the same intensity as the free enzyme, but there is a shift of some 15 mp toward the visible. Chance comments that the identity of the Soret bands of the peroxide complexes with HzOz and methyl or ethyl hydrogen peroxide indicates that the nature of the ironperoxide bond is unaffected by substitution in the Hz02 molecule. Theorell (20,51) has suggested on the basis of a magnetochemical study that the bonding in the pale red secondary complexes is covalent, whereas

390

PHILIP GEORGE

a comparison of the spectra of the green primary complexes with the green peroxidase-fluoride complex, which is known t o contain ionic bonds, indicates that the bonding in these primary complexes is ionic. Chance (50) found that the alkyl hydroperoxides do not give a deep red-colored complex analogous t o peroxidase-hydrogen peroxide 111. Instead a bright green complex with a very strong band at 675 mp and a faint band at 557 mp is formed. This complex which he designates complex IV has a Soret band resembling that of the primary complexes. Since there is considerable destruction of peroxidase at th e large peroxidase concentrations required t o form this complex, hLe suggests th a t it is related t o t h e green oxidation products of hemoglobin and catalase studied by Lemberg and his associates (see Lemberg and Legge, 5). Abrams, Altschul, and Hogness (52) suggested th a t the similarity between the visible bands in the pale red complex given by the peroxidase which can oxidize cytochrome c, and th e bands in oxymyoglobin and oxyhemoglobin indicates th at the secondary complexes may be ferrous compounds in spite of their resistance t o carbon monoxide. However, Chance points out that this spectroscopic analogy does not hold in the Soret region 370-450 mp. I n this region there is some resemblance when the bands of metmyoglobin in relation t o oxymyoglobin are compared with those of peroxidase in relation t o the primary complex. These inconsistencies in spectral analogies when th e visible region is compared with the Soret region thus makes it impossible t o draw any definite conclusions. Some of the early studies of the kinetics of peroxidase oxidations provided evidence t o suggest that both the hydrogen peroxide and the hydrogen donor were joined to the enzyme molecule (Mann, 53). It was found that an excess of hydrogen peroxide inhibited the oxidation of the hydrogen donor, and th at this inhibition could be relieved by increasing the concentration of the hydrogen donor. This could be explained if hydrogen peroxide first combined with the enzyme and then peroxide and the hydrogen donor competed for a second site on th e enzyme. Combination between the hydrogen donor and the enzyme has been accepted by Lemberg and Legge (5) in their theory of peroxidase action which they represent: Fe3+OH

FeJfOH HzA

Fe3+OOH HzA

FeSfO H

I

+ HzO + A

However, the inhibition of peroxidase action a t high peroxide concentrations is capable of other interpretations, and since there is no evidence for the enzyme uniting with t.he hydrogen donor throiigh a group in its

THE SPECIFIC REACTIONS

OF IRON IN SOME HEMOPROTEINS

391

protein in the detailed kinetic studies of Chance (see below) the acceptance of this mechanism must await a more conclusive proof of such combination. Chance has recently extended his original observations on the kinetics of the formation of the pale red peroxidase-H202 complex and its reactions with hydrogen donors like leucomalachite green and ascorbic acid (54) t o include the kinetics of formation of the green primary complex and its conversion into the pale red secondary complex (55) in the case of both hydrogen peroxide and alkyl hydroperoxides. Using a rapidflow apparatus equipped with a monochromator, kinetic data were obtained in agreement with the following basic reaction mechanism : Per. OH

+ HOOR kikz Per. OOR (I) + H2O kl

Per. OOR (I) + Per. OOR (11)

(2)

ka

Per. OOR (11) -+ Per. (OH)

(3)

ks'

Per. OOR (11)

+ AH2-+ Per. (OH) + ROH + A

(3')

The secondary complex decomposes spontaneously as in 3 and can react with hydrogen donors as in 3'. The study of the last reaction entailed using concentrations of hydrogen donors (AH2) in excess of t h a t of the secondary complex such that the kinetics were pseudo first order. The net velocity constant k,' is thus equal t o kd(AH2) where kk is the true bimolecular constant for this reaction. The values of kl were found t o be 9 k 2 X lo6, 1.5 X lo6 and 3.6 X 106 M.-l sec.-l for hydrogen peroxide and methyl and ethyl hydrop.eroxides respectively; and k , the first order constant for the conversion of the primary into the secondary complex was about 4 sec.-l for all three peroxides. The spontaneous decomposition of the secondary complex depends principally upon whether peroxides have been previously added t o the solution. A fresh peroxidase gives velocity constants for k , of about 0.02 set.-', but several additions of peroxide can reduce this velocity constant by a factor of ten. T o explain this effect Chance suggests t h a t a hydrogen donor is present with the enzyme initially which can accentuate the decomposition of the secondary complex as in reaction 3', but with successive additions of peroxide it becomes oxidized. This basic mechanism holds if the reaction is carried out in two stages: first, the formation of the primary green complex followed by its quantitative conversion into the secondary complex, then addition of the hydrogen donor which is oxidized by the secondary complex. However, a more complicated mechanism is possible for Chance also found that the conversion of the primary complex into the red secondary complex is accentuated in the presence of a

392

PHILIP GEORGE

hydrogen donor, and the velocity of formation of the secondary complex from the primary complex can be nearly as rapid as the actual formation of the primary complex. The value of kz, the velocity constant for the decomposition of the primary complex into peroxidase and peroxide cannot be obtained by direct measurement. Measurements of the equilibrium constant for the dissociation of the primary and secondary complexes with methyl M. and 3 X lo-' M. from which hydroperoxide gave values of 3.2 X kz may be estimated to be 2.2 sec.-l or 3.4 sec.-' from the known formation velocity constants. Chance found that pH has little effect on the dissociation constants in the range 3.6 to 8.8, which suggests that neither of the two hemelinked groups found by Theorell and Paul (15) affects the ability of peroxidase to form its enzyme substrate complexes. On the basis of this observation he suggests that the primary combination takes the form: Per. OH

+ HOOR kikl Per. OOR (I) f H1O

Above pH 8.8, where further decrease of hydrogen ion concentration leads to increased formation of alkaline peroxidase, peroxide complex formation is suppressed, which indicates that peroxides cannot replace the covalently bound hydroxyl group. There is an interesting parallel here in the case of methemoglobin where Haurowitz claimed that the alkaline form cannot give a peroxide complex (44). The reactions of the secondary complexes with hydrogen donors are rapid bimolecular reactions. Chance (49) has obtained the velocity constants for horseradish peroxidase, hydrogen peroxide and methyl and ethyl hydroperoxide complexes respectively reacting with ascorbic acid at pH 7.0: k4 = 2.8 X lo3, 2.8 X lo3, and 2.2 X lo3 M.-l sec.-l and reacting with pyrogallol at pH 7.0, k4 = 2.1 X IOs, 2.1 X lo6, and 1.8 X lo6 M.-l set.-'. With the three hydrogen donors, hydroquinone, guaiacol, and pyrogallol, there was no systematic decrease in the values of kb over the pH range 3.6-6.7. Whereas the peroxidase reaction is reversibly and competitively inhibited by cyanide ions (Lemberg and Legge, 5 ; Keilin and Mann, 46; Chance, 56) carbon monoxide has no inhibiting effect (Theorell, 47), which has led to the widely accepted view that the enzyme remains in the ferric state throughout the reaction. A very different type of reaction that peroxidase can catalyze is the autoxidation of dihydroxymaleic acid HOOC.C(O'K)=C(OH)-COOH which has the same oxidizable group as ascorbic acid (see Lemberg and Legge, 5). I n contrast to the usual peroxidase reaction with peroxides,

THE SPECIFIC REACTIONS O F IRON I N SOME HEMOPROTEINS

393

this system can be poisoned with carbon monoxide, and this inhibition is reversed by light. It would seem therefore that a valency change occurs in the course of this reaction. Theorell and Swedin (117) have shown that the brown color of free peroxidase changes to deep red when oxygen is bubbled through a mixture of peroxidase and dihydroxymaleic acid, and that the absorption bands of the red substance correspond to those of peroxidase in the presence of an excess of hydrogen peroxide, i.e., the third comRlex, Per. (HzOz)111. The initiation of the polymerization of vinyl compounds is a good indication of the production of OH radicals (Dainton, 108) particularly in systems containing hydrogen peroxide and ferrous ion (Baxendale, Evans, and Park, 84). It is interesting therefore that peroxidase and catalase with hydrogen peroxide have been found inactive (Dainton and Smith, 107). Another difference between the reactions of the enzymes and ionic ion, where the mechanism is generally accepted t o involve free radicals, appears when the products of hydrogen peroxide oxidation obtained using peroxidase are compared with those obtained using ferrous iron. Mann and Saunders (log), Saunders and Mann (110), and Chapman and Saunders (111) investigated the oxidation of aniline, p-toluidine, and mesidine by these two systems and found very different reaction products in the two cases. V. CATALASE REACTIONS

I . Catalytic Decomposition of Hydrogen Peroxide Many kinetic studies have been made of this reaction which is complicated by the destruction of the enzyme as the reaction proceeds. Typical of the early investigations are those by Yamasaki (57), Morgulis (58), Northrop (59), Williams (60), Nosaka (61), Maximovitsch and Antonomova (62), and Zeile and Hellstrom (63). Under certain experimental conditions the following kinetic equations hold: -d(H202)/dt = k[Catalase] . [H~OZ]

and -d(Catalase)/dt

=

k’[Catalase] . [ H z ~ z ]

I

I1

The enzyme destruction particularly at a temperature of 20°C. or below is a relatively slow reaction compared with the peroxide decomposition, and since all the evidence shows it t o be a true side reaction, it will not be discussed further here. A change in activity of the enzyme observable at high peroxide concentrations but distinct from the slow destruction reaction will be referred to Iater. Measurements of peroxide decomposition under the conditions where Eq. (I) holds have been widely used for estimation of catalase activity

394

PHILIP QEORGE

according to the procedure suggested by Zeile and Hellstrom (63). This entails using very dilute catalase solution about 5 X 10-l1 M , and approximately N HzO2 and estimating the residual peroxide at three minute intervals after the reaction is started. This procedure has recently been criticized by Bonnichsen, Chance, and Theorell (64), who showed that Zeile and Hellstrom's procedure gave values for the enzyme activity about 60% of that obtained when higher enzyme concentrations were used, e.g., 2 X lou9 M , and samples withdrawn. for titration at 13, 28, and 43 seconds from the start of the reaction. 'They suggest that enzyme inactivation arising from the spreading of the catalase over the various interfaces in monolayers or chemical inactivation by peroxide itself in reversible or irreversible reactions is responsible. These results may be compared with those obtained by George (65,66,67). I n this investigation a manometric technique was used which made it possible t o measure the rate of peroxide decomposition at constant peroxide concentration 0.01-5.0 M . There was an initial rapid evolution of oxygen which lasted for about two minutes, depending on the peroxide concentration, followed by evolution at a steady rate which slowly decreased in the course of an hour. This decrease was undoubtedly due t o enzyme destruction. I n the initial reaction the rate of oxygen evolution was found to decrease exponentially with time and this exponential decrease was more rapid the higher the peroxide concentration according to the following expression for the exponential constant:

where P is the peroxide concentration in moles per liter, and a, b, and c are constants having the values 0.072 M.-l sec.-', 0.15 M.-l, and 0.0185 M.-I sec.-l at O'C., respectively. Both initial and steady rates were directly proportional t o the catalase concentration but showed a complex variation with the peroxide concentration, the rates rising t o maximum values at 0.4 M and 0.07 M H202, respectively, and then decreasing markedly with further increases in the peroxide concentration. Peroxide itself thus inhibits its own decomposition in high concentration. These results were obtained with erythrocyte and liver catalase and lysed red cells, and further experiments showed that the po:isible presence of inhibitors in the peroxide or the nature and concentnition of the buffer solution played no part in the reaction. The inhibition of the steady rate decomposition at high peroxide concentrations was shown to be quantitatively reversible by several experiments in which dilute buffer solution or small amounts of concentrated peroxide solution were added during the course of the reaction. This inhibition is therefore not caused

T H E SPECIFIC REACTIONS O F I R O N IN SOME HEMOPROTEINS

395

by enzyme destruction. Dilution during the period of initial high activity also showed a reversibility which suggests that the transition from high to low activity is not caused by partial destruction of the enzyme. The initial high activity was found to be more susceptible to inhibition by azide and cyanide ions than the subsequent low activity for at certain concentrations of these inhibitors the initial high activity could be completely eliminated without affecting the low activity. Foulkes and Lemberg have confirmed that the activity of catalase decreases initially much faster than can be accounted for by its irreversible destruction but using different experimental conditions found that azide ions can enhance the initial falling off in activity (18). George (65) also investigated the kinetics of the decomposition of peroxide by concentrated catalase solutions in the presence of high concentrations of sodium azide (0.6 M ) and found a residual catalytic activity apparently attributable to azide catalase itself. The catalysis by azide catalase did not show the inhibition by peroxide at high peroxide concentrations like the free enzyme. The kinetics of the catalase peroxide reaction over a wide range of concentrations are thus very complicated and the problem is t o decide what are the significant kinetics for the oxygen evolution reaction by the intact enzyme. Bonnichsen, Chance, and Theorell (64) favor the view that only the first order kinetics given by Eq. (I) are significant. The bulk of their experimental data refer to peroxide concentrations of 0.1 M or less where George also found the initial rate directly proportional t o peroxide concentration. At higher peroxide concentrations, 0.3 and 1.0 M , data obtained by Millikan and McLaughlin (see 64) showed no decrease in rate at the higher concentrations in contrast to George's results. Further experimental work is very necessary in this concentration range to settle this point. The values for the velocity constant in Eq. (I) obtained by Bonnichsen, Chance, and Theorell for pure erythrocyte and liver catalase are k = 3.5 X lo7 and 3.0 X lo7 M.-l, set.-' at 22"C., respectively. The activation energy for the reaction with erythrocyte catalase has the remarkably low value of 1,700 & 100 cal. The comparison between the affinity of an inhibitor for catalase as measured spectrophotometrically and the affinity determined from kinetic measurements has revealed several interesting features. Keilin and Hartree (17) showed that azide ions and hydroxylamine are much more effective inhibitors than the affinity of these compounds for ferric catalase would suggest. Their results are given in Table VII, the first column giving the molarity of NaN3, NH20H, and KCN required to produce 50% inhibition, the second column giving the relative affinity

396

PHILIP GEORGE

TABLE VII Inhibition of Catalase Activity of Azide, Hydroxylamine, and Cyanide, the Relative A f i n i t y of These Substances for the Enzyme and the Corresponding Dissociation Constants for the Complexes Inhibitor

Molarity to Give 50% Inhibition

Relative Affinity for Ferric Catalase

IXssociation Constant of Ferric Complex

NaN3 NHiOH KCN

6.3 X lo-* 6.3 x 10-7 4.3 x 10-6

67 40 10,000

6 X 10-4 M. 10-8 M. 4 x 10-6 M.

of the three compounds for catalase determined spectroscopically, and the third column giving the dissociation constants of the azide ion and hydroxylamine complexes calculated from the dissociation constant of the cyanide ion complex as determined spectrophotometrically by Chance (68). If azide ions and hydroxylamine exercised their inhibitory power entirely by combination with ferric catalase, then the molarity required to give 50% inhibition would be numerically equal to the spectroscopically determined dissociation constant. The data shows them to be between lo3 and lo4 times more effective. Keilin and Hartree (17) have shown in a series of experiments that this enhanced inhibiting power of azide and hydroxylamine is connected with the reduction of catalase to a ferrous form when these substances are present, and they suggest that these compounds stabilize the ferrous form. The reduction of the iron is clearly demonstrated in the following way. When sodium azide is added t o a catalase c;olution, the color originally greenish brown becomes slightly more greenish and the absorption band at 622 mp is intensified and moves about 2 p nearer the blue end of the spectrum. Addition of hydrogen peroxide changes the color of the solution t o red, and the absorption bands are replaced by two stronger bands at 587 and 559 mp. If carbon monoxide is passed through the solution, these bands become more distinct and shift to shorter wavelengths, the band maxima now being at about 577 and 546 mp. Both these red compounds, especially the latter, are fairly stable in the absence of oxygen, but in its presence they rapidly change back into the original greenish brown azide catalase. Even in a nitrogen atmosphere enough oxygen is liberated from the hydrogen peroxide t o bring about this change although in a carbon monoxide atmosphere it occurs more slowly. Hydroxylamine catalase is also reduced by hydrogen peroxide to give a spectroscopically similar autoxidizable compound. A quantitative study of the inhibition of the peroxide decomposition by sodium azide in the presence of carbon monoxide showed that the

TEE SPECIFIC REACTIONS O F I R O N I N SOME HEMOPROTEINS

397

resultant inhibition was far more pronounced in CO/N2 than in C O / O n gas mixtures (Keilin and Hartree, 17). The data indicates that in CO/Oz mixtures there is a definite competition between the two gases for the ferrous azide catalase, the enzyme having a greater affinity for O2 than CO. Expressing the partition coefficient of azide catalase between 0 2 and CO in the form k =

[Cat. Oz]CO [Cat. CO1.02

where Cat. 0 2 and Cat. CO represent the concentrations of active and inactivated enzyme respectively, the experimental data were found to agree very closely with calculated values obtained using lc = 9. The inhibition of azide catalase by CO is very largely relieved by radiation from a mercury vapor lamp, which is additional support for the formation of a ferrous complex, for carbonmonoxyhemoglobin can be dissociated in this way, and it also confirms that very little irreversible inhibition is occurring. A further piece of evidence for the reduction of the iron in this system is the observation that although the affinity of cyanide for ferric catalase is about 150 times greater than that of azide (see Table VIIj, when cyanide is added to azide catalase, previously treated with hydrogen peroxide in a pure nitrogen atmosphere, no change in the spectrum occurs. The band at 587 mp remains sharper than the band at 559 mp in contrast to the cyanide-catalase spectrum in which the band at about 585 mp is broader and much less intense. Replacing the nitrogen atmosphere by carbon monoxide shifts the bands toward shorter wavelengths as was found in the absence of cyanide. Although these experiments establish beyond doubt the reduction of catalase hematin when azide or hydroxylamine are present, there is no similar direct evidence for reduction with the free enzyme, for in this system carbon monoxide has no inhibiting effect and no spectroscopic changes have been observed. The inhibition of catalase by cyanide shows none of the characteristics of the azide or hydroxylamine inhibition as is to be expected if cyanide combines with the ferric form. At low peroxide concentrations M the equilibrium constant for the formation of the cyanabout catalase complex (K1) determined from kinetic data using the expression

where R, is the uninhibited rate and R 1 the inhibited rate a t various (HCN) concentrations agrees well with the value of K I determined by direct spectrophotometric measurements (Chance, 68). This shows that under these conditions cyanide inhibits simply by the removal of

398

PHILIP GEORGE

active catalase as the inactive cyan-catalase complex :tnd that no competition occurs between cyanide and peroxide for the hematin iron atoms. This is in agreement with earlier observations that the extent of cyanide inhibition is independent of the peroxide concentration (see Lemberg and Legge, 5 ) . Simple though this appears, it is difficult t o correlate this observation with the formation of a catalase-hydrogen peroxide complex whose existence was revealed by a rapid-recording spectrophotometric technique developed by Chance (69). The results he obtained will now be described.

2. Complexes of Catalase with Hydrogen Peroxide and A1A:yE Hydroperoxides Catalase was found t o form an intermediate compound in the presence of hydrogen peroxide (Chance, 69). The spectrum was measured from 380-430 mp and is slightly shifted toward the visible as compared with free catalase. The complex shows no similarities to1 cyan-catalase or the compound formed when peroxide is added to azide catalase. Its formation is very rapid, the bimolecular velocity constant having a value of about 3 x lo7 M.-l see.-'. In the absence of added hydrogen donors, the complex decomposes slowly according to a first order reaction with a velocity constant of about 0.02 set.-'. This catalase complex thus resembles the green primary hydrogen peroxide complex of peroxidase. A very remarkable property of the complex is tha,t apparently only one of the four hematin groups of the catalase molecule is bound t o peroxide (Chance, 70). Three independent experimental methods provide qualitative evidence for this. First, the spectrophotometric titration of catalase with hydrogen peroxide gives amounts of the complex in excess of the molar amount of peroxide added when the calculation is based upon the combination of all the hematin groups with the peroxide. Secondly, the noncompetitive inhibition of catalase activity by cyanide indicates that not all the hematin groups are combined with peroxide under the conditions of the test for catalase activity. Thirdly, the difference between the extinction coefficient of catalase and the complex is rather small having a value of 40 cm.-l x mM.-l ak 405 mp which is comparable t o the difference between the extinction coefficients of erythrocyte catalase with four hematin groups (380 cm.-' X mM.-l) and liver catalase with three hematin groups (340 cm.-' X mM.-l). This small difference in extinction coefficient becomes even more significant when compared with the large difference found for the alkyl hydroperoxide complexes (180 cm.-' X mM.-' at 405 mp) whose formation and reaction will be described later. Although the peroxide decomposition by catalase shows noncom-

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

399

petitive inhibition by cyanide ions, Chance found that cyanide ions do compete with hydrogen peroxide in the formation of the peroxide complex (70). A quantitative study of the reaction of the catalase hydrogen peroxide complex with cyanide or alkyl hydroperoxides revealed that 1.2 +_ 0.1 hematin groups are combined no matter whether 1.3 of the 4 hematin groups are present as degraded bile pigments, and this composition is independent of large variations in the peroxide concentration. Chance suggests two possible interpretations of these results. If all the catalase hematins react in the same way and independently of each other, it follows that each group is saturated to the extent of (1.2 _+ 0.1) + 4 as a consequence of a steady state being set up, the complex being formed from catalase and peroxide and destroyed by reacting with further peroxide molecules. This interpretation, however, does not account for the constant composition of the complex when some of the hematins have been converted into bile pigment, as in liver catalase, without the additional postulate of an interaction between the catalase hematins. There is, however, no evidence for such interaction in the cases of combination with cyanide and methyl hydroperoxide which have been thoroughly examined for this effect (Chance 68,71). A second interpretation is that the peroxide molecule can only combine reversibly with one of the hematins of the catalase molecule and that having combined, the remaining hematin groups acquire “special properties ” such that they now react rapidly with peroxide and destroy it. This would account for a constant composition of one peroxide per catalase molecule independent of the bile pigment content, and thus conflicts with the observed value of 1.2 -t 0.1 bound hematins. More recent work by Chance and Herbert (72) on bacterial catalase has revealed that with this enzyme 1.6 peroxide molecules are bound to the catalase molecule. This suggests very strongly that the second interpretation entailing the unique combination with one hematin group is wrong and that the first interpretation in terms of a steady state is correct. Possible reaction mechanisms with the peroxide complex will be discussed later in comparison with hydroperoxide complexes. Like peroxidase, catalase forms both primary and secondary complexes with methyl and ethyl hydroperoxide (Chance, 73). The primary complexes are green having a diffuse absorption band in the red starting at 670 mp. The secondary complexes are red and have absorption maxima in the visible region at 572 and 536 mp. The catalase-ethyl hydroperoxide complex found by Stern (74) had maxima at approximately these wavelengths and was thus the secondary complex. The Soret bands of the primary complexes are similar in shape to that of the free enzyme but are shifted toward the red by several millimicrons. At

400

PHILIP GEORGE

405 mp both ethyl and methyl hydroperoxide complexes show a decrease of extinction coefficient as compared to the free enzyme of about 180 cm.-l X mM.-' for erythrocyte catalase, i.e., about 45 cm.-' X mM.-' per hematin group bound t o peroxide. The Soret bands of the secondary complexes are very different from those of free catalase and the primary complexes. They show a large shift toward the red, for instance the maximum for the methyl hydroperoxide complex is :st 422 mp with an extinction coefficient of 242 cm.-' x mM.-', and they appear somewhat similar to the cyan-catalase complex. A striking difference between the primary alkyl hydroperoxide complexes and the primary hydrogen peroxide complex is that in the former all the catalase hematins are combined (Chance, 70). I n the p H range 3.8-9.0 the combination follows the equation Cat. ( O H ) r

+ 4HOOR

Cat. (OOR),

+ 4EIoO

each group being bound independently. The bimolecular velocity constants for their formation although large are a little less than that for the formation of the hydrogen peroxide complex. The values for ethyl and methyl hydroperoxide and hydrogen peroxide are 2 X lo4, 8.5 X lo5, and 3 X lo7 M.-l set.-', respectively, showing a decrease in magnitude with increasing size of the peroxide. The formation of the secondary complexes from the primary alkyl hydroperoxide complexes is not simple (Chance, 71). The secondary complexes do not form until an appreciable amount of the primary complex is already present, yet the reaction is not a first order transformation of the primary complex as in the case of peroxidase. The velocity of formation of the secondary complexes increases with increasing alkyl hydroperoxide concentrations but not enough t o follow a second order equation. A property all the primary complexes have in common is the decomposition giving free catalase which does follow first order kinetics (Chance, 71). The velocity constants for ethyl and methyl hydroperoxides are 0.04 and 0.016 sec.-l as compared with 0.02 set.-' for the hydrogen peroxide complex. The secondary complexes decompose far more slowly, the first order velocity constants for ethyl and methyl hydroperoxides having the values 2.3 X and 4 X SIX.-', respectively. 3. Catalase in Coupled Oxidations

In 1936 Keilin and Hartree (75) showed that addition of ethyl alcohol t o certain enzymatic oxidation systems such as uricase with uric acid and amino acid oxidase with amino acids doubles the oxygen uptake.

