FIFTH REPORT OF THE COMMITTEE ON COYTACT CATALYSIS

F I F T H REPORT OF T H E COMMITTEE ON COYTACT CATALYSIS* BY E. EMMET RElU

Introduction There is among chemists a growing tendency to forget the distinction between catalyzed and non-catalyzed reactions. Catalyzed reactions used to be considered exceptional and therefore remarkable. As we find more and more of such reactions their exceptional nature disappears. Soon all reactions will be catalytic and then this term will cease to be descriptive. We are coming to face two more fundamental questions: I ) Why do chemicals react at all? and 2) Why do they not always react instantly when brought in contact with each other, provided the energy change would be in the right direction? We shall probably not get a clear answer to either of these untilwecometoamuch more complete understanding of atomic structure. The union of one atom with another is due to the transfer of one or more electrons or to the mutual holding of one or more pairs of electrons. Hence the study of reactions is reduced to finding out why and how the electrons readjust their relations with atomic nuclei. We are far enough along to see that this must be so but the investigation of atomic structure has not progressed far enough to answer the many questions that arise. The distinction between homogeneous and heterogeneous catalysis is gradually vanishing. The radiation theory of catalysis in its original restricted form seems to be fast losing ground. The intermediate compound theory appears to be gaining in acceptance but losing in definiteness. Perhaps it would be better to change “intermediate compound” to “intermediate complex” since the stoichiometric relations implied by ‘(compound” may exist but seldom. Combination and adsorption are probably not as different as we used to think them. X-ray photography has shown us that in a crystal of sodium chloride there are no molecules to be represented by Ka-C1 in which a univalent sodium atom is combined with a likewise univalent chlorine atom. We find that each sodium atom is directly joined to eight chlorine atoms and that each chlorine atom is similarly joined to eight sodiums. If the joining of atom to atom is thus indiscriminate within the crystal why should it be limited to strict valence relationships at the surface of a mass? The surface atoms can not be linked with the proper number of atoms and must satisfy their affinities by adsorbing vapors or gases. Whatever we may call the forces that bind the sodium atom to the eight surrounding chlorine atoms, they are nevertheless forces. These forces when not satisfied, as they cannot be at the surface of a crystal, must reach out for something that can be adsorbed. The * Report of the Committee on Contact Catalysis of the Division of Chemistry and Chemical Technology of the Kational Research Council. Written by E. Emmet Reid assisted by the other members of the committee: Messrs. H. Adkins, E. F. Armstrong, W. C. Bray, 0. U‘.Brown, R. F. Chambers, C. G. Fink, J. C. W. Frazer, H. S. Taylor and W.D. Bancroft, Chairman.

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adsorbed layer is held by the same forces that hold the atoms together within the mass. As a matter of fact we find adsorbed films on all surfaces. If the sodium and chloride atoms are joined promiscuously in the solid sodium chlorine, what happens when melting takes place? The relations can not be very different in the molten material, yet there is freedom of motion: the atoms must be able to shift their unions from one to another as they pass. This labile state of the unions may be the explanation of the greatly enhanced reactivity of substances in the liquid state. Adsorptive power becomes solvent power and all manner of substances are dissolved as they come in contact with the liquid. We may picture the solute as being adsorbed on the surface of the liquid in consequence of the unbalanced forces that exist there. On account of the mobility of the liquid thk surface layer would not remain as such but would mix with the rest of the liquid leaving a fresh surface to repeat this process. The adsorptive forces may be so strong that the adsorbed molecule is pulled apart in the process. This may be what happens when a molecule of hydrogen chloride is adsorbed by a water surface. If a polar molecule is thus disrupted, the fragments will carry charges and be ions. Solution and adsorption are both highly selective. We know that the reactivity of any substance is greatly enhanced when it is dissolved: it is probably likewise increased by adsorption. If the two processes are essentially the same, there is no reason for distinguishing between homogeneous and heterogeneous catalysis. The old conception according to which the catalyst contributes only its presence and takes no part in the reaction has been given up. Zelinsky’ remarks: “My observations on catalysis extending over several years have brought me to the same view of catalytic phenomena as was expressed by MendelejefP long ago in such a simple and original form, a view which later Raschig’ and recently it seems Bodenstein4 have adopted. In the contact processes with carbon compounds, the catalyst does not determine the reaction simply by its presence but by taking the ride of an active principle in the process; its surface energy produces far-reaching alterations in the substances which come into contact with it.” The Fourth Report of this Committee written by H. S. Taylor was devoted to the exposition of one point of view and did not include a general survey of catalytic literature. A different plan has been followed in the present report and an attempt has been made to cover the field more generally for 1925 as well as 1926. The references were obtained chiefly by searching Chemical Abstracts. In all, 240 articles have been noticed though there is no claim of completeness. Only abstracts of some of these were available. The most of the space has been devoted to articles on the theories of catalysis but large numbers of others have been catalogued with brief statements of results accomplished. ‘Brennstoff-Chem., 7, 207 (1926);Ber., 59, 15 (1926). * J. Ruea. Phys. Chem. SOC., 18,8 (1886). a Z. angew. Chem., 19, 1985 (1906). Ann., 440, 177 (1924).

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General Articles and Addresses Two important summaries and discussions have been published, one on homogeneous catalysis by Hinshelwoodl who, with his coworkers, has made many important contributions to this field and the other on auto-oxidation, anti-oxygenic and pro-oxygenic activity by Moureu and Dufraisse2 who have made many and remarkable discoveries on the acceleration and retardation of oxidation. These papers cover these fields much more thoroughly than can be done in the present report. Reference must be made to them for a fuller discussion. An important paper by Armstrong and Hilditch' gives a summary of our knowledge of catalysis at that date. The views are very similar to those expounded by Taylor in the Fourth Report. The same authors' have reviewed hydrogenation. Bone and Andrew5 give an excellent statement of results so far achieved in explaining surface catalysis. I n an address, Taylor6 gives a fine view of the field of catalysis. The most useful catalysts for each type of reaction are tabulated. Many important facts as to reactions at boundaries are brought together. In another address the same author' sketches the industrial developments due to catalytic invest igat io n. Two masters of the art describe their own work: Sabatiers tells in an intimate retrospect of his great achievements how he was led to his hydrogenation method and PatartQgives a detailed description of his methanol process. The industrial production of synthetic methanol is discussed by Lormand*O. Mittasch," who has spent a score of years in the catalysis research laboratory of the Badische Co. gives, in an address before the German Chemical Society, a fine account of the history and achievements of catalysis. RidealI2 reviews catalysis from the point of view of Rayleigh, Hardy and Langmuir, who have established that the seat of catalytic activity is limited to the film of reacting substances adsorbed on the surface of the catalyst. TweedyI3gives an excellent statement of the theories of catalysis. Guiselin" covers the same subject with particular emphasis on the importance of porosity Chem. Rev., 3, 227-256 (1926). Chem. Rev., 3, 113-162(1926). 3Chemistry and Industry, 44, 701 (1925). 4 Deuxieme Cons. Solvay, 1926, 492. 3 Proc. Roy. SOC., 109A, 459 (1925). 6 Canadian Chem. Met., 10, 35 (1926). 7 Ind. Eng. Chem., 18,958 (1926). 8 Ind. Ena. Chem , 18, IOOS 1926). 9 Bull. so;. enc. ind. nat., 137, 1417; Chem. Abs., 19,2026 (1925);Chimie et industrie, 13, I79 (192.5). 1OInd. Eng. Chem., 17,430 (1925). "Ber., 59, 13B (1926). 'ZDeuxiBme Cons. Solvay, 454 (1926). ' C h ~ m i s t r yand Industry, 45, 157-9,177-80(1926). 1aMem. Compt. rend. SOC. ing. riv. Fmnoe. 78, 2.5 (1925). 1

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and surface action. He suggests making catalysts with pores of known size and shape. Swientoslawski’ classifies catalytic reactions in two groups: I. Those that are allowed to reach a stable equilibrium. 2. Those in which the reactants are removed from the influence of the catalyst before a stable state is reached. Seymour2 and Pascal3 discuss catalysis in industry. Firth4 writes of catalysts and enzymes. Almquist5 describes catalysis at high pressures. We have sets of lecture demonstrations by Webster6 and by Kolthoff.’ Critical Increment There is a general agreement that the reason that reactions in homogeneous mixtures do not take place instantaneously and completely is that at any one time only a few of the molecules have the necessary amount of energy. Those that have this are said to be in the “excited” state and the amount of energy required to put a molecule in this state is called the “critical” increment.” The theories of catalysis differ from each other in the mechanism which is proposed by which this critical increment can be imparted to a molecule. I n a thorough mathematical paper G. N. Lewis and Smiths consider the possible sources of the energy required for activation. They conclude “that there is no valid argument against the general radiation hypothesis. On the other hand, the special radiation hypothesis, as announced in the first papers by Perrin and W. C. McC. Lewis, in which it is assumed that ordinary chemical reaction is caused by nearly monochromatic radiation, is not tenable. Not only is there no experimental evidence in its favor, but it is certain that the number of collisions between the molecules and such very restricted light quanta would be inadequate to account for the observed rates of reaction. It is shown that both collisions and radiant energy offer opportunities for the activation of molecules far greater than the number required to account for observed reaction rates.” In a similarly thorough paper Tolmang discusses “the various mechanisms of chemical activation that have been proposed, in order to estimate their possible importance for chemical reaction.” He concludes that “activation by collision with molecules of high enough kinetic energy cannot take place fast enough to account for the decomposition of nitrogen pentoxide or other unimolecular gas reaction. Activation by collision also could not take place fast enough to account for the decomposition of nitrous oxide or other bimolecular reactions, if the total energy of activation is brought into the reaction by one of the two molecules. . , ilctivation, in accordance with the

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J. Chim. phys., 2 2 , 73 (1925). * Industrial Chemist, 2 , 226 (1926). Technique moderne, 17, 757 (1925). ‘Chem. Age (London! 14, 324, 348 (1926). 6 J. Chem. Ed., 3, 385 (1926). Chem. Weekblsd, 2 2 , 317 (1925). ‘Chem. Weekblad, 2 2 , 356 (1925). e J. Am. Chem. Soc., 47, 1508 (1925). 9 J. Am. Chem. SOC., 47, 1524 (1925). 1

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simple radiation theory, by the adsorption of the frequency calculated by assuming the energy of activation to be taken up as a single quantum, cannot take place fast enough to account for the decomposition of nitrogen pentoxide, or other unimolecular reactions, and leads to incorrect predictions as to the frequencies that will be active. . . In conclusion it can be stated that activation may occur t o some extent by all the four mechanisms suggested. Indeed, since the great difficulty is to account for the rapidity with which the energy of activation is supplied, we must not despise the assistance afforded by any method of activation. However, it seems at the present as if our main :lopes must be centered on the elaborated radiation theory.” Cofman-Nicoresti’ regards the catalyst as “a separator of heat from other elements in a reaction.’’ Wolfenden2 in a study of the critical potentials of hydrogen in presence of catalytic nickel and copper, concludes that substantial quantities of atomic hydrogen are present on these metals. The ionization curve in presence of the catalytic metal resembles that of an incandescent grid. Theory of Surface Action In the Fourth Report of this Committee, H. S. Taylor3 has developed in fine fashion the theory of catalytic action on a surface. His ideas have gained general acceptance and have had a great influence on recent investigations. The remarkable magnitude of the heat of absorption of oxygen on charcoal observed by Garner4 led him to believe “that the valency of the surface carbon atoms is very far from being satisfied by the neighboring carbon atoms.” He feels that the differencebeween the reactions of carbon and oxygen and a catalytic reaction lies in the fact that the atoms of the catalyst do not leave the surface although their relative positions with respect to one another may change during the reaction. I t is, however, clear that those atoms lying in an exposed position on the surface, by virtue of their high energy content, will play a large part in the chemical process. They will be the most likely to possess the energy necessary for promoting the trigger actionof catalysis. Also, it can be seen why the surface becomes increasingly effective with use. If part of the heat liberated during the heterogeneous reaction is absorbed by the surface, this may result not only in an increase in the total surface but also in an increase in the proportion of exposed atoms on thesurface,z.e.,an increased free energy of surface per sq. cm. “It is not unlikely that a comparatively small number of the total atoms on the surface are the active agents in promoting chemical change.” The action of a catalytic surface has been studied by Hinshelwood and his coworkers in a series of papers. Hinshelwood and Pritchard5 explain the decomposition of nitrous oxide on a hot platinum wire by assuming that “the nitrous oxide gives its oxygen atom to the bare platinum surface . . . the Pharm. J., 115, 345 (1925). Proc. Roy. SOC.,llOA, 464 (1926). J. Phys. Chem., 30, 145 (1926). See also Proc. Roy. SOC.,108A, r o j (1925). Nature, 114,932 (1924). J. Chem. SOC.,127, 327 (1925).

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function of the surface is thus to act as an acceptor for atomic oxygen, thereby rendering possible an unimolecular in place of a bimolecular process. What makes the odd oxygen atom more easily detachable is its affinity either for the platinum or for atomic oxygen already on the platinum.” Green and Hinshelwood’ have studied the decomposition of nitric oxide on platinum and believe that this “provides a further example of a reaction that is bimolecular in the gas phase becoming unimolecular a t the surface of a catalyst. The reaction NO = K 0 is therefore rendered possible without the communication of energy in prohibitive amounts.” Hinshelwood and Pritchard* have investigated the reaction between hydrogen and carbon monoxide on a platinum wire a t 1000’. They say: “We must therefore assume that the catalytic activity of the surface is localized in certain active points forming a small fraction only of the total surface. A simple mechanism which now accounts for most of the facts is that reaction occurs when hydrogen and carbon dioxide become adsorbed adjacent to each other on an active part of the surface. I t must be assumed further that the fraction of the active surface covered by carbon dioxide incremes from zero to nearly unity as the pressure of carbon dioxide increases from o to 400 mm, while the adsorption of hydrogen on those points left free from carbon dioxide is never very great, so that we have to deal with that portion of the hydrogen adsorption isotherm where adsorption is more or less directly proportional to pressure.” I n the second paper they find: “The adsorption of each gas on the active centers of the catalyst is almost independent of the pressure of the other gas. This shows that not all the surface is active, but that only certain parts are able to adsorb hydrogen and carbon dioxide and cause them to react. The parts which adsorb hydrogen in this way are different from those which adsorb and render active the carbon dioxide. Interaction apparently takes place when molecules of the two gases are adsorbed on adjacent centers of the appropriate kind.” The same authorsSfind that the decomposition of hydriodic acid on a gold surface is of the zero order and is uninfluenced by the pressure of the hydrogen. “The function of the surface in the catalyzed decomposition reactions of nitrous oxide on platinum, hydriodic acid on gold and ammonia on tungsten is to permit the occurrence of a unimolecular instead of a bimolecular process requiring an energy of activation about twice as great.” Hinshelwood and Burk4 have studied the decomposition of ammonia and of hydriodic acid on various surfaces and confirm the above views. “Hence we must conclude once more that different reactions are provoked at different points (“active centers”) on the catalytic surface.” Hutchinson and Hinshelwood5 have studied the interaction of hydrogen and nitrous oxide on the surface of a gold wire. “The results can be inter-

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2 J

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J. Chem. SOC., 1926, 17c9.

J. J. J. J.

Chem. Chem. Chem. Chem.

SOC., 127, 806, 1546 (1925). Soc., 127, 1552 (1925). SOC., 127, 1x05,2896 (1925). SOC., 1926, 1556.


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preted by assuming that hydrogen and nitrous oxide are adsorbed on the surface independently of each other, and that interaction can take place between adjacent molecules. . . . A second effect was traced to hydrogen, namely a steady decrease in the activity of the wire, especially when it was heated in mixtures containing hydrogen in excess. . . I t is suggested that hydrogen gradually disssolves in the body of the metal-this solution being distinct from its primary adsorption on the surface-and decreases the adsorptive power towards nitrous oxide.” Burk’ goes further into the details of adsorption by assuming that the two ends of a molecule are independently adsorbed by active points on the catalytic surface. If these points are further apart than the normal length of the molecule, it would be stretched and the energy required to break it up thereby lessened. Faresti? also believes that the surface of the subdivided nickel is constituted of adsorbing centers of various power and that the catalytic activity depends on which centers are capable of adsorbing the gas from the strongest bonds. The activity of a copper catalyst is found by Palmer and Constable3 to vary periodically with the temperature of reduction from the oxide, there being three maxima between 220’ and 420’. They4 observe that “the rates of dehydrogenation of the primary alcohols, ethyl, propyl, butyl, isobutyl and isoamyl are all equal within the limits of experimental error; and the temperature coefficient of the velocity is the same for all. . . . Secondary propyl alcohol reacts with a velocity about five times that of the primary alcohols. “Reaction occurs in an adsorption film covering the copper surface. The film becomes one molecule thick a t 28ooC. and its thickness increases as the temperature falls. All primary alcohols contain the --CH20H group a t the end of the hydrocarbon chain. The rate of dehydrogenation has been shown to remain constant while the length of the hydrocarbon chain is doubled. I t is very improbable, considering the very specific action of catalysts, that reaction could be initiated at a distance from the surface. These observations, therefore, show that the primary alcohols are adsorbed, the -CH20H group in contact with the surface, and the hydrocarbon chains in contact perpendicular to the copper surface.” Constable5 maintains that “chemical reactions occur only when alcohol molecules are adsorbed over a characteristic arrangement of copper atoms.” I n a second paper he8 finds that “if the molecules only react in the unimolecular film next to the catalyst surface, then varying the pressure only affects the change in so far as it affects the life of a molecule in the surface . Hence the reaction velocity should be practically independent layer.

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1 J. Phys. Chem., 30, 1132 (1926). *Gam., 5 5 , 185 (1925), Proc. Roy. SOC.,106.4, 250 (1924). ‘ Proc. Roy. Soc., 107.A, 255 (1925). 5 Proc. Roy. Soc., 107A, 270 (1925). 6 Proc. Roy. SOC., 107A, 279 (1925).