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

401

Similar additions of ethyl alcohol to xanthine oxidase with hypoxanthine or xanthine oxidase with acetaldehyde had no effect; however, a little catalase added to these systems resulted in increased oxygen uptake like that observed in the previous cases. These results showed that alcohol undergoes a secondary or coupled oxidation by the hydrogen peroxide formed in the primary oxidation reaction. It was shown that catalase could utilize hydrogen peroxide for this type of coupled oxidation in a series of experiments in which hydrogen peroxide was gradually formed by the hydrolysis of barium and cerium peroxides. It therefore seemed likely that the difference between the primary oxidation systems was that uricase and amino acid oxidase contained slight traces of catalase whereas xanthine oxidase did not. Later experiments confirmed these results (Keilin and Hartree, 76), and the following set of compounds were found to undergo coupled oxidation: methanol, ethanol, n-propanol, isobutanol, 0-amino-ethanol, and ethylene glycol. Keilin and Hartree suggested that this coupled oxidation is a more likely biological function of the enzyme than the catalytic decomposition of the hydrogen peroxide into water and molecular oxygen. Chance (69,118) showed that this coupled oxidation can be accounted for by the primary hydrogen peroxide complex reacting with the alcohol in a bimolecular reaction liberating t,he free enzyme. The velocity constants for its reaction with methyl, ethyl, n-propyl, n-butyl, and isoamyl alcohols are lo3, lo3, 17, 2, and 0.1 M.-' sec.-l, respectively. Ascorbic acid was also found t o react with an apparent bimolecular velocity constant of 3.4 X lo2 M.-l set.-', but in a later paper Chance (119) has shown this to be a reaction of a different type involving the production of a secondary complex and not the free enzyme. I n addition in this paper, reactions of the primary complex with nitrite and formate are described which are similar to the reactions with alcohols. The primary compounds of the alkyl hydroperoxides also react with hydrogen donors. The methyl hydroperoxide complex reacts at approximately the same rate with methyl, ethyl, and n-propyl alcohol as does the hydrogen peroxide complex; the ethyl hydroperoxide complex reacts somewhat faster, the velocity constants for the three alcohols being 2.1 X lo3, 2.1 X lo3, and 33 M.-l see.-', respectively. The activity of the secondary complexes is negligible in dilute solution, so that formation of the secondary complex in the reacting system has the effect of inhibiting the reaction. This accounts too for the inhibition of the ordinary catalytic decomposition of hydrogen peroxide when ethyl hydroperoxide is added (Chance, 77). It is very interesting that hydrogen peroxide itself can react extremely rapidly with the primary alkyl hydroperoxide complexes in a manner analogous to the alcohols. The velocity constants for its reac-

402

PHILIP GEORGE

tion with the ethyl and methyl complexes are estimaied to be greater than lo6, and about 5 X 107 M.-l set.-', respectively. As in the case of catalytic decomposition of hydrogen peroxide the peroxidatic activity of the enzyme shows no inhibition by carbon monoxide. Chance investigated this in the system primary methyl hydroperoxide complex reacting with ethyl alcohol. Instead of inhibition a slight increase in the rate of disappearance of the complex was noted which could be attributed t o formate being produced by the hydration of the carbon monoxide and acting as an additional substrate (71).

4. Reaction Mechanisms with the Catalase-Peroxide Complexes Chance (78) has discussed this experimental data in terms of the extended Michaelis theory which accounts for the similar peroxidatic action of peroxidase, the only difference being that with peroxidase the main reaction can proceed via the secondary complexes, whereas with catalase these complexes are inactive and the main reaction proceeds via the primary complexes. Representing the primary complexes by FeOOH and FeOOR he suggests the various reactions are: (i) Oxidation of alcohol by the primary hydrogen per'oxide complex:

(ii) Oxidation of alcohol by the primary alkyl hydroperoxide complex:

+ -____.

FcO~I$- ROH +

E'c :67 OR ,

L

4

I n these two reactions the choice of the reactive oxygen atom in the peroxide group is purely arbitrary. (iii) Hydrogen peroxide reacting with the primary alkyl hydroperoxide complex : FerO)]OR ..-d

L--,

;H.!oo;Hj

+

FcOH -I- ROH 4

0:

+

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

403

(iv) Hydrogen peroxide reacting with the primary hydrogen peroxide complex, i.e., the “key” reaction in the catalytic decomposition of hydrogen peroxide : Fc :6: OH

f

.__a L--

IH.;oo!.H;

-

FeOH

+ €I:O + 0 2

Reaction (iii) as written is preferred t o the following alternative F c ; ? OR + 0

,

H i0.i OH

-

E’COII + l i O B

+

0 2

on the grounds that if this occurred an analogous reaction with an alkyl hydroperoxide would be possible : FC 8

€I:O

OR ,

+ H j.Q OR

+

-

FeOH

+ 2ROH + 0 2

and there is no evidence .at all for such a reaction evolving oxygen in the spontaneous decomposition of the catalase hydroperoxides. Important evidence in favor of reaction (iv) as the “key” reaction in the catalytic peroxide decomposition is the fact that the rate of formation of the primary hydrogen peroxide complex is high enough t o account for the entire catalytic activity, provided a mechanism for its rapid decomposition is available. The bimolecular formation velocity constant has a value of about 3 )( lo7 M.-l set.-' as compared with the value of the overall bimolecular constant for the catalytic reaction of 3 - 3.5 x lo7 M.-1 sec.-l. Any simple mechanism with consecutive reactions including the unimolecular decomposition must be excluded, for the decomposition rate of the primary complex is extremely low, the first order constant being about 0.02 set.-' (Chance, 69). The noncompetitive inhibition of the decomposition of hydrogen peroxide b y cyanide is not immediately obvious from the above reaction mechanism for if cyanide can compete in the formation of the peroxide complex which is responsible for the oxygen evolution in step IV, competitive inhibition might be expected. However, under the experimental conditions necessary t o observe peroxide decomposition, a n excess of peroxide is required and this is sufficient t o give the maximal concentration of the peroxide complex, 1.2 or 1.6 moles of bound peroxide for each erythrocyte or bacterial catalase molecule respectively, i.e., the peroxide complex concentration is independent of the peroxide concentration. Analysis of the system under these conditions shows noncompetitive inhibition t o hold.

404

P H I L I P GEORGE

Theorell (79) has suggested an alternative mechan-ism based on the combination of substrate and acceptor molecules with two hematin groups. Chance (78) has pointed out that this mechanism cannot apply t o the primary alkyl hydroperoxide complexes reacting with an alcohol or hydrogen peroxide (as in reactions ii an'd iii) because all hematin groups are attached t o hydroperoxide molecules in these complexes; however, it is applicable to the catalytic decomposition of hydrogen peroxide. The mechanism may be represented: HOFe-FeOOH

HOFe-FeOH

I

HOFe-FeOH

!

HOFe-FeOH

T

HOFe-IFe-OOH

-4 I HzOz

+ Hz0

HOFe-IFe-OOH

+HzO

+

0 2

The essential step evolving oxygen is the intramolecular reaction of the complex containing the two combined peroxide groups. There is no experimental evidence for this complex and for the mechanism to agree with Chance's data the complex would have t o have a very short half life. Chance has added that the following points are in favor of such a mechanism. First, it explains why peroxidase is a feeble catalyst for peroxide decomposition; with only one hematin group per molecule this type of mechanism is excluded. Secondly, it explains why methernoglobin is a poor catalyst, for methemoglobin only gives an inactive secondary peroxide complex. Thirdly, it accounts for the decrease in catalase activity with a decrease in the number of intact hematin groups. But this mechanism has the serious disadvantage that it cannot apply t o the reaction of the primary alkyl hydroperoxide complexes with hydrogen peroxide and in addition it requires further elaboration t o explain the peroxidatic action of catalase. Certain other aspects of these mechanisms will be discussed later in comparison and contrast with the reactions of ionic iron.

VI. THEMECHANISM O F THESEH E M O P R O T E I N REACTIONS A very valuable approach t o the problem of the general mechanism of these reactions has been made by Theorell (2) who has emphasized the manner in which structural differences between the various hemoproteins can be correlated with the appearance of highly specific reactions. An important feature of these particular structural groups is the interaction which exists between the reaction of the group and the specific reaction of the heme or hematin group. For instance, when oxygen is attached to hemoglobin in neutral solution, the pH is lowered

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

405

as a result of the ionization of a neighboring acid group. Theorell has named groups of this kind heme-linked groups and has examined the reactions of a wide range of hemoproteins for this type of interaction. There is not sufficient data yet to decide whether such groups play a primary role in the kind of oxidation-reduction reactions discussed above and undergo oxidation-reduction themselves, or whether their function is secondary in that by ionization the groups affect the speed with which the heme or hematin group reacts. I n the case of the hemoglobin, a mechanism will be advanced later in which the assumption of a n auxiliary electron accepting group on the protein molecule makes it possible to explain the autoxidation kinetics. But first it is necessary t o examine the several mechanisms which have been suggested to explain the various oxidation-reduction reactions. It has been shown above that the catalytic action of catalase and peroxidase is intimately connected with the ability of these hemoproteins to form complexes with hydrogen peroxide (or alkyl peroxides). By choosing the experimental conditions the existence of three different complexes can be demonstrated spectroscopically. The chemical nature of these complexes is as yet unknown, and mechanisms have been represented as bimolecular reactions between substrate and complex. It is difficult to reconcile such mechanisms in which some four strong covalent bonds are broken and then re-formed simultaneously with the high speed and low activation energy of these reactions (Bonnichsen, Chance, and Theorell, 64), and some stepwise mechanism seems more likely. These kinetic features in the case of oxidation-reduction systems are usually associated with very simple reactions such as electron transfer or reactions in which only one bond is broken and another formed. The reactions of hydrocarbon free radicals, as elucidated by the “sodium flame” technique developed by Polanyi and his school (80) and the part they play in the oxidation of hydrocarbons as evidenced by many kinetic studies are cases in point (Pease, 81; Steacie, 82). The mechanisms proposed by Lemberg and Legge ( 5 ) have a n added disadvantage. An acceptor group joined to the protein or the hydrogen accepting substrate linked to the protein is made to react with HzOzor OzH- linked to the hemoprotein iron atom by an intramolecular rearrangement. With the reacting groups attached in this way some distance apart, it is difficult t o see how a reaction involving the breaking of several bonds and transfer of atoms or fragments of molecules can occur, for the gain in energy coming from the partial formation of the new bonds, which can normally offset a large activation energy needed to break the old bonds, would appear to be excluded. Some of these difficulties may be illusory if the structure of the or [Fe,-H202]+ but is peroxide complex involved is not Fe,-OOH

406

PHILIP GEORGE

somewhat simpler (Fe,+ represents the ferriprotoporp hyrin group joined t o the protein). The existence of three complexes with very distinct spectroscopic characteristics formed from the same two molecules suggests that one or two of these are reaction products or degradation products of the two components. At the most two might be accounted for by Fe,-OOH and [Fe,-H20z]+. There is no evidence that any of these complexes has involved an irreversible reaction with a group on the protein, and it is not easy t o see how any such reaction should alter the absorption spectrum so markedly in the visible region. Likewise there is no evidence for irreversible attack on the porphyrin ring itself. The experiments of Chance show the interrelation of two of the complexes. The green complex of peroxidase (complex I) changes t o the red complex (complex 11) according t o first order kinetics. It seems highly likely then that the possibility of a complex itself undergoing a reaction must be taken into account in a complete reaction mechanism, although under some experimental conditions such a reaction may not play a dominant role. In its wider imp1ic;ations this means that the entire question as to whether a valency change occurs and the conjugate problem of the participation of free radicals as intermediates must be reconsidered. I n discussing this a comparison has often been made between hemoprotein reactions and those of ionic iron (Stern, 3). Calculations have been made showing that the relative activity of catalase, hemin and ionic iron as catalysts for hydrogen peroxide decomposition are in the ratio lo6: 10P. However, such calculations can be misleading, for the reactions show different dependencies on the hydrogen ion concentration which make a valid comparison impossible. The 10l0-fold greater efficiency of catalase as compared with ionic iron, which these figures show, could be taken t o suggest an entirely different type of reaction mechanism. Yet if ionic iron were not precipitated in alkaline solution, a t pH 10.46 it would have a catalytic activity equal to the maximum activity of catalase itself. This conclusion follows from the values for the velocity constant k in the rate equation. -dHzOz/dt

=

k[H~O~][Catalyst]

which for catalase and ionic iron (neglecting the formation of hydroxides) are 3.5 x 107 and I .2 x 10-3/[H+]M.-1 set.-', respectively. For catalase and peroxidase there is no direct evidence that a valency change of the iron atom occurs except when azide js added to catalase, and peroxidase is acting as an oxidase in the autoxidat ion of dihydroxymaleic acid, when both systems can be poisoned by carbon monoxide. The normal absence of such poisoning cannot be taken as proof that the

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

407

iron remains in the trivalent state throughout the reaction, for an alternative explanation is that although a valency change occurs, other reactions competing for the ferrous form are more rapid than the combination with carbon monoxide. In this connection the recent experiments of Keilin and Hartree (45) with methemoglobin and hydrogen peroxide are significant, for the system shows a well-developed peroxidatic and catalatic action and reduction of the iron is apparent from the formation of HbOz, orHbCO when carbon monoxide is admitted. In a like manner the different oxidation products formed when peroxidase and ferrous ion are used in peroxidatic reactions (109,110,111) do not prove that the underlying mechanism is different. The mechanism in the case of ferrous ion involves competition between Fe++ and Fe+++ for radical intermediates, and it is quite possible with reactions of such complexity, particularly if they proceed through several distinct oxidation stages, that such competition reactions could lead to different reaction products predominating, depending on the catalyst employed. In the absence of direct proof of a valency change or free radical intermediates participating, it is valuable to survey the data on the reactions of ionic iron in similar systems and then examine the possibility of similar mechanisms occurring in hemoprotein reactions.

1. Free Radical Mechanisms with Ferrous and Ferric Ions The evidence for the free radical mechanisms of the reaction between ferrous and ferric ions and hydrogen peroxide is fully discussed in the article by J. H. Baxendale in this volume, and it is necessary here only to summarize and comment on those features especially relevant to hemoprotein reactions. This evidence is essentially indirect. Experiment shows very reactive intermediates t o be present and extensive kinetic studies reveal competition reactions for these intermediates in that the overall order of the reaction is found t o depend on the reactant concentrations. A free radical mechanism is adopted because it accounts for the chemical reactivity of the system in the oxidation of substrates (Fenton's reaction) and the initiation of the polymerization of vinyl compounds (Baxendale, Evans, and Park, 84) and it provides a set of reactions which largely account for the observed kinetics. The set of reactions which fit best the most recent experimental data is that proposed by Barb, Baxendale, George, and Hargrave (83) : Fe3+ Fez+ Fez+ HO Fez+ Fe*+

+ OzH- Fea++ HOz + Hz02 Fe3++ OH- + HO + HO Fe3++ OH+ HzOa H a 0 + HOz + HOn + Fe3++ OZH+ + Fez+ + O2 + -+ 4

+

02-

(i') (0) (1) (2)

(3) (4')

408

PHILIP GEORGE

(Reactions i and 4 would refer to corresponding reactions with un-ionized

HzOzand HOz.) This reaction scheme is to be preferred to the original mechanism proposed by Haber and Weiss (85) in which oxygen was evolved in the step HOa

+ Ha02 -+

0s

+ Ha0 + HO

(w)

for two main reasons. The Haber and Weiss mechanism is incompatible with the recent kinetic data, and it does not provide a consistent mechanism for the reaction of ferrous salts with hydrogen peroxide and the catalytic decomposition of hydrogen peroxide by ferric salts. George’s stoichiometric experiments with potassium superoxide, KOz, also indicated that reaction w is insignificant when ferric and ferrous ions are present in solution (86). Although superficially similar in so far as some reactions occur in both mechanisms the kinetic differences between the two are fundamental. In the Haber and Weiss mechanism the ferrous and ferric ions only initiate and terminate a reaction chain propagated by steps w and 2 above which involve the two chain carriers, the HO and HOz radicals. I n the new mechanism there are three chain carriers, the radicals HO and HOZ and the ferrous ion itself. The new mechanism is supported by the elucidation of the thermochemistry of the various steps in terms of the ionization potential ( I ) of the ferrous ion in aqueous solution, the electron affinities ( E ) of the radicals HO and HOz plus the heats of solvation (8)of the corresponding ions OH- and OSH- in water, the various 0 * H arid 0 * * 0 bond strengths ( D ) and other thermal quantities such as heats of evaporation (A) and heats of solvation (8) (see 83,84,87,88,89). A411the steps with the exception of (i’) and (0) are exothermic, and being; electron transfer reactions or a simple bond breaking reaction (step 2) are to be expected to proceed extremely rapidly as is required by the kinetic mechanism. Steps (i’) and (0) are endothermic to the extent of 28 and 5 kcal., respectively, and these values are in accord with their activation energies which are 28 5 8 and 9.4 kcal. As with the kinetic evidence this thermochemical evidence is indirect support for the radical mechanism. An important point which will be elaborated later is that thermal data of this kind provide criteria for rejecting certain steps when the endothermicity becomes impossibly large for the reaction to proceed a t a sensible rate. There are some further aspects of the ionic iron-hydrogen peroxide system which have a possible bearing on hemoprotein reactions. Apart from the reactions of the HO radical listed above (1 and 2) and reactions with oxidizable or polymerizable substrates, there is Eome experimental evidence that two additional types of reaction may occur. The first

- -

-

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

409

is electron transfer from anions, e.g., HO.

+ C1- --+

OH-

+ C1.

a reaction which is energetically favored. The second, and perhaps the more significant, is a reaction with the ferric ion itself. The evidence for this is not so clear cut as in the previous case, yet it is difficult to see any other interpretation for the observation that in the presence of higher concentrations of ferric salts than those usually employed (10-1 M rather than M ) the ratio of the velocity constants of steps 1 and 2 has a different value, control experiments having shown that a normal kinetic salt effect was not responsible. Except that it takes part in reactions similar to those of the HO radical, there is no evidence as to the nature of this ferric ion-HO radical reaction. One possibility is the formation of a ferryl ion FeO++, a compound of quadrivalent iron: FeS+

+ HO

--t

FeOHS+ -+ Fe02+

+ H+

The participation of the ferryl ion in the ionic iron-peroxide system was suggested in 1932 by Bray and Gorin (90) although its relationship to, and possible role in replacing, the HO radical was not realized. Cupric ions which for some time have been known to accelerate the ferric ion catalyzed decomposition of peroxide (Bohnson and Robertson, 91) have also been shown to affect the ferrous ion reaction (83). There is now good kinetic evidence that this arises from the cupric ion reacting more rapidly than the ferric ion with the 02-: cu2+

+ oz--t

Cu+

+ Fe3+-very fast

cu+

+

01

followed by Cu2+

+ Fez+

The second step accounts for the fact that the action is purely catalytic and that a limit can be reached above which further additions of cupric ion have no effect (Barb, Baxendale, George and Hargrave, 83). One point of marked resemblance between ferric ions and the hemoproteins in their ferric state is the formation of a complex with hydrogen peroxide. By the appropriate choice of peroxide and hydrogen ion concentrations a deep brown-red colored complex is formed with ferric ions. Evans, George, and Uri (88) showed this to be an “ion pair” complex formed: FeS+

+ OzH-

[FeOZHla+

The heat and entropy of formation of this complex have values in accord with the data on other ferric ion pair compiexes, e.g., FeOH2+ and FeCI2+. In contrast with the apparent behavior of the hemoprotein-

410

PHILIP GEORGE

hydrogen peroxide complexes this ferric ion complex plays no distinct part in the reaction mechanism. A bimolecular reaction between the complex and another hydrogen peroxide molecule does not occur. The kinetics of the peroxide decomposition can be explained if the initiating reaction is (i') in the above set of reactions, and there is no means of telling from the kinetics whether reaction (i') occurs by unimolecular decomposition of the ion-pair complex or simple collisions between Fe3+ and OzH-, or both (83). Recent studies have shown that the oxidation of a substrate by the ferrous ion-hydrogen peroxide system is a reaction of great complexity. I n the absence of oxygen the following general scheme offers an explanation for the chief kinetic features. Kolthoff and Medrtlia (92,93) found it to hold for the oxidation of ethyl alcohol to acetaldehyde and Barb, Baxendale, George, and Hargrave (83) found it to apply to the oxidation of traces of organic impurities in distilled water. H.ZA represents the oxidizable substrate.

+ + + + + + +

+ + + + + + + +

Fez+ HZOZ + Fe3+ OHHO Fez+ HO + Fe3+ OHHzA HO+ HA HzO Fe3++ HA+ Fez+ HA HA+ OH-+ HAOH HAOH HO + HOA HzO Fe3+ A 0 Fez+ H+ HOA --f

The presence of oxygen profoundly affects the reaction. This is similar to the effect of oxygen on the initiation of the polymerization of vinyl compounds by the ferrous ion-hydrogen peroxide system (Baxendale, Evans, and Park, 84). It is due t o the addition. of oxygen t o the organic radical, and when oxygen is present in these oxidation systems the following reactions have also to be taken into account (84,92,93).