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of the pressure, provided the catalyst is completely covered with the adsorbed alcohol film.” In two further papers Constable’ considers the effects of diluents on the rates of surface catalysis. Assuming the theory of catalytic centers he arrives by mathematical analysis a t a formula which he verifies for the dehydration of alcohol over copper. The same author* gives a mathematical treatment of the dynamics of surface action in a closed vessel. The formulae arrived at are verified by experiments. While the activity of a metallic surface depends on the method of preparation, Levi3 finds no difference in the performances of thoria catalysts made in seven widely different ways. Gauger4 found thin films of platinum or nickel on Pyrex glass to be inactive and concludes that the importance of mere extent of surface has been overestimated. “It seems not at all unlikely that a catalyst may have electrons that are on somewhat different energy levels than in the case of the crystalline metal. The so-called active. surface of the catalyst consists, therefore, of those molecules that have electrons in these outer energy levels.” Almquist and Crittenden5find that copper reduced from the oxide is much more effective than massive copper in removing oxygen from mixtures of hydrogen and nitrogen. Hoover and Ridea16have studied the decomposition of ethanol over thoria. They find both dehydration and dehydrogenation. They conclude that “the theory that the two reactions are promoted by different active areas or patches on the surface has been shown to offer a satisfactory explanation of the observed phenomena. Support for this theory is found from the fact that the two reactions require different energies of activation, that poisons cut down one reaction to a greater extent than the other, that chloroform promotes the ethylene reaction, and that the ratio of the reaction velocities of the two reactions is altered a t low pressures. , . I t seems likely that both the orientation of the surface atoms and the adsorptive power of the active patch play a part in determining the type of reaction that proceeds on the active patch. At atmospheric pressure the reaction velocities depend in part on the poisoning effect of the products of reaction.” Constable’ prepared copper in many different ways and measured its activity. A smooth copper surface is probably less than 1/5ooo as active as the same area of reduced copper. “The essential feature necessary to the production of copper catalytically active in the dehydrogenation of alcohols is the sudden liberation of free copper atoms under conditions in which the kinetic energy of the atoms of the structure is insufficient to cause the collapse of the active centres. Thus methods of producing copper under the same

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Kature, 116, 278 (1925); Proc. Camb. Phil. SOC.,23, I 7 2 ( 1 5 ~ 0 . Nature, 117, 230 (1926). 3 Atti. hccad. Lincei, (6) 2 , 419 (1925). 4 J . Am. Chem. SOC.,47, 2278 (192j). 5 Ind. Eng. Chem.. 18, 866 (1926). 8 J. Am. Chem. SOC., 49, 104 (1927). Proc. Roy. SOC.,110.4,283 (19261. 1


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physical conditions should give copper with the same surface activity and showing the same temperature coefficient. In spite of the wide divergence between the spacing of the copper atoms in copper formate and valerate thermal decomposition gives a product showing very nearly the same temperature coefficient and surface activity. “Since the surface fields vary with the nature as well as the arrangement of the surface atoms, the nature and the magnitude of the molecular distortion will vary with the chemical nature and physical state of the surface, and we have an explanation of the specific action of catalysts. It is apparent that very small quantities of added material of suitable physical properties could cause wide variation in the nature and distribution of the centres of activity, and we have a tentative explanation of promoter action. “The persistence of the centres of activity unchanged by chemical reaction requires that the adsorption and desorption of the reactants shall be reversible without alteration of the surface; and that the heat absorption or evolution shall not sinter the existing centres of activity during reaction. Many chemical changes are violent enough to cause marked changes in the surface. I n these cases the activity of the surface increases on use to a maximum, which corresponds with the maximum number of centres capable of existing together at a given instant.” Fryling’ regards some of the surface atoms of a nickel catalyst as being in an almost gaseous condition which he considers as corresponding to a high temperature. He assumes “that active nickel atoms dissociate hydrogen molecules.” Heat treatment of a catalyst reduces the heat of absorption. This reduction is less in the case of a promoted catalyst. An interesting extension of his former theories is given by Taylor2 to explain the dual action of oxide catalysts. “It is significant that, in the case of formic acid decomposition, in presence of metal catalysts the reaction products are exclusively those of dehydrogenation: hydrogen and carbon dioxide, irrespective of the metal. With oxide catalysts, on the other hand, the products may be dehydrogenation as, for example, with zinc oxide, or mainly dehydration as, for example, with aluminium oxide. To what can this variation be attributed? A metal surface is composed of metal ions and electrons, an oxide surface of metal ions and oxide ions. The conclusion seems inevitable that, on the metal ion (positively charged) the dehydrogenation process occurs, whilst on the oxide ion, dehydration occurs. Attachment to the positive ion bas the effect of giving the hydrogen atoms in the formic acid molecule greater freedom, in agreement with the observation made above that substitution by an electrogenative substituent such as chlorine causes the hydrogen ion to have greater freedom. The negative oxide ion causes the hydrogen atoms to be more firmly attached and so the dehydration split is favored. On this basis, an oxide,catalyst surface is to be regarded as composed, not of a single catalyst, but of two catalysts, metal ions and oxide ions and the J. Phys. Chem., 30,818 (1926).

* Fourth Colloid Symposium, 19 (1926).

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nature of the changes induced in the adsorbed reactant is determined by the charge of the ion on which the reactant molecule is adsorbed. The extent of the two alternative changes will be determined by the relative extent of adsorption of reactant on the two ions, on the relative frequency of the two ions in the surface and on their specific individual catalytic activities. These several factors, extent of adsorption, frequency of ions in the surface and catalytic activity will be determined by the degree of saturation of the lattice ions (Le., catalyst structure) and by the extent to which the ions are already covered by poisons (salts, ammonia, water, etc.)” Russell adopts the view that, for chromium, manganese, iron, cobalt and nickel, the number of electrons in the outermost orbit is between one and two and is, therefore, either one or two. The consequence is that these atoms can never be electrically neutral in the sense that other atoms are. Hence each of these metals has an active and a passive state according as the outer orbit contains two electrons or only one. These metals “adsorb gases and act as catalysts when in the passive form. They adsorb because of their property of never being electrically neutral” and they catalyse a reaction “because of their potentiality of passing from the passive to the active form.” Tzentnershver and Steaumanis*have measured the over-voltages of various metals and compared them with the catalytic effects of these metals on rate of solution of zinc in acids. The orders of magnitude turned out to be entirely different. Baudisch and Welo3 have investigated active ferric oxide made by heating in oxygen to 330’ and inactive which had been heated to 550’. The X-ray spectrograms were entirely different. They conclude that the catalytic activity is connected with the “spatial arrangements of the electrons and protons in the atom. The activity seems to depend on the size of the crystals and upon the intramolecular arrangement and not upon the presence of Fe” nor on the adsorptive capacity.” Levi and Haardt4 in a study of the relation of catalytic activity to extent of surface have determined particle size of a platinum catalyst. The area decreased with heat treatment. The following were the areas of 0.01g. catalyst at different temperatures’:60°, 5588 sq. cm.; IIO’, 5 0 5 0 ; I~o’,3300; 180°, 1880;ZI~’, 1385. “The experiments indicate that the catalytic power of a metal is predominatingly a function of its surface area.” On the contrary Rocasolano5 holds that the effectiveness of catalysts depends on composition rather than on physical properties. He thinks that platinum particles have a shell of oxide which is alternately reduced and oxidised. According to Rice6 many of the systems which have been regarded as homogeneous are quite otherwise. As ordinary distilled water contains some

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Kature, 117, 47 (1926). Z. physik. Chem., 118, 438 ( 1 9 2 5 ) . a Chem. Ztg., 49, 661;J. Biol. Chem., 65, 215 (1925). Gam., 56,424 (1926);Atti. Accad Lineei, (6)3, 91, 2 1 5 (1926). 6 Nat. Ges. Wias. Gottingen, 1924,177. J. Am. Chem. SOC., 48, 2099 (1926).

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2 j,ooo suspended particles per cc., some reactions, a t least, may be confined to the surface of these particles. “The thermal decomposition of hydrogen peroxide is an example of a reaction that is to some extent catalyzed by the walls of the vessel but mainly by suspended dust particles; when hydrogen peroxide, free from dust, is kept in a vessel of fused silica it is a remarkably stable substance and may be heated for several days to 60’ without appreciable decomposition.” He suspects that in many reactions the real catalyst is the suspended dust which is promoted or poisoned by other substances. Dust may be responsible for anomalous results in many supposedly homogeneous reactions. Many of the theories previously put forward relative to the decomposition of hydrogen peroxide may require revision. Adsorption and Catalytic Activity Since it is generally admitted that adsorption by a surface is necessary Eo catalytic effect it is natural that many investigations have been made to relate the two quantitatively. While catalysis may not be possible without adsorption, it is certain that adsorption takes place frequently without leading to any chemical change. The conclusion from many measurements is that, while the two are closely related, the one does not follow the other according to any mathematical relation. Alexeyevskii’ has measured the adsorption of 7 2 organic compounds and bromine by animal charcoal and by calcined copper sulphate. Hoskins and Bray2 find that manganese dioxide, copper oxide, and mixtures of the two adsorb larger amounts of carbon monoxide than of any other gas. The rate of adsorption of carbon monoxide by one of these is in general closely related to its activity in oxidizing this gas. Firth and Watson3 conclude that the catalytic activity of charcoal in the decomposition of hydrogen peroxide is not related to its capacity for absorbing iodine from a chloroform solution. Lazier and Adkins4 have measured the adsorption of ethylene and hydrogen by zinc oxide, ferric oxide, nickel, and copper; and also hydrogenation rates with these catalysts. They say that: “The results of the adsorption and catalytic activity measurements indicate that there is a qualitative agreement between the two manifestations of chemical activity, but there is apparently no quantitative relationship between total adsorption and catalytic activity.” Duclaux6 assumes that atoms or molecules which are brought in contact These addition comwith one another give addition compounds. pounds are formed spontaneously, Le., either without activation or, more probably, by autoactivation. Adsorption is but a particular case of the . These addition compounds can underformation of these compounds. go, either without activation or by autoactivation, an internal transposition

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J. R u m Phys. Chem. SOC., 5 5 , 401 (1924); Chem. Abs., 19, 2634 (1927). SOC., 48, 1454 (1926). J. Phys. Chem., 29, 987 (1925). J. Phys. Chem., 30, 353 (1926). “DeuxiPme Cons. Solvay,” 630-45;Chem. Abs., 20, 3614 (1926).

* J. Am. Chem.