+ -+

HA 0 2 + HAOz HAOz Fe2+ H+ -+ HAOzH Fea+ Fez++ HA0 4-Fe3+ OHHAO,H HzA+ HAOH HA HA0 HA HA02 HzA + HAOzH Fe3+ HA0 Fez+ H + + HAOH

+

+

+ + + +

+ + + + +

The use of ferric ion and hydrogen peroxide as an oxidizing system has been very little studied. This is surprising because it provides the simplest model for peroxidase. Walton and Christiansen (94) showed that ethyl alcohol could be oxidized at 30°C. and the reaction was retarded in more acid solutions. The apparently low catalytic activity of this system would appear to be due t o the slowness of the initiating

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

411

reaction (i) arising from its big activation energy Fe3+

+ H2OZ

Fe2+

4

+ HO, + H+

(i)

compared with the corresponding reaction with the ferrous ion which is rapid Fez+

+ H2OZ+ Fe3+ + OH- + HO

(0)

At 25°C. and pH 2.0 the velocity constants ki and k, have the values 9.1 X and 53.0 M.-l sec.-l, respectively. For the reactions with hydrogen peroxide there thus exists a large accumulation of experimental data t o compare with the data on hemoprotein reactions. However, this is not the case for the reaction with molecular oxygen where the experimental data is very scanty. Weiss (95) has suggested the following radical mechanism for the autoxidation of ferrous salts and this has been generally accepted.

+ O2 Fe3++ 02+ H+ HOz 3 Fez+ + HOz Fe3+ + OzHOzH- + H+ 2 HzOz 2Fe2+ + HZO2 2Fe3+ + 20Ha

Fez+

4

02-

-+

-+

From the appropriate stationary state equations it can be shown that the oxidation rate should be:

and initially when Fez+>> Fe3+ -dFeZ+/dt = 4k,FeZ+02

The experimental data of Lamb and Elder (96), however, are not in agreement with this predicted rate expression for they find the initial rate proportional t o the square of the ferrous ion concentration and directly proportional t o the oxygen pressure. This has recently been confirmed by the author (97), and it would appear that either the autoxidation is subject t o a true catalysis by trace impurities (an induced reaction is excluded by the total ferrous ion oxidized being large, about M / 2 0 ) or the actual mechanism is different from t h a t suggested by Weiss. 2. An Examination of the Possibility of Similar Free Radical Reactions

with Hemoproteins The purpose of this discussion is t o see which of the reactions listed above for the free ferrous and ferric ions are possible with the hemo-

412

PHILIP GEORGE

proteins. An examination of some of the experimental data on these reactions to try to decide whether they do in fact occur will follow in the next section. In considering the thermochemistry of the above reactions the only quantity which changes when ionic iron is replaced by another metal or coordination complex is the ionization potentisil I of the reduced form. The following list gives the heats of the reactions (i’--4’) in terms of electron affinities, solvation energies, bond strengths, etc., as given above, using the most recent value for these quantities (83,87,88,89). Reaction (i’): Fe3+

+ OzH-+

+

Fez+ HOZ +I - ( E H o ~ SO~A-)S H O=~+I - 123 kcal. OH Reaction (0) : Fez+ HzOz -+ Fe3+ OHQo = - I - X H ~ O , - S H ~-ODEO-OH ~ (EHO SOH-) SHO== - I Reaction (1): Fez+ OH + Fe3+ OHQ,r

=

+ +

+

+

+

+

+ +

+

+

+ 89.7 kcal.

- I + (EHO + SOH-)- SHO= - I + 137.6 kcal. + HOZ + Fe3+ + OZH= - I + ( E H o+ ~ SOZH-) - S m a = - I + 123 kcal. Reaction (4’) : Fe3+ + + Fez+ + Qr I - (Eoz + Soz-) + Soz = I - 76 kcal. Fe3+ + Reaction (a): Fez+ + QB = - I + (Eoz + Soz-) - So2 = - I + 76 kcal. QI

=

Reaction (3): Fez+ Q3

02-

0 2

=

0 2 -

02-

Unfortunately there is not sufficient thermal data for direct calculation of I in the case of heme, hemochromogens, and reduced peroxidase. In the first three cases, however, estimates can be based on a knowledge of their oxidation-reduction potentials. The ionization potential I for a metal complex M2+ionizing in solution is given by the heat of the reaction: and since

I can be obtained from the heat of the cell reaction ( - A H 0 ) M(aq.)’+

+ H++

M(aq.)’+

+ 3Hz - A H o

from the equation: - A H o = 82 - I

or I = A H o

+ 82 kcal.

Knowing the standard free energy change in the cell reaction, A H o can be calculated if ASo is known. From standard entropy values the change Hf to &Hz corresponds to f15.6 e.u. and so only a n estimate of the entropy change associated with the valency change of the complex is required to give a value for A H o and hence I .

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

413

In'the case of such large ions as the metalloporphyrin complexes with their rigid planar structure, t o a first approximation the entropy change will be given by the difference in the Born charging entropy of the two valency forms. The Born charging entropy of an ion is given by: 8B.C. =

-2R D 2

where Z is the magnitude of the ionic charge, e the electronic charge, R the radius of the ion, D the dielectric constant of the medium, and ( d D / d T ) , the temperature coefficient of the dielectric constant at constant pressure (Eley and Evans, 98). Taking heme as a sphere of radius 5.8 A. equivalent in volume to its platelet size of 15 X 15 X 3.7 A.3, the change in entropy from the reduced to oxidized form amounts to - 1.7 e.u., considering only the central iron atom (i.e., change in charge 0- +1) and +5.0 e.u., taking into account the ionized propionic acid side chains of the protoporphyrin (i.e., change in charge - 2 + -1). The total entropy change in the cell reaction is thus 13.9 or 20.6 e.u., which give for TAXo, 4.1 or 6.1 kcal. As the difference between these values is small considering the assumptions made in the calculation, a n approximate value of TASO of 5.0 kcal. has been used in the calculation of I listed in Table VIII, which also includes values of I for Fe2+, Co2+,Cr2+, Cu2+and Fe (CN)64-for comparison. TABLE VIII Calculations of Ionization Potentials in Aqueous Solution

Eo OT Eo'

AG

Complex or Metal Ion

Volts

ICcal.

Heme.. . . . . . . . . . . . . . . . .. a Cyanide hemochromogen . . a Nicotine hemochromogen . .Q Globin hemochromogen . . .a Hemoglobin. . . . . . . . . . . . . .b Myoglobin . . . . . . . . . . . . . . .c Fez+. . . . . . . . . . . . . . . . . .d coz+. . . . . . . . . . . . . . . . . . e Cr2+. . . . . . . . . . . . . . . . .. * c u * + .. . . . . . . . . . . . . . . . .I Fe(CiY)sa-. . . . . . . . . . . . . I

+0.110 +O .183 -0.184 $0.098 -0.144 -0,046 -0.77 -1.84 + O . 41 -0.167 -0.47

- 2.5 - 4.2 4.2 - 2.3 3.3 1.1 +17.8 $42.5 - 9.5 3.8 +10.8

+ + +

+

TASO Kcal.

Kcal.

+ 5.0 + 5.0 + 5.0 + 5.0 + 5.0 + 5.0

+ 2.5 + 0.8 + 9.2 + 2.7 + 8.3 + 6.1

4.9 4.9 4.9 7.4 +14.2

$12.9 +37.6 -14.4 - 3.6 +25.0

-

A€€O

I Kcal. 84.5 82.8 91.2 84.7 90.3 88.1 94.9 119.6 67.6 78.4 107.0

~

a

ED'values at pH 7.0 and 30'C. (Guznian Barron, 99).

Taylor and Hastinga (100). Taylor (101), Wyrnan and Ingalls (102). Taylor and Morgan (103). Evans, Baxendale, and Uri (89). * Eo values from Latimer (104). Entropy changes assumed to be identical with Fez+ systems. I EDand entropy of Cu*+from Latimer (104). Entropy of Cu+ assumed identical with Na+. 0 Data given in Butler (105) corrected for AH0 = -82 kcal. for H + e+ 1Hz (89). In a, b , c the free energy changes are not strictly standard free energy changes, for the redox potentials refer to unit concentrations and not unit activities of the reactants. b

+

414

PHILIP GEORGE

The only available check with experiment for the heme compounds is in the case of hemoglobin where a t ph. 7.0, Eo' = -0.152 at 24°C. (Conant and Pappenheimer, 106) and -0.144 or -0.14!3 (Taylor, 101) at 30°C. A H o of the cell reaction is thus 12.6 or 6.9 kcal. and the corresponding value of I , 94.6 or 88.9 kcal. The calculated value is in tolerable agreement with these experimental values, but it is clear that a thorough investigation of the variation of Eo' over a much wider temperature range on the same sample of hemoglobin and using the same experimental method is needed to establish a reliable value. The fairly close agreement between calculated and experimental values does show, however, that the assumption regarding the entropy change awociated with the valency change is not far wrong, and the entropy change is in fact small. It would appear from this that no marked reorganization of the protein structure accompanies the valency change. These calculations of the ionization potential show that heme, hemochromogens, and hemoglobin should in their reduced form react as rapidly or more rapidly than ferrous ions. Reactions (3) and (4) are again exothermic and should be very rapid. Reaction (0)is about thermoneutral instead of 5 kcal. endothermic, and so this direct reaction of the ferrous form with hydrogen peroxide should be faster than with ferrous ion itself. Reaction (a), a step to be considered in an autoxidation mechanism, is also less endothermic than in the case of ferrous ion and again a higher rate is to be expected. Reaction (i'), between the ferric form and OzH- is conversely more endothermic and should be much slower than with ferric ions because of this and also because the temperature independent factor should also be less than lo2*,the value found for the ferric ion reaction. This high value is attributable to reaction taking place between oppositely charged ions, one carrying a high charge. If this reaction is considered as taking place between the ferric form and the H202 molecule (instead of the OzH- anion) it will be 8.2 kcal. more endothermic on account of the heat of ionization of the hydrogen peroxide. A value of the bimolecular velocity constant for a hemin derivative reacting with hydrogen peroxide at ph. 7.0 and 20"C., assuming I = 90 kcal., the activation energy equal to the endothermicity and a normal temperature independent factor of lo", is given by: k =

1O1I . exp . (90 - 123) X lo3/&" X 1.8 X lo-'* = 3.6 X 10-7

where 1.8 X 10-l2 is the dissociation constant of hydrogen peroxide. I n spite of the very approximate nature of this calculation arising from uncertainty as to a good value for the temperature independent factor there can be no doubt that the magnitude of the velocity constant will

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

415

be very small indeed, and it is very unlikely that this type of reaction can play any significant part in heme or hemoprotein reactions. This is especially true in the case of catalase which has not been mentioned in this section because all attempts to prepare the ferrous form have failed. The most likely reason for this is that its ionization potential is very low and the ferrous form is able to react with any oxidizing entity present, even H+ ions. It may be concluded from this discussion that heme, hemochromogens, and the hemoproteins could undergo similar free radical and electron transfer reaction as the free ferrous and ferric ions. On thermochemical grounds reactions (l), (3), and (4)should be very rapid and reactions (0) and (a) should be more rapid than with ferrous ion in contrast to reaction (i), the ferric form reacting with hydrogen peroxide which should proceed extremely slowly. 3. The Applicability of These Free Radical Mechanisms in the Hemoprotein

Reactions a. Reactions of Peroxidase and Catalase. The observation (Dainton and Smith, 107) that neither catalase nor peroxidase can initiate the polymerization of vinyl compounds in the presence of hydrogen peroxide cannot be taken as evidence that free radicals are not formed when they react with peroxide. This observation resembles the absence of carbon monoxide poisoning in the reactions of peroxidase and catalase. It could be that other reaction steps compete for the radicals in the first case and the reduced form of the enzyme in the second case to the exclusion of the test reactions with vinyl compounds or carbon monoxide and so lead to a negative result. However, the above estimate of the heat of the reaction when a hemin derivative reacts with hydrogen peroxide in a manner analogous to the ionic ferric ion reaction: Fe3+

+ OzH-

+

Fez+

+ H01

makes it extremely unlikely that a reaction of this type is the initial redox reaction with peroxidase and catalase. The primary and secondary complexes that these enzymes give with peroxide appear to play the part of intermediates in an alternative reaction sequence by which the peroxide molecule undergoes reduction. In this respect this process is diametrically opposed to the ionic ferric ion reaction above which is the first step by which the peroxide molecule can undergo oxidation to molecular oxygen. Chance has not discussed in detail possible structures for the primary and secondary complexes except to point out that the Soret bands of the primary complexes with catalase resemble those of open chain porphyrin degradation products like degraded hemoglobin and

416

PHILIP GEORGE

specimens of catalase containing a high bile pigment content. There can be no doubt that peroxide combines directly with the iron atom of the hematin group and he suggests that “speculation a,s t o whether the porphyrin ring is actually oxidized on formation of the primary complex by electron transfer from the iron peroxide complex and is then reduced on reaction with the reducing substrate or acceptor, affords very interesting possibilities” (73). I n the absence of direct evidence all arguments a t present as t o the nature of these peroxide complexes can only be based on analogy with other hemoprotein derivatives and their reactions. Chance’s suggestion t hat the primary catalase complexes involve attack on the porphyrin ring does not seem very likely when comparison is made with peroxidase. The primary peroxidase complex resembles the primary catalase complex particularly in the Soret region, but the structure with a modified porphyrin ring is attributed to complex I V formed with peroxidase in the presence of excess alkyl hydroperoxide. I n the visible region the spectra of the primary complexes resemble the free enzymes, the very rapid formation of the complexes, and in the case of catalase the competitive formation in the presence of cyanide, all suggest that the formation involves a simple bimolecular displacement. If this is so then either H202or 02H- is joined to the iron atom. The high affinity of the hematin group of the enzymes for hydrogen peroxide tends t o exclude the bonding of the H202 molecule as such, for this type of reaction is in the category of replacement of a bound OH- anion by H202 of solvation and this reaction would appear unfavorable both on considerations of heat and entropy changes. Bonding of the OzH- anion is more likely t o account for the high affinity particularly if an OH- anion and not a H 2 0 molecule is to be replaced. It is interesting to note t h a t a compound of this type is not restricted to hemoproteins. The complex obtrained with ionic ferric iron is of this type having an “ion pair’’ structure Fe02H2+(Evans, George, and Uri, 88). Nevertheless there is an important distinction in the function of such complexes in the two systems. With ionic iron the compl ex plays no independent kinetic role, electron transfer predominates and there is no indication whether it occurs by unimolecular fission of t h e complex or direct bimolecular reaction hetween Fe3+ and 02H-, but the net result is one step toward the oxidation of the peroxide molecule. With the hemoproteins the peroxide complexes play a dominant kinetic role connected with a direct overall reduction of the peroxide molecule. It is true that with excess peroxide catalytic decomposition occurs in all cases which may be regarded as a mutual oxidation reduction of one molecule by the other. The distinction in the two cages is in the fate of the peroxide molecule involved in the complex.

THE SPECIFIC REACTIONS

OF IRON IN SOME HEMOPROTEINS

417

The conversion of the green primary complex into the pale red secondary complex appears t o be a reduction process even though it occurs in the absence of any added hydrogen donors. The most definite evidence for this is the case of peroxidase where the speed of the conversion is increased in the presence of all compounds with which the peroxide system reacts (Chance, 5 5 ) . For catalase, where the conversion can only be obtained with alkyl hydroperoxides, the evidence is not so clearcut, but a t least the velocity of formation of the secondary complexes increases as the hydroperoxide concentration is increased. An alternative explanation for these effects would be that the primary and secondary complexes are in some sort of equilibrium where removal of the latter would have the effect of increasing the rate of conversion. There is no indication of any such equilibrium, however, and direct reduction of the primary complex appears t o be the most likely explanation. One possible formulation for this change involves the production of a “ferry1 ion’’ type of compound by the removal of an OH radical by the hydrogen donor from the OzH- anion bound t o the iron atom: Prot Fe,OOH

+ AH2-

Prot Fe,O

+ HzO + AH.

(green primary complex) (pale red secondary complex)

The “ ferrylperoxidase ” being nominally a compound of quadrivalent iron would have oxidizing properties corresponding to one equivalent. This raises a difficulty encountered in any attempt to formulate a more detailed reaction mechanism taking account of the increased rate of formation of the secondary complex in the presence of the hydrogen donor. Chance (54)has shown that the secondary complex when formed takes part in a bimolecular reaction involving two equivalents of the hydrogen donor. As hydrogen peroxide initially only has an oxidizing capacity of two equivalents, this is incompatible with the conversion itself being a reduction process. * Further experimental studies are required to resolve this difficulty. In view of the complications which can occur with the ferrous ion-hydrogen peroxide system when it is used as an oxidizing agent in the presence of molecular oxygen, a test of the effect of removing oxygen from these enzyme reactions appears desirable. It is abundantly clear that further discussion of the possible free radical nature of subsequent reactions in these enzyme systems must await a more detailed knowledge of the structure of these complexes. The recent experiments by Keilin and Hartree (45) showing most clearly that methemoglobin and metmyoglobin are reduced when they

* This difficulty may only be apparent, for the bimolecular reaction was largely studied in the presence of an excess of hydrogen donor where the kinetics become pseudofirst order and the underlying stoichiometry can be obscured.

418

PHILIP GEORGE

react with hydrogen peroxide in sufficiently high concentration to cause catalytic decomposition, taken in conjunction with their previous demonstration (1G,17) of the reduction of azide catalase by hydrogen peroxide, are indisputable evidence that a reaction path for peroxide decomposition by hemoproteins involving a change of valency does exist. It is difficult, however, to formulate a mechanism for the methemoglobin and metmyoglobin reactions because again a peroxide complex of uncertain structure takes part. The experiments of George and Irvine (1 16) on the metmyoglobin complex showed t h a t it is probably formed in a reaction of the type MetMb

+ H 2 O 2 + Complex + X

(2 oxid. equiv.) (1 oxid. equiv.) (1 oxid. equiv.)

and that a transient oxidizing entity X is produced along with the complex, which can react with added reducing agents if present initially, but which disappears if the complex formation is allowed to proceed independently. A tentative mechanism which can account for these results is the formation of an OH radical and a “ferTy1’’ compound of myoglobin in the reaction

+ Prot Fe,O + OH + €LO+ + H 2 0 2 + Complex + X

Prot FeP+(H20) i.e.,

MetMb

H202+

The reduction of the iron t o the ferrous form in the presence of excess hydrogen peroxide which was observed b y Keilin and Hartree (4.5) could then occur by ferrylmyoglobin reacting in a simple bond breaking reaction Prot Fe,O

+ 02H-

Ferrylmyoglobin

-+

Prot FepO*

+ OH-

Oxymyoglobin

The same reaction in the case of peroxidase would suggest that the deep red complex peroxidase-H202 (111) has the oxyheinoglobin type of structure. The participation of ferry1 compounds in this manner offers an alternative way for a ferric compound to react in a series of simple steps comparable t o the simple free radical reactions. In a more recent series of experiments, George (97) has shown that the secondary hydrogen peroxide complexes of horseradish peroxidase and cytochrome-c pcroxidase can be titrated with ferrocyanide or ferrous ions and alvo appear to take part in a one oxidizing equivalent reduction t o the ferric form of the enzyme. In this important chemical property they thus rcsemblc the metmyoglobin complex in spite of marked spcctroscopic differences in the visible region of the spectrum (Keilin and Hartree, 48). If the peroxide molecule is not a component part of the structure i t

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

419

was to be expected that other oxidizing agents might, be found to yield the secondary complex; HOCl, HOBr, Br03-, IOa- and (310, form a compound spectroscopically indistinguishable from the secondary H z 0 2or MeOOH complexes of horseradish peroxidase. Cytochrome-c peroxidase forms its secondary complex with all but HOCl and HOBr, which cause degradation of the enzyme. In a search for other redox systems capable of reacting with either the free enzyme or the secondary complex it was found that chloriridite ions reduce both secondary complexes whereas ferrous tris-dipyridyl ions or ferrous tris-ortho phenanthroline ions do not. This indicates that the redox potential for the reaction Complex I1

+ electron

4

Ferric form of the enzyme

lies between -1.02 and -1.06 volts, these being the potentials for the chloriridite ion and ferrous tris-dipyridyl ion couples respectively. These experiments suggest that the secondary complexes should no longer be regarded as Michaelis-Menten enzyme substrate complexes but as reaction intermediates in the same sense that free radicals and semiquinones are reaction intermediates, for all three classes of compounds provide a path for stepwise reactions. As a consequence the accepted mechanism for peroxidase action needs revision. There are several structures which may be considered for the metmyoglobin complex and the secondary complexes of the peroxidases: a) A simple quadrivalent ion structure containing Fe4+ b) A derivative of quadrivalent iron such as the ferry1 ion, Fe02+(see George and Irvine, 116, 120) c) A diradical structure in which trivalent iron forms “one end of the radical,” the other end being a normal radical grouping at a methene carbon atom, a pyrollic carbon atom, a some other atom within the conjugated network of porphyrin ring and heme-protein linkage d) A higher oxidation state of the hematin group in which the electron has been removed from a n-orbital common to the ring as a whole. The existence of this type of oxidation state has recently been demonstrated by Cahill and Taube in the case of the structurally similar compounds-copper, iron, cobalt, zinc, and aluminum tetrasulfonated phthalocyanines (121). Structures a and d arise by simple electron transfer and only differ as to which is the lower lying energy level, whereas structures b and c require bond breaking reactions both in the formation and subsequent reaction of

420

PHILIP OEOROE

the complex. Further experiments are required to enable a choice t o be made. Preliminary measurements of the paramagnetic susceptibility show the iron in the metmyoglobin complex to be essentially ionjcally bound with p = 5.10 & 0.05 Bohr magnetons, close to the theoretical value of 4.90 calculated on a “spin only” basis for four unpaired electrons (George and Irvine, 121). This would be in accord with ionic structures of type a or b provided there was no large orbital contribution to the magnetic moment. I n the case of the secondary horseradish peroxidase complex, Theorell showed that the bonding is essentially covalent with xtn5 4,800 X c.g.s. units and it has been generally accepted that it is a covalent ferric compound with one unpaired electron. This value is somewhat uncertain since the solution examined contained in addition to the secondary complex some of the primary complex. More recently Theorell, Ehrenberg c.g.8. units for the and Chance (122) report a better value of 3,500 X methyl hydroperoxide secondary complex. This value is close to the susceptibility of the peroxidase-cyanide complex, 2,970 X c.g.s. units, and they conclude by analogy that the complex containsi ferric iron essentially covalently bound like the cyanide complex. A susceptibility value of this magnitude is very high for a compound with only one unpaired electron where the contribution arising from electron spin is 1,270 X c.g.s. units at 20°C. Contributions from orbital moment are believed to account for such deviations observed generally with covalent ferric complexes of the hemoproteins (Pauling, 4). It is interesting however that c.g.s. units for the complex is even a little the value of 3,500 X greater than the theoretical value of 3,390 X 1 0 P c.g.8. units required for two unpaired electrons calculated on a “spin only” basis. This is the number required for a covalent hexacoordinated compound of quadrivalent iron or its derivatives. Since there is no a priori reason for expecting the orbital contribution to be large or small for the secondary complex the susceptibility value quoted does not exclude a quadrivalent iron type of structure. b. Oxidative Reactions of Hemoglobin and Myoglobin Involving Oxygen. There is no evidence that the autoxidation of these compounds is complicated by the participation of complexes of unknown structure and since the overall chemical change occurring is the simple oxidation of a ferrous compound to a ferric compound, these reactions are more suitable for kinetic analysis. The intramolecular reaction mechanism for hemoglobin oxidation proposed by Lemberg and Legge in which only the intermediate Hb4(02)z undergoes a reaction has been discussed above. The fact that the mechanism cannot account for the oxidation at high oxygen pressures makes it

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINR

421

extremely unlikely that it is correct. It is, however, possibIe t o extend this mechanism and permit reaction between free heme groups and heme groups combined with oxygen in the intermediates Hb4(02) and Hba(O2)~as well as in Hb4(02)2, discarding the idea that in addition a hydrogen donor AH2 is necessary. Such a mechanism with “reacting pairs” of heme groups fits the experimental data far better than that suggested by Lemberg and Legge, and with the additional assumption that the pairs in Hb4(02)3 react a little faster than the pairs in the other intermediates a very close fit can be obtained. Thus it is not possible to reject this type of intramolecular mechanism outright, but there remain the inherent difficulties in understanding how such a reaction can take place between separated groups and why an identical reaction occurring intermolecularly and favored by collision should be excluded. The experimental studies of George and Stratmann (37) showing that the kinetics of myoglobin autoxidation are similar in every respect to those of hemoglobin suggest very strongly that a common reaction mechanism is at work. Since myoglobin contains only one heme group this would exclude all intramolecular mechanisms and focus attention on possible intermolec,ular mechanisms. A rate dquation for the autoxidation appropriate to an intermolecular mechanism which agrees well with Brooks’ experimental data has been discussed above. It has the form:

&- k a ( a - 2) . (1 - a ) ( a - ). dt

a--2

where (a - x) and x are the concentrations of unoxidized and oxidized hemoglobin or myoglobin and a the fraction of unoxygenated hemoglobin or myoglobin. The numerator in this equation is reminiscent of the corresponding term in the rate equation derived from the free radical mechanism for autoxidation referred to earlier, though this simple mechanism cannot apply in this case because the denominators in the rate equations are different, which means in effect that the free radical reaction would not be first order in unoxidized compound. It is possible to extend the free radical mechanism so that the derived rate equation has the required first order dependence by assuming that there is an auxiliary electron accepting group in hemoglobin and myoglobin that can act as cupric ions do in the ionic iron-hydrogen peroxide system catalyzing the reaction between Fe3+ and 0 2 - . To illustrate this the following symbols are necessary: 0-Fe2+, 0-Fe2+.02, and 0-Fe3+ represent hemoglobin, oxyhemoglobin, and methemoglobin with the electron accepting group in its oxidized state.

422

PHILIP GEORGE

When it accepts an electron, it is represented by @--Fez+ or e)-Fe2+*0z. The reaction scheme is then: 1

0-Fe2+.0z

O-Fe3+ 3

{ (

O-FeZ+ 0-Fe2+

+

02-

+ HOz 2 O-Fe3+ + OZH-

O-Fez+Oz @-Fez+ @-Fe2+02

--*

+

4 0 2 - 4

+ 0-~e3+

{

@-Fez+ + 0 2

@-Fe*+ 5 O-Fez+ .-+ fast

(

O-Fe2+.0z

+ 0--Fez+

followed by the rapid reaction of HzOz with two more hemoglobin molecules. The solution of the stationary state equations written in Brooks’ terminology is :

where kl, kz, ka, and k d are velocity constants of steps 1, 2, 3, and 4, respectively. If steps 4 and 5 predominate, i.e., the catalysis of step 3 which is the back reaction regenerating hemoglobin, then Eq. (2) reduces to

which is identical in kinetic form with Eq. (1). The fate of the HzOz is not kinetically significant, if it reacts with the ferrous iron of heme the number 4 appears in Eq. (3) ; if it reacts entirely with other groups the number 4 is replaced by 2: it merely affects the stoichiometry, not the kinetics. Essentially this mechanism is a competition for HO2 radicals; only a fraction of those produced in the initial electron transfer from the ferrous iron of heme to 0 2 lead to an overall oxidation because of the back reaction regenerating heme. The predominant back reaction is that catalyzed by the electron-accepting group in the hemoglobin, and it is by this reaction that the iron is protected against oxidation. It is held in the ferrous state, not in a static sense by its structural environment, but in the dynamic sense that the primary step giving rnethemoglobin is obscured by catalyzed back reactions regenerating heme, thus giving a slow net oxidation. There is another type of reaction of hemoglobin which can readily be explained by a similar radical mechanism. Acid or pyridine denature Hbh(02)k and liberate two of the four 0 2 molecules, a third one being evidently required for the oxidation of four iron equivalents, while the

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

423

fourth is used in unspecified oxidation reactions. Lemberg and Legge (5) have taken this as evidence for the existence of the two XH2 groups per molecule and represent the reaction : Hb4(Oz)4

+ O 2 + 2H20 2HzX +

+

Hb4* 402 Hb4*(OH)4i.e., oxidation of the Fe 4 2Hz0 2X

-+

Hb4*

-+

0

2

+

Hbr(On)r(HzX)z+ Hba*(OH)r(X)z

+ 202

where the asterisk indicates denatured globin. Lemberg, Legge, and Lockwood (41) and later Lemberg (40) studied the oxidizing action of acidified Hb4(02)4 on ascorbic acid and biliverdin. In the case of ascorbic acid, oxidation occurs instantaneously on acidification if the ascorbic acid is already mixed with the Hb4(02)4, but far less is oxidized if it is added shortly after acidification. With biliverdin there was little evidence for the instantaneous oxidation, but a slow oxidation was observed which was attributed to the formation of H102 in the initial reaction. The results were explained by assuming the formation of active oxygen which reacted directly in an oxyhemoglobin-ascorbic acid complex. These observations may also be interpreted by the free radical mechanism, 02- again being formed as in step 1 of the autoxidation mechanism 402-

+ 4H+-+ 2HzOz + 202

where MetHb4* signifies the acid hematin protein complex. 0 2 - reacts rapidly with water giving HzO2 and 0 2 (George, 86), and this will predominate, for even if some acid hematin-denatured protein complex reacts with 02-the corresponding heme complex will autoxidize regenerating 02-. The oxidizing action of the system will depend on whether the substrate can react with 0 2 - or the HOz radical with which it is in ionic equilibrium. If it does not, oxidation can still occur through the peroxidatic action of the H 2 0 2and acid hematin-denatured protein complex. The fact Ascorbic acid unlike biliverdin apparently reacts with 0 2 - . that 02-(H02) is not an extremely powerful oxidizing agent has been shown in the recent experiments of Barb, Baxendale, George, and Hargrave (83). The action of pyridine on Hb4(02)4 can be explained in a similar way for it denatures the protein and parahematin derivatives are formed.

+

Hbr(O2)r 402-

pyridine

Globin and pyridine parahematin

4H+ + 2Hz02

+ 202

+ 402-

424

PHILIP GEORGE

These mechanisms explain immediately why usually only 50% of the total combined Oz appears as 0 2 gas. The fact is often expressed by saying 50% rather than the theoretical 75% is liberated. On this new interpretation the theoretical amount would be SO%,, and the actual amount obtained might be less if 0 2 - reacted with a substrate or more if the HzOZ formed were catalytically decomposed. An advantage of these mechanisms is that the stoichiometry is explained without the assumption of two XH2 groups. Other reactions of hemoglobin also permit a free radical interpretation, notably the coupled oxidation with ascorbic acid by molecular 0 2 which yields choleglobin, but further discussion requires a full kinetic analysis. Even though the denaturation reactions described above have not been examined kinetically it is worth emphasizing that their chief features can be explained by the formation of 0 2 - as in the mechanism advanced for the autoxidation. The liberation of an activated lo2molecule is no longer required-02- is the active oxygen.

VII. SUMMARY In comparing the reactions of hemoglobin, myoglobin, peroxidase, and catalase with molecular oxygen or hydrogen peroxide in relation t o the similar reactions of ionic iron the following conclusions can be drawn: 1. Many experimental studies now suggest a free radical mechanism in the case of ionic iron, but the essential nature of this evidence is indirect. Most important are the kinetic and thermochemical data. 2. Consideration of the possibility of similar free radical reactions occurring with the hemoproteins according to this thermochemical data suggests that the ferrous compounds should react as rapidly or more rapidly than free ferrous ions. The ferric compounds reacting with HOz radical or 0 2 - should still be rapid reactions but in the case of hydrogen peroxide or OzH- the reaction should proceed extremely slowly. 3. Kinetic studies of the autoxidation of hemoglobin and myoglobin do not favor intramolecular mechanisms involving reaction between separated groups on the protein molecule, but the form of the rate equations resembles that found when two valency states of a metal ion compete for the same radical intermediate. A free radical mechanism based on competition for the HOz radical can be developed to account for the observed results. 4. It is the reaction of the ferric forms of these hemoproteins with hydrogen peroxide that provides the sharpest contra,& to ionic iron. Complexes are formed in both systems. Although ferric ion only gives an ion pair complex Fe02H2+,whereas peroxidase can give three com-

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

425

plexes, depending on experimental conditions, the first of which probably has this ion pair structure, the contrast is in the type of reaction undergone by the complex. The ferric ion complex plays no independent kinetic role, the dominant reaction being electron transfer resulting in the partial oxidation of the peroxide molecule: in the hemoprotein peroxide complexes the peroxide is utilized in reactions involving the reduction of peroxide t o water. So little is known about the structure of these complexes t h a t a detailed discussion of reaction mechanisms must await further experimental evidence. I n the reaction of ionic iron there is some indication that a “ferryl” ion, Fe02+ may take part under certain experimental conditions. A compound with this structure would adequately explain part of the known behavior of the secondary complexes with peroxidase and catalase and the single complex formed with methemogIobin or metmyoglobin. I n the light of these conclusions and the data on which they are based, the effect of coordination within the porphyrin ring and t o the protein molecule on the catalytic action of the iron atom is potentially fourfold. 1. This coordination makes the complexes water soluble over a large pH range. This effect is important in itself for if ionic iron were not precipitated as the hydroxides, a t p H 10.46 i t would have a catalytic activity comparable t o that of catalase in the straightforward decomposition of hydrogen peroxide. 2. The ionization potential of the ferrous compounds in aqueous solution is lowered by the coordination, which has the effect of making any endothermic reaction of free ferrous ion less endothermic and hence potentially faster in the case of the hemoproteins. 3. Conversely the reaction of the ferric compounds will be potentially less exothermic, or endothermic and hence slower. I n this respect the coordination of the iron atom appears t o make possible a new reaction path in the case of hydrogen peroxide and alkyl hydroperoxides involving at first complex formation with the peroxide molecule and then the production of further complexes of unknown structure but which function as electron or hydrogen acceptors. 4. Another possible result of the coordination is t o allow electron accepting or donating groups on the protein molecule t o enter into competition reactions which could lead t o the predominance of certain reactions resulting in a selective or specific overall reaction.

REFERENCES 1. Wyman, J., Advances in Protein Chem. 4, 410 (1948). 2. Theorell, H., Adv. in Enzymol. 7, 265 (1947). 3. Respiratory Enzymes, University of Wisconsin Press, Madison, 1942.

426

PHILIP GEORGE

4. Haemoglobin, Barcroft Memorial Conference. Butterworths, London, 1949. 5. Lemberg, R., and Legge, J. W., Haematin Compounds .and Bile Pigments. Interscience, New York, 1949. 6. Theorell, H., and Akeson, A., Arkiv Kemi, Mineral. Geol. A16, No. 8, (1942). 7. Bonnichsen, R. K., Arch. Biochem. 12,83 (1947). 8. Roughton, F. J. W., Harvey Lectures 39,96 (1944). 9. Shack, J., and Clark, W. M., J . Biol. Chem. 171, 143 (1947). 10. Keilin, J., Biochem. J. 46,448 (1949). 11. Haurowita, F., Enzymologia 4, 139 (1937). 12. Haurowitz, F., Brdicka, R., and Kraus, F.,Enzymologia 2, 9 (1937). 13. Clark, W. M., Cold Spring Harbor Symposia Quant. Biol. 7, 18 (1939). 14. Lemberg, R., Cortis-Jones, B., and Norrie, M., Nature 140,6!i (1937);Biochem. J . 32, 171 (1938). 15. Theorell, H., and Paul, K. G., Arkiv Kemi, Mineral. Geol. MA, No. 12 (1944). 16. Keilin, D., and Hartree, E. F., PTOC. Roy. SOC.(London) l21B, 173 (1936). 17. Keilin, D., and Hartree, E. F., Biochem. J. 39, 148 (1948). 18. Foulkes, E. C., and Lemberg, R., Enzymologia 13, 302 (194'3). 19. Agner, K., and Theorell, H., Arch. Biochem. 10,321 (1946). 20. Theorell, H., Arkiv Kemi, Mineral. Geol. 16A, No. 3 (1942). 21. Hartree, E. F., Ann. Repts. on Progress Chem. (Chem. SOC.London) 43,287 (1946). 22. Millikan, G., Proc. Roy. SOC.(London) 166A,277 (1936). 23. Eley, D. 1) , Trans. Faradoy SOC.39, 172 (1943). 24. Haurowitz, F., Z. physiol. Chem. 264,266 (1938). 25. Theorell, H., Arkiv Kemi, Mineral. Geol. 14B,No. 20 (1940:l. 26. Theorell, H., Ergeb. Enzymforsch. 9, 239 (1943). 27. Lemberg, R., Legge, J. W., and Lockwood, W. H., Biochem. J. 36, 328 (1941). 28. Brooks, J., Proc. Roy. SOC.(London) 109B, 35 (1931). 29. Brooks, J., Proc. Roy. Soe. (London) 118B, 560 (1935). 30. Coryell, C. D., Stitt, F., and Pauling, L., J. Am. Chem. SOC.69,633 (1937). 31. Conant, J. B., and Fieser, L. F., J . Biol. Chem. 62, 595 (19245). I Biol. . Chem. 63,479 (1'325). 32. Neill, J. M., and Hastings, A. B., . 33. Legge, J. W., J . Proc. Roy. SOC.N . S. Wales 76, 47 (1942). 34. Pauling, L., Proc. Nail. Acad. Sci. U.S. 21, 186 (1935). 35. Brooks, J., J . Physiol. 107,332 (1948). 36. George, P., Abstracts 1st Intern. Congr. Biochem. Cambridge, England, 385, 1949. 37. George, P., and Stratmann, C. J., Biochem. J. In press. 38. Holden, H. F., Austrahm J . Ezptl. Biol. Med. Sci. 14,291 (1936). 39. Lemberg, R., Legge, J. W., and Lockwood, W. H., Biochem. J. 36, 339 (1941). 40. Lemberg, R., Australian J . Exptl. Biol. Med. Sci. 20, 111 (1942). 41. Lemberg, R., Legge, J. W., andLockwood, W. H., Biochem. J . 36, 363 (1941). 42. Kobert, R., Arch. Gen. Physiol. 82,603 (1900). 43. Keilin, D., and Hartree, E. F., Proc. Roy. SOC.(London) 117B, 1 (1935). 44. Haurowitz, F., 2. physiol. Chem. 232, 125 (1935). 45. Keilin, D., and Hartree, E. F., Nature 166, 513 (1950). 46. Keilin, D., and Mann, T., Proc. Roy. SOC.(London) l22B, 119 (1937). 47. Theorell, H., Enzymologia 10,250 (1942). 48. Keilin, D., and Hartree, E. F., Biochem. J. 49,88 (1951). 49. Chance, B., Science 109,204 (1949). 50. Chance, B., Arch. Biochem. 21,416 (1949).

T H E SPECIFIC REACTIONS O F IR ON I N SOME HEMOPROTEINS

427

Theorell, H., and Akeson, A., Arkiv Kemi, Mineral. Geol. A16, No. 1 (1942). Abrams, R., Altschul, A. M., and Hogness, T. R., J . Biol. Chem. 142,303 (1942). Mann, P. J. G., Biochem. J. 26, 918 (1931). Chance, B., J. Biol. Chem. 161, 553 (1943). Chance, B., Arch. Biochem. 22, 224 (1949). Chance, B., J. Cellular Comp. Physiol. 22, 33 (1943). Yamasaki, E., Science Repts. Tohoku Imp. Univ., First Ser. 9, 13 (1920). Morgulis, S., J. Biol. Chem. 47,341 (1921). Northrop, J. H., J . Gen. Physiol. 7, 373 (1924-25). Williams, J., J. Gen. Physiol. 11, 309 (1927-28). Nosaka, K., J. Biochem. (Japan) 8, 275, 301 (1927). Maximovitsch, S. M., and Antonomovn, E. S., 2.physiol. Chem. 174,233 (1928). Zeile, K., and Hellstrom, H., 2.physiol. Chem. 192, 171 (1930). Bonnichsen, R. K., Chance, B., and Theorell, H., Acta Chem. Scand. 1, 685 (1947). 65. George, P., Nature 180, 41 (1947). 66. George, P., Biochem. J. 43,287 (1948). 67. George, P., Biochem. J. 44,197 (1949). 68. Chance, B., J. Biol. Chem. 179, 1299 (1949). 69. Chance, B., Acta Chem. Scand. 1, 236 (1947). 70. Chance, B., J . Biol. Chem. 179, 1311 (1949). 71. Chance, B., J. Biol. Chem. 179, 1341 (1949). 72. Chance, B., and Herbert D., Biochem. J . 46, 402 (1950). 73. Chance, B., J. Biol. Chem. 179, 1331 (1949). 74. Stern, K. G., J . Biol. Chem. 114,473 (1936). 75. Keilin, D., and Hartree, E. F., Proe. Roy. Soc. (London) 119B, 141 (1936). 76. Keilin, D., and Hartree, E. F., Biochem. J. 39, 293 (1945). 77. Chance, B., J. Biol. Chem. 180, 865 (1950). 78. Chance, B., J. Biol. Chem. 180,947 (1950). 79. Theorell, H., Ezperientia 4, 100 (1948). 80. Polanyi, M., Atomic Reactions. London, 1932. 81. Pease, R. N., Equilibrium and Kinetics of Gas Reactions. Princeton University Press, Princeton, 1942. 82. Steacie, E. W. R., Atomic and Free Radical Reactions. Reinhold Publishing Corporation, New York, 1946. 83. Barb, W. G., Baxendale, J. H., George, P., and Hargrave, K. H., Nature 163, 692 (1949); Trans. Faraday Soc. 47, 462 and 591 (1951). 84. Baxendale, J. H., Evans, M. G., and Park, G. S., Trans. Faraday SOC.42, 155 (1946). 85. Haber, F., and Weiss, J., Proc. Roy. SOC.(London) 147A,332 (1934). 86. George, P., Faraday SOC.Discussions No. 2, 196,218,219 (1947). 87. Evans, M. G., and Uri, N., Trans. Faraday SOC.46, 224 (1949). 88. Evans, M. G., George, P., and Uri, N., Trans. Faraday SOC.46, 230 (1949). 89. Evans, M. G., Baxendale, J. H., and Uri, N., Trans. Faraday SOC.46,236 (1949). 90. Bray, W. C., and Gorin, M. H., J . Am. Chem. SOC.64, 2124 (1932). 91. Bohnson, V. L., and Robertson, A. C., J . Am. Chem. SOC.46, 2512 (1923). 92. Kolthoff, I. M., and Medalia, A. I., J . Am. Chem. SOC.71, 3777 (1949). 93. Kolthoff, I. M., and Medalia, A. I. J . Am. Chem. SOC.71, 3784 (1949). 94. Walton, J. H., and Christiansen, C. J., J. Am. Chem. SOC.48,2083 (1926). 95. Weiss, J., Naturwissenschaften 23, 64 (1935). 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64.

428

PHILIP GEORGE

96. Lamb, A. B., and Elders, L. N., J . Am. Chem. SOC.63, 137 (1931). 97. George, P., unpublished experiments. 98. Eley, D. D., and Evans, M. G., Trans. Faraday SOC.34, 1093 (1938). 99. Gusman Barron, E. S., Cold Spring Harbor Symposis Quant. Biol.7 , 154 (1939). 100. Taylor, J. F., and Hastings, A. B., J . Biol. Chem. 132, 649 (1939). 101. Taylor, J. F., J . Biol. Chem. 144, 7 (1942). 102. Wyman, J., and Ingalls, E. N., J . Biol. Chem. 139,877 (1!341). 103. Taylor, J. F., and Morgan, V. J., J . Biol. Chem. 144, 15 (1942). 104. L a t h e r , W. M., Oxidation Potentials. Prentice-Hall, New York, 1938. 105. Butler, J. A. V., Electrocapillarity. Methuen, London, 1940. 106. Conant, J. B., and Pappenheimer, A. M., J . Biol. Chem. 98, 57 (1932). 107. Dainton, F. S., and Smith, P. In press. 108. Dainton, F. S., J . Phys. Chem. 62, 490 (1948). 109. Mann, P. J. G., and Saunders, B. C., Proc. Roy. SOC.(London) 119B, 47 (1935). 110. Saunders, B. C., and Mann, P. J. G., J. Chem. Soc. 769 (1940). 111. Chapman, N. H., and Saunders, B. C., J. Chem. Soc. 496 (1941). 112. Keilin, D., and Hartree, E. F.,Nature 164, 254 (1949). 113. Theorell, H., and Maehly, A. C., Acta Chem. Scand. 4, 42:! (1950). 114. Maehly, A. C., in Enzymes and Enzyme Systems. Harvard University Press, Cambridge, Massachusetts, 1951. 115. Theorell, H., Arkiv Kemi, Mineral. GeoE. M A , No. 14 (19412). 116. George, P., and Irvine, D. H., Nature 168, 164 (1951). 117. Theorell, H., and Swedin, B., Nature 143, 71 (1940). 118. Chance, B., J. Biol. Chem. 182,643 (1950). 119. Chance, B., J. Biol. Chern. 182, 649 (1950). 120. George, P., and Irvine, D. H., Biochem J . In press. 121. Cahill, A. E., and Taube, H., J . Am. Chem. SOC.73, 2847 (1951). 122. Theorell, H., Ehrenberg, A., Chance, B., Arch. Biochem. Biophys. I n press.

Author Index Nunihers in parentheses are reference numbers. They are included to sesist in locating references in which the authors' names are not mentioned in the text. Numbers in italics refer t o the page on which the reference is listed a t the end of each article. Ezample: Agner, K., 376(19), 486 indicates t h a t this author'e article is reference 19 on page 376 and is listed in the bibliography on page 426.