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which can in turn be followed by dissociation. Under these conditions the function of the catalyzer consists essentially in allowing of a transposition which is equivalent to a reaction which, in its absence, would take place with difficulty and in low yield, or else a t a higher temperature.’’ Pearce and Alvarado‘ have studied the adsorption of vapors of water, ethanol, acetic acid and ethyl acetate by thoria and alumina a t 99.4’. They believe the most plausible explanation of catalytic esterification to be that “in the passage of the vapors of ethyl alcohol and acetic acid over alumina or thoria the alcohol molecule is strongly adsorbed, primarily through the residual valencies of the hydroxyl group. Under the prevailing conditions of stress the less highly attracted ethylidene group splits off from the alcohol molecule and there combines with the acid molecule to form the ester.” Bischoff and Adkins2 measured the adsorption by three varieties of titania of certain products of the catalytic reactions in order “to discover the relationship that has been supposed to exist between the adsorptive power of substances and their catalytic activity,” but find that “the adsorption measurements of the products of reaction at the surface of titania show no connection with either the activity of the catalyst or its selective effects.” “It is important to note that the adsorbing powers of the three catalysts are not in the same order as their catalytic activity.” Bone and Andrew3 find that the catalytic activity of a gold surface for the union of carbon monoxide and oxygen is increased by exposure to either gas. Both gases seem t o be activated a t the catalytic surface. That the relation between adsorption and catalytic activity is not as simple as might be supposed appears from the observations of Remy and Gonningen4 who studied the synthesis of water over the metals of the iron group and ruthenium and their alloys. Previous treatment with oxygen makes iron alloys more active, while hydrogen increases the activity of nickel alloys; but pure iron and iron-ruthenium alloys are less active after treatment with oxygen. Veil6 finds that certain metallic hydroxides alter their magnetic properties progressively while catalyzing the decomposition of hydrogen peroxide. Boswell and Bayley6 consider the normal platinum catalyst to consist of particles of platinum with interior content of oxygen, the particles being surrounded by a layer of dissociated water, the hydrogen and hydroxyl ions alternating. This layer they believe to be the seat of the catalytic action. X nickel catalyst is similar. Boswell and Dilworth’ extend this idea to alumina. “All the above reactions catalysed by aluminium oxide involve the elements of water, and in 1

J. Phys. Chem., 29, 256 (1925).

J. Am. Chem. SOC.,47, 807 (1925). 3Proc. Roy. SOC., 109A,459 (1925). 4 Z.anorg. allg. Chem., 149, 283 (192 5). 6 Compt. rend., 182, 1028 (1926). 6 J. Phys. Chem., 29, 11 (192j). J. Phya. Chem., 29, 1489 (1925). 9


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the two former either the addition or removal of water to or from the other reacting compound. Associated with this is the fact that aluminium oxide catalyst contains water, and the further fact that when this water content is diminished by heating at a high temperature the catalytic activity is likewise diminished. This all points to the conclusion that this catalyst functions by means of a surface film of water and that this film is the real seat of the catalysis. The marked stability of the film indicates that it is present in a special condition such as positively charged hydrogens and negatively charged hydroxyls, alternating with each other and completely enveloping each aluminum oxide particle. “That is, the initiation of the reactions comes from the charged hydrogens and hydroxyls on the aluminium oxide particles. These react with the hydrogens or hydroxyls, or both, in the compounds brought in contact with the heated surfaces, thus setting up the reactions which always result in the formation of as much water as has been removed from the catalytic surface. The surface film is thus restored and the cycle continues. “The catalyst from this point of view does not accelerate a reaction already in progress . . . but actually initiates the change. . . .” Intermediate Compounds There have been two extreme views as to the action of a catalyst. One, that the catalyst simply speeds up a reaction by its presence without taking any part in it. The other, that the catalyst is actually a reactant but is regenerated in the end. The latter view, or intermediate compound theory goes back to Desormes and Clement.’ This has been strongly advocated by Sabatier all along and has been “the guiding star” in all of his investigations as he says in his book and in his address before the American Chemical Society.* He believes that the temporary hydrides S i H l and S i H z are formed and give their hydrogen to the organic compounds in catalytic hydrogenation. Kubota and Toshikawas regard reduced nickel catalysts as containing many active unstable hydrides. They divide these into three classes, the first or most active being able to hydrogenate benzene and to be poisoned by thiophene. The second can reduce ethylenic groups and is poisoned by ethyl sulphide while the last is only able to reduce nitro-groups and is poisoned only by hydrogen sulphide. The same theory is promulgated by Nittasch‘; but he is somewhat more liberal in his definition of compound. He feels that in homogeneous systems, particularly in solutions, there is no doubt that catalysis takes place by causing intermediate reactions. After enumerating the difficulties that stand in the way of exact measurements he says: “Yet it remains, according to my judgment, that here also the zntermedzate reactzon hypothesis zs the most suggestzve and most sntzsfactory explanation, provzded one takes a broad enough mew;i.e.,when it is taken into account that inaddition to Ann. Chim. Phys., 5 9 , 329 (1806). Chem., 18, IOOS (1926); Chem. Listy, 20, 45 (1926). Sci. Papers Inst. Phys. Chem. Research. 3 . 223 (192j). Ber., 5 9 , 13B (1926).

* Ind. Eng. 3

1’34

E. EMMET R E I D

the chemical processes with their velocities there are also physical processes with their own velocities.” There is not always agreement among investigators as to the exact hypothetical intermediate processes involved in a catalytic reaction. As is known, some follow Sabatier in considering hydrogenation with nickel as more chemical, involving labile nickel hydride, possibly with the participation of oxygen (Schlenk, Willstatter, Waldschmidt) , while others regard it as more of a physical process in which adsorption plays the leading role: “This adsorption may be simply physical, depending on capillarity and the critical constants of the gases,or may be specific and chemical involving valence forces.” Person’ reviews the theories of catalysis from the standpoint of intermediate compounds. He finds it possible to approach catalytic reactions through the electron theory. Child and Adkine,2 in a study of the condensation of an aldehyde to an ester, assume that an aldehyde-catalyst complex is formed rapidly and that this decomposes slowly into the ester and catalyst. Duparc, Wenger and UrferS regard catalytic oxidation in the presence of the metals of the platinum group as depending on the formation of oxides of these metals. “With platinum we see that the yields [of sulphur trioxide] are small while the metal is PtO, they increase suddenly when the platinum goes to the condition PtO,, and they come down again [with rise of temperature] as the oxide begins to dissociate, till at 800” the yields are zero.”

E. and F. Muller4 suppose a series of intermediate compounds in the decomposition of formoldehyde by the metals of the palladium group. Benton and Emmett5conclude that the union of oxygen and hydrogen in contact with reduced nickel is accounted for by the alternate oxidation and reduction of the metal. Andrussov6 postulates intermediate compounds in the catalytic oxidation of hydrocyanic acid and ammonia. Travers’ assumes an intermediate complex in the oxidation of manganese. Yost8 supposes the trivalent silver ion to be the active agent in oxidation of ammonia by persulphates. Intermediate reactions are discussed by Rosenmund and Jordan9 and by Gulevichloin the catalytic reaction of oximes. Rosenmund” assumes an intermediate complex in the formation of secondary amines. Euler, Olander and RudbergI2 deduce formulae for velocity of the mutarotation of glucose from assumptions that glucose forms salts amphoteri1 2

J. Russ. P h p - C h e m . Soc., 57, 189 (1925). J. Am. Chem. Soc., 47, 798 (1925).

3

Helv. Chem. Acta., 8, 609 (I92jj.

4

Z. Elektrochemie, 31, 41 ( I 9 2 j ) . J. Am. Chem. SOC.,48, 632 (1926).

6

Ber., 59, 458 (1926). Compt. rend., 182, 972 (1926). 8 J. Am. Chem. SOC.,48, 374 (1926). 9 Ber., 58, 51 (1925). ‘@Ber.,58, 798 (1925). ” 2 . angew. Chem., 38, 145 (1925). 1%. anorg. allg. Chem., 146, 45 (1925). 6

7

FIFTH REPORT O F THE CUMMITTEE ON CONTACT C.4TALYSIS

I I3 j

cally. Euler and Olander’ say: “Among the large number of catalyzed hydrolyses, not a single case is met with where it is objectionable to assume the formation of a salt between the catalyst and the reacting substances.” Spitalsky* develops equations for homogeneous catalysis on the basis of the formation of intermediate compounds between catalyst and reactant. All cases are thoroughly worked out and it is shown that the order of the reaction between catalyst and substrate may be entirely different from the apparent order of reaction deduced from the curve. The order of a catalyst reaction is determined by the affinity constant of the catalyst to the substrate so that a reaction of the nth order may run according to any order between o and n. ?rlerezhkorskii3regards variable valence as a prime requisite for a catalyst. “For a given reaction a catalyst is a body capable of modifying the valence of the reactants, able to combine reversibly with the reactants, at the same time having at least two stages of oxidation.” llanchot and Gall‘ find that the platinum acts as hydrogen acceptor in the reversible dehydrogenation of hydroquinol. Job5 explains a number of catalytic reactions by the formation of unstable electronic complexes which decompose, giving the final product of the reaction and regenerating the catalyst. Two articles appeared practically simultaneously on the transformation of maleic into fumaric acid, by Neerwein and Weber6 and Terry and Eichelberger.’ I n both, the assumption is made that the catalyst does not actually add to the double bond but only activates it. If one component of the double bond is disrupted there can be free rotation about the other. The latter authors believe the activation to be due to the addition of the catalyst to the unsaturated carboxyl groups. Lebeder, Koblinasky and Yabubchik8 assume that hydrogenation depends on an unstable intermediate product (of the nature of an adsorption compound) between catalyst and unsaturated molecule. Zelinskiig stresses the differences that must exist between the motions of particles in the surface and of those in the interior of an object. There are also peculiar conditions at the interface when two bodies touch. These may approach chemical phenomena which take place only on contact. He concludes: “The mechanism of catalysis does not demand the formation of intermediate compounds but is dependent upon the degree of tension of the chemical system, which tension is necessary for the initiation of chemical action and is brought about by the contact of bodies at ordinary or at elevated temperature. Z. anorg. allg. Chem., 152. 113 (1926). Z. phgsik. Chem., 122, 257 (1926). Bull., (4) 39, 41-3 (1926). Ber., 5 8 , 486 (1925). “DeuxiPme. Cons. Solvay,” 417 (1926). Ber., 5 8 . 1266 (1925). 7 J. h m . Chem. SOC.,47, 1402 (19~51. J. Chem. Soc., 127, 417 (1925). 8 Ber., 5 8 , 2 7 j j (1925). I