Bailey, K. C., 74, 85 Bancroft, W. D., 34, 85 Abel, E.r 33, 35y 36(7), 41, 42, 45, 56, Barb, W. G., 47(43), 48,49,53,54,55,56, 57(59), 59(59), 75, 83, 84, 86, 349, 57(43, 62), 58(66), 59(43), 60,61,64, 364 65, 66, 72, 84, 85, 345(6), 352(6), Abrams, R., 390, 427 364, 408(83), 409, 410(83), 412(83), Ackermann, P., 306, 309(51), 315(72), 423, 427 333, 334, 338, 339, 341 Barrer, R. M., 233, 234, 235, 241, 250, Adcock, W. A., 317, 334, 340, 341 252, 256, 257 Agner, K., 376(19), 426 Barrett, E. P., 99, 147 Aicher, A., 309(56), 339 Barron, G. E. S., 413, 417 Akeson, A., 370, 371, 389(51), 4"~ Bartlett, p. D., 167, 177, $08, 209 Alberts, L., 321(94), 340 Basford, P. R., 221(21), 222(21), 655 Albrecht, O., 169, 208 Bashkirov, A. N., 329, 341 Alden, R. C., 310(61), 339 Baudisch, O., 70, 85 Alder, M. G., 359, 365 Bauer, S. H., 147 Alexander, J., 8, 29 Bauermeister, H. O., 92, 97, 98, 135, 149 Allmand, A. J., 357, 565 Baughan, E. C., 159, 172, 207, 209, Almquist, J. A., 323, 840 362(56), 365, 365 Altschul, A. M., 390, 427 Baxendale, J. H., 31, 47(43), 48, 49, 51, Andersen, V. S., 57, 59, 60, 61, 85, 348, 52, 53, 54(43), 55(43), 56, 57(43, 621, 58(66), 59(43), 60(43), 61(43, 661, 349, 364 Anderson, J. S., 20, 30 63(73), 64, 65(66, 73), 66(43), 70, Anderson, R. B., 144, 147, 229, 256, 316, 72(62), 740061, 84, 86, 345(6), 317, 330, 334(123), 340, 341 347(10), 352(6), 361(10), 362, 364, Andrews, A., 317, 340 365, 393(84), 407, 408(83), 409, Antheaume, A., 297, 338 410(83), 412(83, 89), 413, 423, 427 Antonomava, E. S., 393, 467 Bean, E. H., 317, 360 Armbruster, M. H., 244, 257 Beebe, R. A.,139, 147, 225, 256 Arndt, F., 165, 208 Beckwith, J. B., 139, 147 Arnold, J. R., 233, 256 Beeck, O., 96, 147 Aston, J. G., 195, 210, 243, 246, 247, $57 ~ ~ 1F., 1 ,36(17), 84 Audibert, C., 297, 338 Bell, R. P., 151, 152, 153, 154, 157, 159, Austin, J. B., 244, 267 161, 164, 165, 169, 171, 172, 186, 187, 191, 194, 195, 198,202, 203, 206, B 207, 207, 208, 209, 210 Bahr, H. A,, 332(114), 334(120, 122), 541 Benedict, W. S., 316(75b), 340 Benson, H. E., 306(45), 314(71), 338, 33.9 Bahr, Th., 284, 334(20), 338, 341 429

A

430

AUTHOR INDEX

Bergmann, M., 172, 209 von Berneck, Muller, 352(20), 364 v. Bertalan, J., 46, 56, 57, 61, 84,348, 364 Bevan, E. J., 50(48), 84 Birkhimer, E. R.,12, 29 Black, C. A., 323, 340 Blau, F., 64, 86 Bloch, H. S., 88, 147 Bobtelsky, M., 72, 73, 78, 80(121), 83, 86, 86 Bobtelsky-Chaikin, L., 72,78(114), 86,86 Boedeker, E. R., 14, 30, 88, 148 Bogdanov, G. A., 82, 83, 86 Bohnson, V. L., 46,50,56(72), 57(63), 61, 84, 86, 352(19), 364, 409, 427 Bolinger, E. D., 180, 209 Bonhoeffer, K. F., 169, 187, 193,608,210, 355, 366 Bonnichsen, R. K., 394, 395, 405, 426, 427 Booth, V. H., 171, 209 Born, M., 225, 226, 266 Bosch, C., 327(99), 340 Brackin, C. W., 89, 148 Bray, W. C., 33, 35, 36, 37, 40, 41, 43, 44, 45, 51, 56(6), 58(69), 83, 84, 86, 348, 351, 360, 364, 366, 409, 427 Brdicka, R., 373, 436 Bredig, G., 352, 353(24), 357(33), 364 Brescia, F., 185, 209 Briggs, S. H. C., 68, 86 Brill, R., 334(125), 341 Brode, J., 36(12), 50, 83, 84 Brodie, B. C., 360, 366 Bronsted, J. N., 152, 154, 156, 157, 158, 162, 193, 198, 200, 205, 206,207, 210 Brooks, J., 382, 383, 384, 385, 421, 426 Brown, C. L., 3, 29 Brown, J. A., 130, 143, 147 Brown, J. F., 164, 808 Browning, L. C., 335, 941 Brunauer, S., 88, 90, 95(7), 97, 104, 107, 127, 129, 147, 148, 213, 219, 227, 233(45), 266, 666, 267, 260, 263, 669, 316, 326(96), 339, 340 Bruner, F. H., 312, 313, 339 Buff, F. P., 227, 266 Buffleb, H., 286, 289(25), 338 Burgstaller, A., 359, 360, 566 Burnett, R., 161, 808

Burt, W. F., 3, 29 Butler, J. A. V., 169, 185, 208, 209, 413, 428 Byck, H. T., 128, 1/t? C

Cahill, A. E., 419, 428 Caldin, E. F., 186, 209 Callow, A. E., 43(29), 84 Cambi, L., 171, 208 Cassel, H. M., 218, 230, 231, 232, 242, 248, 265, 266, 267 Cassie, A. B. D., 95, 147, 227, 244(47), 256 Caulkins, A. L., 43, 44, 45, 84 Chance, B., 379, 389, 390, 391, 392, 394, 395, 396, 397, 398, 399,400, 401, 402, 403, 404, 405, 406, 415, 416(73), 417, 420, 426, 427, 4S8 Chang, T. L., 74(103), 86 Chapman, N. B., 393, 407(111), 428 Christiansen, C. J., 58(68), 86, 410, 427 Clark, A., 317, 340 Clark, W. M., 373, 426 Clunk, J. C., 191, 620 Cochran, C. N., 138, 149 Cohan, L. H., 98, 147 Cohen, H., 187, 209 Cohn, E. M., 317, 333, 334, 340 Cohn, M., 172, 209 Conant, J. B., 166, 108, 382, 414, 426, 428 Conn, A. L., 89, 103, 148 Conn, M. E., 12, 29 Connick, R. E., 38, 40,41, 43, 84 Connolly, G. C., 12, 29 Cook, M. A,, 214, 229, 248, 266, 266 Coolidge, A. S., 242, 243, 267 Cordes, C., 4(8), 26(3), 29 Cortes-Jones, B., 374, 386, 426 Coryell, C. D., 382, ,426 Craig, A. M., 161, 208 Craxford, S. R., 309, 329, 332, 339, 340, 341 Criddle, D., 221(22), 266 Crittenden, E. D., 3:34(125), 341 Cross, C. F., 50(48), 84 Crowell, I. H., 306, :314(71), 538, 339 310(61), 339 Crutchfield, J. Curtis, H. A., 323(96), 3.40

m.,

431

AUTHOR INDEX

D Dainton, F. S., 393, 415, 428 Darwent, B. de B., 171, 191, 194, 209 Davidson, D., 71, 85 Davidson, R. C., 89, 112, 148 Davies, C. W., 154, 207 Davis, I. D., 310(60), 339 Dawitt, T. W., 333, 341 Dawsey, L. H., 354, 365 Dawson, H. M., 158, 165, 173, 189, 191, 107, 209, 210 Day, J. N. E., 173, 209 DeBoer, J. H., 257 DecarriBre, E., 297, 338 Dechamps, G., 327(99), 340 De La Mare, P. B. D., 195(139), 210 Deloney, J. E., 13, 30 Deming, L. S., 95(7), 104, 107, 127, 129, 1-47 Deming, W. E., 95(7), 104, 107, 127, 129, 147 Denton, S. W., 13, 30 Devonshire, A. F., 220, 241, 255 DeWitt, T. W., 13, 30, 98, 148, 335(130), 341 Deyrup, A. J., 157(10), 207 Dieckmann, W., 172, 209 Dienst, W., 286 Diets, V. R., 98, 148 Dingwall, A., 166, 186, 208, 209 Dole, M., 229, 256 Dowden, D. A., 352, 364 Drain, L. E., 216, 225, 238, 248, 155 Drake, L. C., 14, 18, 24(43), 28(43), 30, 100, 105, 143, 148, 149 Drenan, J. W., 217, 256 Duclaux, J., 56(60), 84 Duncan, R. D., 147 Dwyer, R. J., 361, 365

E Eckstrom, H. C., 317, 334, $40, 341 Egloff, G., 309(57), 339 Ehrenberg, A., 420, 428 Ehrhardt, C. H., 112, 149 Eidus, Ya. T., 318, 335, 3.40, 341 Elder, L. W., Jr., 411, 428 Eley, D. D., 379, 413, 486,428

Eliot, T. Q., 312(67), 339 Elkin, P. B., 99, 100, 149 Elvin, 0. C., 309(53), 339 Emmett, P. H., 13, 30,88, 90, 95, 97, 98, 167, 148, 219, 227, 233(45), 250, 255, 256, 257, 263, 316(73), 326(91), 333(115a), 335, 339, 340, 341 Engelhardt, F., 329, 341 Erickson, H., 94, 131 Erlenbach, E., 4(8), 26(8), 19 v.Euler, H., 72(94), 85, 178, 209 Evans, A. G., 22, 30 Evans, L. P., 3, 6, 29 Evans, M. G., 32(2), 40(24), 43(28), 48, 51, 52, 60, 70(45), 83, 84, 347, 361, 362(56, 57), 363, 364, 365, 393, 407, 408(84), 409, 412(87, 89), 413, 417, 427, 428 Everett, D. H., 240, 243, 245, 256, 257

F Fajans, K., 193, 110 Farr, J., 204, 210 Faulkner, I. J., 188, 210 Feisst, W., 279(14), 305(43), 337, 338 Fenton, H. J. H., 46, 50, 52, 54, 84, 343, 361, 407 Fergusson, R. R., 233, 234, 235, 256 Field, I. H., 306(45), 314(71), 338, 339 Fieldner, A. C., 310(61), 339 Fieser, L. F., 382, 426 Finck, B., 171, 208 Fischer, F., 272, 273, 274(5), 275, 276, 277, 278, 279(13), 280, 281(18), 284(19), 286, 291(26, 27), 296, 297, 298, 299, 302, 305(43), 310, 320, 323(95), 329, 331, 332, 337, 388, 340, 341 Fischer, H. G. M., 3, 29 Fleming, H. W., 317, 340 Flexser, L. A., 166, 186, 208, 209 Fortner, M., 353(24), 364 Foster, A. G., 98, 148 Fouinat, F., 74, 75, 85 Foulkes, E. C., 376, 395, 426 Fowler, R. H., 222, 223, 227, 229, 256, 260, 269 Francis, A. W., 329(102), 840 Franck, J., 343, 364

.

432

AUTHOR INDEX

Franke, W., 50, 84, 349(16), 364 Frankenburg, W. G., 90, 148, 262, 269, 316(73), 339 Frederikse, H. P. R., 242, 257 Frenkel, J., 236, 238, 239, 240, 241, 242, 252, 267 Fricke, H., 357(35), 358, 966 Friedel, R. A., 317, 330, 340, 341 Friedman, B. S., 147 Frost, A. V. J., 4, 29 Fujimura, K., 310 Funck, A., 80,81, 86 G

Galbreich, A., 78, 81(126), 86 Gayer, F. H., 4, 29 Gelles, E., 195, 202, 210 George, P., 32(2), 47(43), 48, 49, 53, 54(43), 55(43), 56, 57(43, 62), 58(66), 59(43), 60(43), 61(43, 66)) 63(73), 64, 65(66, 73), 66(43), 72(62), 83, 84, 85, 345(6), 352(6), 364, 367, 384, 385, 387, 394, 395, 407, 408(83, 88)) 409, 410(83), 411(97), 412(83, 88), 417, 418, 419, 420, 421, 423, 426, 467, 428 Gilfert, W., 337(135), $41 Gill, R., 36(17), 84 Glasner, A., 78(114), 79(117), 86 Gleysteen, L. F., 98, 148 Gluud, G., 327(99), 340 Goddin, C. S., 312(67), 339 Gohr, E. J., 3, 29 Golden, P. L., 310(60), 339 Goldschmidt, S., 50(49), 84 Goldsmith, H. L., 202, 210 Good, G. M., 21(48), 23, 26, 30, 88, 148 Gorin, M. H., 51, 56, 58, 84, 409, 427 Gorter, C. J., 242, 267 Green, €1. S., 225, 226, 256 Greensfelder, B. S., 21, 23, 26, 27, 30, 88, 148 Gregg, S. J., 221(21), 255 Grennall, A., 15, 30 Griffith, R. H., 279(15), 337 Griffith, R. O., 38, 43, 84 Gring, J. L., 94 Gross, P., 185, 209 Grosse, 15

Grote, H. W., 13(25), 30 Grummer, M., 317,, 340 Guggenheim, E. A., 154, 157, 207, 223, 243, 253, 256, 257 Gunness, R. C., 310(61), 939 Gurry, R. W., 154(5), 807 Gurwitsch, L., 4, 2:1 Gwinn, W. D., 216,256

H Haber, F., 33, 34, 47, 48, 50, 51, 52, 53, 54, 56(4), 58, 59, 64, 83, 83, 343, 344, 345, 346, 348(3b), 351, 352, 354(3b), 364, 408, 437 Hagg, G., 308, 334, 338, 341 Haensel, V., 137, l g 8 von Halban, H., 193, 210 Hall, C. C., 309(58;1, 316, 317, 339, 341 Hall, W. K., 144, 1.47 Halle, F., 308, 317, 334, 340 Halsey, G., 224, 223, 230, 236, 238, 239, 240, 241,242, 244,252, 266, 267, 259, 261, 262, 263, 2!65, 269 Hamill, W. H., 185, 187, 209, 210 Hammel, E. F., 27(52), 30, 88, 148 Hammett, L. P., 157, 161, 166, 168, 177, 186, 207, 208, b!09 Hansen, R. S., 242, 252, 257 Hansford, R. C., l(l), 16, 18, 19(42), 21(42, 47), 23, 24(43), 28(43), 29, 30, 88, 89, 117, I 4 8 Hantcsch, A., 170, 171, 193(126, 127, 128)) 195, 208, 210 Hargrave, K. R., 47(43), 48, 49, 53, 54(43), 55(43), 56, 57(43, 62), 58(66), 59(43), 60(43), 61(43, 66), 65(66), 66(43), 72(62), 84, 86, 345(6), 352(6), 364, 407, 408(83), 409, 410(83), 412(83), 423, 427 Harkins, W. D., 95, 99, 132, 139, 148, 221(21), 222(21), 227, 238, 250, 265, 256, 257 Harkness, R. W., 333(115a), 341 Hartner-Seberich, R., 280(16), 337 Hartree, E. F., 370, 371, 374, 375, 377, 380, 387, 395, 396, 397, 400, 401, 407, 417, 418(lti), 426, 427, 428 Hartridge, H., 378 Haskell, V. C., 161, 208

433

AUTHOR INDEX

Hastings, A. B., 383, 413, 426, 428 Haufe, W., 58, 63, 64, 65, 85, 348(15), 364 Haurowitz, F., 373, 374, 380, 387, 392, 426 Hawk, C. O., 310(60), 339 Heiberg, T., 50(48), 84 Heidt, L. J., 355, 366 Heinemann, H., 101, 118, 148 Heitler, W., 362 Hellstrom, H., 393, 394, 427 Herbert, D., 399, 427 Herbst, M., 317, 334, 340 Herbst, H., 4, 29 Herington, E. F. G., 325, 332(111), 340, 341 Higginson, W. C. E., 171, 172, 206, 207, 209 Hill, G. R., 359, 365 Hill, T. L., 95, 148, 211, 214, 216, 217, 218, 222, 223, 224, 227, 228(48), 229, 230, 232(48, 54), 233(65), 234(48), 235, 236, 239, 240(54), 241, 242, 246, 247, 248, 250(83), 251(83), 252, 253(83), 254, 655, 256, 257, 262, 265, 269, 269 Hindin, S. G., 17, 18, 19, 24, 26, 30, 101, 118, 148 Hinshelwood, C. N., 173, 191, 209, 210 Hirst, J. P. H., 172, 209 Hoekstra, J., 13(25), 30 Hofer, L. J. E., 144, 147, 316, 317, 333, 334(122, 123), $40, 341 Hofmann, K. A., 68, 86 Hogness, T. R., 390, 427 Holden, D., 36(17), 84 Holden, H. F., 386, 426 Holmes, J., 97, 148 Holt, R. B., 362, 566 Honig, J. M., 139, 147 Honig, R. E., 18, 24(43), 28(43), 30 Horn, O., 280(16), 337 Hornberg, C. V., 3, 29 Houdry, E., 3, 4, 5, 29 Howard, H. C., 92, I49 Hsii, S. K., 167, 190(119), 208, 210 Hiinecke, H., 193, 210 Hiittig, G. F., 233, 234, 235, 236, 256, 257 Hughes, E. D., 195, 210 Humphrey, C. W., 49, 84, 345(7, 8), 346, 349(8), 3/74

Hunter, E. C. E., 171, 208 Hush, N. S., 43(28), 84

I Iffland, D., 169(47), 208 Iimori, S., 68, 85 Ingalls, E. N., 413, 428 Ingold, C. K., 167, 168, 173, 190, 208, 209, 210 Innes, W. B., 242, 257 Irvine, D. H., 387, 418, 419, 420, 428 Ives, D. J. G., 193, 210

J Jack, H., 309, 339 Jacobsen, B., 334, 341 Jacques, A. G., 84 Jahr, K. F., 80, 81, 86 Jansson, B., 72, 85 Jessen, H., 334, 341 Jessen, V., 334(122), 334, 341 Johnson, F. B., 103, 148 Johnson, M. F. L., 90, 92, 95, 97, 98, 99, 100, 110, 131, 132, 135, 138, 141, 147, 148 Jones, P., 191, 210 Jones, W. A., 187, 209 Jope, E. M., 371, 380 Joyner, L. G., 99, 127, 147, 148, 250, 257 Junell, R., 169, 208 Jura, G., 95, 99, 132, 139, 148, 221, 222, 227, 238, 248, 250, 255, 655,256, 25? Juza, R., 309(51), 333, 334, 338, 341

K Kagan, Y. B., 329, 341 Kalb, M., 195, 210 Karzhavin, W. A., 327(99), S4O Kaufmann, B., 223, 224, 256 Kazarnovskii, I. A., 363, 365 Keilin, J., 370, 371, 373, 374, 375, 380, 387, 389, 392, 395, 396, 397, 400, 401, 407, 417, 418(16), 426, 427, 428 Keith, P. C., 310(61), 311, 339 Kemball, C., 216, 217, 248, 656, 257 Kendrew, J., 371

434

AUTHOR INDEX

Kilham, I. K., 52 Kilpatrick, J. E., 328 Kilpatrick, M., 193, 210 Kimball, R. H., 186, 209 King, C. V., 180, 209 Kington, G. L., 225, 243, 246, 247, 256, 257

Kirkwood, J. G., 222,225,226,227,266 Kirson, B., 73, 86 v. Kiss, A., 71, 72, 85 Kistiakowsky, W., 67, 68, 85 Kistler, S. S., 94, 148 Kita, G., 310 Klar, R., 187, 210 Kobayashi, K., 4, 29 Kobert, R., 387, 426 Kobosev, N . I., 77(112), 78, 79, 81, 86 Koch, H., 278, 280(16), 299, 329, 332, 337(135), 337, $40, 341

Kodama, S., 310 Kolbel, H., 306, 309, 315(72), 329, 333, 334, 338, 339, 341

Kohler, E. P., 165, 208 Kolthoff, I. M., 50(53), 56, 58(53), 59(53), 63(72), 64, 65(75), 84, 85, 112, 137, 149, 348, 349, 361, 364, 410, 427 Konovalova, B. A., 49, 57(51), 82, 84, 86 Kornblum, N., 169, 208 Kornfeld, G., 355, 364 Kortum, G., 171, 208 Kosterlits, H., 70, 86 Kraemer, E. O., 98, 148 Kramers, H . A., 223, 224, 266 Kraus, F., 373, 426 Kreger, W. E., 90, 97, 99, 100, 107, 109, 110, 112, 128, 136, 147, 149 Krieg, A., 316, 317, 330, 3.40, 341 Kryokov, Y . B., 329, 341 Kiister, H., 280, 281(18), 305(43), 337, 338 Kuhn, R., 58(65), 62, 63, 64, 65, 85, 169, 208

Kuhn, W., 357(33), 865 Kummer, J . T., 333, 335, 341

L Lachmann, A., 171, 208 Lal, B. B., 68, 69(84, 85), 70, 86

Lamb, A. B., 84, 411, 428 La Mer, V . K., 185, 187, 209, 210 Langheim, R., 309(51), 333(119), 338,341 Langmuir, I., 220, 222, 227, 229, 255, 256, 263

Lapworth, A., 165, 208 Latimer, W . M., 413, 428 Latta, J . E., 311, 339 Lamer, H. F., 74(104), 85 Lea, D. E., 355, 357, 365 Leamon, W. G., 4(8), 26(8), 29 LeClerc, G., 297, 338 Lederer, E., 71, 72, E'5 Lee, J. A., 310(61), 339 Lee, T . S., 63(72), 65(75), 85 Lefebvre, H., 297, 3% Legard, A. R., 173, 1?09 Legge, J. W., 369, 37.3, 375, 378, 380, 381, 383, 386, 387, 390, 392, 398, 405, 420, 421, 423, 426 Lehmann, G., 47, 58, 62, 84, 343, 344, 357(33), 364, 366 Lemberg, R., 369, 373, 374, 375, 376, 378, 380, 381, 383, 386, 387, 390, 392, 395, 398, 405, 4!10, 421, 423, 426 Lennard-Jones, J. E., 213, 220, 241, 255 Leuchs, H., 168, 208 Leum, L. N., 12, 29 Leussing, D. L., 63(i'2), 65(75), 85 Leva, M., 317, 340 Levesque, C. L., 161, 208 Lewis, G. N., 192, 21'0 Ley, H., 193, 210 Li, Y . Y., 223, 256 Lichtin, N . N . , 169(47), 208 Lidwell, 0. M., 165, 186, 202, 203, 208, 209, 210 Liebhafsky, H. A., 35, 36(13), 36(16), 37(14), 38, 40, 41, 44, 45, 83, 84 Lindberg, R. I., 92 Livingston, H. K., 219, 255, 257 Livingston, R. S., 35, 36, 37, 38(15), 83, 84 Lockwood, W. H., 380, 386, 387, 423, 426 Loebl, H., 361, 365 Loeser, E. H., 221(21), 222(21), 255 Lohmar, W., 286, 291(26), 338 London, F., 214, 256 Los, J. M., 216, 225(9), 238(9), 265

435

AUTHOR INDEX

Lowry, T. M., 157, 158, 169, 188, 190, 191, 207,208, 210 Lowson, W., 173, 209 M McBain, J. W., 98, 148 McCartney, J. T., 144, 147 McEwen, W. K., 198, 202, 210 McKenzie, A., 173, 209 McKeown, A., 38, 43, 84 McLane, C. K., 365, 365 McLaughlin, P., 395 MacMillan, W. G., 228, 236, 238, 239, 241, 256, 257, 264, 265, 269 McReynolds, H., 13, 30 Macuga, S. J., 12, 29 Maehly, A. C., 372, 428 Magnanini, G., 36, 83 Makower, B., 37, 38, 41, 55, 84 Manchot, W., 46, 47, 50, 58, 62, 84,343, 344, 345, 364 Mann, P. J. G., 390, 393, 407(109), 427, 488 Mann, T., 389, 392, 426 Marisic, M. M., 7, 29 Mark, H., 334(125), 341 Maron, S. H., 169, 203, 208, 210 Martin, F., 298(41), 338 Mason, L. S., 330, 341 Matheson, G. L., 13, 30 Matsumura, S., 310 Maximovitsch, S. M., 393, 427 Mayer, J. E.,222, 225, 226, 866 Medallia, A. I., 50(53), 56, 58(53), 59(53), 64, 84, 348, 349, 361, 364, 410, 427 Meehan, W F., 103, 148 Melik, J. S., 95, 99, 100, 110, 112, 132, 138, 141, 147, 1.69 MerkeI, H., 286, 288, 297, 309, 317, 333, 334, ,336(134), 338,S4l Merz, J. H., 361, 365 Meusel, A,, 286 Meyer, K., 277, 279(13), 284, 297, 337, 338 Michael, W., 305 Michaelis, L., 71, 85,402, 419 Mickeley, A., 172, 209 Miller, A. R., 223, 256, 257