2

I

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E. EMMET REID

Alterattons o j the dynamic state of molecules during catalytzc process are not doubted. These give the chief impulse to catalytic action.’’ He quotes with approval Bodenstein’s‘ statement that the formation of water from oxygen and hydrogen on a platinum surface concerns only the surface and that the adsorbed hydrogen is present not RS hydride nor as atomic hydrogen but as deformed hydrogen molecules. The deformation results from the field of force on the surface of the catalyst which induces the adsorption and effects the deformation of the molecule.” Partington2 quotes Graham3 as saying that the hydrogen molecule is adsorbed by one end and thereby polarized, the outer end having its affinity for oxygen increased. Other chemists are outspoken in their opposition to the theory of intermediate compounds. Larson and Smith,4 after a study of the catalytic synthesis of water, say: “With both nickel and copper the reaction was accompanied by the gradual formation of a stable oxide on the catalyst surface. This oxide decreased the catalytic activity of the metal.” Chapman, Ramsbottom and Trotman5 studied the union of hydrogen and oxygen. They conclude: “Where silver is heated to dull redness in oxygen a t a pressure exceeding 0.00j mm. of mercury, it becomes almost completely covered with a film of oxide, which considerably impairs the efficiency of the metal as a catalyst for the reaction between hydrogen and oxygen. The film of oxide can be removed by heating the metal to redness in a vacuum or in oxygen at a much lower pressure. The catalytic activity of the metal is then the same as that of silver which has been heated in hydrogen. These hypotheses furnish a consistent explanation of the discovery of Bone and Wheeler that silver, after it has been heated in hydrogen, becomes a much more efficient catalyst for the reaction between hydrogen and oxygen than silver which has not been thus treated. . . The behaviour of gold is similar to that of silver . . although . . . not so pronounced.”

.

.

Bone and Andrew6 investigated the union of carbon monoxide and oxygen over nickel and copper and their oxides. They conclude “that, in cases of surfaces composed either of readily oxidizable metals or reducible oxides, the true catalytic combination of the gases does not involve any alternating oxidation and reduction of the surface, as many have supposed; indeed, it was shown that any such purely chemical view is quite inadequate to explain the phenomenon. Thus, in the case of copper oxide, the truly catalytic process was shown to depend primarily upon the condensation of a film of “activated” oxygen on the surface; indeed, so far from the CuO participating directly in the formation of steam, it was actually protected by the film of “activated” oxygen from the attacks of the hydrogen, which otherwise would have energetically reduced it. -

’ A n n , 440. 1 7 7 (1924).

Nature, 115. 534 (192s). Proc Roy. SOC , 16, 422 (1868). J. Am. Chem. S O C ,47, 346 (1925). 5Proc. Roy. Soc., 107.4, 92 (1925). 6 Proc. Roy. S O C . , 110.4, 16 (1926). 9


1137

“The facts observed led to the conclusion that the real phenomenon of surface combustion “depends primarily on a condensation of one or other (and in some cases, possibly both) of the reacting gases on the heated surface,” and that the prime cause is an “activation” of the combining gases (certainly of the hydrogen, and possibly also of the oxygen) by “association” with the surface, the question being left open whether such “association” is a mere surface “condensation” or some deeper “occlusion.” “The catalyzing power of the surface, when i n “normal” activity, can be highly stimulated by previous exposure to either carbon monoxide or oxygen a t the said experimental temperature; in neither case, however, was the full extent of such stimulus immediately manifested when the stimulating gas was rapidly removed and a fresh charge of the reacting mixture (nCOfO2) reintroduced into the apparatus. “Although the catalytic combustion was subsequently found to depend on the presence of moisture in the system, no hydrogen was ever detected in the reacting gases, even when a large excess of carbon monoxide was present. “A microscopic examination of the gold surface, at the conclusion of the experiments failed to reveal any signs of “pitting” having occurred; on the contrary, neither the lustre nor the smoothness of the metal seemed to have suffered any diminution.” Boesekenl combats, with vigor and effect, the idea of intermediate compounds. It has been shown that aluminum chloride combines with acid chlorides and this has been considered an explanation of the Friedel and Crafts reaction. I t is here brought out that molecules such as chlorine, ethyl chloride, carbon tetrachloride, which either do not combine with aluminum chloride or do so with difficulty, react on aromatic hydrocarbons much more vigorously and with the aid of much less of the catalyst than the acid chlorides which form molecular compounds with the aluminum chloride. Benzyl chloride “does not give an addition product but reacts rapidly a t oo while its nitro derivative forms a well-crystallized addition product N02C3HaCH2Cl.AlCl~ which reacts with benzene but does so only a t a higher temperature, Le., under the conditions under which the addition compound begans to dissociate. The formation of addition products may inhibit the reaction entirely and we may conclude that the formation of addition products never gives a satisfactory explanation of catalytic action. If the catalyst and one of the reactants combine they usually do so with a liberation of energy which leaves the system with less energy than it had before. This means that a larger amount of energy will be required to start the reaction.” “The explanation of catalytic phenomena must be sought in what happens during the contact between the catalyst and the activated molecules. I assume that particular bonds are modified and I have called this dislocation. We can imagine that the catalyst without combining intimately with the bond which it is to activate, changes the orbits of the electrons in such a way that the molecule can act much more rapidly than without this change.” Rec. Trav. chim., 45, 458 (1926).

1138

E. EMMET REID

Camers Rosenmund and Joithe’ think that the carrier plays a part in the reaction. In the reduction of borneol to iso-camphane, it has been assumed that the alumina caused water to split off and then the nickel effected hydrogenation. They consider that the splitting off of the w a t u would be too slow at the low temperature and assume formation of complexes. “The borneol and H,are bound to the catalyst by partial affinities. This using up of the affinities on the outside weakens the internal bonds and prepares for a rearrangement of the forces. . . The r61e of the alumina is a double one; in consequence of its powerful adsorptive power it binds the borneol to the catalyst, secondly it splits mater from the labile complex and helps finish the process. The reduction process appears as condensation between ROH and H2 in which the nickel loosens up the H, while the alumina aids the splitting off of the water. . . . The complex-theory of catalyst gains ground more and more as we have investigated it. Wieland,* who has formerly proposed complex-formation in certain reduction processes, now holds it,3 as do we, of general application.”

.

Modification of Catalysts In promoted action and in negative catalysis we have a change in the velocity of a catalyzed reaction, here we have to do with a modification of the direction of the reaction. I t is quite common that a catalyst facilitates two or more reactions as thoria dehydrates and also dehydrogenates an alcohol, the two reactions going on simultaneously. Adkins has shown that alumina prepared in different ways gives very different ratios of ethylene and hydrogen. If we add to a double-acting catalyst a selective promoter, one which will speed up the one of the reactions but not the other, almost the whole product will be that of the favored reaction. Extremely important results have been obtained by Adams and his COworkers on the modification of the platinum oxide catalyst by the addition of various substances, particularly iron salts. The hydrogenation of a compound, such as citral or cinnamic aldehyde in which a double bond and an aldehyde group are present, can thus be directed so as to yield the unsaturated alcohol or the saturated aldehyde practically completely. For details of procedure and many examples reference is made t o the papers of Carothers and Kern, Shriner and Adams6 and Adams and Garvey.6 Faillebin? also has studied the modification of a platinum catalyst by the addition of iron and other substances. Xomatsu and Kuratas have experimented with modified copper catalysts. Harag prepared copper catalysts one of which was deBer., 58,

(192j). (1912). 3 Ergeb. Physiol., 20, 516 (1922). J. Am. Chem. SOC.,47, 1047 (1925). J. Am. Chem. SOC.,47, 1147 (1925). J. Am. Chem. S O C ,4 8 , 477 (1926). ‘Ann. C h m . , 4 , 156, 410 (1925); Compt. r e n d , 182, 132 (1926). 8 M e m . Coll. Sc. Kyoto, 8A, 35, 147 (1925). Mem. Coll Sc Kyoto, 9A, 405 (1926) l