Milligan, W. O., 109, 128, 148 Millikan, G., 378, 395, 426 Milliken, T. H., Jr., 15, 16, 17, 19(39), 20, 23, 30 Mills, G. A., 4, 5, 11, 14, 15, 16(39), 17, 18, 19(39), 23(39), 24, 26(44), 29,30, 88, 89, 101, 118, 148 MilIs, I. W., 8, 29 Mitchell, A. G., 173, 209 Mittasch, A., 274(6), 316(73), 337, 339 Moller, E.,202, 210 Moelwyn-Hughes, E. A., 187, 210 Mohammad, A., 36, 37(14), 83 Montgomery, C. W., 127, 148, 329, 340 Morgan, V. J., 413, 428 Morgulis, S., 393, 427 Morikawa, K., 316(75b), 340 Morris, J. C., 40(25), 84 Morrison, J. A., 216, 225, 238, 248, 255, 266

Mummery, C. S., 46, 47, 48, 50, 53, 84 Murata, Y., 310 Murphree, E. V., 3, 29, 310(61), SS9 Murphy, N. F. J., 34(5), 83 Myddleton, W. W., 309, 329(102), 339, 340

N Nash, A. W., 309(53), 339 Natta, G., 318, 340 Neill, J. M., 383, 426 Nelson, W. E., 169, 185, 208, 209 Neuss, J. D., 86 Newman, M. S., 172, 209 Nikolev, L. A., 73, 85 Nordmann, J. B., 169(48), 208 Norrie, M., 374, 386, 426 Northrop, J. H.,393, 427 Nosaka, K., 393, 427 Noyes, A. A., 36, 83 0

Oblad, A. G., 14, 15, 16(39), 17, 18, 19(39), 23(39), 24, 26(44), 30, 88, 148,214, 248, 255 Ogston, A. G., 195, 210 Olander, A., 178, 209 Oldenberg, O., 361, 362, 365 Olson, L. E., 147

436

AUTHOR INDEX

O’Neill, W. F., 316, 340 Ono,S., 226, 229, 240, 241, 250, 252, 666 Onsager, L., 223, 224, 256 Orr, W. J. C., 185, 209, 215, 255 Osterrieth, J. W., 327(99), 340 Ostwald, W., 152, 207 Oulton, T. D., 100, 119, i48 Owen, B. B., 154, 207 P Pace, B. S., 312(67), 339 Pack, D. H., 214, 248, 256 Pacsu, J., 187, 209 Pappenheimer, A. M., 414, 428 Park, G. S., 48, 51, 70(45), 84, 347(10), 361(10), 364, 393, 407, 408(84), 427 Parravano, G., 27(52), 30, 88, 148 Partington, J. R., 171, 208 Parts, A. G., 58(70), 85 Patart, Ch., 297(33), 338 Patton, J. T., 169(47, 48), 208 Paul, H., 298(41), 338 Paul, K. G., 375, 392, 426 Pauling, L., 377, 382, 383, 426 Paunce, S., 50(49), 84 Payne, J. W., 3, 29 Pearce, M. S. B., 72(95), 85 Pearson, T. G., 355, 366 Pease, R. N., 405, 427 Pedersen, K. J., 158, 162, 165, 169, 171, 174, 186, 198, 200, 203, 205, 207, 208, 209, 210 Peebles, W. C., 316, 317, 334(122), 840, 341 Penney, W. G., 19, 30 Pernoux, E., 297, 338 Perutz, M. F., 371, 379 Peters, K., 305(43), 306, 338 Peters, W. A., Jr., 3, 89 Peterson, S., 58(69), 85, 351, 364 Petin, N. N., 46, 56(41), 83, 84, 86 Petow, H., 70, 85 Pew, A. E., Jr., 3, 29 Pflaum, W., 50, 84 Phinney, J. A., 310(61), 339 Pichler, H., 271, 280, 281(18), 284, 285, 286, 287, 288, 289(25), 291(26, 27), 296, 297, 298, 302, 305(43), 308,311, 317,323(95), 333,334,336(134), S37, 338,339, 340

Pier, M., 274 Pierce, C., 95, 148, 244, 267 Pines, H., 24, 30 Pitzer, K. S., 216, :?55, 328 Plank, C. J., 20, 30 Polanyi, M., 22, 30, 236, 267, 356(56), 365, 405, 427

Porter, R. W., 7, 2:? Powis, F., 158(14), 165(14), 207 Preissler, R., 348(1!i), 364 Prettre, M., 297, 398 Prue, J. E., 154, 207 Purett, H. T., 310(61), 339

Q Qureshi, M., 70, 85

R Rachford, H. H., Jr., 109, 128, I48 Rakowski, A. V., 1‘77, 209 Raney, M., 279, 285, 337 Rao, A. R., 67, 68, 69, 71, 85 Reete, Th., 348(15), 364 Reiff, 0. M., 44, 84 Reinders, W., 353(:!4), 8/34 Reineau, A., 297, 5.78 Reite, O., 167, 169, 185, 188, 198, 203, 208, 209, 210

Rescorla, A. R., 8, 19(14), 29 Reyerson, L. I-I., 195, 210 Reynolds, D. A., 310(60), 339 Rhodin, Jr., T. N., 268, 269, 269 Rice, F. O., 44, 84, 354, 364 Richards, E. M., 181, 210 Richardson, R. W., 103, 148 Rideal, E. K., 248, 667, 263, 269, 309, 332, 339, 34i

Rieman, W., 86 Ries, H. E., Jr., 13, 30, 87, 90, 92, 95, 97, 98, 99, 100, 107, 109, 110, 112, 120(52), 128, 131, 132, 133, 135, 136, 138, 141, 1’48, 149 Riesenfeld, E. H., 74, 75, 86, 86 Risse, O., 357(34), 366 Ritchie, P. D., 173, 209 Ritter, H. L., 14, 50, 100, 143, 149 Rittner, E. S., 92, 149 Rius, A., 353(24), 364 Robbins, 1,. V., Jr., 103, I48 Roberts, G., 310(61), 339

437

AUTHOR INDEX

Robertson, A. C., 46, 50, 56(42), 57(63, 641, 61, 80(120), 84, 86, 86, 352(19), 364, 409, 427 Robins, A. B., 241, 250, 252, 267 Roelen, O., 279, 299, 305(43), 307, 337, 338 Roess, L. C., 99, 100, 149 Rolfe, A. C., 173, 209 Roller, P. S., 13, 30 Ronay, G., 128, 149 Ross, S., 233, 235, 266 Rossini, F. D., 328 Rothmund, V., 359, 360, 366 Roughton, F. J. W., 171, 209, 372, 378, 381, 426 Rowley, H. H., 242, 267 Ruckensteiner, K., 286 Rumer, G., 362 Rumpf, M., 79(116), 86 Rushbrooke, G. S., 223, 266 Russell, A. S., 138, 149 Rybicka, S. M., 191, 210 Ryland, L. B., 7, 29 S

Sabatier, P., 274, 296, 297, 326, 338 Sachsse, H., 327(99), 340 Sands, A. E., 279(15), 337 Saunders, B. C., 393, 407(109, l l l ) , 428 Schaefgen, J. R., 172, 209 Scheuermann, A., 308, 334(125), 338, 341

Scheumann, W. W., 8, 14(19), 29 Schmid, H., 346, 364 Schmidt, L. E., 279(15), 337 Schmitkons, G. E., 12, 29 Schneider, C. H., 233, 266 Schroeder, W. C., 310(61), 339 Schults, 316 Schuman, S. C., 335, 341 Schumb, W. C., 92, 149 Schutza, H., 324, 3.40 Schwarzenbach, G., 165, 208 Seaborg, G. T., 192, 210 Seitz, F., 267 Seligman, B., 316, 317, 340 Semenoff, N., 220, 266 Senderens, J. B., 274, 296, 538 Sennewald, K., 359, 366 Shack, J., 373, 426

Shankland, R. V., 12, 29, 103, I48 Shapiro, I., 112, 137, 149 Sherred, J. A., 169, 208 Shimp, H. G., 8, 29 Shriner, R. L., 169, 208 Shull, C. G., 99, 100, 149 Sidgwick, N. V., 79(115), 82(115), 86 Simon, A., 58, 63, 64,65, 86, 348, 364 Simpson, T. P., 3, 29 Singhal, C. P., 69, 86 Sinnatt, F. S., 309(52), 338 Sips, R., 224, 266, 261, 269 Smedley-Maclean, I., 72(95), 86 Smith, D. F., 310, 328, 339, S4O Smith, G. F., 157, 207 Smith, P., 393, 415, 428 Smith, R. C.,Jr., 92, 149 Smith, R. D., 202, 210 Smith, R. K., 27, 30, 88, 149 Smith, R. N., 95, 148, 244, 257 Smith, S. L., 335(132), 341 Sokolev, N. N., 81, 86 Sokolov, N., 362, 366 Spitalsky, E. I., 32, 36(1), 46, 49, 56(41), 57(51), 75, 76, 77, 79, 80, 81, 82, 83, 84,86

Spivey, E., 189, 191, 210 Srikantan, B. S., 67, 68, 69, 71, 86 Stauffer, C. H., 167, 208 Steacie, E. W. R., 405, 427 Stein, G., 361, 366 Steiner, H., 185, 209 Sterba, M. J., 137, 148 Stern, K. G., 399, 406, 427 Stitt, F., 382, 426 Storch, H. H., 273(3), 306, 310(61), 314, 316, 337, 338,339 Stratmann, C. J., 384, 385, 421, 426 Style, D. W. G., 357, 366 Suess, H., 185, 209 Sullivan, F. W., 312(67), 339 Swain, C. G., 189, 191, 210 Swedin, B., 393, 428 Sweeney, W. J., 3, 29 Szego, L., 171, 208

T Tamele, M. W., 7, 14, 15, 16, 18, 19, 90, SO, 128, 149 Tantram, A. D., 165, 208

438

AUTHOR INDEX

Taube, H., 360, 365, 419, 428 Taylor, G. T., 74, 86 Taylor, H. S., 15, 27, 30, 88, 148, 224, 266, 267, 261, 262, 263, 269, 316, 332(112), 340, 341 Taylor, J. F., 413, 414, 428 Taylor, T. W. J., 178, 209 Taylor, W. J., 328 Teeter, C. E., 154, 207 Teichner, S., 297, 338 Teletov, J., 352(20), 364 Teller, E., 95(7), 104, 107, 127, 129, 147, 227, 230, 233(45), 236, 238, 239, 241, 256, 257, 263, 264, 265, 269, 269, 316(75a), 339 Tentschert, H., 309(51), 333, 334, 338, 341 Ter-Nedden, W., 274, 275 Teter, J. W., 90, 98, 120(52), 133, 147, 149

Theimer, O., 235(67), 267 Theorell, H., 368,370, 371, 372, 375, 376, 377, 380, 389, 392, 393, 394, 395, 404, 405, 420, 426, 426, 4.27, 428 Thiele, F. C., 4(8), 26(8), 29 Thiele, J., 171, 208 Thomas, C. L., 7, 15, 16, 19, 21, 23, 29, 88, 147,149 Thompson, D., 165, 208 Titsenthaler, E., 296, 338 Tobiasson, G. T., 13(25), 30 Tolman, R. C., 234(66), 256 Tompkins, F. C., 74(105), 85, 225, 256 Trambouze, Y., 297, 338 Tramm, H., 298(41), 338 Trapnell, B. M. W., 263, 269 Trenner, N. R., 316(75b), 340 Tristram, G. R., 370 Troitzkii, K. V., 288, 338 Tropsch, H., 272, 273, 274, 275, 276, 280(16), 281(18), 297, 298, 320, 328,

331, 337, 340, 341 Trotman-Dickenson, A. F., 171, 208 Tsuneoka, S., 310 Turkevich, J., 15, 27, SO, 88, 149 Turnbull, D., 169, 203, 208, 210

U Urey, H. C., 172, 209, 354, 365 Uri, N., 32(2), 40(24), 43(28), 60, 80, 83,

83, 84, 86, 362 (57), 363,365,408(88), 409, 412(87, 89), 413, 416, 427

'V Van Nordstrand, R.. A., 90, 92, 97, 98, 107, 109, 112, 120(52), 128, 133, 135, 136, 147, 149 Veit, A., 170, 208 Verhoek, F. H., 172, 209 Vernon, J. L., 195(139), 210 Vinograd, J. R., 128, 149 Voge, H. H., 21(48), 23, 26, 30, 88, 148 Volman, D. H., 355,365 Volmer, M., 220, 255

W Wackher, R. C., 24, SO Wagman, D. D., 328 Waind, G. M., 154, 207 Waldo, P. G., 18, 24(43), 28(43), 30 Walker, E., 70, 74(106), 85 Walker, J. J., 309(56), 339 Walker, S. W., 311, 339 Walker, W. C., 137, 149, 228, 266 Walling, C., 15, SO, 161, 208 Walsh, A. D., 362, 365 Walters, W. D., 169, 187, 208 Walton, J. H., 36(ll), 58(68), 83, 85, 410, 427 Wannier, G. H., 223, 224, 256 Ward, A. F. H., 247, 257 Wassermann, A., 58(65), 62, 63, 64, 65, 85, 161, 208 Waters, W. A., 361, 365 Watson, C. W., 330, 341 Webb, G. M., 13, 30, 112, 149 Weigert, F., 67, 85 Weil-Malherbe, H., 15, SO Weinberger, E. B., 127, 148, 329, 340 Weinrotter, F., 286, 287 Weintraub, M., 317, 340 Weiss, J., 15, 30, 34, 42, 47, 48, 49, 50, 51, 52, 53, 54, !i6(4), 58, 59, 64, 83, 85, 84, 85, 343,, 344, 345(7,8), 346, 348(3b), 349(8), 351, 352, 354(3b), 355(29), 357(36,37), 358(38), 359(38), 361 (47,4811, 362(55), 363, 364, 365, 408, 411, 4.87

439

AUTHOR INDEX

Weller, S., 316, 317, 334(123), 335, 3.40, 341 Wenzel, W., 307, 338 Westgren, A., 334, 341 Westheimer, F. H., 79(119), 86, 187, 200, 209, 210 Wheeler, A., 13, SO, 99, 149, 226, 229, 240, 241, 250, 252, 256 Wheland, G. W., 166(33a), 204, 208, 210 White, L., Jr., 233, 256 Whitmore, F. C. J., 4, 29 Wiegel, B., 353, 364 Wieland, H., 50, 72(96), 84, 86, 349, 364 Wiley, J. T., 13, 30 Wiley, J. W., 95, 148 Wilhelms, O., 46, 50, 84 Wilhelmy, L. F., 151, 207 Wilkins, F. J., 244, 251 Wilks, G. C., 193, 210 Williams, J., 393, 427 Willstatter, R., 33, 34, 51, 52, 54, 56, 59, 83, 343, 36.4 Wilson, C. L., 167, 185, 190(118), 208, 209, 210 Winkler, K., 274

Winther, C., 67, 86 Wittwer, C., 165, 208 Woodward, L. A., 202, 210, 325(98), 340 Wright, J., 186, 209 Wutke, J., 168(39), 208 Wykoff, R. W. G., 334(125), 341 Wyman, J., 368, 370, 372, 381, 413, 426, 428 Wynne-Jones, W. F. K., 36, 84, 185, 187, 209, b10

Y Yamamota, K. J., 4, 29 Yamasaki, E., 393, 427 Yang, C. N., 223, 266 Ykeda, K., 352(20), 364 Yost, D. M., 74(104), 85 Young, J. H., 169, 208 2

Zeile, K., 393, 394, 42T Zettlemoyer, A. C., 137, 149, 228, 266 Ziesecke, K. H., 286, 296, 338 Zucker, L., 177, 186, 209

Subject Index A

Acid catalysis, 173-174, 176, 181, 198200, 206 Acetaldehyde, 171, 191, 206 Acid centers, 15-21, 27, 28 Acetamide, 178, 188 Acid hematin protein complex, 423 Acetic acid, anhydrous, dissociation in, 160 Acid strength and reaction rate, 193, 203-205 Acetic acid catalysis, 188 Acetoacetic ester, 198, 203 Acid strengths and dissociation constants, 161, 182 Acetone condensation, 187 Acetophenone, 198, 202 “Acid Theory,” 15 Acetylacetone, 202, 203 Activated alumina, 95 Aci form conversion, 169-170 Activated clay cataly:sts, 5-6 Acid, definition of, 157-158 Activation energies for halogens, 37, 43 Acid-base catalysis, 151-210 and reaction velocities, 197-198 in absence of a solvent, 161 Active radicals, lifetime of, 357 acid-base strength and catalytic power, Addition of hydroxyl compounds to a 16I- 164 carbonyl group, :171-172 and catalyst structure, 204-209 Adiabatic heat of adsorption, 246, 247, 258 classical theory of, 152-157 concentration and reaction velocity, Adsorbed atoms, 216-217 152-156 Adsorbed molecules, definition, 157, 158, 164 degree of freedom for, 217 detection of, 162-164 interaction between, 217-219 dissociation in, 153-155, 161, 182 one-component system of, 243-248 distinction between, 165-169, 185 and solid adsorbent,, 213-214 empirical laws of, 153-164 Adsorbent-adsorbate, attraction between, en01 formation in, 166-167 238 kinetic steps in, 174-192 Adsorption, see also under Multilayer mechanisms of, 164-192 adsorption and F’hysical adsorption molecular mechanism of, 164-192 apparatus for measuring gas, 90-92 molecular structure in, 151-210 of argon on rutile, 21.6,225,238,242,250 nature of, 164 cooperative, on a nonuniform surface, of nitro compounds and nitramide, 265, 267, 268 169-171 on crystalline plane, 215-218 in nonaqueous solvents, 159-161 dividing surface in, 252 prototropic change in, 188 effect on pore structure, 133-134 reaction rate, 151-156, 192-201 energy of, 212-216, 262-263 reaction types, 192 entropy of, 249, 250, 254, 260, 262 reversible addition in, 171-173 of gas, &8,90-92 salt effects, 153-157 heats of, 246-248, 254, 258 single proton transfer, 174178 of heptane on liquid mercury, 248 substrate structure, 201-204 horizontal interaction in, 229-230, 232 ternary mechanism, 188-192 of hydrogen on tungsten, 262, 263 two proton transfer in, 178-182 of independent molecules, 212-218 440

SUBJECT INDEX

isotherm data, 95-99 iaotherm for B large, 236-239 isotherms, 90-99, 104-134 liquid-like, 264-265 localized, 216, 218, 222-225, 228-229, 263-264

measurement of, 90-92 from a mixture of gases, 232-233 mobile, 216, 218-222 of a molecule on a nonporous, nonpolar adsorbent, 236-239 of a monatomic gas on a nonpolar solid, 212-215

monolayer, 212-225, 265 multilayer, 225-242, 263-268 of nitrogen on metallic surface, 96-97 noncooperative, 268 physical, 212-259, 263-268 on a porous adsorbent, 239-240 potential energy effects in, 236-238 and solution thermodynamics, 248,250 surface area in, 219,220,244,246,248252, 258

surface heterogeneity in, 217-218,224225, 228-230, 269-268

thermodynamic properties of monolayer, 222-223 thermodynamics of, 220-225, 242-255 for three forms of silica, 133-135 uniform surface in, 222-224, 228 vertical interaction in, 230 Adsorption-desorption, of aerogels, 128-129 of alumina, 137-138 of clay catalysts, 119-123 of diatomaceous earth, 140-141, 146 of Linde catalyst, 132-133 of Nalco catalyst, 104 of silica-alumina catalyst, 104, 105-120 of silica-magnesia catalyst, 103-105 of titanium dioxide, 139-140 of xerogel, 127-128 Adsorption-demrption hysteresis experimenfs, 9&91 Adsorptiondesorption isotherm technique, 89-90 Adsorption-desorption isotherms, 89-141, 146

Aerocat catalyst, 93-94, 103, 114-119 microspheres, 114-118

441

Aerogel “S”, 94, 128-130, 135, 136 Aerogels, 99, 126-132, 135, 136 Air-elutriation method, 13 Alcohol, 307-308, 402 coupled oxidation of, 401 conversion of, 337 dissociation in, 159-160 synthesis of, 304, 307-308, 318, 336 Aldehyde, 307-308 conversion of, 337 synthesis of, 307-308 Aldol condensation, 168, 187 Alkali effect, on hydrocarbon synthesis products, 285

on iron catalyst, 275-276, 280, 285288, 305, 313, 320 on thorium precipitation catalyst, 294296 on zinc oxide and thorium oxide, 320321 Alkalized iron filings catalyst, 307 Alkyl hydroperoxide, 389-391, 404 Alumina boria catalyst, 5 Alumina catalyst, 4-5, 126, 137-138, 293, 294, 296 Aluminum-chromium oxides catalyst, 296 Aluminum oxide precipitation catalyst, 293 Aluminum oxide-thorium oxide catalyst, 294 Ammonium hydroxide effect on cupric ion catalysis, 73 Andersen’s empirical equation, 351 Aprotic solvents, 160, 162, 164, 182184 Aquo salts, 68-69, 70, 71 Area-temperature curves of catalysts, 102-103, 117, 122-124, 135-136 Argon-graphite sorption system8, 243246, 251 Ascorbic acid, 386, 424 Autoxidation, of dihydroxymaleic acid, 392-393 of hemoglobin, 381-388, 4-24 of myoglobin, 381-388, 420-424 Azide catalase, 375-376, 395-397, 418 methemoglobin, 380 Azomethines interconversion, 190-191

442

SUBJECT INDEX

B Base, definition of, 157, 158 Basic catalysis, 171, 173, 177, 179-180, 185-187, 200 Bauxite catalysts, 5 Bead aerogel, 130-131 Benzene hydroxylation, 361 Benzene-rubber sorption system, 243245,251 Benzene synthesis, 292 Benzoic acid hydroxylation, 361 Benzoylacetone, 202 Beta rule, 23 Bilirubin, 387 Biliverdin, 387,423 Boiling range of hydrocarbons, 2 Bond type changes, 379 Born charging entropy, 413 Bronsted acid, 15-16, 19, 20 Bronsted-Lowry definition, 158 Bronsted equation, 200, 201 Bronsted relation, and acid-base catalysis, 162-164, 196-198, 206-207 derivations from, 206-207 molecular basis of, 197-198 molecular interpretation of, 196-200 Brunauer-Emmett-Teller and Hutig equations, 233-236 Brunauer-Emmett-Teller isotherms, 231, 236-239 Brunauer-Emmett-Teller method, 95-97, 219 Brunauer-Emmett-Teller model, 227-230 Brunauer-Emmett-Teller theory, criticism of, 263-265 and modifications, 227-228, 230-232, 240-242 recent refinements in, 240-242 t-Butyl and isobutyl carbonium ions, 23 C Calorimetric heats of adsorption, 246247, 254-255 Calorimetric and isothermic measurements, 255 Capillary condensation and hysteresis, 239-240

2-Carbethoxycyclohe:~anone,202-203 2-Carbethoxycyclope:ntanoneI 202-203 Carbide theory, 276, 331-337 Carbided cobalt, 332, 333, 336 Carbided iron, 332, 333-334, 336 Carbided nickel, 334, 336 Carbided and uncarbided catalysts, 329 Carbon balances a t different carbon monoxide conversions, 331 Carbon dioxide formation, 328-331 Carbon dioxide reduction, 274-275 Carbon monoxide-heme reaction, 373 Carbon monoxide-hemoglobin reactions, 378-379 Carbon monoxide-hydrogen catalytic conversion, 272-4341 Carbonium ion mechttnism, 21-24 Carbonium ion theory, 27-28 Carbonyl group, 171-173, 187, 199-200 Carboxylate anion structure, 206-207 Carboxylate ion, 200 Carboxylic acids as catalysts, 119, 206 Carboxylic esters, 178-174 Catalase, 369-374, 393-404, 415-420 activity of, 393-399, 404 arginine content of: 370 complexes of, 398-400, 402404, 415417 compounds of, 374--384 in coupled oxidatio as, 400-402 decomposition of hydrogen peroxide, 367-368, 393-398 inhibition of, 395-399 reaction mechanism of, 402-404 reactions, 393-400, 415-420 structure, 369-372, 375, 377-378 Catalase-alkyl hydroperoxide complexes, 398-400 Catalase-iron protoporphyrin linkage, 371-372 “Catalatic” activity, 368 Catalysis of nitram ide decomposition, 200 Catalysis of nitroprusside, 70 Catalysis by undissociated acids and bases, 160 Catalyst activity test “ A , ” 8, 9, 14 Catalysts, general, see also Cracking catalysts acid centers in, 15-21, 27, 28

SUBJECT INDEX

acid properties of, 14-17 acidity of, 21, 22 acidity measurements, 14-21 activation of, 18-19 activity of, 8-9, 12, 13, 14, 18-19, 285, 321-327 and surface area, 88, 99, 100 and temperature, 278, 281 tests, 8-13 analysis of, 5 apparatus for testing activity of, 8-13 area of, and sintering, 112-117, 121123, 127-128 area temperature curves of, 102-103, 117, 120-124, 135-136 behavior of, 308-309,316,317,328-330 for branched hydrocarbon synthesis, 292-296 calcination of, 16, 18, 20 carbided and uncarbided, 329 chemical nature of, 14-21 chemical tests of, 13-14 commercial, 4-8 composition, 295-296 concentration and reaction velocity, 153-154, 158-159 cracking, 4-8, 15-24, 27, 87-149 degassing of, 91 deterioration of, 88-89, 323 experiments with iron, 287-288 for high pressure carbon monoxide hydrogenation, 289-296 for higher hydrocarbon synthesis, 320321 in homogeneous aqueous solution, 3186 for hydrocarbon synthesis, 281-284, 293-296, 308-309, 328-330, 333 hydrocarbon synthesis, effect on, 275276, 284, 295-296, 308 hydrogen acceptor-donor centers in, 27 for hydrogen peroxide decomposition, 31-86 for hydrogenation of carbon monoxide, 275-277, 282-284,289-296 for isosynthesis, 293-296, 336 life of, 282, 284 for methanol synthesis, 292, 318 molecular structure of, 204-207 nonporous, 139-140

443

for oxygenated compounds, 320-321 for paraffin synthesis, 289-292 particle size of, 13 physical properties of, 88, 95-99 physical tests of, 13-14 platelet structure of, 98, 99, 110-112 poisoning of, 88-89, 321-327 pore structure of, S9, 93, 95-138, 140146 preparation of, 5-7, 278, 291, 294 pretreatment of, 285-288, 336 properties and conditions for synthesis, 319-321 and reaction products, 320-321 reduction of, 275-277, 300-301 regeneration of, 313 sintering effects, 100-147 structure, 7-8,89, 93, 95-138, 140-146, 201-209 structure and sintering properties of, 88-147 and substrate equilibrium, 181 support for, 140, 336 suspended in oil slury, 305-306 thermodynamic studies, 297 titration tests, 14 for water g+s reaction, 329-330 Catalytic activity, 4, 8-14, 62-69, 352, 353 Catalytic conversion of carbon monoxide and hydrogen, 272-341, see Hydrocarbon synthesis Catalytic cracking, 1-30 carbonium ion in, 21-24, 27-28 catalyst evaluation in, 8-14 chemical concepts of, 1-30 chemistry of, 14-29 first commercial unit, 3 fluid process, 3, 5, 6 mechanisms, 21-28 protonic acidity, 28 TCC process, 3, 5, 6 theories, 27-28 vs. thermal cracking, 2, 3 Catalytic power and acid-base strength, 161-164, 182, 193, 199-200 Cation acids, 158 Cationic reactions, 21-24 Celite, 95, 126, 143-145, 269 Cerium oxide catalyst, 294

444

SUBJECT INDEX

Chain-breaking process, 355, 356, 359, 360 Chain mechanism, 72, 343-364 Chain reactions, 60-61, 64-66, 344-364 Chemical potential, 232, 234, 235, 241, 244, 248-249, 251-253, 258 Chemisorption, 259-263 Chromate catalysis, 75-80 Chromium ion complexes, 75-79 Clay catalysts, 14, 101, 121-122 activated, 5-6, 15, 99 Clay type catalysts, 85, 100, 119-124 Cobalt catalyst conversion to carbide, 332, 333, 336 Cobalt catalysts, 275, 277, 317, 320, 328336, 381-384 reaction products on, 280, 282 Cobalt-chromium oxide catalysts, 275 Cobalt-copper-manganese catalyst, 310 Cobalt-manganese-thorium-kieselguhr catalyst, 299 Cobalt precipitation catalysts, 278-279, 282 Cobalt standard catalyst, 278, 299-301 Cobalt-thorium-kieselguhr catalyst, 278279, 282 Cogelation of silica-alumina, 6-7 Combinatorial entropy, 264 “ Compensating reaction” mechanism, 32-34, 67, 77, 79 Compensating reactions, 35, 39, 58 Complexing agent, 62-64 Conductivity of peroxide decompositionchromium mixture, 77-78 Conductivity of solution, 153-154 Consumption ratio, 47-56 and acidity, 47-49, 52 and ion concentration, 48-49 and kinetics of oxidation, 54-56 Continuum theory, 264-265 Conversion of keto to enol, 186 Conversion of para to ortho hydrogen, 332-333 Copper catalysts, 71-73 Copper effect on iron Catalysts, 286-288 &precipitation of silica-alumina, 6-7 Coupled oxidations, 400-402, 424 Covalent iron, 377-378 Cracking catalysts, 4-14, 87-149 acid theory of, 15

alumina, 137-139 area measurementn, 93-97 diatomaceous earth, 140-146 experimental procedures, 90-99 gas adsorption merisurements of, 90-92 pore structure, 89, 93, 95-138, 140-146 pore structure and surface area, 95-99 representative types, 99-119 silica types, 124-1:37 sintering curves for, 135-137 sintering treatments, 92 structure and sintering properties, 87146 titania, 139-140 in virgin, steam-treated, vacuum sintered and used states, 89-93, 101124 Cracking processes, 2!-3 Cresol as a solvent, 190-191 Cuprammonium-peroxide mixtures, 73 Cupric ion, 71-73, 409 Cupric ion catalysis, 71-73 Cyan-catalase complex, 397-399 Cyanide ion-heme complex, 373 Cyanide ion-hemin complex, 373 Cycloversion process, 5 Cytochrome-c peroxi.dase complex, 418419

D DA-5 silica magnesium catalyst, 93, 94, 100 Davison silica gel, 104, 134, 135 Davison silica xerogel, 125-126 Decomposition of :hydrogen peroxide, 31-86, 352-358, 393-398, see under Hydrogen peroxide Decomposition of nitramide, 171, 174, 206 Degree of freedom, 216-217 Dehydration of acetaJdehyde, 206 Dehydration of silica-alumina catalyst, 18 Density in molecules per cubic centimeter, 241, 258 Depolymerieation of dihydroxy-acetone, 172 Deuterium, use of in kinetic mechanisms, 184-185

445

SUBJECT INDEX

Deuterium exchange and ionization rates, 193 Deuterium exchange rate vs. temperature, 24-25 Deuterium exchange reactions, 18, 24-26, 165, 169, 184-185, 193 Deuterium reaction mechanism, 26 Deuterium transfer vs. proton transfer, 184 Devenshire theory, 220-222 Diacetone alcohol, 187 Diakel silica alumina catalyst, 94, 103, 118-120 Diatomaceous earth, 126, 140-146 Diatomic molecule, 217 Diesel oil, 283, 307, 315 Differential entropy, 242-250, 253, 258, see also under Entropy Differential heat of adsorption, 246, 258 Differential quantum yield, 354-356 Dihydroxymaleic acid, 392-393, 406 Dipyridyl, 58, 62, 63 Dissociation constant, of acetoacetic ester, 198 of acetophenone, 198 and acid strengths, 161, 182 of ferric catalase complexes, 396 of nitromethane, 198 Dissociation of strong electrolytes, 153

E Electrolytic conductivity and degree of dissociation, 152, 153 Electron affinities, 362-364 Electron micrographs of catalyst structure, 141-145 Electron transfer process, 344-345, 348, 352, 357 Electron transfer reactions, 34, 45, 344345 Endothermic protolysis, 195 Energy of adsorption, 212-216, 262, 263 Enol formation mechanism, 166-167 Enolization, 165-167, 185-186 Enols of 8-diketones, 206 Entropy, of adsorbate on surface, 242-250, 258 of adsorption, 249-250, 254, 260, 262 of activation, 60

change, 60, 249-250, 412-414 combinatorial, 264 communal, 220 differential, 242, 244-247, 250, 253 differential and integral, 245-246 Integral, 245-250, 252, 258 Environment effect and reaction velocity, 156-157 Equilibrium constants, and heats of reaction, 197-198 in hydrocarbon synthesis, 328 and reaction rate, 196-198 Equilibrium of ferrous and ferric salts, 349-351 Equilibrium gas pressure, 219, 249-252, 258, 261, 262, 263, 264 Esterification of carboxylic esters, 173174 Ethyl carbonate, 187-188 Exchange reactions, of deuterium, 24-26, 169, 184-185, 193 between esters and alcohols, 172-173 Exothermal reaction heat, 305-307, 315

F Fergusson-Barrer statistical derivation, 234-235 Ferric catalase complexes, 396 Ferric dipyridyl complex, 58 Ferric ion catalysis, 56-61 Ferric ion complexes, 63-65, 409-410, 424-425 Ferric ion-hydrogen peroxide system, 63-66, 410-411 Ferric ion hydrolysis constant, 57 Ferric ions, 56-65, 346-348 Ferricyanide catalyst, 67-71 Ferricyanide hydrolysis, 68 Ferricyanide-peroxide reaction, 67-69 Ferrocyanide catalysis, 67-71 Ferrocyanide-ferricyanide redox system, 67-71 Ferrocyanide-peroxide mixtures, 69-71 Ferrocyanide solution, 68-69 Ferrous-ferric complexes, 64, 65 Ferrous-ferric ion reactions, 64-65, 343351, 407-411 difference between, 348-350 Ferrous ion catalysis, 46-66

446

SUBJECT INDEX

Ferrous ion complexes, 58, 62-65 Ferrous ions-peroxide system, 46-66, 410, 417 Ferry1 ion, 51, 56, 58, 409, 417, 425 Ferrylmyglobin, 418 Ferrylperoxidase, 417 Filtrol SR clay catalyst, 93, 94, 123 Fischer-Tropsch, early petroleum synthesis, 272-273, 275-276 medium pressure synthesis, 281-289 normal pressure synthesis, 274-275 pre-normal pressure synthesis publications, 274 reaction, 274-275, 277-296 Flow diagram of CAT “ A ” apparatus, 8-9 Fluid Filtrol clay catalyst, 93, 94, 119124 Fluid Filtrol clay isotherm plots, 119124 Fluid process, 3, 5, 6 Fluidized catalysts, 310-311, 313, 318 Free radical concept, 51-52 Free radical mechanism, of ferrous and ferric ions, 407-411 with hemoproteins, 407-420 in hydrogen peroxide reactions, 42, 343-364, 393 in ionic iron reactions, 368, 407-411, 424 Frenkel-Halsey-Hill theory, 236-240, 242 Freundlich isotherm, 259-262, 265

G Gasoline, aviation, 3 boiling range, 2 distillation apparatus, 11 manufacturing processes, 2-3 synthesis, 272-341 General and specific catalysis, 157-159, 176, 179-181 Gibbs equation, 230-231 Gibbs integral, 231, 238 Gibbs method, 251-254 Glossary of svmbols. 258 Glycol aldehyde, 172

Haber-Weiss mechanism, 348, 408 Haber-Willstiitter chain reactions, 34 Halide catalysis of hydrogen peroxide, 35-43 Halide catalysts, 35-46 Halide-halogen concentrations, 35-40 Halides and halogenci, 38-43 Halogenation of acetone, 165 Halogenation, of ketones, 165-167, 185186, 201-205 Halogenation of nitroparaffins, 169 Heat of adsorption and compression, 246 Heat capacity of argon, 216 Heats of adsorption, 246-248, 254, 258 Heats of hydration, :362-364 Heats of immersion, 255 Helium for dead space measurements, 91 Helium-mercury displacement method, 92, 97 Helium-mercury pore volume, 92-94 Hematin, 369, 372, 3,73, 416 Hematin compounds, 372-373 Heme compounds, 3‘72-373, 412-415 Heme group reactions, 373, 421-422 Heme-linked groups, 373, 392, 405, 412415 Hemochromogens, 373-375, 386, 412-414 Hemochromogens-hy drogen peroxide reaction, 386 Hemochromogens-parahematin system, 373 Hemoglobin, amino acid content of, 370 autoxidation of, 381-388, 420-424 compounds of, 374-384 mechanism of reaction, 382-384, 386387, 404-425 oxidation of, 382-386 oxidation reactionti of, 420-424 reaction with acids, 386, 422-423 with biliverdin, 423 with carbon monoxide, 378-381 with hydrogen peroxide, 422 with oxygen, 367-369, 374, 382-386 with oxyhemoglubin, 374 with pvridine. 4 2 2 4 2 3 rates ii,379,382, 384, 385

SUBJECT INDEX

structure of, 369-372, 379-382 x-ray studies of, 371 Hemoglobin-iron protoporphyric linkage, 371-372, 382 Hemoglobin-methemoglobin system, 381 Hemoproteins, 367-428, see also under Catalase, Hemoglobin, Myoglobin and Peroxidase absorption spectra of, 378 amino acid content of, 370 autoxidation of, 381-388 catalase reactions, 393-404 chemistry of, 369-381 compounds of, 374-381, 409, 410 cystine or cystein content of, 370 denaturation of, 380 differences in properties and reactions of, 381 free radical mechanism of, 411-420 heme groups in, 369-374, 379-380 histidine content of, 370-371 lysine content in, 370-371 oxidative attack on, 380 peroxidase reactions, 388-393 reaction with hydrogen peroxide, 367368, 393-398, 404, 410-417, 422425 reaction with iron, 367-428 reaction mechanism of, 404-425 stability of, 374, 381-382 structural differences of, 369-370, 404406 structural differences and reactions of, 404-405 velocity constants in, 391-392 Hemoproteins-peroxide complexes, 416417, 425 Hemoproteins-prosthetic group linkage, 371-372 n-Heptane on ferric oxide curves, 222 Heterogeneity of surface, 267-269 Hexagonal iron carbide, 334 High area gels, 99 Horseradish peroxidase complexes, 418420 Hot gas recycle process, 305-306 Houdry pelletted catalyst 5-46, 94, 114 Houdry porous beads, 94, 112-114 Houdry process, 3-6 Houdry silica alumina catalyst, 112-114

447

Houdry Type I catalyst, 93, 94, 103 Hutig equation, 233-236 Hydration of acetaldehyde, 171, 191 Hydration of carbonyl compounds, 171 Hydrocarbon synthesis, 272-341 by-products in, 307-308 carbide theory, 331-337 carbon dioxide formation in, 328-331 catalyst properties of, 319-321 catalysts for, 281-290, 293-296, 303313, 317-318, 328-330, 333, 336 composition of synthetic products, 280 developments in France, 296-298 developments in Germany, 298-309 developments in Great Britain, 309310 developments in Japan, 310 developments in other countries, 318319 developments in United States, 310318 equilibrium constants, 328 ethylene in, 310 with fluidized catalysts, 310-313 German operating plants, 298 high pressure, 289-296 hot gas recycle process, 305-306 hydrogen-carbon monoxide ratio in, 301-302 liquid reaction products, 280-281 medium pressure, 281-289, 301, 305307 modification of medium pressure, 305309 oil recycle in, 306-307, 314 oil slurry in, 305-306 olefin content in, 301-302 pilot plant experiments, 279-280 procedures used in German plants, 298-299 reaction products and catalyst properties, 319-321 removal of reaction heat, 279 solved and unsolved problems, 319-337 typical conversions, 327-330 water formation in, 328, 331 water- and oil-soluble products, 313314 yields of products, 280-287, 294, 299300

448

SUBJECT INDEX

Hydrocol process, 311-313 Hydrogen-charcoal sorption system, 243245 Hydrogen-deuterium exchange, rate of, 165, 184-185, 193 Hydrogen ion concentration, in buffer solution, 155-156 and protein transfer, 176 and reaction velocity, 177 Hydrogen ion in solution, 152, 155-158 Hydrogen-palladium sorption system, 243, 245, 251 Hydrogen para-ortho conversion, 332333 Hydrogen peroxide, reaction rate, 346-347, 352-353, 359360 as reducing agent, 34 reduction of, 51-52 reduction of halogens by, 37 x-ray decomposition of, 357-358 Hydrogen peroxide-catalase complexes, 398-404, 415-417 Hydrogen peroxide-catalase reaction, 367-368, 393-398 Hydrogen peroxide complex, 398-402, 415-41 6 Hydrogen peroxide-cupric ion reaction, 351-352 Hydrogen peroxide decomposition, 3186, 352-354 active intermediates, 50-53 of aquopentacyanoferrite, 69 by catalase, 393-399 catalysts for, 31-86 chain terminating reactions in, 60-61, 64, 66, 72 chemical kinetics of, 35 chromate catalysts for, 75-80 chromium ion complexes, 75-79 compensating reactions in, 35, 39, 58, 77, 79 concentration of peroxide in, 394-395, 397-399 concentrations and consumption ratio, 4 7 4 9 , 79 copper catalysts for, 71-73 cupric ion mechanism in, 54-55, 72 cupric ions in, 71-73 decomposition by cyanide, 403

dependence of acidity of catalysts, 74-77, 80-82 ferrocyanide and ferricyanide catalysis, 67-71 free radical concept,, 51-52 halide catalysts in, 3 5 4 6 by hemoproteins, 393-398, 401, 406, 408-411, 418 in homogenous aqueous solution, 3186 hydroxyl radicals in, 58-60 inhibition of, 396-397, 401 Iodate catalyst for, 43-46 by ionizing radiations, 357-358 intermediate oxidation reactions in, 52 intermediate polymerization in, 52 intermediates in, 3:1-34, 58, 65, 81, 82 iron catalysts for, 46-67 iron salts in presence of complexing agents, 62-67 on metal surfaces, 352-354 molybdate catalysts for, 80-82 noble metal catalysts, 353 permanganate catalysts for, 73-75 by peroxidase, 367-368, 383-384, 388392, 404, 410-411, 415417 photochemical, 34, 354-357 quantum yield, 354-356 rate of, 36, 56-58, Ci7, 76-78, 81, 82 reaction mechanisms, 39-43 theories of, 32-34 tungstate catalyst, 82-83 Hydrogen peroxide-ferric ion reaction, 56-61, 347-351, 368, 407, 424-425 Hydrogen peroxide-ferrous ion reactions, 47-50, 343-347, 368, 407, 424-425 Hydrogen peroxide-heme group reaction, 421-422 Hydrogen peroxide-hemochromogens reaction, 386 Hydrogen peroxide-ionic iron system, 408-410, 421 Hydrogen peroxide-methemoglobin reaction, 387 Hydrogen peroxide-metmyoglobin reaction, 417418 Hydrogen peroxide-ozone reaction, 359361 Hydrogen peroxide-parahematin reaction, 373, 374

SUBJECT INDEX

Hydrogen peroxide-permanganate system, 73-75 Hydrogen peroxide-peroxidase reaction, 383-384, 388-392 Hydrogen transfer mechanism, 23, 26, 27-28 Hydrogen-tungsten sorption system, 243-245 Hydrogenation of carbon monoxide to hydrocarbons, 272-341 Hydrogenation mechanism in cracking, 27 Hydrolysis of acetamide, 178, 188 Hydrolysis of carboxylic esters, 173-174 Hydrolysis of ethyl ortho-carbonate, 187188 Hydrolysis of halogens, 39-40, 45 Hydrous oxides of aluminum, 4-5 Hydroxyl ion catalysis, 156, 187-188, 195 Hydroxylamine catalase, 396 Hypohalous acid, 39-40 Hypohalite ion reaction, 39, 41 Hysteresis isotherms, 95-146 Hysteresis loops in adsorption or desorption, 240 and pore structure, 99-100

I I. G. Farbenindustrie hydrocarbon synthesis experiments, 305-306 Illumination and ferricyanide-ferrocyanide catalysis, 67-69 Incompletely dissociated salts, 154 Inert adsorbent, 251-253, 254 Inhibitor, azide ions of catalase activity, 395-397 cyanide ion of catalase activity, 395399 cyanide ions of peroxidase activity, 392-393 ethyl hydroperoxide of peroxide decomposition, 401 hydroxylamine of catalase activity, 395-396 sodium aaide of peroxide decomposition, 396-397 Integral entropy, 245, 246-250, 252, 258 Interaction energy of molecules, 218, 220 Intermediate product theory, 32-33

449

Intermediate reduction of ferric ions, 348-349 Intermediates, of carbonium ion, 21-22 in chromate catalysis, 79 in hemoglobin-oxygen reaction, 381, 383 in hydrogen peroxide-iron system, 58, 65, 66, 74, 81, 82 in manganese dioxide catalysis, 74 Intermolecular mechanism, 386, 421 Intermolecular potential energy, 213214 Intramolecular mechanism, 381, 383 Iodate catalysis, 43-46 Iodate reduction, 44, 45 Iodination of acetone, 165, 189-191 Ion displacement reaction, 40-41 Ionic activity coefficients, 154 Ionic-iron compounds, 378 Ionic iron reactions, 368, 404-411, 421, 424 Ionic strength, 154-156 Ionization of aliphatic amines in alcohol, 195 Ionization of p-diketone, 206 Ionization potentials of hemoprotein complexes, 412-413 Ionization radiations on hydrogen peroxide, 357-358 Iridium catalyst, 292 Iron-alumina-lime-potassium carbonate fused catalyst, 303-305 Iron carbide hexagonal catalyst, 308-309, 317, 318 Iron catalysts, 46-67, 275-277, 284-289, 308-310, 317, 320, 328-330, 331 behavior of, 304, 305, 328, 330 conversion to carbide, 333-334, 336 experiments with, 287, 302-305 for hydrogen peroxide decomposition, 46-67 precipitation type, 284-288, 303-305 pretreatment of, 288, 305 raw and reduced, 315-316 reaction products on, 285-287 sintered, 305 of synthetic ammonia type, 288, 314316 Iron complexes, 62-66, 409-410, 424-425