205

* Ber., 145, 484


1139

hydrating and another dehydrogenating. Komatsu and Masumoto‘ compared nickel, copper and thoria catalysts at different temperatures and found that the activity of thoria is most influenced by temperature. hrmstrong and Hilditch* investigated the various factors in selective hydrogenation particularly the relative location of the double bonds. Richardson and Snoddl” have studied the selectivity of platinum and nickel catalysts in the hydrogenation of oils containing more than one double bond. Rosenmund and Jordan4 hydrogenated aromatic aldehydes. By partially poisoning the catalyst they were able to stop at the alcohol. Zetzsches with Arnd, Enderlin, Flutsch, Menzi, and Loosli as coworkers, has used partially poisoned or “regulated” catalysts for the hydrogenation of acid chlorides. According to Paal and Poetke6 the action of palladium on formic acid is different according to whether the solution is acid or neutral. Galecki and Mlle. Bincer’ and Galecki and Mlle. Krzeczkovskas have experimented with gold sols prepared in various ways and found great differences in their activities. Promoters Under this head are grouped a variety of phenomena. Tuly and -&damsg find that zinc and iron salts are truly remarkable as promoters with the platinum oxide catalyst. I n this way cinnamic acid can be hydrogenated to the corresponding alcohol without reactivation of the catalyst. Bray and DossIO show that a mixture of manganese dioxide and copper oxide is more efficient in the catalytic oxidation of carbon monoxide than either of these oxides alone. Robertson” in a study of the decomposition of hydrogen peroxide by ferric and cupric salts and mixtures of these, defines promotion provisionally “as a change in the path of a reaction with a concomitant displacement of thesteady state.” The absorption spectra indicated the presence of cupric acid when ferric iron was promoted by copper. Quartaroli12 has investigated the same reaction with the same two catalysts. The activity of cupric hydroxide is increased so-fold by ferric hydroxide which alone is only slightly active. The most powerful catalyst is a mixture of the two in presence of alkali. He finds wide variations in the activity of catalysts according to the proportions of the reactants. Mem. Coll. Sc. Kyoto, QA, 15 (1926). Proc. Roy. SOC., 108A, IZI (1925). 3 Ind. Eng. Chem., 18, 570 (1926). * Ber., 58, 160(1925). Helv. Chem. Acta, 8, 591 (1925);9, 177, 182 (1926). Ber., 59, 1511(1926). Bull. Acad. polonaise, 1925,A , 93. 8 Bull. Acad. polonaise, 1925 A, I I I. J. Am. Chem. SOC.,47,306 (1925). l a J . Am. Chem. SOC.,48, 2060 (1926). llJ. Am. Chem. Soc., 47, 1299 (1925). ’~Gaae.,55, 252, 619 (1925).

’

I 140

E. EMMET REID

Russell and Taylor’ observe that 10% of thoria increases the activity of a nickel catalyst ten times. “Therefore, it appears impossible to explain the action of the promoter either by a quantitative extension of surface, or by a change in the relative concentrations of the adsorbed reactants. I t seems probable, however, that the surface of the promoted catalyst is such that it holds adsorbed in a reactive condition a larger fraction of the total amount of gas adsorbed than does the unpromoted catalyst.” Casse12 considers it necessary to assume that molecules absorbed on the gas-solid interface have two-dimensional heat motion. “Each of the two reacting gases is preferentially adsorbed on one of the two phases of the solid catalyst.” Kab3 investigated the effect of gases on the activity of platinum in the decomposition of formic acid. Platinum prepared in high vacuum was found to be active. That saturated with hydrogen or carbon monoxide was inactive while that containing oxygen was moderately so. The catalyst could be reactivated by evacuation. ‘ ( T h eassumption made by m a n y that p l a t i n u m without oxygen is not able to effect the catalytic transfer of hydrogen is rendered improbable by this experiment. . . . T h e most active black appears to be one without a n y adsorbed gas. A superficial adsorption of oxygen after disintegration appears to have only the advantage of avoiding loading with the harmful gases, hydrogen] carbon monoxide, etc., and to keep these away until the black enters the reaction; yet it is not necessary for catalytic action and is lost, as the following electromotive experiments show, immediately on contact with hydrogen.” Waldschmidt-Leitz and Seit~~rnaintain that oxygen-containing and oxygenfree platinum black are to be regarded as two different catalysts of which only the first is capable of adding hydrogen to unsaturated compounds. “The influence of the electrical force, emanating from the positive metal and acting in the intermediate layer, on the stability of the hydrogen molecule will make itself felt to different degrees according to the greater or less possession of the metal surface by oxygen; so the significance is seen of the dependence of the specific activation of hydrogen on the oxygen content of the catalyst, as has been observed in the hydrogenation of napthalene.” Kuhn’ takes quite a different view of the action of oxygen. He assumes that: “In the interior of the platinum there are hydrogen nuclei and free electrons. As these leave the platinum the recombination begins passing through excited and normal hydrogen atoms, Hz ions and excited Hz molecules finally to the normal inactive H1. The path of these various recombinations lies on the surface of the platinum and in its immediate vicinity where the substance to be hydrogenated is. On pure platinum and palladium the recombination is so catalyzed that only normal HZ molecules reach the subJ. Phys. Chem., 29, 1325 (1925).

* Naturwiss., 14, 103 (1926). 5

2. physik. Chem., 115, 224 (1925). Ber., 58,563 (1925). Naturwiea., 13, 169 (1925).

FIFTH REPORT OF THE COMMITTEE

os

CONTACT CATALYSIS

1141

strate and no hydrogenation takes place.” If the surface is poisoned by oxygen or air then active hydrogen reaches the substrate. The oxygen is a poison for the catalyst that would destroy the active atomic hydrogen on which the hydrogenation depends. Schmidt’ suggests that the presence of a second metal in a catalyst may serve as a promoter by preventing sintering which would diminish its activity. Wyckoff and Crittenden2made an X-ray examination of some ammonia catalysts. They find that the presence of potassium and aluminium oxides maintains a large iron surface by preventing the growth of iron crystals. Stadnikov, Gavrilov and Vinogradov3 observe that carbon deposited on finely divided iron is extraordinarily active in reductions at relatively low temperature. According to Rideal and Wright4 nitrogen and iron have great influence on the catalytic power of charcoal in the oxidation of oxalic acid. “The incorporation of 10% of urea in sugar and activation of the resultant charcoal by slow combustion until some two-thirds have been removed by oxidation, will produce a charcoal of specific surface of 2 0 0 sq. in. per g. . . The incorporation of these “promoters” in the carbon results in the extension of the total surface and possibly a small extension in the fraction of the surface which is catalytically active. At the same time, two new types of catalytically active surface make their appearance: an iron-carbon-nitrogen complex surface with a specific activity some 800 times that of the original activecarbon surface, and an iron-carbon surface with a specific activity some 50 times that of the original surface.” Sandonnini5has found that carbon is a more active catalyst with copper and that the effect of the two is moderately increased by light. Chirnoaga6 observes that alumina is without action in the decomposition of sodium hypochlorite but that it is a promoter for nickel and cobalt peroxides. Mixtures of these two oxides are more effective than either alone, the maximum being at 3070 nickel. Medreder7 has studied the catalytic oxidation of methane with free oxygen and finds that the addition of 4y0 of hydrogen chloride raised the precentage of methane oxidised from 2 to 27’. With a small amount of hydrogen chloride, formaldehyde is the chief product but much causes the formation of carbon monoxide. The best yield of formaldehyde was with a mixture of the phosphates of tin, iron and aluminium as catalyst and 0.13% hydrogen chloride. Taipale8 observes that the nature of the solvent has much influence on the rate of hydrogenation of amines, acetic acid being the best solvent, ether, methanol and ethanol being poorer. Z physik. Chem., 118, 193 (1925). Chem Soc., 47, 2866 (192j). Ber., 58,242 (192j). J. Am. Chem. Soc., 1926, 1813. Atti Accad. Lincei, (6) 2, 427 (1925). J Chem. Sac., 1926, 1693. Trans. Karpov Inst. Chem., 1925,KO.4, I 17. J RUM.Phys.-Chem. Soc., 57, 487 (1925).

* 3. Am.

1142

E. EMMET REID

Taylor and Close‘ have studied the acid catalysis of lactone formation. The reaction is much faster in met ether than in dry. This looks like a case of the influence of water but they believe that the velocity is not dependent on the solvent but depends “upon the thermodynamic activity of the hydrogen ion which niay correspond to the concentration of the unhydrated ion.” Abderhalden and Iiomm? have investigated the oxidation of glycine anhydride by hydrogen peroxide under the action of sunlight aided by ferrous sulphate. Theeffect of hydrogen ion concentration has been shown by Witzemann3 who finds that the oxidation of butyric acid by hydrogen peroxide is slower in presence of monosodium phosphate than with the disodium salt. Clark, McGrath and Johnson4 have studied the effect of X-rays on the catalytic oxidation of sulphur dioxide. They find no effect in dry air but a slight increase of oxidation in presence of moisture. Smith5 has found the effect of gum arabic and other colloids on the rate of hydrolysis of esters in heterogeneous systems to be quite irregular sometimes in one direction and sometimes in the other. I n homogeneous systems the effects were slight. Bronsted and King6 have investigated salt effects in hydroxyl ion catalysis. Induced Reactions. These may be considered cases of catalysis provided we are not strict in holding to the idea that the catalyst suffers no change. Reference must be made to the admirable review of oxidation catalysis by Moureu and Dufraisse’ for much material that, can not be discussed here. Dey and Dhars oxidise sulphur, sugars, starch and various other substances by passing air at room temperature into solutions containing these and finely divided copper, cuprous chloride or oxide or yellow phosphorus. They regard induced reactions as due to activation by emit’ted ions. Palit and Dharg add a number of observations on the oxidations of carbohydrates in presence of ferrous hydroxide and sodium sulphite. They find that “substances which are difficultly oxidised by passing air can be usually more readily oxidised in presence of reducing agents than in presence of a feeble oxidising agent like ferric salt or cupric salt.” According to Spoehr and Smithlo sodium ferropyrophosphate which is itself readily oxidised by air is a catalyst for the air oxidation of carbohydrates and hydroxy acids by air. Horiuchi” finds that safrol is isomerized by heating with 207~caustic . is not affected by this treatment when alone potash at 1 8 0 ~ - 2 0 0 ~Eugenol but is isomerized if mixed with safrol. With I jY0 alkali at 1 6 0 ~ - 1 7 0neither ~ is changed when alone, but both are if they are mixed together. l

J. Phys. Chem., 29, 1085 (1925).

* Z. physiol. Chem.,

144, 234 (19zj). J. Am. Chem. SOC.,48, 202 (1926).

‘ Proc. Nat. Acad. Sci., 11, 646 (1925). J. Chem. SOC.,127, 2625 (1925). J. Am. Chem. SOC.,47, 2523 (1925).