450

SUBJECT INDEX

Iron-copper catalyst, 309 Iron-copper-alkali catalyst, 303-305 Iron-copper-kieselguhr catalyst, 304-305 Iron-copper-zinc-alkali catalyst, 303-305 Iron decomposition catalysts, 288 Iron fused ammonia synthesis type catalyst, 288 Iron kieselguhr catalysts, 284, 304-305 Iron magnesium catalyst, 309 Iron oxide quintquevalent, 344 Iron-percarbide catalyst, 309-310 Iron peroxide complex, 415-416 Iron protoporphyrins, 369, 372-374 Iron salts catalysis, 46-67 Iron salts in presence of complexing agents, 62-67 Iron turnings catalyst, 320 Iron-zinc oxide catalyst, 275 Isomorphous substitution in silica lattice, 16 Isosteric heats of adsorption, 244-247, 258, 261, 262, 268 Isosynthesis, 292-296, 321 Isotherm plots, of Aerocat microspheres, 115-116 of alumina, 137-138 of DA-5 catalyst, 105 of Diakel silica alumina, 118-120 of diatomaceous earth, 140-141, 146 of Fluid Filtrol clay catalysts, 119-124 of Houdry silica alumina, 113-1 14 interpretation of, 95-99 of Linde nonporous silica, 132-133 multilayer adsorption, 266-268 of Nalco catalyst, 104 of silica aerogels, 128-130 of silica aluminum TCC, 105-113 of silica magnesia, 103-105 of silica xerogel, 127-128 of titania, 139-140 Isotherm types I1 to V, 129, 132, 240, 241 Isothermal heat of adsorption, 246-247, 258 Isothermal process, 247 Isotope exchange, 167, 185-186, 193

K “Katalasesto~s,’~ 62, 64, 66 Kelvin equation, 98-99

Keto-eno tautomerization, 165-167, 178, 181 Ketone, acid-catalyzed prototropy of, 194 conversion of, 166-168 enolization of, 165-169 halogenation of, 167-168 racemization of, 167-168, 170 reversible aldol condensation of, 168169 Kieselguhr, efficiency of, 297-298 Kieselguhr, origin of, 299-300 Kinetic analyses, examples of, 185-188 Kinetic mechanism, in non-aqueous solvents, 182-184 involving deuterium, 184-185 Kinetic studies, of chromium ion catalysis, 75-79 of cupric ion catal:ysis, 72 of hemoglobin oxidations, 382-386 of hemoproteins, 379,382-386,390-391 of hydrogen peroxide decomposition, 35, 393-395 of molybdate catalysis, 80-82 of myoglobin oxidrrtion, 384-386 of permanganate, ‘74 of peroxidase oxidntions, 390-391 of tungstate catalyst, 82 Kinetics, of aldol condensation, 187 of carbiding, 335 dependence on acidity, 37-41, 47-49 of ferric ion decomposition, 56-62 of ferrous ion oxidntion, 54-56 of halogen reactions, 37-43 of iodate-iodine renction, 44-45 of proton transfer, 175-180 of reaction, 35-43

L Labilization of hydroxyl hydrogen, 19 Lewis acid, 15-19, 28, 29 Linde nonporous silica, 94, 126, 132-136 Liquid “communal ” entropy, 220 Liquid density method, 219 Liquid hydrocarbona, composition of, 295-296 Liquid state, 222, 225-227, 233 Localized adsorption, 263-264

SUBJECT INDEX

with interactions, 222-225 without interactions, 228-229 London Law, 266 Low temperature distillation of C& fractions, 292-293 Low temperature gas adsorption technique, 88

M Magnetic measurements, 377 Manganese dioxide catalysis, 74-75 McMillan-Teller continuum theory, 264265 Mean consumption ratios in hydrogen peroxide reaction ions, 300, 344-347 Mechanisms, acid-base catalysis, 164-192 for autoxidation of ferrous salts, 411 of carbonium ion catalysis, 21-24 of catalase-peroxide reaction, 402-404 of catalysis in nonaqueous solvents, 182-184 of catalytic cracking, 21-27 cationic, 21-24 chain, 72, 343-364 of chromium ion catalysis, 78 compensating reactions, 32-34, 67, 77,

79 of deuterium reaction, 26 of enol formation, 166-167 of ferric and ferrous ion catalysis, 5361, 407-411 free radical, 343-364, 368, 393, 407420, 424 of hemoprotein reactions, 382-384, 386-387,404-425 of hydrocarbon synthesis, 276 of hydrogen transfer, 28 of hydrogenation, 27 intermolecular, 386, 421 intramolecular, 381, 383 kinetic, 182-185 metal surface decomposition, 352-354 multilayer adsorption, 266-268 of mutarotation, 169-170 of nitramide decomposition, 171 of ozone decomposition, 358-360 photochemical reaction, 354 racemisation, 167-168, 170

451

radical, 42-43, 408 Taylor’s, 27 ternary, 188-192 Medium pressure synthesis with iron catalysts, 303-305 Mesityl oxide oxalic ester, 191 Mesomeric structure of ions, 200, 203 Metalloporphyrin complexes, 413 Methanol synthesis, 292, 296-297, 318 Methemoglobin, 375, 377-380, 382, 383, 386, 404, 421-422 Methemoglobin-hydrogen peroxide reaction, 387, 417-418 Methyl hydroperoxide-catalase complex, 398-402, 420 Methyl isocyanide-heme reaction, 373 Metmyoglobin, 384, 385 Metmyoglobin complex, 387-388, 417420 Michaelis-Menten enzyme substrate complexes, 419 Michaelis theory, extended, 402 Mobile adsorption with interactions, 218222 Molar integral entropy, 246 Molecular structure and acid-base catalysis, 151-210 Molybdate catalysis, 80-83 Molybdate catalyst, 80-83 Moving bed processes, 3-6 Multilayer adsorption, 225-242, 263-268 BET theory and modifications, 227233 capillary condensation and hysteresis, 239-240 Hutig equation, 233-236 isotherm for 0 large, 236-239 mechanism of, 266-268 from a mixture of gases, 232-233 surface area of, 267-268 surface tension, 230-232 Mutarotation of camphor, 169-170 Mutarotation of esters, 173 Mutarotation of glucose, 158, 172, 187, 190, 191 Myoglobin, 369, 381-388 amino acid content of, 370 autoxidation of, 384-388, 420-424 lysine content of, 370-371 oxidation of, 367-368, 382-386

452

SUBJECT INDEX

oxidative reactions, 420-424 reaction rates of, 379, 385 structure of, 369-372 x-ray studies of, 371 Myoglobin-carbon monoxide reaction, 384 Myoglobin-iron protoporphyrin linkage, 371-372

N Nalco microspheres, 93-94 Nalco silica-magnesium catalyst, 93-94, 100, 103, 104 Naphtha reforming process, 5 Nearest neighbor interactions, 222-223, 232, 236 Nickel catalysts, 275, 277-278, 297, 320 Nickel-cobalt silicon alloy, 279 Nickel-manganese-aluminum kieselguhr catalyst, 277-278 Nickel precipitation catalysts, 277-279 Nickel-thorium-kieselguhr catalyst, 277, 279 Nickel-thorium oxide catalyst, 277 Nitramide, 169-171, 206 Nitrobenzene hydroxylation, 361 Nitrogen adsorption isotherms, 95-146, see under Isotherm plots Nitrogen data on anatase, 231, 238 Nitrogenous bases-heme reactions, 373 Nitromethane, 192, 198, 203, 204 Nitroparaffins, 169, 206 Nitrourethane, 205-206 Nonaqueous solvents, 182-184 Nonporous catalysts, 94, 126, 132-136, 139-140 Normal pressure synthesis, 274-275, 277, 279, 284 Number of molecules adsorbed, 217, 234, 258 Number of moles of adsorbed gas, 245, 246, 247, 249, 258 using Gibbs method, 252-253, 258 Number of moles of adsorbent, 244, 251, 258 0

Octahedral coordination complexes, 378 Olefin increase in hydrocarbon synthesis, 301, 305

Olefins, proton affinities of, 22-23 Oxidation of alcohol, 402 Oxidation of ascorbic acid, 386, 424 Oxidation of ferrous compounds, 46-49, 420-423 Oxidation of halides, 36-43 Oxidation of hemoproteins, 382-386, 420-424 Oxidation of iodine, 42-45 Oxidation states in chromate reactions, 79-80 Oxygen-ferric ion reaction, 345 Oxygen-hemoglobin reaction, 378-379, 420-424 Oxyhemoglobin, 374, 380, 421-422 Oxyhemoglobin-ascorbic acid complex, 387, 423 Oxyporphyrin hemoc:hrome, 374 Oxysynthesis of alcohols, 307-308 Oxysynthesis of aldehydes, 307-308 Ozone decomposition,, 358-361 Ozone-hydrogen peroxide reaction, 355, 358-361

P Palladium catalyst, 292 Paraffin, 280-283, 285, 287, 289-290 Paraffin synthesis, high pressure, 289-292 Parahematins, 373 Paraperoxidase, 380 Particle size, 13-14 “Perhydrol,” 50 Perhydroxyl ion, 58 Periodic potential barriers, 215-216 Permanganate catalysis, 73-75 Peroxidase, activity of, 392-393 arginine content of, 370 complexes of, 388-393, 419-420, 424425 with cyanide ion, 420 with ethyl and inethyl hydroperoxide, 389-393 compounds of, 374--378, 380, 384 decomposition of hydrogen peroxide, 367-368, 383-384, 388-392, 404, 410-411, 415417 inhibition of action of, 390-393 oxidation reaction of, 390-391

SUBJECT INDEX

reactions of, 388-393, 415-420 structure of, 369-372, 380, 419-420 Peroxidase-dihydroxymaleic acid reaction, 392-393, 406 Peroxidase-hydrogen peroxide reaction, 367-368, 404, 410-411, 415-417 Peroxidase-iron protoporphyrin linkage, 369-372 Peroxidic intermediates, 33-34 Petroleum synthesis, early history, 272276 Phenanthroline, 62-64 Phenols as catalysts, 206 Photochemical bromide-bromine catalysis, 43 Photochemical decomposition of hydrogen peroxide, 354-357 Photochemical reaction mechanism, 354 Phthalocyanines, 369 Physical adsorption, 212-250, 263-268 cooperative, 265 liquid like, 264-265 localized, 263-264 multilayer, 265-268 nonuniform surface, 265 surface heterogeneity, 265-268 Planar rotation, degree of, 217 Platelet structure, 98, 99, 110-112 Platinum catalyst, 292 Poisoning, by carbon dioxide, 323-324 by carbon monoxide, 393, 406 of catalyst and area measurement, 8889 of cobalt catalyst, 321-325 by cyanide, 386 of hydrocarbon synthesis, 275, 321-327 of iron catalysts, 323-327 of nickel catalysts, 323-327 of platinum or palladium, 353 of ruthenium catalyst, 325, 327 by steam, 323-327 by sulfur, 323-327 thorium oxide resistance to, 327 Polycrystalline copper, 268-269 Polymerization, 71, 361, 393, 415, 416 Pore radius measurements, 89, 93, 98-99 Pore size, 13-14, 99-100 Pore structure, 89, 95-100 adsorption effect on, 133-134

453

of Aerocat, 115-118 of Aerogel S, 126, 128-132 of clay catalyst, 100, 120-123 of Diakel, 11&126 of diatomaceous earth, 141-145 of gels, 99 of Linde nonporous silica, 126 of magnesia catalyst, 103-104 of Santocel “ C ” silica, 94, 126 of silica alumina catalysts, 105-120 of silica bead aerogel, 126 of silica catalysts, 94, 126-132, 135 of silica xerogel, 127-128 of titania, 126 Pore volume, 13-14 of cracking catalysts, 13-14 helium-mercury, 92-94 measurement of, 93, 97-98 Porphyrin ring, 371-372, 416, 425 Porphyrins, 387 Potential energy curves, 197, 204-205 Potential energy of interaction of molecules, 213-214, 237, 241, 258 Potential intermolecular force, 213, 258 Precipitation catalysts, 276-279, 282, 293-296, 303-305 Preequilibrium, 179, 182, 184 Pressure-area curves of n-heptane, 222 Pressure effect on synthesis with thorium oxide catalyst, 293-294 Pressure-temperature diagram for synthetic processes, 319 “Primiirstoss,” 349 Primary salt effect, 156-158 Principle of microscopic reversibility, 233-234 Product distribution of Synol synthesis, 307-308 Products of hydrocarbon synthesis, 279, 280, 282-287, 289-294, 315 on cobalt catalyst, 282-284 on iron catalyst, 285-287 on ruthenium catalyst, 291 Promoters, alkali on iron catalysts, 305, 313, 320 alkali of carbide formation, 333 of aluminum chloride catalyst, 19, 29 aluminum oxide of thorium oxide catalyst, 294 of boron fluoride catalyst, 19, 29

454

SUBJECT INDEX

cerium ion, 80 chloro-copper, 80, 82-83 citrate copper, 80, 82-83 of chromate catalysis, 80 cobalt ion, 80, 82-83 copper, 80, 286, 287, 288, 333 cupric ion, 82-83, 352 of cupric ion catalysis, 72, 82-83 cupric salts, 57, 61 ferric ion, 82-83 of ferric ion catalysis, 57, 61 of hydrocarbon synthesis, 24, 275, 276, 336 of hydrogen halides, 18-19, 29 of hydrogen peroxide decomposition, 46-50, 61, 352 iodine ion, 82-83 of iron catalyst, 276, 286, 287, 288 iron salts, 46-50 manganese ions, 82-83 manganous sulfate, 72 of molybdate catalysis, 82-83 molybdate of tungstate catalysis, 83 nickel ion, 80, 82-83 of nickel catalysts, 77 olefins, 24 of silica aluminum catalyst, 18-21, 29 and supports, 88 of tungstate catalysis, 82-83 tungstate of molybdate catalysis, 83 water, 18-21, 29 Protolytic reaction, 164, 170 Proton affinities, 22, 23 Proton transfer, in aprotic solvents, 182-183 to a base, 205 and potential energy curves, 204-205 and rate determination, 152, 182-183, 185 and reaction velocity, 176-179 single, 174-178 between substrate and catalyst, 174177 two, 178-182 Prototropic isomeriration, 178-179 Prototropy, 180-181, 194 Pseudo acid, definition, 195 and normal acids, 204-207

and pseudo bases, 192-195 and ‘‘true” acids, 193 Pyridine as a solvent, 190, 191

R. Racemization, 185-1136, 190 and ionization rates, 193 mechanism, 167-168, 170 Radical mechanism, 42-43, 408 Radicals, free OH, detection of, 361 thermodynamics of free, 361-364 Radioactive C14 in carburizing of catalysts, 333 Raney type nickel-cobalt catalyst, 279 Raney type skeleton catalysts, 279 Rate constants and activity energies of halogens, 37, 40 Rate constants and ionic strength of halogens, 36 Rate of decomposition, hydrogen peroxide-ferric ion reaction, 349-351 Rate dependence and iodine concentration, 45 Rate determining step in catalysis, 152, 182-183, 185, 194-195 Rate equation, for autoxidation of hemoproteins, 421 hydrogen peroxide-ferrous ion reaction, 346-348 hydrogen peroxide photochemical decomposition, 354-356 Rate of ferrous ion oxidation, 49, 54-55 Rate of formation of ferrous complexes, 63-66 Rate of molecules adsorbed, 229, 231233, 241, 247, 258, 260-266, 268 Rate of oxygen evolution in hydrogen peroxide-catalase reaction, 394 Reaction mechanisms, 39-42 Reactions involving two proton transfers, 178-182 single proton transfer, 174-178 Reactions at surface of metal catalysts, 352-354 Reaction velocity, of acid-base catalysis, 152-156, 19220 1 and activation energy, 197-198

SUBJECT INDEX

and catalyst concentration, 153-154, 156, 158-159 and conductivity of solution, 153, 154 and environmental effect, 156-157 and hydrogen ion concentration, 177 transition state of, 156 Recycle experiments, 274-281, 301-302 Recycle operations, 301-302, 305-307 Recycling oil in hydrocarbon synthesis, 314-3 15 Reduction of halogens, 37, 42 Reduction of iodate, 44, 45 Reduction of iron, 397 Reversible dimerization of aldehydes and ketones, 172 Ruthenium catalyst, 289-292, 320 S

Salt effects, 36, 38, 153-157, 178 Santocel “ C ” silica, 94, 126 Schiff base isomerism, 181 Screen analyses for particle size, 13 Secondary salt effects, 154-158 Sedimentation methods, 13 Silica, forms of, 125 structure of various forms of, 124137 Silica aerogel, 94, 115, 128, 133-134 Silica Aerogel S, 126, 128-132 Silica Aerogel Santocel “C,” 94, 126, 131, 135 Silica alumina catalysts, 14, 18, 24, 27, 88, 99, 100-101, 105-120 acid centers in, 15-21 active centers of, 20-21 alumina content in, 16-17, 20 calcined, 18, 20 preparation of, 6-7 promotion by water, 18-21 stability of, 6, 7 structure of, 107-119 synthetic, 6-7, 15-20, 100 temperature effect on surface area, 102-103, 117 Silica-alumina gel, structure of, 20 Silica-alumina hydrogel, 6-7, 16, 20 Silica-alumina-zirconia catalysts, 5 Silica bead aerogel, 126

455

Silica gel, Davison, 94, 100, 104 Silica lattice, isomorphous substitution in, 16 Silica-magnesia catalyst, 5, 14, 88, 99-104 Silica-magnesia DA-5 catalyst, 94, 100, 103 Silica-magnesia catalysts, sintering of, 101 structure of, 20, 103-104 Silica xerogel, 94, 125-128, 133-134 sintering of, 127-128 structure of, 127 Silica-zirconia catalysts, 5, 14, 19 Sintering, and area of catalyst, 92, 112-117, 121123, 127-128 and catalytic activity, 300 and deterioration of catalyst, 100-101 effect of clay catalyst, 121-124 effect of heat on, 101-102 effect on pore structure, 100-101, 127128, 136 effect on pore .structure of silicas, 101103, 104, 127-128, 136 effect on pore volume, 100-101 effect in steam, 89, 92, 101-122 effect on surface area, 93, 100-101 effects in vacuum, 93, 101-102, 106112, 116, 123-124, 127 factors affecting, 137 geometry of, 110-112 and grinding of catalyst, 112 nonuniform, 112 properties of cracking catalyst, 87-149 studies on silicas, 105-119, 135-137 and surface area, 100 treatments of cracking catalysts, 92 Sintering curves, 92, 101-102, 117-118 for Filtrol clay catalysts, 122-124 for silicas, 135-137 techniques, 89 Solution thermodynamics, 243-246, 254 Specific catalysis, 159, 163, 181 Specific hydrogen ion catalysis, 176 Stationary concentration, 38-39, 60, 63, 65 Stationary ferric-ferrous level, 349-351 Stationary state of radicals, 349-351, 353-354 Statistical BET theory, 231-234

456

SUBJECT INDEX

Statistical deduction of the B E T equation, 227-228 Statistical derivation of B E T model, 228 Statistical mechanical partition function, 234, 246, 258 Steady state, 37-39, 43 Steam deactivation, 89 Steam deterioration of catalysts, 89, 117, 121-122 Steam poisoning of catalysts, 323-324 Steam sintering, effect on clay catalysts, 121-122 effect on silica alumina catalyst, 105118 effect on structure, 105-118 Steam stripping and catalyst hydration, 89 Steam treated and vacuum sintered catalysts, 92-124, 135-138 Substrate equilibrium, 181 Substrate structure, 20 1-204 Surface area in adsorption, 219, 220, 244-252, 258 Surface area measurements, 88-89, 9597, 219, 227 Surface concentration, using Gibbs method, 253, 258 Surface heterogeneity in physical adsorption, 259-268 Surface or spreading pressure, 217, 220, 230-231, 247-251, 261 Surface or spreading pressure, Gibbs method, 252, 253, 258 Surface tension and B E T theory, 230-232 Synol process, 307-308 Synthesis of branched hydrocarbons, 292-296 Synthesis conditions and catalyst properties, 319-321 Synthesis of high melting paraffins, 289292 “Synthesis of Synthol,” 307, 320

T Tautomerism, 178, 181 Taylor’s mechanism, catalytic cracking, 27 TCC clay type catalysts, 93, 94, 100, 103, 105-109, 119

TCC process, 3, 5, 6 Temperature, dependence of, o n steady state function, 38 effect on hydrocarbon synthesis, 293294,296 effect on monolayer adsorption, 220, 223-224 effect on surface area, 102-103, 117120, 122-124, 135-136 Temperature-area curves, 102-103, 117, 122-124, 135-1.36 Temperature-area plot technique, 89-90 Termolecular hypothesis, 189-190 Ternary mechanism, 188-192 Theory of, acid-base catalysis, 152 liquid state, 222, ‘225-227, 233 multilayer adsorption, 225-233 peroxidase action, 90 physical adsorption, 212-258 Thermal cracking, 28 Thermal deactivation, 89 Thermal destructive distillation, 1-2 Thermal stability, 135-138 Thermochemical evidence for radical mechanism, 408 Thermochemistry, 412 Thermodynamic calculations, 335 Thermodynamic data on free radicals, 361-364 Thermodynamic studies on catalysts, 297 Thermodynamic system, 253-254 Thermodynamic treatment, 251-254 Thermodynamics, of adsorption, 220-225, 242-255 alternative treatments, 251-254 calorimetric heats of, 246-247 solution thermodynamics, 243-246 three dimensional, 248-249 Thermofor process, :3 Thermomagnetic studies, 333-334 Thorium oxide catalysts, 292-296, 327 Thorium precipitation catalysts, 294 Thorium-zinc-alumiiium oxides catalyst, 296 Titania, 95, 126, 139-140 Titanox-A MO catalyst, 95 Tri-(p-nitropheny1)-methane, 192-193

457

SUBJECT INDEX

Tripyridyl, 65 Tungstate catalysts, 80-83

U Uncharged acids as catalysts, 158 Undissociated acid molecules as catalyst, 173 Unitary theory of hydrocarbon catalysis, 27 Uricase and amino acid oxidase. 400401

v Vacuum sintered catalysts, 92-93, 101124, 127, 135-138 van der Waals equation, 218-220, 241 van der Waals forces, 212-213, 218, 268 Verdohemochrome, 374, 386

W Water formation in hydrocarbon synthesis, 328, 331

Water gas reaction, 327-331 Water-sodium chloride sorption system, 243-245 Water-sulfuric acid sorption system, 243245, 251 Wetted aerogel, 131

X Xanthine oxidase, 401 Xerogels, 94, 96, 125-128, 133-134 X-ray decomposition of hydrogen peroxide, 357-358 X-ray studies on carburized iron, 334-335 X-ray studies of hemoglobin molecule, 379 Z

Zinc-copper catalyst, 293 Zinc oxide catalyst, 296, 318, 320 Zinc oxide-thorium oxide catalyst, 296 Zirconium oxide catalyst, 294 Zirconium structure, 19-20

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