’ Chem. Rev., 3, 113 (1926). * Z.anorg. allg. Ch., 144,307

(rgz5j.

J. Phys. Chem., 30, 939 (1926).

1oJ. Am.

Chem. SOC.,48,236 (1926).

”J. Chem. SOC.Japan, 45, zog (1925).

FIFTH REPORT OF THE COMMITTEE O S COSTACT CATALYSIS

I

I43

Catalyst Poisons and Negative Catalysts Xaxted‘ finds that the adsorption of lead and mercury ions by a platinum catalyst is linear up to saturation. So is the poisoning curve. “Thus the activity of the catalyst in the presence of such a poison is, at any rate for the first stage, a linear function of the actual concentration of poison on the surface of the catalyst.” He restricts poisoning to heterogeneous systems as he can see no meaning t o this word when applied to a homogeneous system. Kubota and Yoshikawa2 have studied the toxicity of thiophene for nickel of thioand copper catalysts. Sickel is rapidly poisoned at 300’ by phene in benzene yet its activity for hydrogenating phorone was maintained. The rapidity with which nickel is poisoned depends on the temperature at which it is reduced, the higher the temperature, the more rapid the poisoning, Copper is not affected by thiophene. Boswell and Bayley3 consider the normal platinum or nickel catalyst to consist of particles of the metal with interior content of oxygen, the particles being surrpunded by a layer of dissociated water. this layer being the seat of catalytic oxidation and reduction. They say: “Experinients on the poisoning of nickel and platinum catalysts by chlorine seem to indicate that the poisoning is accomplished by the destruction of the surface film on the catalytic particles, which film is the seat of the normal catalytic action, thus rendering the interior oxygen content accessible to free hydrogen. This interior oxygen so vita! to the maintenance of this surface film and hence of catalytic action in the normal catalyst, is thus quickly removed.” Bakj4 has investigated the poisoning of palladium catalysts for the reiction:NaHZP02 H%O-+NaH2PO3 Hz. He thinks the poisoning to be due partly to chemical combination with thc .)alladous chloride and partly to adsorption on the palladium black. For 1 atom of palladium, 1.75 molecules of potassium cyanide, I of mercuric chloide, 1.5 of thiourea and more than I O of quinine hydrochloride were required. Charrion5 dehydrated alcohol at z 50’ with alumina containing various additions. Calcium oxide, phosphoric acid, cobalt and copper oxides, tungsten trioxide and sulphuric acid and mercury inhibit the reaction in the order given. Wieland and Fischer6 have found a curious case. The rate of oxidation of oxalic acid by hydriodic acid is reduced 8 8 7 by 0.0001mole. hydrocyanic acid and is stopped by twice this amount. They regard the hydrocyanic acid as an anticatalyst but the most painstaking search failed to reveal the catalyst. Quartaroli’ observes that one part of hydrogen peroxide in ~oo,ooo,ooo parts of water can be detected by its effect on the color change of cupric hy-

+

+

J. Chem. SOC.,127, 73 (1925); Ind. Chemist, I , 449 (1925). Sci. Papers Inst. Phys. Chem. Res. (Japan), 3, 33 (192s). 3 J. Phys. Chem., 29, 11 (1925). 1 Trans. Karpov. Inst. Chem. (Russia), 1925, So. 4, I I . 5 Compt. rend., 180, 213 (1925). 8 Ber., 59, 1171 (1926). 7 G a m , 55, 264, (1925). 1

2

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E. EMMET REID

droxide. This reaction is hindered by salts of magnesium and other electrolytes even in extremely small amounts. Zetzsche and Xrnd’ have investigated the purification of solvents used in hydrogenation. Lamb and Vail? have made an elaborate study of the effects of different amounts of moisture on Hopcalite and “conclude that the activity of the catalyst is primarily determined by its water content.” Constable3 finds that the temperature coefficient for the dehydrogenation of alcohol remains the same during gradual poisoning. Cusmano4 concludes that camphor does not act as a contact catalyst in the union of sulphur dioxide and chlorine but rather by virtue of the residual valence of its oxygen. Organic compounds containing the groups =CO, -COIH or -0- act as positive catalysts for this reaction but those containing the groups -h’O?, -SO*H or a halogen inhibit the reaction. Goldschmidt and AZathiesenj have made a comparison of the catalytic action and the conductivities of hydrochloric, hydrobromic and hydriodic acids in butyl alcohol and find water to be an anti-catalyst. Toda6 observed the oxidation of cystein by methylene blue. This reaction is inhibited by hydrocyanic acid from which it is supposed to be due to an unknown catalyst. I t is speeded up by a trace of iron. Palit and Dhar’ dissolved metals in nitric acid in presence of catalysts. Reducing agents, except formic acid, retard the solution. According to Gault and Trauffault3 chloroform is readily chlorinated without a catalyst but this is prevented by ferric chloride. The subject of catalysis and auto-oxidation is handled so extensively and in in such a masterly way by Moureu and Dufraisseg that it need not be treated here except for a few references. Moureu, Dufraisse and Lotte’Ofind that “practically all catalyzers of autooxidation (either positive or negative) are easily oxidisable substances.” They usually contain hydroxyl, iodine or sulphur. Methoxyl is inactive. Sulphones and sulphoxides are inactive. Moureu and Dufraisse” regard catalytic oxidation as dependent on the formation of a sort of peroxide. “The antioxygens act by decomposing catalytically the peroxide A [O,] which results from the union of the auto-oxidisable substance with a molecule of free oxygen.” In a later article1*they find nitrogen compounds to be important. Helv. Chem. Acta, 9, 173 (1926). Sac.. 47, 123 (19~j). 3Proc. Camb. Phil. SOC., 22, 738 (1925). G a m , 55, 218 (1925). 2. physik. Chem., 121, I j3 (1926). 6 Biochem. Z., 172, (1926). J. Phys. Chem., 30, 1125(1926). BCornpt.rend., 179,467 (1924). g Chem. Rev., 3, 113 (1926);“ D e u x i h e Cons. Solvay,” 524 (1926). ’OCompt. rend., 180, 993 (1925). *’J.Chem. Sac., 127, I (192j). Wompt. rend., 182, 949 (1926). 1

* J. Am. Chem.

’


I145

Dharl combats Luther’s view2 that a negative catalyst only counteracts a positive. He maintains that negative catalysts are much more numerous than positive. The most of the substances that hinder oxidation are reducing agents. “It is very probable that the phenomena of negative catalysts of induced reactions and of formation of molecular complexes are in close relation to each other.” Underwood3 follows Taylor’s theory that inhibitors function by forming molecular compounds with the active molecules of the substance which is being preserved. The stabilizing actions of 2 5 substances towards chloroform were studied without finding serious disagreement with this theory. Fugitive dyes may be protected against fading in light by treatmenf with a solution of phenol or resorcinol. Granting that the active molecules are thus taken care of, the question arises why are new active molecules not formed? They naturally would be if the assumption is correct that, under given conditions, a certain portion of the molecules is active. Over ahundredinhibiting agents for the oxidation of oils have beenstudied by Smith and Wood4. “The most effective are basic unsaturated compounds as the amines, aromatic phenols, and inorganic basic reducers.” They conclude that : I. “The antioxidant being basic, combines with the acidic products of oxidation and prevents them from acting as auto-catalysts toward oxidation. “The triple-bonded nitrogen atom with two partial valencies or ele2. ments with free valences forms intermediate compounds with the easily oxidised ethenoid carbon. 3. “This temporary compound controls the rate of reaction for a definite, but limited period of time.” Observations on Catalysis Under this head are noted numerous investigations in which the theory of catalysis is not the prominent part. They contain a host of observations of catalytic effects of many classes. On account of their number little space can be given to any one. Dehydration. Adkins and Perkins, Lazier and Adkins, Adkins and Lazier5 have made a comprehensive study of the dehydration and dehydrogenation of alcohols over alumina and zinc oxide. Bonham6 and Kesting’ have passed propyl and ethyl alcohol over alumina, clay, etc. Clark, Graham and Winte?, and Jatkar and Watsong have found alumina and alum to be excellent catalysts for the preparation of ether from the alcohols, while Plusslohas used Z. anorg. allg. Chem., 144, 289 (1925). Chem., 45, 662 (1903). Proc. S a t . Acad. Sci., 11, 78 (192j). Ind. Eng. Chem., 18, 691 (1926). J. Am. Chem. SOC.,47, 1163(192j);47, 1719(192j);48, 1671(1926). J. Am. Pharm. Asso., 14, 114 (1925). 2 . angew. Chern., 38, 362 (1925). * J. Am. Chem. SOC.,47, 2748 (1925). J. SOC.Chem. Ind., 45, 23, 1681(1926). ‘OHelv. Chem. Acta, 8, 507 (192j).

* Z. physik.

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mixtures of alcohols. Takagi and Ishimasal demethylahe veratrol over a Japanese clay, lead sulphate and alum. Mailhe2 decomposes the esters of secondary alcohols over alumina, thoria and titania. Senderens3 has made a comprehensive study of the preparation of ethers from both primary and secondary alcohols by heating the alcohols with sulphuric acid containing some water. The amounts of acid required decrease from 100% by volume for ethyl to 30y0 for heptyl. Less sulphuric acid is required for secondary. The temperatures used are from 100' to 1 4 j O . Excellent yields were obtained. Darrel14 finds 300' to be the optimum temperature for the preparation of ethyl amine from alcohol and ammonia over alumina. The influence of other factors was studied. Cleminson and Briscoej show that the reaction, 2CO - 4 0 s C, does not take place below 300' in clean glass but does as low as 250' in presence of magnesia and alumina. Hydrogenation and dehydrogenation. Pierce and ddams, Heckel and iidams, and Hiers and iidams6have applied hydrogenation with platinum oxide catalyst to various substances. The last article is a comparison of the oxide catalyst with colloidal platinum. Zelinskii and Turowa-Pollak'have compared the activity of platinum, iridium, rhodium, ruthenium, palladium and nickel for the hydrogenation of benzene at 7 temperatures, 100' to 300". Kegoshis has determined the opt,imum conditions for the hydrogenation of acetaldehyde with nickel. Kluyrer and DoukerO have studied the catalytic transfer of hydrogen. Ryerson and Thomas'O have hydrogenated ethylene and phenol with palladium and nickel in silica gel. Kash" has prepared cobalt, copper and manganese catalysts for the hydrogenation of carbon monoxide. Carothers and JonesI2 prepared primary amines by the hydrogenation of nitriles with platinum. Chakravarty and G h ~ s h ' .used ~ nickel on sugar charcoal in the hydrogenation of carbon monoxide. For this reaction at 300°, Franz Fischer, Tropsch and Dilthey14 arrange the catalysts in this order: ruthenium, iridium, rhodium, nickel, cobalt, osmium, platinum, iron, molybdenum, palladium and silver. Zetzsche and ZalaI5 dehydrogenated alcohols in the presence of cupric and cuprous oxides and manganese dioxide. Ghosh and Chakravarty have redetermined the equilibria: Hz+HCHO%CH3OH and HCHO -

+

J. Pharm. Soc. Japan, No. 517, 266 (1925). Caoutchouc and Guttapercha, 22, 12937 (1925). 3 Compt. rend., 180, 790 (1925); 181, 698 (1925); 182, 612 (1926). J. Chem. Soc., 127, 2399 (1926). J. Chem. Soc., 1926, 2148. 6 J. Am. Chem. Soc., 47, 1098 (1925); 47, 1712 (1925); Ber., 59, I62 (1926). Ber., 58, 1298 (1925). 6 Report Osaka Ind. Res. Lab. Japan, 5 , S o . 6, I (1924). eProc. head. Sci. Amsterdam, 28, 605 (1925). 1U"Third Colloid Symposium," 99 (1925 ) . ""Fuel in Science and Practice," 5 , 263 (1926). 125. Am. Chem. Soc., 47, 305 (1925). '?Quarterly J. Ind. Chem. S o c . , 1, I j o (1925). "Brennst. Chem., 6,265 (1925). "Helv. Chem. Acta, 9, 288 (1926). I

2


+

1147

%CO H1and obtained values different from the accepted. Escourron* discusses hydrogenation under reduced pressure. Komatsu and Tanaka3 hydrogenated aniline over copper and nickel. The transfer of hydrogen from one part of a molecule to another is effected by nickel or copper according to Delaby and Dumoulin.4 Vinyl-alkyl carbinols are thus transformed into ethyl ketones. Xckel is used by Mailhe5 to decompose acid chlorides. X1kenes, carbon monoxide and hydrogen chloride are produced. Benzoyl chloride gives chlor-benzene and carbon monoxide. Oxidation. Sinozaki and Hara6 oxidised hydrocyanic acid to nitric oxide over a variety of catalysts. They believe cyanic acid to be first formed. Horiuchi and Cyeda' find that the nitrates of mercury and lead improve the yield in the oxidation of anethol to anisaldehyde. Langs observes the accelerating effect of iodine and iodides on the oxidation of arsenious acid by permanganate. I n the titration of oxalic acid with permanganate, Ridleye believes that both manganese ions and manganese sulphate molecules have catalytic effect. Miscellaneous. Briner, Pluss and Paillardloobtained hexamethyl-benzene from pbenol, etc., and methanol over alumina a t 400'. Wibaut, Diekmann and Rutgers" find bismuth chloride to be a catalyst for the addition of hydrobromic and hydrochloric acids to ethylene and propylene. Tzentnershver and StraumanW observed an increase in the rate of solution of zinc in mineral acids in the presence of various salts. Stadnikov and I v a n ~ v s k i idecompose '~ fatty acids over iron at 400'. Yamaguchi'4 used reduced copperinBeckmann's rearrangement. The same author'5 tried the action of the same catalyst on pinacols. Koraczynski and KierzeklGfind nickel, cobalt and copper powders and cobalt and nickel chlorides t o be catalysts for indol synthesis. The formation of acetals is comprehensively studied by Adanis and Adkins." Ferric and calcium chlorides were used. To be a catalyst the salt must form an alcoholate. The useful salts have acid reaction but the ew-lence is against catalysis by the hydrogen ion. Hara and KomatsuIs pass an al.!ohol Quarterly J. Ind. Chem. SOC.,2, 142 (1925). "Parfums de France," So. 26, 86 (1925.) J M e m . Coll. Sc. Kyoto, 8A, 135 (1925). Compt. rend., 180, 1277 (1925). 5Compt. rend., 180, 111 ( I g z j ) . Tech. Rep. Tohoku I. U., 6, 95 (1926). ' J. Chem. SOC.Japan, 45,203 (1924). 8 Z . anorg. allg. Chem., 152, 197 (1926). Chem., Sews, 130,3 0 j (1925). '@Helv.Acta. Chem., 7, 1046 (1924). l'Proc. Acad. Sci. Amsterdam, 27, 671 (1924). 'IZ. physik. Chem., 118,415 (1925). 13TTrans.Karpov Inst. Chem. (Moscow), 1925,S o . 4, 175. '4Mem. Coll. Sc. Kyoto, 9A, 33 (1925). ':Bull. Chem. Soc. Japan, 1, 64 (1926). '*Gam., 55,361 (1925). "J. Am. Chem. Sac., 47, 1358 (1925). 1sMem. Coll. Sci. Kyoto, SA, 241 (1925). l

1148

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and ammonia over copper at 300' and get a nitrile. Mailhe' decomposes anides a t 400' over nickel, nitriles being formed in some cases. Oparinaz condensed isovaleric aldehyde with ammonia over alumina a t 360' and obtained pyridine derivatives. Ipatiev and Xlinkvin3 condensed ethylene in an iron vessel with alumina. 0. and C. A. Silberrad and Parke4have chlorinated toluene with several agents in presence of a large number of catalysts which they have classified. Nelson and Engelder5 and Westcott and Engelder have decomposed formic acid by passing over copper, alumina, thoria and nickel at various temperatures from 3 jo" up. Simonss has determined the decomposition products of triacetin and tripropionin over thoria at joo'. Blumbergn' has studied the decomposition of diazo compounds by copper, Paal and Boeters8 have perfected the preparation of collodial cobalt. The sulphonation of anthraquinone in presence of mercury has been thoroughly investigated by cop pen^.^ Zelinskiilo has shown that 1-pinene passed over palladium asbestos gives a mixture of cymene and dihydropinene which shows that change of structure may be effected by a hydrogenation catalyst. Richter and Wolf" isomerize beta pinene to the alpha by shaking wit palladium containing hydrogen. If the metal does not contain hydrogen there is no change. According to Favorsky and Mlle. Chilingaren12 certair ketones are isomerized by zinc chloride at 3 joo. Veil13has found .hat the activity of nickel 1 ixide in the decomposition of hydrogen peroxide changes greatly with time jd use. The magnetic properties were roughly parallel to the catalyti. activity. The following five papers pertain to homogeneous solutions. Livingstone and BrayI4 have studied the chlorine-chloride decomposition of hydrogen peroxide, Hammick, Hutchinson and Snell find the rate of oxidation of formic acid by bromine to be inversely proportional to the concentration of the hydrogen ion. The reaction is between the formate ion and molecular bromine. Euler and OlandeP have investigated the splitting of acetoacetic ester by acids and bases and find that the rates are not proportional to hydrogen and hydroxyl ions. Brbnsted and Duus16 have used this reaction to Bull.. 37, 1394 (1925). J. Russ. Physik. Chem. SOC.,57, 319 (1925) 3 Khim. Promushlennost', 3, 57 (1925). J. Chem. SOC., 127, 1724 (1925). J. Phys. Chem. 30, 470, 476 (1926). 6 J. Am. Chem. SOC., 48, 1991 (1926). Chem. Weekblad., 22, 599 (1925). Ber., 58, I 542 (1925). Rec. Trav. chim., 44, 907 (1925). 'OBer., 58, 864 (1925). "Ber., 59, 1733 (1925). W o m p t . rend., 182, 221 (1926). I3Compt. rend., 180, 932 (1925). 14J.Am. Chem. SOC.,47, 2069 (1925). l5Z. anorg. allg. Chem., 147, 295 (1925). leZ. physik. Chem., 117, 299 (192j).

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FIFTH REPORT O F THE COMMITTEE O S CONTACT CATALYSIS

1149

estimate the concentration of the base present. Kilpatrick' has carried on reactions in buffer solutions, and discussed salt effect. As compared with the multitude of studies in which the regulation inorganic catalysts appear we have far too few investigations that involve organic catalysts. This is all the more true when we consider that the bodily functions of animals, growth and decay are probably initiated and regulated by organic substances which may be classed as catalysts. Morgulis, Beber and Rabkin2 find catalase to be most active in the decomposition of hydrogen peroxide between o and IO'. According to WatsonS the acid bromide is a powerful catalyst in the bromination of an aliphatic acid. Krasonskii and Kiprinov4 condense phenyl-acetylene to s-triphenylbenzene using a primary amine as catalyst. Secondary and tertiary do not cause this reaction. 1

J. Am. Chem. SOC., 48, 2091 (1926).

* J. Biol. Chem., 68,521 (1926). a 4

J. Chem. SOC., 127,2067 (1925). J. Rum. Phys. Chem. SOC., 56, I, (1925).

FIFTH REPORT OF THE COMMITTEE ON COYTACT CATALYSIS

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